Thermodynamic Data for the Speciation and Solubility of Pd, Pb, … · 2004. 12. 21. · JNC TN8400 99-011 JP9955327 42 Thermodynamic Data for the Speciation and Solubility of Pd,

JNC TN8400 99-011 JP9955327 42

Thermodynamic Data for the Speciation and Solubility

of Pd, Pb, Sn, Sb, Nb and Bi in Aqueous Solution

January, 1999

3 3 0 0 2 9 3 6 .

JAPAN NUCLEAR CYCLE DEVELOPMENT INSTITUTE

3 1 - 0 9

319-1194

Inquiries about copyright and reproduction should be addressed to :Techinical Information SectionAdministration Division Tokai WorksJapan Nuclear Cycle Development Institute4-33 Muramatsu, Naka-gun, Ibaraki 319-1194,Japan.

JNC TN8400 9 9 - 0 1 1

January,1999

Thermodynamic Data for the Speciation and Solubility

of Pd, Pb, Sn, Sb, Nb and Bi in Aqueous Solution

Barbara Lothenbach**, Michael Ochs**, Hans Wanner*** and Mikazu Yui*

Abstract

This report provides thermodynamic data for predicting concentrations of

palladium Pd, lead Pb, tin Sn, antimony Sb, niobium Nb and bismuth Bi in

geologic environments, and contributes to an integration of the JNC chemical

thermodynamic database, JNC-TDB (previously PNC-TDB), for the performance

analysis of geological isolation system of high-level radioactive wastes. Besides

treating hydrolysis in detail, this report focuses on the formation of complexes or

compounds with chloride, fluoride, carbonate, nitrate, sulfate and phosphate.

Other important inorganic ligands (sulfide for lead and antimony, ammonia for

palladium) are also included. In this study, the specific ion interaction theory

(SIT) approach is used to extrapolate thermodynamic constants to zero ionic

strength at 25 °C.

*: Waste Isolation Research Division, Tokai works, Japan Nuclear Cycle Development Institute (JNC)

**: BMG ENGINEERING Ltd., Switzerland

**: HSK, Switzerland

JNC TN8400 9 9 - 0 1 1

1999 £fU£

Pd, Pb,Sn,Sb, Bi

Barbara Lothenbach**, Michael Ochs**,

Hans Wanner***, M # H f P *

PNC-TDB) ^ f O — l i t L t i l L / : ,

, - * 7* (Nb) iJ «t O* H'X-7^

mtzo ttz, &

(10

£&(Pb), ^ X ( S n ) ,

interaction theory ) , 25°C, -<

r { i , SIT (specific ion

0 Hi3tt£-*libtf>!tftfc L

**: BMG ENGINEERING Ltd. (X -f 7,)

***: HSK ( X - f X )

> ^ -

JNC TN8400 9 9 - 0 1 1

Preface

The PNC-TDB project was initiated by PNC in 1996 with the aim to establish, by March 31,1998, an internationally acknowledged equilibrium database for 20 elements that were definedas potentially important by PNC for the safety of a nuclear waste repository. I accepted tocoordinate the project at an international level, and Dr. Dhanpat Rai (Battelle PNNL, Richland)as well as Dr. Gregory Choppin (Florida State University, Tallahassee) accepted to participateas experts. Their task was the setting-up of an equilibrium database on the actinide elements.The tedious job of administrative project management, including the contracting business withall participants was assumed by Dr. Michael Ochs (BMG Engineering, Schlieren, Zurich). Inthe 2nd year of this project, he and his colleague Dr. Barbara Lothenbach also were put incharge to develop the chemical thermodynamics of some non-actinide elements.

The kick-off meeting was held at Tokai-mura, May 27-30, 1996, in the presence of theJapanese experts: Dr. O. Tochiyama, Dr. H. Moriyama, Dr. S. Nakayama and Dr. T.Yamaguchi. Nevertheless, the various contracts were not in place until October 1996. A secondmeeting involving only PNC, BMG and myself was held at BMG, Zurich, November 18 - 20,a third meeting was held at Richland, Washington, May 27 - 29, 1997, without participation ofthe Japanese experts, and a fourth one on October 29, 1997, on the occasion of the Migration'97 conference at Sendai, again with the presence of the Japanese experts.

It was clear from the beginning, and stated explicitly at the kick-off meeting, that the timescaleof the project (1.5 years) was extremely ambitious. Obviously, it was impossible to achieve, forthe final product, a quality level that would be comparable to that of the NEA-TDB. Wenevertheless decided to choose a procedure that resembles the one of NEA-TDB at least in thebasic principles. This will allow immediate improvements of the pre-selected data in a possiblefollow-up project.

PNC decided from the start to accept different procedures for the establishment of the actinidedatabase on one hand, and of the non-actinide databases on the other hand. For the actinidecompounds and complexes of the +III and +IV oxidation states PNC preferred theestablishment of a complete Pitzer model to the adoption of the NEA database. Less weight wasput on the development of databases for penta- and hexavalent actinide species. The motivationbehind this decision was PNC's conviction that only the +III and +IV oxidation states of theactinides will be relevant in their performance assessment analyses due to their choice of near-field components and geological sites.

For the non-actinide elements, PNC preferred initially to have its own staff perform a review ofthe available literature and establish selected data sets, while the members of the expert teamshould review these data sets and make constructive comments in such a way that the PNC staffcould then implement the improvements and thus accomplish the various data sets according tointernational standards. The state-of-the-art analysis of the project during the second meeting inNovember 1996 revealed that this procedure was unlikely to result in satisfactory data sets bythe final deadline of March 31, 1998. It was then decided to split the non-actinide elements upinto two sets, one to be treated by BMG according to the guidelines of the NEA-TDB, the otherto be treated by PNC staff according to their own procedures.

in

JNC TN8400 9 9 - 0 1 1

The tasks that had to be accomplished by the PNC-TDB project team were the following:

D. Rai/PNNL Database on actinide compounds and complexes in the +III and +IVoxidation states

G. Choppin/FSU Database on actinide compounds and complexes in the +V and +VIoxidation states, plus the actinide redox potentials

M. Ochs/BMG Database on the elements Bi, Nb, Pb, Pd, Sb and Sn

As I have indicated above, the data sets presented in this report cannot be considered tocorrespond to the quality level of the NEA-TDB due to the reasons outlined below. However, Iconsider them as an excellent, partly high-quality basis that will require some quality-checkingwith independent experimental literature in the case of the actinide models, and a more detailedreview of some of the reported literature in the other cases.

My personal assessment of the quality of the database presented in this report and the reports byD. Rai, G. Choppin and co-workers, and of the improvements and refinements that may berequired to obtain ,,final", high-quality data sets, is briefly summarized below:

Actinide elements:The actinide +III model is based on the extensive experimental experience of D. Rai with thechemical behavior of lanthanides and actinides. The evidence for analogous treatment of the +IIIoxidation states is compelling and convincing. A complete Pitzer parameter set is provided. It isimportant that the Pitzer parameters be used in modeling studies, because in some cases (e.g.,the sulfate complexes) complex formation is expressed by Pitzer parameters rather thanequilibrium constants. This data set has, to my knowledge, not yet been checked againstindependent experimental studies from other laboratories. This task, which was not part of thepresent project but would improve the credibility of the data set, is strongly recommended forthe near future.

For the +IV oxidation state of the actinides, the analogy concept is less evident than for the +IIIoxidation state. The recommendation to cross-check the data against independent experiments isalso valid here.

The reason why the +V and +VI oxidation states of the actinides have been treated in a lesssophisticated way is due to the low weight PNC has assigned to them. Without expressing anycriticism of the selected values themselves, I feel that the review procedure here is notsufficiently transparent to cope with international standards. These data are based on expertjudgment rather than on detailed analysis of the available experimental papers. The samecomment applies to the redox potentials. However, due to the low weight PNC had assigned tothis part of the project, the corresponding budget was quite limited, and a detailed analysis ofthese chemical systems could not be carried out.

Non-actinide elements (Bi, Nb, Pb, Pd, Sb and Sn):For these elements all the experimental data have been compiled from the available literature.Detailed reviews of the experimental papers have in general not been performed due to time and

IV

JNC TN8400 9 9 - 0 1 1

budget constraints, an exception being the hydrolysis papers of the palladium review. Theprocedure chosen is consistent with the guidelines of the NEA-TDB: Experimental data are usedto perform SIT plots (SIT = Specific Ion Interaction Theory, cf. NEA-TDB). In cases wherethe experimental data from different sources show satisfactory agreement on the SIT plot, anextrapolation to zero ionic strength may provide confidence in the resulting thermodynamicconstants. It will nevertheless be necessary to review the experimental papers in detail in orderto achieve international quality standards for the selected data sets. In addition, important gapshave been identified in many cases, and experimental programs should be performed to providethe information required for credible predictive modeling.

As a conclusion, I believe that the objectives of the PNC-TDB project, as defined at the kick-offmeeting and re-defined at the second meeting in June 1997, have been reached. It is somewhatunfortunate that part of the database comes with complete Pitzer parameter sets, others with SITparameter sets, and yet others with no ion interaction parameters at all. However, as I havementioned in the beginning of this Preface, the procedure that unavoidably had to lead to thisdifference in activity factor treatment was consciously chosen by PNC. In my view, thedatabase presented here provides a good basis for further refinement. It could therefore carrythe attribute ,,Phase I". I strongly recommend to follow the above suggestions for independentverification on one hand, and for more detailed review on the other hand, in addition to thepossible definition of experimental programs to fill the gaps that are considered critical forperformance assessment purposes.

Wiirenlingen, May 18, 1998 Hans WannerScientific Project Coordinator

JNC TN8400 9 9 - 0 1 1

Executive Summary

PNC is planning to submit a new Performance Assessment Report by March 2000. Besidesinformations on geology, repository design etc., basic data on radionuclide behavior arerequired for the safety analysis. These basic data include chemical thermodynamic data inaqueous media, requiring the establishment of a chemical thermodynamic database, PNC-TDB,for the interaction of key radionuclides with significant ligands under relevant conditions.

The subject of the present report are the key elements:

• tin• antimony• lead• bismuth• niobium• palladium

For these elements, element-specific datasets have been developed based exclusively onexperimental studies published in the literature, rather than relying on existing compilations.This report focuses on the formation of complexes or compounds with:

• hydroxide• chloride• fluoride• carbonate• nitrate• sulfate• phosphate

Other important inorganic ligands (sulfide for lead and antimony, ammonia in the case ofpalladium) are also included. The number of experimental studies varies from one element to theother, and the data set selected here cannot in general be considered 'complete' for geochemicalmodeling applications. Where data for different ionic strengths are available, the specific ioninteraction equation (SIT) is used to extrapolate formation constants to zero ionic strength. Theselected formation constants are listed in Section 3 of this report.

Sections 2 and 3 of this report contain

• the data selection criteria• an outline of the extrapolation procedure to I = 0 with the specific ion interaction equation

(SIT)• the selected data for the key elements Sn, Sb, Pb, Bi, Nb, Pd

In the following Sections 4 - 9 , the data selection for the individual key elements is discussed indetail. For each element, tables containing the complete data compilation as well as allcalculations used in the data evaluation for the individual elements, are made available to PNC.

VI

JNC TN8400 99-011

Table of contents

1 Introduction 1

1.1 Background 1

1.2 Organization of the JNC-TDB project 1

2 Standards and Conventions 3

2.1 Symbols, units, notations and conversion factors 3

2.1.1 Symbols and notation 32.1.2 Compilation of thermodynamic data 42.1.3 Phase designators 52.1.4 Physical constants 52.1.5 Equilibrium constants 52.1.6 Redox reactions 7

2.2 Data selection criteria 7

2.3 Ionic strength corrections 7

2.4 Auxiliary Data 82.4.1 Selected thermodynamic data for auxiliary species 9

2.4.2 Conversion of AfG°values to equilibrium constants 9

3 Selected data : 12

3.1 Selected data for Sn, Sb, Pb, Bi, Nb andPd 12

3.2 Gaps and uncertainties 223.2.1 Uncertainties 22

3.2.2 Gaps 22

4 Tin 27

4.1 Hydrolysis of tin(TV) 27

4.1.1 Hydrolysis of tin(IV) under acidic conditions: Sn4+, SnOH3+,Sn(OH)2

2+and Sn(OH)3+ 28

4.1.2 Hydrolysis of tin(IV) under neutral and alkaline conditions:Sn(OH)5-and Sn(OH)6

2" 294.2 Solid tin(IV) oxides/hydroxides 32

4.2.1 Freshly precipitated Sn(OH)4(am) 334.2.2 SnO2(precip) 33

Vll

J N C T N 8 4 0 0 99 - O i l

4.2.3 SnO2(cassiterite) 344.2.4 Additional data compiled for the formation of tin(IV)

hydroxide/oxide compounds 35

4.3 Other tin(TV) complexes and compounds 36

4.4 Redox reactions 40

4.4.1 Sn2+/Sn(cr) 404.4.2 Sn2+/Sn4+ 414.4.3 Sn(OH)4° 424.4.4 Additional data compiled for the tin redox system 43

4.5 Hydrolysis of tin(II) 45

4.5.1 SnOH+ 474.5.2 Sn(OH)2° 484.5.3 Sn(OH)3" 494.5.4 Sn3(OH)4

2+ 504.5.5 Sn2(OH)2

2+ 504.5.6 MSn(OH)3+ 504.5.7 Additional equilibrium data compiled for Sn(II) hydrolysis 50

4.6 Solid tin(H) oxide/hydroxide 54

4.6.1 SnO(cr) and Sn(OH)2(precip) 554.6.2 Additional equilibrium data compiled for tin(II) hydroxide/oxide

compounds 56

4.7 Tin(II) chloride system 58

4.7.1 SnCl+, SnCl2°, SnCl3-, and SnCl42- 60

4.7.2 SnOHCl0 : 624.7.3 SnCl2(s) 634.7.4 SnOHCl(s) 634.7.5 Additional equilibrium data compiled for the tin(II) chloride

system 64

4.8 Tin(II) fluoride system 68

4.8.1 SnF+, SnF2°, andSnF3- 694.8.2 SnF2(s) 714.8.3 Additional equilibrium data compiled for the tin(II) fluoride

system 71

4.9 Tin(II) carbonate system 73

4.10 Tin(II) nitrate system 74

4.10.1 SnNO3+, Sn(NO3)2°, Sn(NO3)3-and Sn(NO3)4

2- 754.10.2 Tin(II) nitrate compounds 774.10.3 Additional equilibrium data compiled for tin(II) nitrate system 77

4.11 Tin(II) phosphate system 79

via

JNC T N 8 4 0 0 99 - O i l

4.12 Tin(II) sulfate system 80

4.12.1 Tin(n) sulfate complexes 80

4.12.2 SnSO4(s) 81

4.13 Tin(II) sulfide system 82

4.13.1 SnS(herzenbergite) 824.13.2 Sn2S3(s) and Sn3S4(s) 82

4.14 Comments on selected references 83

5 Antimony 87

5.1 Hydrolysis of antimony (III) 87

5.1.1 Sb3+ 895.1.2 SbOH2+ 905.1.3 Sb(OH)2

+ 905.1.4 Sb(OH)4- 925.1.5 Sb2(OH)4

2+ 935.1.6 Sb2(OH)6° 935.1.7 Additional equilibrium data compiled for the hydrolysis of

antimony(m) 93

5.2 Solid antimony(m)-oxide/hydroxide 95

5.2.1 ct-Sb2O3 (valentinite) 965.2.2 £-Sb2O3 (senarmontite) 97

5.3 Antimony(in) chloride system 99

5.3.1 Reactions of Sb3+: SbCl2+,.SbCl2+, SbCl3°, SbCLr, SbCl52-

andSbCl63- 100

5.3.2 SbCl4~ and SbOHCl3- 1015.3.3 SbCl3(s) and SbOCl(s) (or Sb4O5Cl2(s)) 1025.3.4 Additional equilibrium data compiled for the antimony (III)

chloride system 102

5.4 Antimony(ITf) fluoride system 104

5.4.1 Reactions of Sb3+:SbF2+, SbF 2 \ SbF3o, and SbF4- 1045.4.2 Reactions of SbF3°:SbF4-, and SbF3OH- 1055.4.3 SbOF0 or Sb(OH)2F° 1055.4.4 SbF3(s) 105

5.5 Sb(III) sulfate system 107

5.5.1 SbOSO4- 1075.5.2 Sb2(SO4)3(s) 107

5.6 Antimony(III) sulfide system 108

5.6.1 Sb2S42-, HSb2S4-, and H2Sb2S4

0 1085.6.2 Sb2S3(stibnite) 109

IX

JNC TN8400 9 9 - 0 1 1

5.7 Redox reactions 110

5.7.1 Sb(cr)/Sb(OH)3° I l l5.7.2 Sb(in)/Sb(V) .....1125.7.3 Sb(cr)/Sb(OH)5° 1135.7.4 Additional data compiled for the redox potential of antimony 113

5.8 Hydrolysis of antimony(V) 115

5.9 Sb2O5(precip) 118

5.10 Sb2O4(cr)andSb6O13(cr) 120

5.11 Other antimony(V) complexes and compounds 122


6 Lead 126

6.1 Hydrolysis of lead 126

6.1.1 PbOH+ 1306.1.2 Pb(OH)2° 1316.1.3 Pb(OH)3- 1326.1.4 Pb(OH)4

2- 1326.1.5 Pb2OH3+ 1336.1.6 Pb4(OH)4

4+ 1346.1.7 Pb3(OH)4

2+ 1356.1.8 Pb3(OH)3

3+and Pb3(OH)5+ 1366.1.9 Pb6O(OH)6

4+ , 1376.1.10 Additional equilibrium data compiled for the lead hydroxide

system 138

6.2 Solid lead-oxide/hydroxide phases 144

6.2.1 PbO(litharge) and PbO(massicot) 1446.2.2 Precipitated lead hydroxide 1456.2.3 Pb(OH)2(s) 1456.2.4 Additional data for lead hydroxide/oxide compounds 145

6.3 Lead chloride system 149

6.3.1 Lead chloride complexes 1496.3.2 PbCl2(s) 1566.3.3 Mendipite: Pb2(OH)3Cl(cr) or Pb4(OH)6Cl2(cr) 1566.3.4 Laurionite and paralaurionite 1576.3.5 Additional equilibrium data compiled for the lead chloride system 158

6.4 Lead fluoride system 163

6.4.1 PbF+andPbF2° 1646.4.2 PbF2(s) 165

JNC TN8400 9 9 - 0 1 1

6.5 Mixed lead fluoride chloride system 168

6.5.1 PbFCl0 169

6.5.2 Matlockite (PbClF(cr)) 169

6.6 Lead carbonate system 171

6.6.1 PbCO3° 1736.6.2 Pb(CO3)2

2- 1746.6.3 Pb(CO3)3

4-andPb(CO3)46- 1756.6.4 PbCO3OH- 1756.6.5 PbHCO3

+, Pb(HCO3)2° and Pb(HCO3)3- 1756.6.6 PbCO3(cerrusite) 1766.6.7 Pb3(CO3)2(OH)2(hydrocerrusite) 1776.6.8 Plumbonacrite 1786.6.9 PbCO3PbO(s) and PbCO3(PbO)2(s) 1786.6.10 Phosgenite 1786.6.11 Additional data for the lead carbonate system 179

6.7 Lead nitrate system 184

6.7.1 Lead nitrate complexes 1866.7.2 Pb(NO3)2(s) and PbOHNO3(cr) 1886.7.3 Additional data compiled for the lead nitrate system 188

6.8 Lead phosphate system 191

6.8.1 PbH2PO4+and PbHPO4° 192

6.8.2 PbHPO4(cr) and Pb3(PO4)2(s) 1926.8.3 Pb(H2PO4)2(s), Pb4(PO4)2O(s) and plumbogummite

(PbAl3(PO4)2(OH)5(cr)) 1936.8.4 Pyromorphites: Pb5(PO4)3Cl(s), Pb5(PO4)3F(s), and

Pb5(PO4)3OH(s) 1936.8.5 Pb10(PO4)6(OH)2(hydroxylapatite) 1936.8.6 Additional data compiled for the lead phosphate system 194

6.9 Lead sulfate system 197

6.9.1 Lead sulfate complexes 1986.9.2 PbSO4(anglesite) 1996.9.3 Hinsdalite: PbAl3PO4SO4(OH)6(cr) 200

6.10 Lead sulfide system 204

6.10.1 Lead sulfide complexes 2056.10.2 Galena (PbS) 2056.10.3 Additional data compiled for the lead sulfide system 206

6.11 Redox equilibria 208

6.11.1 Pb2+/Pb(cr) 2086.11.2 Pb2+/Pb4+ 2086.11.3 PbO2andPb3O4 2086.11.4 Data compiled for the lead redox system 209

XI

JNC TN8400 9 9 - 0 1 1

6.12 Lead(IV) 211


7 Bismuth 219

7.1 Hydrolysis of bismuth 219

7.1.1 BiOH2+ 2227.1.2 Bi(OH)2

+ 2237.1.3 Bi(OH)3° 2247.1.4 Bi(OH)4- 2257.1.5 Bi6(OH)i26+ 2267.1.6 Bi9(OH)2o7+, Bi9(OH)2i6+, and Bi9(OH)22

5+ 2277.1.7 Bi3(OH)4

5+ 2277.1.8 Bi6(OH)15

3+ 2277.1.9 Additional equilibrium data compiled for the bismuth hydroxide

system 228

7.2 Solid bismuth-oxide/hydroxide 232

7.2.1 a-Bi2O3(cr) 2327.2.2 Precipitated Bi(OH)3(s) 2337.2.3 Additional data compiled for solid bismuth oxides/hydroxides 233

7.3 Bismuth chloride system 236

7.3.1 Bismuth chloride complexes 2367.3.2 BiOCl(s) and Bi(OH)2Cl(s) 2427.3.3 BiCl3(s) 2437.3.4 Additional equilibrium data compiled for the bismuth chloride

system 243

7.4 Bismuth perchlorate 247

7.4.1 BiC1042+ or Bi(H2O)6

3+ C1O4- 2477.4.2 BiOClO4(precip) 247

7.5 Bismuth fluoride system 248

7.6 Bismuth carbonate system 250

7.7 Bismuth nitrate system 251

7.7.1 Bismuth nitrate complexes 2517.7.2 BiONO3(s) 2557.7.3 Additional equilibrium constants compiled for the bismuth(III)

nitrate system 256

7.8 Mixed bismuth nitrate and chloride system 259

7.8.1 Bismuth chloride nitrate complexes 259

xn

JNC TN8400 9 9 - 0 1 1

7.8.2 Additional equilibrium data compiled for the bismuth(III) chloridenitrate system 263

7.9 Bismuth phosphate system 264

7.10 Bismuth sulfate system 265

7.10.1 Bismuth sulfate complexes 2657.10.2 Bi2(SO4)3(s) 2657.10.3 Equilibrium data compiled for the bismuth sulfate system 265

7.11 BP+/Bi(cr) 267


8 Niobium 275

8.1 Hydrolysis of niobium(V) 275

8.1.1 Monomeric Nb(V) species 2758.1.2 Polymeric Nb(V) species 277

8.1.2.1 Protonation at pH > 8: N^Ch 98-, Nb6Oi 9H7-,

Nb6O19H26-, and Nb6Oi 9H35- 277

8.1.2.2 Very alkaline solutions: N b ^ 2(OH)48-, Nb4O[6

12-and NbO2(OH)43- 280

8.1.2.3 Neutral and acidic solutions: H^Nb] 2C>366~,H5Nb120367-, H4Nb12O368-, andH3Nb12O369- 280

8.1.3 Additional equilibrium data compiled for the niobium(V)hydroxide system ' 280

8.2 Solubility of solid niobium pentoxide ..- 283

8.2.1 Additional equilibrium data compiled for Nb2C>5(s) 285

8.3 Solid niobium phases: redox equilibria 286

8.4 Other niobium(V) complexes and compounds 287


9 Palladium 290

9.1 The redox pair Pd27Pd(s) 290

9.2 Hydrolysis of palladium(n) 292

9.2.1 Hydrolysis of palladium(II) 2929.2.2 Additional equilibrium data compiled for the hydrolysis of

palladium(II) 293

9.3 Solid palladium(II)-oxide/hydroxide 296

9.3.1 Pd(OH)2 (precip) 296

X l l l

JNC TN8400 9 9 - 0 1 1

9.3.2 PdO(cr) 296

9.4 Chloride complexes of palladium(II) 298

9.4.1 Chloride complexes 2989.4.2 Additional equilibrium data compiled for palladium chloride

complexes 299

9.5 Amino complexes of palladium(II) 302

9.6 Conclusions 302

9.7 Comments on selected references: 303

10 References 309

XIV

JNC TN8400 9 9 - 0 1 1

1 Introduction

1.1 Background

JNC is planning to submit a new Performance Assessment Report by March 2000. Besidesinformations on geology, repository design etc., basic data on radionuclide behavior arerequired for the safety analysis. These basic data include chemical thermodynamic data inaqueous media, requiring the establishment of a chemical thermodynamic database, PNC-TDB,for the interaction of key radionuclides with significant ligands under relevant conditions.

JNC has defined two groups of key elements for the next Performance Assessment Report.They are termed '1st priority elements' and '2nd priority elements', respectively, and arecomprised of the following elements:

• 1st priority elements: Pu, U, Np, Th, Am, Ra, Sn, Zr, Ni, Pd, Tc, Se, Pa, Cm and Cs.• 2nd priority elements: Sm, Ac, Po, Pb, Nb, Bi and Sb.

Ligands considered by JNC include hydroxide, carbonate, chloride, fluoride, sulfate andphosphate. In addition, solid-solution modeling may be considered for Ra, and possibly otherelements. Since the aqueous chemistry of Cs is very simple, this element is not treated withinthe JNC-TDB development.

1.2 Organization of the JNC-TDB project

The JNC-TDB development project involves JNC staff members, an external expert group, anda Japanese expert advisory group. The expert group consists of Dr. H. Wanner, HSK; Dr. D.Rai, Batelle PNNL; Prof. G. Choppin, Florida State University; and Dr. M. Ochs, BMG. Theroles and responsibilities of each member of the expert group are defined in Table 1.1. TheJapanese expert advisory group comprises Prof. Tochiyama, Tohoku University, Prof.Moriyama, Kyoto University, and Dr. Nakayama, JAERI. The treatment of elements by thedifferent groups is detailed in Table 1.1.

JNC TN8400 9 9 - 0 1 1

Table 1.1: Assigned responsibilities for the JNC-TDB project for the complete 2-year period ofFY96-97:

Elements Responsibility Timeframe

Pu,Pd,Sm:

Ra,

Ac,Pb,

Zr,

Th,Sb,

, SeTc,

u,Bi,

Po

Np,Nb,

Pa, Am, CmSn

PNNL iBMG2

JNC (review byJNC (review by

PNNL 1)BMG2)

FY96-97FY97FY96-97FY96-97

1 This includes Dr. D. Rai and his co-workers at Batelle PNNL, as well as Dr. G.Choppin and his co-workers at Florida State University

2 This includes Dr. M. Ochs, Dr. B. Lothenbach, and other personnel at BMGEngineering Ltd, with Dr. H. Wanner functioning as external advisor.

The key elements Sn, Sb, Pb, Bi, Nb, Pd are the subject of the present report. For theseelements, element-specific datasets are developed in this report based exclusively onexperimental studies published in the literature, rather than relying on existing compilations.

Seperate reports on actinide elements and Ac are being prepared by Dr. D. Rai and co-workers.The elements Sm, Ni, Se, Ra, Zr, Tc, and Po are treated by JNC.

JNC TN8400 9 9 - 0 1 1

2 Standards and Conventions

2.1 Symbols, units, notations and conversion factors

2.1.1 Symbols and notation

The symbols for physical and chemical quantities used in this report are summarized in Table2 . 1 .

Table 2.1: Symbols and notation

AfG° the standard molar Gibbs energy of formation [kJ mol"] ]

ArG° the molar Gibbs energy of a reaction [kJ mol"1]

AfH° the standard molar enthalpy of formation [kJ mol"1]

S° the standard molar entropy [J K"1 m o F ]

T absolute temperature [K]

R molar gas constant (8.3145 J K"1 m o F ]

n number of electrons involved in a redox reaction

CB concentration of a solute B in [mol/L]

me concentration of a solute B in [mol/kg solvent]

YB activity coefficient of a substance B

I ionic strength [mol/L]

I m ionic strength [mol/kg solvent]

p Factor for the conversion of molarity, CB, to molality, m s , of substanceB [dm3 solution per kg H2O]

VB stoichiometric coefficient of a substance B (negative for reactants,positive for products)

J N C T N 8 4 0 0 99 - O i l

2.1.2 Compilation of thermodynamic data

In the Sections 4 -9 of this report, thermodynamic data are compiled in tables. The followingconventions have been used throughout:

'Reference': The references are ordered chronologically and alphabetically by the first twoauthors within each year. A more detailed description is given in Chapter II. 1.8in [1992GRE/FUG]. The references are given in Section 3 of this report.

'Comments': T indicates the temperature (K) to which the given constant refers. Manycomments were imported directly from the NEA database.I indicates the conditions under which the constant was determined, e.g., I =0.1-1.

T : I indicates the ionic strength to which the given constant refers. This value isoften, but not always, identical with the ionic strength given under 'Comments'.

'Medium': indicates, where available, the electrolyte in which the given constant ismeasured or refers to.

'Method': Abbreviations for the method of measurement are listed in Table 2.2.

Table 2.2: Abbreviations for experimental methods

cat = cation exchangecol = colorimetric analysiscon = conductivity measurementsel = electrophoresisemf = electromotive force measurements at high temperaturesextr = extractionfe = fluoride selective electrodekin = kinetic measurementsn/a = method not known to the reviewersNMR= "FNMRpol = polarographypot = potentiometryse = sulfide electrodeSO2 = pSO2 measurements at high temperaturessol = solubility measurementssp = spectrophotometry, NMRtit = titration (evaluation of equilibrium through pH only)

JNC TN8400 9 9 - 0 1 1

2.1.3 Phase designators

Chemical formulae may refer to different chemical species and are often required to be specifiedmore clearly in order to avoid ambiguities. For example, PbCC>3 occurs as a solid and asaqueous complex. The distinction between the different phases is made by phase designatorsthat immediately follow the chemical formula and appear in parentheses. The only formulae thatare not provided with a phase designator in this report are aqueous ions. The use of the phasedesignators is described below:

The designator (1) is used for pure liquid substances, e.g.,

The designators ('name'), (cr), (precip) and (s) is used for solid substances. When the solidhas a common name, ('name') is used, e.g. PbCO3(cerrusite). (cr) is used when it is knownthat the solid is crystalline, e.g. PbOHCl(cr). (precip) is used when it is known that the solidwas precipitated from solution. Otherwise, where no such information is available, (s) isused.

For aqueous species no designators are used, e.g. PbOH+ or PbCO3°.

2.1.4 Physical constants

The fundamental physical constants are taken from [1992GRE/FUG] and are listed in Table2.3.

Table 2.3: Fundamental physical constants. These values were taken from [1992GRE/FUG].

R molar gas constant 8.3145 J R-1 mol"1

F Faraday constant 96 485 C moW

2.1.5 Equilibrium constants

The IUPAC has not explicitly defined the symbols and terminology for equilibrium constants ofreactions in aqueous solutions. In this report the conventions (based on the work of Baes andMesmer, [1976BAE/MES]) given below have been used throughout.

Formation of an aqueous complex:

JNC TN8400 9 9 - 0 1 1

• Formation of a solid:

Conventionally, equilibrium constants involving a solid are denoted as 'solubility constants'rather than as formation constants of the solid. An index 's ' to the equilibrium constantindicates that the constant refers to a solubility process, as shown below:

MmLn(s)^mM + nL KSo=[M}m[L]n

Kso is the conventional solubility product and the subscript '0' indicates that the equilibriumreaction involves only uncomplexed aqueous species.

In this report, the formation of solid is treated analogously to the formation of aqueousspecies. I.e., the inverse solubility product is used, denoted with an asterisk:

mM + nL<^> MmLn(s) K*So =

Further notations used in this report are compiled in Table 2.4.

Table 2.4: Reactions

P measured cumulative formation constant for a reaction1 (in molarityunits)

(3m measured cumulative formation constant for a reaction corrected from

molarity to molaliry units

(3° cumulative equilibrium constant valid at I = 0

P b measured cumulative equilibrium constant for a reaction involving O H ~

instead of H + l.

K consecutive (stepwise) equilibrium constant for a reaction

Kso Solubility product of a solid 2.\Pb2+]

Example: PbO(cr) + 2H+ <=» Pb2+ + H2O; Kso = l -^

K*so Formation constant of a solid 2.

Example: Pb2+ + H2O <=> PbO(cr) + 2H+; K so = }—=S[Pb2+]

Kbso Solubility product of a solid involving OH~ instead of H+ '.

K*bso Formation constant of a solid involving OH~ instead of H+ ].

' Hydrolysis reactions are formulated in this report using H+ as component and not OH".2 In this report, reactions involving a solid are normally formulated as formation reactions.

JNC TN8400 99-011

2.1.6 Redox reactions

Redox reactions are usually quantified in terms of their electrode (half cell) potential E. E isidentical to the electromotive force (emf) of a redox reaction that involves the standard hydrogenelectrode as an electron donor or acceptor. In this review, electrode potentials are given asreduction potentials relative to the standard hydrogen electrode which acts as an electron donor.The standard redox potential, E°, is related to the change of the Gibbs energy ArG° and theequilibrium constant K as outlined below:

(at 298.15 K)ArGnF nF

2.2 Data selection criteria

To assure a traceable data selection a list of selection criteria is given below. All studies selectedin this review should fulfill the following criteria.

• Experimental study• I = constant• No complex formation with electrolyte (or corresponding correction possible)• All relevant species included (e.g. polymers)• Experimental details reported• T - 2 9 8 K

For more detailed information of the individual papers and additional criteria, see discussion ofthe individual species in Section 4 - 9 of this report.

2.3 Ionic strength corrections

Thermodynamic data always refer to a selected standard state. The standard state for a solute Bin a solution is a hypothetical solution at the standard state pressure (0.1 MPa) and the standardtemperature (298.15 K) in which me = mo = 1 mol/kg, and in which the activity coefficient JQis unity. However, for many reactions, measurements cannot be made accurately (or at all) indilute solutions from which the extrapolation to the standard state would be simple. In thisreport, thermodynamic data were extrapolated to the standard state (1=0) using the specific ioninteraction equation (SIT) as described in [1992GRE/FUG]. An extensive description of thismethod and its use can be found in Appendix B of [1992GRE/FUG] and [1995SIL/BID].Shortly, the correction consists of an extended Debye-Hiickel expression, in which the activitycoefficients of the reactants and products depend only on the ionic charge of the reactants and

JNC TN8400 99 - Oil

the ionic strength of the solution, but it accounts for the medium specific properties byintroducing ion pairing between the medium ions and the species involved in the equilibriumreactions. The correction of the measured data to 1=0 is made with the following equation:

Debye - Hueckel term

1 + 1.5X J/.~\lAm

Im = / xp

where I ionic strength [mol/L]Im ionic strength [mol/kg solvent]p Factor for the conversion of molarity, CB, to molality, me, of substance B [dm3

solution per kg H2O]P measured cumulative formation constant for a reaction expressed in molarity

units|3m measured cumulative formation constant for a reaction corrected from molarity

to molality unitsXv sum of the stoichiometric coefficients of the reaction.

The factors for the conversion of molarity, CB, to molality, mg, of a substance B for thedifferent electrolytes at 298.15 K were taken from Table II.5 in [1995SIL/BID] and from TableII-l in [1976BAE/MES].

2.4 Auxiliary Data

In this section the thermodynamic data for auxiliary compounds and complexes are compiled.Many of the these auxiliary species are used in the evaluation of the recommended data given inSection 2. It is therefore essential to always use these auxiliary data in conjunction with theselected data. The use of other auxiliary data can lead to inconsistencies and erroneous results.Table 2.5 contains the selected thermodynamic data of the auxiliary species considered in thisreview (all taken from the recent NEA publication of [1995SIL/BID]).

JNC TN8400 9 9 - 0 1 1

2.4.1 Selected thermodynamic data for auxiliary species

Table 2.5: Selected thermodynamic data for auxiliary species taken from [1995SIL/BID]

complex AfG° [kJ/mol] Reference

ci-F-HF°SO4

2-HS-H2S°H2S(g)NO3-CO32-HCO3-CO2°CO2(g)PO43-HPO42-H2PCVH2OCa2+

Na+K+

-131.217-281.523-299.675-744.004

12.243-27.648-33.443

-110.794-527.899-586.845-385.970-394.373

-1025.491-1095.985-1137.152

-237.140-552.806-261.953-282.510

[1995SIL/BID][1995SEL/BID][1995SIL/BID][1995SEL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SDL/BID][1995SIL/BID][1995SIL/BID][1995SIL/BID][1995SEL/BID][1995SIL/BID]

2.4.2 Conversion of AjG ° values to equilibrium constants

Experimental papers measure and indicate in most cases log (3 values. In a few cases the resultsof experimental studies are given as AfG° values. To be able to compare these values with thelog B or log K values selected in this report, these AfG° values were converted to B or K valuesaccording to

Arc°

using the AfG° values of the auxiliary species (Table 2.5) and the AfG° values of the masterspecies Sn2+, Sn(OH)4°, Sb(OH)3°, Sb(OH)5°, Pb2+, Bi3+ selected in this report (Table 2.6).In many compilations, which were only used for comparison in this report, AfG° values arecompiled instead of log p or log K values (e.g. in [1982WAG/EVA]). In general, these AfG°values were converted to log P or log K values using the AfG° values of the auxiliary species(Tables 2.5) and the AfG° values of the master species Sn2+, Sn(OH)4°, Sb(OH)3°, Sb(OH)5°,

JNC TN8400 9 9 - 0 1 1

Pb2+, Bi3+, Pd2+ (Table 2.6) selected in this report. However, some compilations also reportAfG° values for the master species (Table 2.7). In that case the AfG° values given in therespective compilation is used.

Table 2.6:

complex

Thermodynamic data for the master species Sn2+, Sn(OH)4°, Sb(OH)3°,Sb(OH)5°, Pb2+, Bi3+, Pd2+ selected in this report. For Nb(OH)50 nothermodynamic data was selected.

AfG° [kJ/mol] Reference

Sn2+

Sn(OH)4°Sb(OH)3°Sb(OH)5°Pb2+

Bi3+

Pd2+

1 tentative value

Table 2.7: Ther

-26.42-944.16-643.0-992.62-24.2495.55187.6 *

modvnami

this reportthis reportthis reportthis report[1989COX/WAG]this reportthis report

Thermodynamic data for the master species Sn2+, Sn(OH)4°, Sb(OH)3°,Sb(OH)5°, Pb2+, Bi3+, Pd2+, given in previous compilations

complex AfG° [kJ/mol] Reference

Sn2+

Sn(OH)4°

Sb(OH)3°

-26.29-27.89-29.89-27.2-27.14-27.2-27.23-27.62

-934.8-950.6

-647.16-644.8-644.66-643.63-643.9

[1952LAT][1978COD][1980BEN/TEA][1982WAG/EVA][1985BAB/MAT][1985GAL][1988PHI/HAL][1989COX/WAG]

[1984KEL/HOU][1988PHI/HAL]

[1985BAB/MAT][1985PAS][1986ITA/NIS][1990SHI/ZOT][1994AKI/ZOT]

10

JNC TN8400 9 9 - 0 1 1

Table 2.7: continued

Pb2+

Pb4+

Bi3+

Nb(OH)5°

HNbO30

Pd2+

-24.24-24.34-24.44-24.34-24.00-24.4-24.43-24.06-24.39-24.34-24.43-24.42-24.42-24.4-24.42-24.34-24.23-24.00

302.43-302.57302.44

82.7982.8091.7991.82

-1448-1448-985.3

190.34176.56177.37176.53176.5176.53176.5176.47

[1989COX/WAG][1952LAT][1969HEL][1977PAU][1978COD][1978ROB/HEM2][1980BEN/TEA][1981HEL/KIR][1981STU/MOR][1982PAU][1982WAG/EVA][1983LAN][1983SAN/BAR][1984 VIE/TAR][1985BAB/MAT][1985GAL][1985RAI/RYA][1988PHI/HAL]

[1952LAT][1983LAN][1985GAL]

[1968 ROB AVAL][1982WAG/EVA][1985BAB/MAT][1985LOV/MEK]

[1985UDU/VEN][1982WAG/EVA][1985BAB/MAT]

[1952LAT][1967IZA/EAT][1968GOL/HEP][1980BEN/TEA][1982WAG/EVA][1985BAB/MAT][1985COL][1988PHI/HAL]

11

J N C T N 8 4 0 0 99 - O i l

3 Selected data

This chapter presents the thermodynamic data set for bismuth, niobium, lead, palladium,antimony, and tin species selected in this review. The Tables 3.1 to 3.6 contain therecommended thermodynamic constants of chemical equilibrium reactions by which bismuth,niobium, lead, palladium, antimony, and tin compounds and complexes are formed. Besidestreating hydrolysis in detail, this review focuses on the formation of complexes or compoundswith chloride, fluoride, carbonate, nitrate, sulfate and phosphate. Other important inorganicligands (sulfide for lead and antimony, ammonia in the case of palladium) are also included.The present report does not include any compounds or complexes containing organic ligands.

It should also be noted that the data set presented in this section may not be 'complete' for allconceivable systems and conditions. Gaps and uncertainties are pointed out in Section 2.2 andin various paragraphs in the Sections 4 -9.

3.1 Selected data for Sn, Sb, Pb, Bi, Nb and Pd

The Tables 3.1 to 3.6 contain the recommended thermodynamic data and should be read asfollows:

Each element has at least one master species whose identity is given in the heading of eachtable. Redox sensitive elements may have more than one master species, allowing to separatethe modeling of the different oxidation states, e. g. in case of kinetic inhibition. Each tablecontains the formation constants of all the complexes and solids proposed in this report.

Example for reading the contents of the database:

• Species SnOH+ (see Table 3.1) is composed of the following components: 1 x componentSn2+, and -1 x component H+. Addition of 1 x H2O1 gives the correct formation reactionfor SnOH+

Sn2++ H2O(1) = SnOH++ H+

with a formation constant log P of -3.75.

1 H2O has an activity of 1 in aqueous solutions according to the conventions used and is therefore not included in

the tables.

12

JNC TN8400 9 9 - 0 1 1

Table 3.1: Selected thermodynamic data for reactions involving tin compounds and reactionsas selected in Section 4 of this report. All ionic species listed in this table areaqueous species. The data refer to the reference temperature of 298 K and to thestandard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinitedilution (1=0).

components =>species

SnOH+

Sn(OH)2«

Sn(OH)3"

Sn3(OH)42+

SnCl+

SnCl2°SnCl3'SnOHCl0

SnF+

SnF2°SnFf

SnNO3+

Sn(NO3)2»

Sn(NO3)3-

Sn(NO3)42"

SnSO4°Sn(SO4)2

2"

Sn(OH)2(precip)

SnO(cr)

SnOHCl(s)

Sn(cr)

Sn4+

Sn(OH)4°

logP

-3.75

-7.71

-17.54

-6.51 1

1.652.312.09-2.27

4.46 1

7.74 19.61 i

1.251.741.370.30 i

2.91 i

2.83 1

-2.82

-2.41

2.42

-4.63

- 5 1

-5.4 i

Sn 2 +

1

1

1

3

111

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

H+

-1

-2

-3

-4

-1

*

-2

-2

-1

-4

ci-

123

1

1

F-

1

23

NO3- SO42"

1

2

3

4

1

2

e~

2

-2

-21 tentative values

13

JNC TN8400 9 9 - 0 1 1



Sn4+

Sn(OH)5"

Sn(OH)62"

SnO2(precip)

SnO2(cassiterite)

logP

0.4 i-7.97

-18.40

7.468.0

Sn(OH)4°

1

11

11

4

-1-2

00

tentative value (see Section 4.4.3); values for other protonated species are missing.

14

J N C T N 8 4 0 0 9 9 - O i l

Table 3.2: Selected thermodynamic data for reactions involving antimony compounds andreactions as selected in Section 5 of this report All ionic species listed in this tableare aqueous species. The data refer to the reference temperature of 298 K and tothe standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinitedilution (1=0).


11

Sb3+

SbOH2+

Sb(OH)2+

Sb(OH)4-

Sb2(OH)6°

SbCl2+

SbCl2+

SbF2+

SbF2+

SbF3°

Sb2S42"

HSb2S4~

H2Sb2S4°

S b2 0 , (valentinite)

Sb2S,(stibnite)

Sb(cr)Sb(OH)5

logP

-0.73 i0.83 1

1.30

-11.93

0.08 i

2.78]

3.27'

6.48'

12.65 i

18.36'

42.53

52.1857.00

8.72

55.14

11.99-21.84

Sb(OH)3°

1

1

1

1

2

1

1

1

1

1

2

22

2

2

11

H+

321

-1

033333234

0'3

3-2

ci-

1

2

F-

1

2

3

HS~

444

3

e~

3-2

1 tentative values

15

JNC TN8400 9 9 - 0 1 1



II

Sb(OH)6"

Sb12(OH)644"

Sb12(OH)655-

Sb12(OH)666-

Sb12(OH)677"

Sb2O5(precip)

logP

-2.7220.3416.7211.896.07

7.40

Sb(OH)s°

1

12

12

12

12

2

H+

-1

-4

-5

-6

-7

0

16

JNC TN8400 9 9 - 0 1 1

Table 3.3: Selected thermodynamic data for reactions involving lead compounds andreactions as selected in Section 6 of this report. All ionic species listed in this tableare aqueous species. The data refer to the reference temperature of 298 K and tothe standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinitedilution (1=0).

components =$species

11

PbOH+

Pb(OH)2°Pb(OH)3"Pb2OH3+

Pb4(OH)44+

Pb3(OH)42+

Pb3(OH)5+

Pb6(OH)84+

PbCl+

PbCl2°

PbCl3"

PbCl42"

PbFPbF2°PbFCl0

PbCO3°Pb(CO3)2

2-PbNO3

+

Pb(NO3)2°Pb(NO3)3"PbHPO4°PbH2PO4

+

PbSO4°Pb(SO4)2

2"Pb(HS)2°Pb(HS)3"

logp

-7.51-16.95-28.02-7.18-20.63-22.48-30.72-42.68

1.552.002.011.352.273.013.557.3010.131.061.480.76 i

15.45 i21.05 i

2.822.37 i

12.34 i13.59 '

p b 2 +

1

112433611111111111111111

1

-1-2-3-1-4-4-5-8

12

ci-

1234

1

F- CO 32 - NO3-

121

12

123

PO43-SO4

2" HS-

11

12

23

17

JNC TN8400 9 9 - 0 1 1


components => Pb 2 + H+ Cl~ F~ CO32" NO 3 - P O 4

3 - S O 42 - HS~ e~

species log p

PbO(red, litharge)PbO(yellow, massicot)Pb(OH)2(precip)PbCl2(s)PbOHCl(cr)PbF2(s)PbFCl(matlockite)PbCO3 (cerrusite)Pb3(CO3)2(OH)2

(hydrocerrusite)Pb10(CO3)6(OH)6O(plumbonacrite)PbOHNO3(cr)PbHPO4(s)Pb3(PO4)2(s)Pb4(PO4)2O(s)Pb(H2PO4)2(s)Pb5(PO4)3OH(hydroxy pyromorphite)Pb5(PO4)3Cl(chloro pyromorphite)Pb5(PO4)3F(fluoro pyromorphite)PbSO4(anglesite)PbS (galena)

-12.68-12.96-13.054.81-0.627.528.8213.2317.64 1

41.21 l

-2.94 123.78 144.40 137.09 148.94 162.80 1

84.40 1

71.60 1

7.8112.17

111111113

10

113415

5

5

11

-2-2-2

-1

-2

-8

-11

-24-1

-1

21

1

1

21

1

12

6

12223

Pb(cr) -4.25 1 2

PbO2(s) -48.98 ] 1 -4 -2Pb3O4(s) -70.98 J 3 -8 -21 tentative values

JNC TN8400 9 9 - 0 1 1

Table 3.4: Selected thermodynamic data for reactions involving bismuth compounds andreactions as selected in Section 7 of this report. All ionic species listed in this tableare aqueous species. The data refer to the reference temperature of 298 K and tothe standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinitedilution (1=0).

components =>spec ie s

BiOH2+

Bi(OH)2+

Bi(OH)3

Bi(OH)4-

Bi6(OH)126+

Bi9(OH)207+

Bi9(OH)216+

Bi9(OH)225+

Bi3(OH)45+

BiCl2+

BiCl2+

BiCl3

BiCl4"BiCl5

2"BiCl6

3"BiNO3

2+

Bi(NO3)2+

Bi(NO3)3

Bi(NO3)4-

BiClNO3+

BiCl(NO3)2°BiCl2NO3°BiCI2(NO3)2-BiCl3NO3"

a-Bi2O3(cr)

BiOCl(s)(BiO)2CO3(cr)(BiO)4(OH)2CO3(cr)BiONO3(s)

Bi(cr)

logP

-0.92-2.56-5.31-18.71

1.34-1.36-3.25-4.86-0.80 13.655.857.629.068.33 i7.64 l

1.972.953.623.095.165.286.865.758.09

-0.768.47

14.27 18.68 i2.75

16.74

Bi 3 +

111169993111111111111111

21241

1

H+

-1-2-3-4

-12-20-21-22-4

-6-2-4

-10-2

ci-

123456

11223

1

CO32" NO3- e-

123412121

11

1

3! tentative values

19

JNC TN8400 99 - Oil

Table 3.5: Selected thermodynamic data for reactions involving niobium compounds andreactions as selected in Section 8 of this report. All ionic species listed in this tableare aqueous species. The data refer to the reference temperature of 298 K and tothe standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinitedilution (1=0).

components => Nb(OH)5° H+

species log P_J

Nb(OH)6- -6.6 1 -1

Nb2O5(s) 16.0 i 21 tentative value (see Section 8.2)

Nb(0H)5° and Nb(OH)6 are hypothetical species (see Section 8.2).

20

JNC TN8400 9 9 - 0 1 1

Table 3.6: Selected thermodynamic data for reactions involving palladium compounds andreactions as selected in Section 9 of this report. All ionic species listed in this tableare aqueous species. The data refer to the reference temperature of 298 K and tothe standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinitedilution (1=0).


pdcrPdCl2°PdCl3"PdCl4

2-PdCl3OH2~PdCl2(OH)2

2-Pd(NH3)2+

Pd(NH3)22+

Pd(NH3)32+

Pd(NH3)42+

Pd(cr)

log [3

5.18.310.911.72.5-7.0 i9.618.526.032.8

32.9 '

Pd2+ H+

11111 -11 -21111

1

cr

l23432

NH3

1234

e~

21 tentative values

21

JNC TN8400 9 9 - 0 1 1

3.2 Gaps and uncertainties

3.2.1 Uncertainties

Some authors report the standard deviations of their experimentally determined formationconstants, but these do not represent the quality of the reported values in absolute terms. It isimportant not to confuse the statistical standard deviation with the accuracy. The standarddeviation is calculated from the dispersion of equally weighted points, while the accuracyreflects the reliability and reproducibility of an experimental value and also includes all kinds ofsystematic errors. The estimation of systematic errors of a method is difficult and can only bemade by a person who is familiar with the experimental method. In many cases thedetermination of the standard deviations is not possible because either only one or two datapoints are available, or the authors did not report the individual values. In this report no attemptwas made to calculate systematically the uncertainties connected to experimentally determinedformation constants. For further indication of the reliability of different methods see discussionsfor the individual species in Sections 4 - 9 .

3.2.2 Gaps

For the user it is important to consider that the selected data set presented in Section 3.1 may notbe 'complete' with respect to all conceivable systems and conditions; there are gaps in theinformation, particularly concerning complex formation of tin(FV), antimony and niobium withinorganic ligands as well as the hydrolysis and the solubility of the oxides of niobium andpalladium. For each individual key element discussed in this report, gaps are listed below.While some missing data may not be important from a practical point of view (e.g. the exactsolubility of a very soluble solid, e.g. SnCl2(s); or the formation constants for complexes thatwill only be important in very concentrated solutions, e.g. the formation of tin(IV) nitrates),other missing data (e.g. the formation of solids with sulfide) may be more important. Gapsconsidered important by the authors of this review are printed in bold.

The gaps also are shortly discussed in the respective sections of the key elements in this report.This information may be used as a basis for the assignment of research priorities.

22

JNC TN8400 9 9 - 0 1 1

Tin(II): Gaps in the tin(II) thermodynamic database. Gaps considered important by theauthors of this review are printed in bold.

complex or solid Comments

SnCl2(s) very solubleSnF2(s) solubleSn(NC>3)2(s) very solubleSnSC>4(s) moderately solubleSnCO3(s), SnCO3° and Sn(CO3)2

2-Sn3(PO4)2(s)SnSO4(s)SnS(s)

Tin(IV): Gaps in the tin(IV) thermodynamic database. Gaps considered important by theauthors of this review are printed in bold.


- chloride complexes and solids '- fluoride complexes and solids l

- nitrate complexes and solids J l

- sulfate complexes and solids 1

- Sn 4 + , SnOH3+, Sn(OH)22+, Sn(OH)3

+

- carbonate complexes and solidsphosphate complexes and solidssulfide complexes and solids- redox equilibrium

Based on the data listed in Tables 4.7 to 4.9, it is our feeling that the complex formation ofSn(IV) with these ligands is rather weak as compared to the hydrolysis of Sn(IV).

23

J N C T N 8 4 0 0 99 - O i l

Antimony(III): Gaps in the antimony(III) thermodynamic database. Gaps considered importantby the authors of this review are printed in bold.


sulfate complexes and solids weak complexesnitrate complexes and solids no information availablecarbonate complexes and solids no information availablephosphate complexes and solids no information available

Antimony(V): Gaps in the antimony(V) thermodynamic database. Gaps considered importantby the authors of this review are printed in bold.


- Sb5+, SbOH4+, Sb(OH)23+, Sb(OH)3

2+, or Sb(OH)4+ stable at pH < 1

- chloride complexes and solids weak complexes- fluoride complexes and solids ?- nitrate complexes and solids ?- sulfate complexes and solids ?- carbonate complexes and solids ?- phosphate complexes and solids * ?- sulfide complexes and solids ?

24

JNC TN8400 9 9 - 0 1 1

Lead: Gaps in the lead thermodynamic database. Gaps considered important by theauthors of this review are printed in bold.


- PbC>2 (s), Pb3O4(s) very oxidizing conditions necessary- Pb(IV) hydrolysis very oxidizing conditions necessary

more information needed:- hydrocerrusite- phosphate complexes and solids

Bismuth: Gaps in the bismuth thermodynamic database. Gaps considered important by theauthors of this review are printed in bold.


B1CI3 (s) very solublefluoride and sulfate complexesBiF3 (s)

J ?BiPO4(s)Bi2(SO4)3(s)Bi2S3(s)

25

JNC TN8400 9 9 - 0 1 1

Niobium: Gaps in the niobium thermodynamic database. Gaps considered important by theauthors of this review are printed in bold.


chloride complexes and solidsfluoride complexes and solidsnitrate complexes and solidssulfate complexes and solidscarbonate complexes and solidsphosphate complexes and solidssulfide complexes and solidsniobium hydrolysis in the neutral and acidic pH range,

formation of polynuclear complexesunder alkaline conditions ?

Palladium: Gaps in the palladium thermodynamic database. Gaps considered important bythe authors of this review are printed in bold.


- fluoride complexes and solidsnitrate complexes and solidssulfate complexes and solidscarbonate complexes and solids

- phosphate complexes and solidssulfide complexes and solids- hydrolysis

- Pd (OH) 2 (s)

weak complexesweak complexes?

probably not important inaqueous solutions

at very low Pd(II) cone (> lO1 0 M) toprevent formation of polymers

26

JNC TN8400 99 - O i l

4 Tin

Tin exists in several oxidation states from -IV to +IV in its compounds, but only the +11 and+IV states are important in its aqueous chemistry. The +IV state is particularly stable in naturalenvironments, and cassiterite SnC>2(cr), is the major source of this element in nature[1976BAE/MES, 1985GAL, 1995WIB]. Little quantitative information exists on the solubilityof SnC>2(cr), but it is rather low (< 10"8 M) in pH region 1 to 7. The stannous ion (Sn2+) iseasily oxidized, and SnO(s) is much more soluble than SnC>2(s). The redox reactions of Sn(0)to Sn(II) and Sn(IV) are discussed in Section 4.4.

Reliable thermodynamic data concerning the complexes formation of Sn(OH)4° with inorganicligands are quite limited. Data for the formation of complexes and solids of tin(IV) with thefollowing inorganic ligands can be found in the literature: water, chloride, fluoride, phosphate,carbonate, sulfate and sulfide. Sufficient experimental data, however, to calculate equilibriumconstants were only available for the hydrolysis of tin(IV) and for the solubility products ofSnO2(precip) and SnO2(cr) (Sections 4.1 and 4.2).

For the formation of Sn2+ complexes or compounds with water, chloride, fluoride, nitrate,phosphate, sulfate and sulfide thermodynamic data are available. Equilibrium constants for thehydrolysis of tin(II) and the complex formation with chloride, fluoride, nitrate, and sulfate arecalculated from experimental data and are given in the Sections 4.5 to 4.12. Also the solubilityproducts of Sn(OH)2(precip), SnO(cr), and SnOHCl(s), are calculated from experimental datagiven in the literature (Sections 4.6 and 4.7).

4.1 Hydrolysis of tin(IV)

The knowledge on tin(IV) hydrolysis is limited. Amaya and coworkers [1997AMA/CHI,1998ODA/AMA] measured the solubility of Sn(IV) in dilute NaC104 solutions between pH 2and 13.5. They observed an increase of the solubility of Sn(IV) above pH = 7, while itremained constant at = 10"8 M between pH 2 - 7 . This indicates the predominance of theuncharged Sn(0H)4° species. In strong acid or base tin(IV) solubility increases. As thehydrolysis of tin(IV) in acidic medium is not well defined, all log p values given in thefollowing paragraphs refer to the Sn(OH)4° complex.

The data used for calculating the formation constants of tin(IV) hydroxide complexes arecompiled in Table 4.1. Additional data for the tin(TV) hydroxide system data which were notchosen for the calculation of log (3° values in this report are compiled in Table 4.2 and 4.3.

27

JNC TN8400 9 9 - 0 1 1

Table 4.1: Experimentally determined equilibrium data compiled for the hydrolysis ofSn(OH)4°, according to the equilibrium: mSn(OH)4° + nH2O «=> Snm(OH)4m+n-n

+ nH*. These data were chosen for the evaluation of recommended values in thepresent report. Additional information see Section 4.14: 'Comments to selectedreferences'. Method: sol = solubility measurements.

log 6m 4m+n Reference Comments I (M) Medium Method

log (3li5: Sn(OH)4° + H2O <=> Sn(OH)f + H+

-1.15 i [1998ODA/AMA] T= 298.15 K, 1=0.1 0.1 NaClO^ sol

log P1>6: Sn(OH)4° + 2H2O <=> Sn(OH)62- + 2H+

-17.74 * [1998ODA/AMA] T= 298.15 K, 1=0.1 0.1 NaC104 sol1 recalculated by [1998ODA/AMA] based on the experiments of [1997AMA/CHI]

4.1.1 Hydrolysis oftin(IV) under acidic conditions: Sn4+, SnOH3+, Sn(OH)22+ and Sn(OH)3+

[1958JOH/KRA] and [1959JOH/KRA] have established, using ultracentrifugation, that Sn(IV)exists in the form of monomeric species in acidic and basic solutions. The only study carried outat 298 K is by [1971NAZ/ANT] (Table 4.2) who studied the hydrolysis of tin(IV) under acidicconditions in 1 M KNO3. [1971NAZ/ANT] used a spectrophotometric method in which thecompetition between the hydrolysis reaction and tomplexation with salicylfluorone wasmeasured. The log (3 values obtained by [1971NAZ/ANT] for the protonation of Sn(OH)4° toSn(OH)3+, Sn(OH)2

2+, Sn(OH)3+, and Sn4+ are given in Table 4.2. The total Sn(IV)concentration in these experiments was 10-5 M. Considering the possible interaction of Sn(IV)with the nitrate ions of the electrolyte solution, these log 6 values can be considered asestimates, but not as exact values for the stability constants of Sn(IV) hydrolysis at I = 1.Nevertheless, based on the observation of [1971NAZ/ANT] and [1997AMA/CHI] it is clear thatprotonation of Sn(OH)4° will occur at pH < 2. [1934HUE/TAR] observed the predominance ofthe Sn4+ cation in solution containing > 0.5 M HC1. This indicates that the protonation ofSn(OH)4° to Sn4+ takes place at slightly higher pH values than indicated by the data of[1971NAZ/ANT] (Table 4.2). The only other experimental values determined for Sn(OH)4°protonation is by [1970KUR/BAR] at 100 °C. In contrast to the observations at 25 °C,protonation of the Sn(OH)4° ion was observed by [1970KUR/BAR] already at a pH of 7.

From the different redox equilibria (cf. Section 4.4.3) a tentative log K° value of 0.40 can becalculated for the reaction Sn(OH)4° + 4H+ <=> Sn4+ + 4H2O. For the other hydrolyzed Sn(IV)complexes it is not possible to propose any reliable thermodynamic values based on theexperimental data available.

28

JNC TN8400 99-011

4.1.2 Hydrolysis oftin(IV) under neutral and alkaline conditions: Sn(OH)f and Sn(OH)(,2-

Hydrolysis of tin(IV) at 25 °C was investigated by [1970BAR/KLI] and Amaya and co-workers[1997AMA/CHI, 1998ODA/AMA]. In contrast to [1997AMA/CHI], who proposed theformation of Sn(OH)5~ and Sn(OH)g2" under alkaline conditions, [1970BAR/KLI] proposedonly the formation of Sn(OH)5-. However, [1970BAR/KLI] calculated pH from the amount ofNaOH added initially. These calculated pH values are probably significantly larger than the realpH in solution, making their results difficult to interpret.

Based on solubility measurements, Amaya and co-workers [1997AMA/CHI, 1998ODA/AMA]proposed the formation of Sn(OH)5~ and Sn(OH)62" under alkaline conditions. The log P15 andPi56 values given by [1998ODA/AMA] for I = 0.1 (these values were based on the experimentsof [1997AMA/CHI], later recalculated by [1998ODA/AMA]) are given in Table 4.1.Extrapolation to 1=0 was made using SIT (see Section 2). Based on the interaction coefficientsgiven in Appendix B.3 of [1992GRE/FUG] an As near or equal 0 can be assumed for bothreactions.

Sn(OH)4° + H2O ^ Sn(OH)5- + H+ log p l i5° = -7.97, As = 0Sn(OH)4° + 2H2O ^ Sn(OH)6

2" + 2H+ log Pi'6° =-18.40, Ae = 0

Some compilations [1984HOU/KEL, 1985GAL, 1992PEA/BER] list thermodynamic data forthe species SnO32~. This is only a different notation for Sn(OH)62\

Table 4.2: Additional, experimentally determined equilibrium data compiled for the hydrolysis of Sn(OH)4°,according to the equilibrium: mSn(OH)4° + nH2O <=> Snm(OH)4m+n-

n + nH+. These data werenot chosen in the present report for the evaluation of recommended stability values. Reasons fornot selecting these references are given at the end of the table or in Section 4.14. Method: sol =solubility measurements, sp = spectrophotfometry, and pot = potentiometry.

log (3m,4m+n Reference

log P!X

0.87 2

log Pi,,

21.82 ]

1.44 2

log Pi,:

14.691.55 2

,.- Sn(OH)4° + 4H+ <=* Sn4+ +

[1971NAZ/ANT1

,: Sn(OH)4° + 3H+ <=> SnOH3

[1970KUR/BAR][1971NAZ/ANT1

Comments

4H2O

T= 298.15 K, 1=1

+ + 3H2O

T=373 K, I=dilT= 298.15 K, 1=1

,.• Sn(OH)4° + 2H+ <=> Sn(OH)22+ + 2H2O

1 [1970KUR/BAR][1971NAZ/ANTJ

T=373 K, I=dilT= 298.15 K, 1=1

I(M)

1

1

1

Medium

KNO,

NaOHKNO,

NaOHKNO,

Method

sp

solsp

solsp

log Sn(OH)4° + H+ Sn(OH)3+ + H2O

7.54 '1.22 2

2.50 •2.06 3

[1970KUR/BAR][1971NAZ/ANT][1981DAD/SOR][1981DAD/SOR1

T=373 K, I=dilT= 298.15 K, 1=1T= 573 K, I=n/aT= 298 K, I=n/a

NaOH1 KNO3

self mediumself medium

solspsolsol

29

JNC TN8400 9 9 - 0 1 1


log Pi,5-

-12.4 4

-9.06 >-7.65 5

-7.97 6

Sn(OH)4° + H2O <=> Sn(OH)5- + H+

[1970BAR/KLI][1970KUR/BAR][1997AMA/CHI]fl998ODA/AMAl

T=298.15K, 1=0.2-2.5T=373 K, I=dilT= 298.15 K, 1=0.1T= 298.15 K, 1=0.1

00

NaOHNaOH

NaC104

solsolsol

log Pi,

-18.30-20.77-20.43-17.31-18.38

6: Sn(OH)4° + 2H2O <=> Sn(<

1 [1970KUR/BAR]1 [1973KLI/BAR]7 [1973GAB/SRI]5 [1997AMA/CHI]

[1998ODA/AMA]

OHtf- + 2H+

T= 373 K, I=dilT= 473 K, I=dilT= 298.15 K, 1=0.25T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1

0.2500

NaOHself

NaOHNaC104

solsolpotsol

temperature in the experiments not 298 Knitrate (used as electrolyte solution) forms complexes with tin(IV)calculated from values determined at 200-400 °Clinear extrapolated by [1970BAR/KLI] from different I. Individual values not given.corrected to 1=0 with Davies equation by [1997AMA7CHI] using the wrong sign for the correction,these values havebeen corrected by [1998ODA/AMA].corrected to 1=0 with Davies equation by [1998ODA/AMA]. For log[3 values at 1=0.1, see Table 4.1estimated from potentiometric data assuming a log Kw of -13.76, log K (Sn(OH)4/Sn(cr)) of -0.77 (at 1=0.25)

Table 4.3: Thermodynamic data for the hydrolysis of Sn(OH)4° taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log (3m,4m+n Reference Comments I (M) Medium

log Pi,0-

0.31 '-1.022.41 2

2.611.271.531.731.940.86 3

1.26 ]

0.31 1

-0.32-0.64

log Pi,;:

3.101.792.072.331.411.43 3

1.59-0.18

Sn(OH)4° + 4H+ 4* Sn4+ +

[1952LAT][1955DEL/ZOU][197OKUR/BAR][1973KLI/BAR][1979VAS/GLA][1979VAS/GLA][1979VAS/GLA][1984HOU/KEL][1984HOU/KEL][1985BAB/MAT][1985GAL][1987BROAVAN]fl988PHI/HAL]

Sn(OH)4° + 3H+ <=? SnOH3

[1973KLI/BAR][1979VAS/GLA][ 1979V AS/GLA][1979VAS/GLA][1984HOU/KEL][1984HOU/KEL][1987BRO/WAN][1988PHI/HAL]

4H2O

T= 298.15 K, 1=0T= 298.15 K, 1=0T=373 K, I=dilT= 298 K, 1=0T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K( 1=0T= 298.15 K,I=0T=298.15K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

+ + 3H2O

T= 298 K, 1=0T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T=298.15K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

00

0234000000

NaOH

HC1O4

HC1O4

HC1O4

02340000

HC1O4

HC1O4

HC1O4

30

JNC TN8400 9 9 - 0 1 1


log Pi,-4

2.911.882.182.481.381.49 3

2.21-0.95

2: Sn(OH)/ + 2H+ tt Sn(OH)22+ + 2H2

[1973KLI/BAR][1979V AS/GLA][1979VAS/GLA][1979VAS/GLA][1984HOU/KEL][1984HOU/KEL][1987BRO/WAN][1988PHI/HAL]

T= 298 K,T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

0

1=0K, 1=2K, 1=3K, 1=4K,I=0K,I=0K, 1=0K, 1=0

02340000

HC1O4

HCIO4

HC1O4

log Pu: S?i(OH)4° + H< <=> Sn(OH)3+ + H2O

2.03 [1973KLI/BAR] T=298K, 1=01.31 [1979VAS/GLA] T= 298.15 K, 1=21.22 [1979VAS/GLA] T= 298.15 K, 1=31.22 [1979VAS/GLA] T= 298.15 K, 1=40.73 l [1980BEN/TEA] T= 298.15 K, 1=00.72 l [1982WAG/EVA] T= 298.15 K, 1=0-0.52 [1984HOU/KEL] T= 298.15 K, 1=01.22 3 [1984HOU/KEL] T= 298.15 K, 1=01.66 [1987BRO/WAN] T= 298.15 K, 1=0-0.24 H988PHI/HAL1 T= 298.15 K, 1=0

0234000000

HC1O4

HC1O4

HC1O4

log PliS: Sn(OH)4° + H2O <=> Sn(OH)s- + H+

-2.77-11.18-8.85 4

[1987BRO/WAN][1988PHI/HAL][1992PEA/BER]

T=rp

298.298.298.

151515

K,K,K,

1=01=01=0

000

log PL6: Sn(OH)4° + 2H2O <=> Sn(OH)62' + 2H

-20.95 •-24.5-21.60 4

-19.11-20.96 '-6.57-22.19-20.86 4

[1952LAT][1955DEL/ZOU][1975KRA][1984HOU/KEL][1985GAL][1987BRO/WAN][1988PHI/HAL][1992PEA7BER]

T= 298.15 K, 1=0T=298.15K, 1=0T=298.15K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

00

00000

log Pl6: Sn(OH)4° <=> SnO32- + 2H+ + H2O

-21.42-23.15 '-23.05 4

[1984HOU/KEL][1985GAL]f!992PEA/BERl

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

000

1 calculated with a A,GO of -944 kJ/mol for Sn(OH)4° (Section 4.4.3)2 calculated by [1970KUR/BAR] from thermodynamic data and own measurements.3 values given in the text; they do not correspond to the values calculated from the A,G0 values given in Table 1 of

[1984HOU/KEL],4 calculated with a log (310 of 0.40 (value from Section 4.4.3).

31

JNC TN8400 9 9 - 0 1 1

4.2 Solid tin(FV) oxides/hydroxides

Fresh precipitates of tin(IV) salt from solutions are amorphous at room temperature. After agingX-ray patterns show broadened reflexes of poorly crystalline SnOj [1963FEI/SCH,1997AMA/CHI]. [1970KUR/BAR] found from X-ray diffraction measurements that the solidphase precipitated from Sn(IV) solution varied from cassiterite at 300 *C to a structure similar tovarlamoffite (Sn(>2 • XH2O) at 100 CC. Their starting material was freshly precipitatedSn(OH)4(s) and equilibrium was attained in 24 hours only at temperatures above 100 *C

Both, the crystalline SnO2(cassiterite), and precipitated SnO2(precip) are quite insoluble, whilea badly defined, freshly precipitated Sn(OH)4(am) is quite soluble (Table 4.5). This solid,however, transforms to SnC>2(precip) within a month.

Only a few experimental measurements of Sn(IV) solubility can be found in the literature. Thedata used for the calculations of the formation constants of tin(IV) hydroxide/oxide compoundsare compiled in Table 4.4. Additional data are given in Table 4.5 and 4.6.

Table 4.4: Experimentally determined equilibrium data compiled for the formation of tin(IV)hydroxide/oxide compounds. These data were chosen for the evaluation ofrecommended values in the present report. Additional information for the differentreferences see Section 4.14: 'Comments to selected references'. Method: sol =solubility measurements.

log K*so Reference Comments I (M) Medium Method

log K*so: Sn(OH)40 t=> SnO2(precipitated) + 2H2O

7.46 ' [1997AMA/CHT] T= 298.15 K, 1=0.1 0.1 NaClO. sol

log K*so: Sn(OH)40 <=> SnO2(crystalline) + 2H2O

8.0 2 [1997AMA/CHTJ T= 298.15 K, 1=0.01 0.01 NaClO4 sol1 In [1997 AM A/CHI], this value is reported as log K*°so with the assumption that the activity coefficient of

Sn(OH)4° is unity.2 calculated from experimental data of [1997 AM A/CHI] in this report

32

JNC TN8400 9 9 - 0 1 1

4.2.1 Freshly precipitated Sn(OH)4(am)

The compilation of [1969FEI/SCH] proposes for freshly precipitated Sn(OH)4(am) a log K*Soof 2.0 for the reaction Sn(OH)4° <=> Sn(OH)4(freshly precipitated). A similar solubility productis determined by [1970KUR/BAR] at 100 °C. However, Sn(OH)4(am) will turn into a poorlycrystalline SnO2(precip) with a constant solubility within a month of aging [1963FEI/SCH,1997AMA/CHI]. No solubility constant for Sn(OH)4(freshly precipitated) is recommended inthe present report.

4.2.2 SnO2(precip)

Only a few experimental measurements of the solubility of precipitated, amorphous SnO2 can befound in the literature. Besides the data given in the compilation of [1969FEI/SCH] which lackexperimental detail, the solubility of precipitated SnC>2 has been determined recently by[1989BAY/EWA] in cement equilibrated water and by [1997AMA/CHI] and [1998ODA/AMA]in 0.1 M NaC104. [1989BAY/EWA] measured at a pH of 9 a Sn(IV) concentration of 6xlO"8

M. [1997AMA7CHI] measured in the pH range 2 to 7 a constant solubility of about 3xlO"8 M.[1997AMA/CHI] calculated a log K*So = 7.46. The additional experiments of [1998ODA/AMA]in 0.1 MNaC104 confirmed the solubility data determined by [1997AMA/CHI] (Table 4.4).[1998ODA/AMA] also showed that the presence of chloride and sulfate influences the SnO2solubility (Table 4.5). The reaction Sn(OH)4° <=» SnO2(precip) + 2H2O is expected to have aAe ~ 0, because no charged species are involved in the reaction. Thus, the value determined by[1997AMA/CHI] (Table 4.5) may be used to calculate the solubility of precipitatedSnO2(precip):

Sn(OH)4° o SnO2(precip) + 2H2O log K*°so= 7.46

33

JNC TN8400 9 9 - 0 1 1

4.2.3 SnO2(cassiterite)

Cassiterite crystallizes at elevated temperatures (« 300 °C), while at lower temperatureprecipitates are less crystalline [1970KUR/BAR]. From the early studies of [1929GRU/LIN] itcan be concluded that the solubility of cassiterite in water is smaller than 106 mol Sn(IV) L 1 .[1970BAR/KLI] measured the solubility of synthetic cassiterite in water, dilute HNO3 anddilute NaOH. They observed a constant, pH-independent (between pH 2 - 11) solubility of 4x 10 7 mol Sn(IV) L"1. No detection limit for the Sn(FV) is indicated which makes it difficult todecide whether the measured concentrations below pH 11 correspond to the detection limit orare real concentrations. Thus these data were not used to calculate cassiterite solubility

More recently, [1997AMA/CHI] determined cassiterite solubility in 0.01 M NaC104. Theyobserved a constant, pH-independent (between pH 2 - 8) solubility of about 9 x 10-9 molSn(IV) L"1. From the average of the experimental values, log K*so = 8.0 can be calculated.Extrapolation to 1=0 of the data given by [1997AMA/CHI] assuming a Ae of 0, as no chargedspecies are involved, gives:

Sn(OH)4° <=> SnO2(cassiterite) + 2H2O logK*°S0=8.0

This value is in fair agreement with the log K*°So of 8.50 given by [1963FEI/SCH].

At higher temperature, cassiterite solubility was determined in different references (Table 4.5).Extrapolation of the measurements of [1981DAEVSOR] to 298 K gives a log K*So of = 7.5 forcassiterite, a value which is in fair agreement with the log K*so of 8.05 as determined by[1997AMA/CHI] for cassiterite (SnO2(cr)). [1973KLJ/BAR], however, determined in thetemperature range 473 - 673 K a higher solubility of cassiterite than [1981DAD/SOR] and[1988BAR/SHA] (Table 4.5). Also at 298 K, the solubility measured by [1973KLI/BAR] ismuch higher than observed in other references [1926GRU/LIN, 1963FEI/SCH,1997AMA/CHI] (Table 4.5). For further comments see Section 4.14.

34

JNC TN8400 9 9 - 0 1 1

4.2.4 Additional data compiled for the formation of tin(FV) hydroxide/oxide compounds

Table 4.5: Additional experimentally determined equilibrium data compiled for the formation of tin(TV)hydroxide/oxide compounds. These data were not chosen in the present report for the evaluation ofrecommended stability values. Reasons for not selecting these references are given at the end of thetable, in Section 4.2 and in Section 4.14: 'Comments to selected references'. Method: sol =solubility measurements.


logK*S0:Sn(OH)4° <=> Sn(OH)4(freshly precipitated)

2.00 l [1963FEI/SCH]2.21 [1970KUR/BAR]

T= 293 K, I=dilT=373 K, I=dil

00

n/aNaOH

solsol

log K*so: Sn(OH)4° <=> SnO2(precip) + 2H2O

5.00 '7.46 2

7.46 2

7.65 3

6.94 3

6.98 3

6.88 3

7.04 3

7.29 3

7.35 3

[1963FEI/SCH][1997AMA/CHI][1998ODA/AMA][1998ODA/AMA][1998ODA/AMA][1998ODA/AMA][1998ODA/AMA][1998ODA/AMA][1998ODA/AMA][1998ODA/AMA]

T= 293 K, I=dilT= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.53; contains 0.4 M NaClT= 298.15 K, 1=0.57 (0.04 M NaCl)T= 298.15 K, 1=0.54 (0.004 M NaCl)T= 298.15 K, 1=0.54 (0.4 M NaC104)T= 298.15 K, 1=0.43 (0.1 M Na2SO4)T= 298.15 K, 1=0.56 (0.01 M Na2SO4

T= 298.15 K, 1=0.53 (0.001 M Na,SO4

000

0.530.570.540.530.430.560.53

n/aNaC104

NaClO4

NaClNaClCVNaClNaClOyNaCl

NaClO4

Na2SO4

NaC10VNa2S04

NaClOVNajSO,

solsolsolsolsolsolsolsolsolsol

logK*S0:Sn(OH)4° <=> SnO2(cassiterite) + 2H2O

> 6 [1926GRU/LIN] T= 298 K, I=n/a8.50 ' [1963FEI/SCH] T= 298 K, 1=06.40 4 [1970BAR/KLI] T= 298.15 K, I=dil6.44 5 [1973KLI/BAR] T= 298.15 K, 1=06.04 [1973KLI/BAR] T= 373 K, 1=05.49 [1973KLI/BAR] T= 473 K, 1=05.25 [1973KLI/BAR] T=573K,I=O5.07 [1973KLI/BAR] T= 673 K, 1=06.23 [1981DAD/SOR] T=473K, 1=05.86 [1981DAD/SOR] T=573K, 1=05.60 [1981DAD/SOR] T= 673 K, 1=07.49 6 [1981DAD/SOR] T= 298 K, 1=08.07 [1988BAR/SHA] T= 573 K, I=dil

000000

0.100000

n/an/anonononono

NaC104

nonononono

solsolsolsolsolsolsolsolsolsolsolsolsol

from unpublished experimental results [1957EGG]. In this report, dissolved species is assumed to be Sn(OH)4°, seealso comment in Section 4.14extrapolated to 1=0 with Davies equation by [1997 AM A/CHI]; [1998ODA/AMA] gives the same valuechloride and sulfate form complexes with tin(IV)linear extrapolated from different I, no detection limit indicatedsee comments in Section 4.14calculated from values determined at 200-400 °C

35

JNC TN8400 9 9 - 0 1 1

Table 4.6: Thennodynamic data for the formation of tin(IV) hydroxide/oxide compounds taken from previouscompilations. As pointed out in Section 2 of this report only experimental data were used for thepresent evaluation. The following table serves only for comparison.

log K*so Reference Comments I (M)

log K*s0: Sn(OH)4° <=> Sn(OH)4(freshly precipitated)

1.33 >-1.80 ]

3.003.101.33 '1.33 '

log K*S0:

8.72 ]

8.75 »6.668.79 l

8.72 >8.76 '8.77 '8.76 l

8.77 l

8.76 ]

8.78 l

8.71 '9.688.72 '8.71 '14.857.585.40 2

[1952LAT][1968SUS/KHO][1975KRA][1984HOU/KEL][1985BAB/MAT][1985GAL1

T= 298.15 K, 1=0T= 298 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

Sn(OH)4° <=> SnO2(cassiterite) + 2H2O

[1952LAT][1954COU][1955DEL/ZOU][1963WIC/BLO][1968SUS/KHO][1971NAU/RYZ][1978COD][1978ROB/HEM2][1979KUB/ALC][1980BEN/TEA][1982PAN][1982WAG/EVA][1984HOU/KEL][1985BAB/MAT][1985GAL][1987BRO/WAN][1988PHI/HAL][1992PEA/BER1

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, I=n/aT=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0

0

000

0

0

0

0

01 calculated with a AfG° of -944 kJ/mol for Sn(OH)4° (Section 4.4.3)2 calculated in this report with a log P, 0 of 0.40 (value from Section 4.4.3)

4.3 Other tin(IV) complexes and compounds

For other tin(IV) species and compounds the amount of experimentally determined dataavailable in the literature is rather limited. From the available data it can be expected that Sn(IV)forms stable complexes and compounds with chloride, fluoride, carbonate, sulfate and sulfide.In Table 4.7 to 4.9 formation constants for tin(IV) species and compounds are compiled.

No data are recommended in this report as not sufficient data are available to judge the reliabilityof these data and to extrapolate formation constants to 1=0. The following tables serve only togive a general idea of the strength of complex formation with tin(IV).

36

JNC TN8400 9 9 - 0 1 1

Table 4.7: Experimentally determined equilibrium data compiled for the tin(IV) system. These data were notchosen in the present report for the evaluation of recommended stability values because notsufficient data are available to extrapolate the formation constants to 1=0. Method: pot =potentiometry, sol = solubility measurements, se = sulfide selective electrode.

logp Reference Comments I (M) Medium Method

log Pu: Sn4+ + Ci <=> SnCl3+

3.71 [1978FAT/ROU1 T= 298.15 K, 1=5 HC1O, pot

log Pu: Sn4+ + 2CI- <

6.46 [1978FAT/ROU1 T= 298.15 K, 1=5 HC1O. pot

log Pu: S?i4+ + 3CI <=> SnCl3+

8.78 ri978FAT/ROUl T= 298.15 K, 1=5 HC10d pot

log Ph4: Sn4+ + 4Ci <=? SnCl4°

9.48 ri978FAT/ROU] T= 298.15 K, 1=5 HC1O_ pot

log P!5: Sn4+ + 5CI <=> SnCl5-

11.23 ri978FAT/ROU1 T= 298.15 K, 1=5 HC1O4 pot

log Ph6: Sn4+ + 6CI- <=> SnCl62-

12.40 ri978FAT/ROUl T= 298.15 K, 1=5 HC1O, pot

log PUJ: Sn(0H)4° + CO32- + H+ <=> Sn(OH)3CO3- + H2O

7.71 ' [1971KUE7BAR] T= 298.15 K, 1=0.1-0.5 NaHCO, sol

log Pi,oX- SnO2(aq) + 2H2SO4 <=* Sn(SO4)2+ + SO4

2- + 2H2O

-1.30-1.55

[1955BRU]ri955BRUl

T= 303 K, I=n/aT=291 K, I=n/a

H2SO4 solsol

log Pi.o.i-' SnS2(s) + S2- <=> SnS32-

5.31 [1968HSE/REC] T= 298 K, 1=0.1 0.1 NaNO, se

log Ks0: Sn(OH)4° + 2HS- + 2

36.45 2 [1984KOC/TOP]

SnS2(s) + 4H2O

T= 298.15 K, 1=1-4 HCIO4 sol1 extrapolated to 1=0 with Debye-HUckel equation by [1971KUR/BAR]2 calculated with a AfG° of -944 kJ/mol for Sn(OH)4° (Section 4.4.3)

37

JNC TN8400 9 9 - 0 1 1

Table 4.8: ThermodynamJc data for tin(IV) system taken from previous compilations. As pointed out inSection 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

log K so

log Pl,0,4-

- 11.52

log Pi,o,6-

0.2111.23

log PIJJ-

6.52 '5.39

log Pi,o,6--

12.64 '12.64 '

log Pi,o,i:

-2.69 !-2.70 '6.25

Reference

Sn4+ + 4Ci <=> SnCl4

[1984HOU/KEL]

Sn4+ + 6CI <=> SnClt

[1984HOU/KEL][1984KOC/TOP]

Sn(OH)4°+ F~+ H+

[1982 WAG/EVA][1988PHI/HAL1

Sn(OH)4° + 6F- + 4H"

[1952LAT][1985GAL]

Comments

0

T= 298.15 K, 1=0

T= 298.15 K, 1=0T= 298.15 K, I=n/a

<x> SnOOHF0 + 2H2O

T= 298.15 K, 1=0T= 298.15 K, 1=0

" <=> SnF62- + 4H2O

T= 298.15 K, 1=0T= 298.15 K, 1=0

Sn(OH)4° + SO42' + 4H+ <=> SnSO4

2+ + 4H2O

[1980BEN/TEA] T= 298.15 K, 1=0[1982WAG/EVA] T= 298.15 K, I=n/a[1988PHI/HAL1 T= 298.15 K, I=n/a

I(M)

0

0

00

00

0

log Pi.0,2- Sn(OH)4° + 2SO/- + 4H+ <^> Sn(SO4)2° + 4H2O

-0.32 l [1982WAG/EVA]-1.44 ri988PHI/HALl

T= 298.15 K, I=n/aT= 298.15 K, I=n/a

log Kso: Sn(OH)4° + 2SO42- + 4H+ <=> Sn(SO4)2(s) + 4H2O

-5.73 '-15.55 '-15.55 '-5.73 '-5.73 [

-6.84

[1952LAT][1977BAR/KNA][1979KUB/ALC][1985BAB/MAT][1985GAL]ri988PHI/HAL]

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/a

log Kso: Sn(OH)4° + 2HS- + 2H+ <=> SnS2(s) + 4H2O

30.51 ! [1974MIL]30.51 ' [1977BAR/KNA]30.51 ' [1979KUB/ALC]36.53 ' [1985GAL]35.40 fl988PHI/HALj

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/a

calculated with a Afi0 of -944 kJ/mol for Sn(OH)4° (Section 4.4.3)

38

JNC TN8400 9 9 - 0 1 1

Table 4.9: Thermodynamic data for Sn4+ predicted by [1984BRO/WAN]. This table serves only forcomparison.

log P u logPi i 2 log Pi,3 log Pi,4 log Pii5 logPi,6 References Comments

log Pmn: mSn4+ + nCt & SnmCln4-"

3.88 7.12 9.84 12.10 13.93 15.34 [1987BROAVAN1 T= 298.15 K, 1=0

log pmin: mSn4+ + nF <* Snm¥n4'n

8.89 17.29 25.34 33.07 40.53 47.72 [1987BROAVAN1 T= 298.15 K, 1=0

log pmn: mSn4+ + nCO32- « Snm(CO3)n

4-2"

11.72 21.84 30.62 38.14 44.47 49.65 [1987BROAVAN1 T= 298.15 K, 1=0

log pm_n: mSn4+ + nHCOf <=> Snm(HCO3)n4n

6.85 12.18 16.22 19.09 20.84 21.50 [1987BROAVAN] T= 298.15 K, 1=0

log pmin: mSn4+ + nNOy <=> Snm(NO3)n4-"

1.83 2.16 1.25 -0.81 [1987BROAVAN1 T= 298.15 K, 1=0

log /3min: mSn4* + nHPO42- « S

12.88 24.86 36.21 47.02 57.35 67.24 [1987BROAVAN] T= 298.15 K, 1=0

log £„,,„: mSn4+ + nH2PO/ <=> SnJH2PO4)n4-"

4.75 8.54 11.62 14.10 16.04 17.46 [1987BROAVAN] T= 298.15 K, 1=0

log Pm,n: mSn4+ + nSOf <=> Snm(SO4)n4-2"

4.95 8^61 11.21 12.88 13.65 13.57 [1987BRO/WAN] T= 298.15 K, 1=0

39

JNC TN8400 9 9 - 0 1 1

4.4 Redox reactions

Tin exists in several oxidation states from -IV to +IV in its compounds, but only the +11 and+IV states are important in its aqueous chemistry. The +IV state is particularly stable in naturalenvironments, and cassiterite is the major source of this element in nature [1976BAE/MES,1985GAL, 1995WIB]. The stannous ion (Sn2+) is easily oxidized. The data used for thecalculations of the redox reactions of tin are compiled in Table 4.10. Additional data for the tinredox system are compiled in Table 4.11 and 4.12. These data were not chosen for thecalculation of log K° values in this report.

Table 4.10: Experimentally determined equilibrium data compiled for redox reactions of tin.Calculated in this report from the respective E° values. These data were chosen forthe evaluation of recommended values in the present report. Additionalinformation for the different references see Section 4.14: 'Comments to selectedreferences'. Method: pol = polarography, pot = potentiometry.

logK Reference

log K: Sn2+ + 2e~ <=> Sn(cr)

-5.06 i-5.11 !-5.11 1-5.24 2

-5.13 3

[1928PRY][1928PRY][1928PRY][1949RIC/POP][1970BON/TAY]

Comments

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=0.08K, 1=0.2K, 1=0.35K, 1=1K, 1=1

KM)

0.080.2

0.3511

Medium

HC1O4

HC1O4

HC1O4

KC1NaC104

Method

potpotpotpotpol

1 calculated in this report from the E° (0.4532, 0.434 and 0.4252 V, respectively) and the respective Sn(II)concentrations as given by [1928PRY] using E°(HgCl/Hg) = 0.2444 V [1996STU/MOR; saturated KC1]

2 calculated by [1949RIC/POP] from data measured with a Calomel electrode in KC1 and KC1O4 medium.Corrected by [1949RIC/POP] for the interaction with Cl, extrapolated to 1=1 with Debye-Huckelapproximation.

3 calculated in this report from the E° value (0.374 V) given by [1970BON/TAY] using E°(AgCl/Ag) = 0.2223V[1996STU/MOR;I=1]

4.4.1 Sn2+/Sn(cr)

The redox potential of the reaction Sn2+ + 2e" <=> Sn(cr) has been measured by several authorsin acidic media. [1928PRY] and [1949RIC/POP] determined the potential against a calomelelectrode in perchlorate solutions and potassium chloride solutions, respectively.[1970BON/TAY] measured in 1 M HCIO4 the potential by polarography against a silver- silverchloride reference electrode. Extrapolation of these values given in Table 4.10 to I = 0 is shownin Figure 4.1 and results in:

40

JNC TN8400 9 9 - 0 1 1

Sn2+ Sn(cr) logK° =E° =

-4.63-O.137V

Sn2+ + 2e~ <=> Sn(cr)

0.5 1 1.5L. molal

Figure 4.1: Plot of log K + 4 D vs. Im for the reaction : Sn2+ + 2e- <=> Sn(cr) at 25 tC. Thestraight line shows the result of the linear regression: Ae = -0.25; log K°= -4.63.

4.4.2 Sn2+ISn4+

Sn2+ oxidizes readily to Sn4+ under oxic conditions [1934HUE/TAR, 1979VAS/GLA].[1934HUE/TAR] measured with a standard hydrogen electrode the redox potential of thisreaction in HCl (Table 4.11). Down to a HCl concentration of 0.5 M, [1934HUE/TAR]observed no hydrolysis of the Sn4+ cation present in concentrated acid, while in 0.1 and 0.2 MHCl [1934HUE/TAR] found a strong shift in their potentiometric measurements indicating thehydrolysis of Sn4+. Unfortunately, [1934HUE/TAR, 1979VAS/GLA] had chloride present intheir measurements. Sn(II) and probably also Sn(IV) (cf. Table 4.9) has a strong tendency toform chloride complexes. As no data are available for the complex formation between Sn(IV)and chloride, it is not possible to calculate a log K° for the reaction Sn4+ + 2er <=> Sn2+ fromthese data. [1934HUE/TAR and 1979VAS/GLA] extrapolated from their measurements a log K°~ 5 for the redox equilibrium Sn4+ + 2e~ <=> Sn2+ (Table 4.11). For the lack of any better data,a tentative log K° value of 5 corresponding to E° = 0.148 V may be used. This value, however,is debatable as the tin chloride complexes are not, or only partly, considered and as the activitycorrections made by [1934HUE/TAR and 1979VAS/GLA] are somewhat doubtful.

41

JNC TN8400 9 9 - 0 1 1

Combining log K° value of 5 with the redox potential of Sn2+ as defined in Section 4.4.1 givesa tentative log K° of 0.37 tor the reaction Sn4+ + 2 H2(g) <=> Sn(cr) + 4H+, corresponding toan E° of 0.005 V.

4.4.3 Sn(OH)4°

No measurements of the redox reaction between aqueous Sn(II) and Sn(IV) at neutral oralkaline conditions exist. Reliable data for the protonation of Sn(OH)4° to Sn4+ would beneeded for the calculation of the redox equilibria between the species Sn(OH)4° and Sn2+ usedas master species in this report. However, such data are not available at present (cf. Section4.1.1).

However, another possibility is to calculate the standard molar Gibbs energy of formation ofdissolved species from solubility measurements and the standard molar Gibbs energy offormation of the respective solid. From the AfH° and S° data given in the review of[1989COX/WAG] a AfG° of -515.826 kJ/mol for cassiterite SnO2(cr) is calculated by[1995SEL/BID]. This value is accepted in this report as the standard molar Gibbs energy offormation for cassiterite. Using this value, the solubility of cassiterite as defined in Section4.2.3, and a AfG° of -237.14 kJ/mol for H2O, a tentative standard molar Gibbs energy offormation for Sn(OH)4° is calculated to be -944.156 kJ/mol resulting in (also these values aretentative):

Sn(OH)4° + 2 H2(g) <=> Sn(cr) + 4H2O logK0 =0.77E° =0.011 V

or Sn(OH)4°+ 2H++ H2(g) « Sn2++ 4H2O log K° = 5.40E° = 0.160 V

Only two values for this reaction are reported in the literature (Table 4.12). The value given by[1984HOU/KEL] is based on the hydrolysis data of [1971NAZ/ANT].

Also for the reaction from Sn(OH)4° to Sn4+ a tentative constant can be calculated from theabove values:

Sn(OH)4° + 4H+ <=» Sn4+ + 4H2O log K° = 0.40

42

JNC TN8400 9 9 - 0 1 1

4.4.4 Additional data compiled for the tin redox system

Table 4.11: Experimentally determined equilibrium data compiled for the tin redox system. These data werenot chosen in the present report for the evaluation of recommended stability values.. Method: pot= potentiometry.

logK Reference Comments I (M) Medium Method

K: Sn2+ + 2e- <=> Sn(cr)

-4.59 l

-4.94 2

-5.17 2-5.10 2

[1928PRY][1949RIC/POP][1949RIC/POP][1949RIC/POP]

T= 298.15 K, 1=0.08-0.35T= 298.15 K, 1=0.1T= 298.15 K, 1=1T= 298.15 K, 1=1.8

0111

HC1O4

KC1KC1KC1

potpotpotpot

log K: Sn2+ <=> Sn4+ + 2e~

-5.21 3

-2.37 4 '5

-4.23 4 '5

-4.88 4

-4.75 4

-4.67 4

-4.48 4

-5.16 6

-5.41 6

-5.67 6

-5,09 6

-4,69 6

-3.85 6

-6.78 7

-6.58 7

-6.58 7

-6.38 7

-5.15 7.8

[1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUEyTAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1934HUE/TAR][1979VAS/GLA][1979V AS/GLA][1979 V AS/GLA][1979V AS/GLA][1979 V AS/GLA]

T=298.15,1=0.5-2T=298.15,1=0.1T=298.15,1=0.2T=298.15,1=0.5T=298.15,1=0.8T=298.15,1=1.1T=298.15,I=2T=298.15,1=0.1T=298.15,1=0.2T=298.15,1=0.5T=298.15,1=0.8T=298.15,1=1.1T=298.15,1=2 .T= 298.15 K, 1=2.1T= 298.15 K, 1=3.1T= 298.15 K, 1=3.1T= 298.15 K, 1=4.1T= 298.15 K, 1=2-4

00.10.20.50.81.12

0.10.20.50.81.12

2.1163.1043.1424.086

0

HC1HC1HC1HC1HC1HC1HC1HC1HC1HC1HC1HC1HC1

HC1O4

HC1O4

HC1O4

HC1O4

HC104

potpotpotpotpotpotpotpotpotpotpotpotpotpotpotpotpotpot

1 extrapolated to 1=0 with Debye-Hiickel term by [1928PRY]. Calculated from E°=0.1359 V.2 calculated from [1949RIC/POP] from data measured with a Calomel electrode. Extrapolated to 1=1 with Debye-Hiickel

approximation.3 extrapolated to 1=0 after correction for activity by [1934HUE/TAR], Calculated in this report from E°=0.154 V.4 corrected for activity of H+ by [1934HUE/TAR]. Calculated in this report from corrected E° vbalues.5 the observed potential is a function of tin concentration, indicating hydrolysis of tin.6 calculated in this report from the E° values and the respective Sn(IV)/Sn(II) concentrations as given by

[1934HUE/TAR], Mean of 5 measurements.1 paper in Russian, no experimental details available.8 extrapolated to 1=0 with Debye-Huckel term by [1979VAS/GLA]. Calculated in this report from E°=0.1522 V.

43

JNC TN8400 9 9 - 0 1 1

Table 4.12: Thermodynamic data for tin redox system taken from previous compilations. As pointed out inSection 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

logK Reference Comments I (M) Medium

log K: Sn2+ + 2e- <=> Sn(cr)

4.604.614.60 '4.60 l

4.884.894.885.244.774.614.76All4.774.844.77

[1941GAR/HEI][1952LAT][1955DEL/ZOU][1975KRA][1976VAS/KOK][1978COD][1979VAS/GLA][1980BEN/TEA][1982WAG/EVA][1984HOU/KEL][1985BAB/MAT][1985GAL][1988PHI/HAL][1989COXAVAG][1992PEA/BER]

T=298.15K, I=n/aT= 298.15,1=0T= 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15,1=0T= 298.15 K, 1=2-4T= 298.15,1=0T= 298.15,1=3T= 298.15 K, 1=0T= 298.15,1=0T= 298.15,1=0T= 298.15,1=0T= 298.15,1=0T= 298.15 K, 1=0

00

00003000000

HClOa

HC1CL

log K: Sn2+ <=> Sn4+ + 2e

-5.07-5.10 :

-5.21 '-5.08-4.27-5.25-5.05-5.10

log K:

7.024.41

[1952LAT]1 [1955DEL/ZOU]1 [1975KRA]

[1984HOU/KEL][1985BAB/MAT][1985GAL][1988PHI/HAL][1992PEA/BER1

Sn(OH)4° + 2H+ + H2(g) ±

[1984HOU/KEL][1988PHI/HAL]

T= 298.15,1=0T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15,1=0T= 298.15,1=0T= 298.15,1=0T= 298.15 K, 1=0

=> Sn2+ + 4H2O

T= 298.15 K, 1=0T= 298.15,1=0

00

00000

00

calculated in this report from E° values.

44

JNC TN8400 9 9 - 0 1 1

4.5 Hydrolysis of tin(II)

Tin(II) is easily oxidized by air to tin(TV) and hydrolyzes readily. In solutions containing morethan 10"5 M Sn(II), the formation of polymeric Sn(II) species (Sn3(OH)4

2+ and Sn2(OH)22+)

can be observed under acidic conditions [1958TOB, 1976BAE/MES]. Between pH = 4 - 8 thesolubility of SnO(s) does not depend on pH, indicating the predominance of the unchargedSn(OH)2° species. In strong acid or in base tin(II) solubility increases [1976BAE/MES].

The data used for the calculations of the formation constants of tin(II) hydroxide complexes arecompiled in Table 4.13. Additional data for the tin(II) hydroxide system which were not chosenfor the calculation of log (3° values in this report are compiled in Table 4.14 and 4.15.

Table 4.13: Experimentally determined equilibrium data compiled for the hydrolysis of Sn2+,according to the equilibrium: mSn2+ + nH2O <=> Snm(OH)n

2m-n + nH+. Thesedata were chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 4.14: 'Comments toselected references'. Method: sol = solubility measurements.

log (3m,nReference Comments I (M) Medium Method

log fiu: Sn2++ H20 <=>

-3.92 [1958TOB]-3.70 [1976GOB]-4.10 [1981PET/MIL]-3.80 [1981PET/MEL]-4.10 [1981PET/MEL]

SnOH+ + H+

T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=0.1T= 298.15 K, 1=0.5T= 298.15 K, 1=1

-3.77 [1997SAL/FER] T= 298.15 K, 1=3

33

0.10.5

13

NaC104

NaC104

NaNO3

NaNO3

NaNO3

NaCIO,

potpotpotpotpotpot

log pli2: Sn2+ + 2H2O o Sn(OH)2° + 2H+

-1-7-7

.90

.90

.80

[1981PET/MIL][1981PET/MIL][1981PET/MIL]

T=T=T=

298.298.298.

151515

K,K,K,

1=01=01=1

.1

.500

.1

.51

NaNO3

NaNO3

NaNO,

potpotpot

45

JNC TN8400 9 9 - 0 1 1


log Pu: Sn2+ + 3H2O ^>

-17.96 [1977MAR]-17.5 [1981PET/MIL]-17.7 [1981PET/MIL]-17.6 [1981PET/MEL]

Sn(OH)3- + 3H+

T= 298.15 K, 1=3 3 NaC104

T= 298.15 K, 1=0.1 0.1 NaNO3

T= 298.15 K, 1=0.5 0.5 NaNO3

T= 298.15 K, 1=1 1 NaNO,

potpotpotpot

log P3

-6.11-6.81-6.87

3Sn2+ + 4H2O <=> Sn3(OH)42+ + 4H+

[1958TOB][1976GOB][1997SAL/FER]

T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3

3

3

3

NaClO4

NaClO4

NaCIO,

potpotpot

[1973JOH/OHT] have concluded from X-ray investigations in 3 M Sn(II) perchlorate solutionsthat hydrolysis leads to the formation of Sn3(OH)42+ in concentrated Sn(II) solutions: theirobservation confirm the potentiometric data of [1958TOB], [1958TOB] studied in his carefulwork the hydrolysis of Sn(II) in 3 M NaC104. He determined both the Sn2+ and the H+

concentration. Titration experiments were conducted in 2.5 - 40 mM of Sn(II) and down to apH of 1. He interpreted his results with four species: Sn2+, SnOH+, Sn3(OH)42+ andSn2(OH)22+. His results were recalculated by [1964LIN/TU] with least-squares analysis.[1964LIN/TU] doubted the existence of SnOH+. In place of this species [1964LIN/TU]proposed the existence of Sn2(OH)3+. [1976GOB] and [1997SAL/FER] confirmed later in theirwork the existence of Sn3(OH)42+ and SnOH+ and rejected the formation of the speciesSn2(OH)3+ proposed by [1964LIN/TU].

Many of the earlier measurements of Sn(II) hydrolysis [1939GOR, 1941GAR/HEI,1952VAN/RHO] neglected the formation of polynuclear tin(II) complexes in their solutions.Thus, their data (given in Table 4.14) were not chosen for the calculation of log (3° values in thisreport.

46

JNC TN8400 9 9 - 0 1 1

4.5.1 SnOH+

Under acidic conditions and in solutions containing more than 10"5 M Sn(II), the formation ofpolymeric Sn(II) species (Sn3(OH)4

2+ or Sn2(OH)22+), can be expected besides Sn2+ and

SnOH+. [1981PET/ML] studied with anodic stripping voltammetry the hydrolysis of tin(II) in0.1 - 1 M NaNO3, NaCl and artificial seawater with a Sn concentration of 10"7 M and observedonly the formation of mononuclear Sn(H) complexes. Further data for the hydrolysis of Sn2+ toSnOH+ are given by [1958TOB] and [1976GOB]. From these experimental results (given inTable 4.13), the following formation constant can be calculated (see Figure 4.2)

Sn2+ + H2O SnOH+ + H+ 1,1 = - 3.75

Sn2+ + H2O <> SnOH+ + H+

lm, molal

Figure 4.2: Plot of log (3U + 2 D vs. Im for the reaction : Sn2+ + H2O o SnOH+ + H+ at 25°C. The straight line shows the result of the linear regression: As = - 0.13; logP°ii = - 3.75.

47

JNC TN8400 9 9 - 0 1 1

4.5.2 Sn(OH)2°

Formation constants for Sn(OH)2° from Sn2+ have been determined by [1941GAR/HEI] and[1981PET/MIL]. [1941GAR/HEI] measured Sn(II) solubility in 0 - 0.4 M HC1 and NaOHassuming only the presence of Sn2+, SnOH+, Sn(OH)2° and Sn(OH)3- at variable I.[1941GAR/HEI] neglected the formation of polymeric species. Under acidic conditions also theformation of tin(II) chloride complexes cannot be excluded. An additional difficulty of the workof [1941GAR/HEI] is that the free acidity was not measured but was calculated indirectly fromassumed equilibria. [1981PET/MIL] studied with anodic stripping voltammetry the hydrolysisof tin(II) with an Sn concentration of 10"7 M. From the data of [1981PETVMIL] in NaNO3

(given in Table 4.13), the following formation constant can be calculated (see Figure 4.3)

Sn2+ + 2H2O <=> Sn(OH)2°+2H+ log =-7.71

Sn2+ + 2H2O <=> Sn(OH)2° + 2H+

lmi molal

Figure 4.3: Plot of log (31,2 + 2 D vs. Im for the reaction : Sn2+ + 2H2O <=> Sn(OH)2° + 2H+at 25 °C. The straight line shows the result of the linear regression: As = - 0.31;log J3°i,2 = -7.71.

48

JNC TN8400 9 9 - 0 1 1

4.5.3 Sn(OH)3-

Formation constants for Sn(OH)3- from Sn2+ have been determined by [1941GAR/HEI,1973GAB/SRI, 1977MAR, 1978DIC/LOT1 and 1981PET/MIL]. The value of[1941GAR/HEI] was not chosen for extrapolation to I = 0 (cf. Section 4.5.2: Sn(OH)2°).Extrapolation of the data of [1977MAR] in 3 NaC104 and of [1981PET7MIL] in NaNO3 (givenin Table 4.13) gives (see Figure 4.4):

Sn2+ + 3H2O ^ Sn(OH)3-+3H+ log(3°u =-17.54

Sn2+ + 3H2O <=> Sn(OH)3- + 3H+

-15

-15.5 --

-16

-18

-18.5

-19

-19.5 •

-20

= -0.12x-17.54

L, molal

Figure 4.4: Plot of log p u + 0D vs. Im for the reaction : Sn2+ + 3H2O <=> Sn(OH)3" + 3H+at 25 °C. The straight line shows the result of the linear regression: Ae = 0.12; logP \ 3 = - 17.54.

[1973GAB/SRI] and [1978DIC/LOT1] determined the redox potential of the reactions Sn(cr) +3OH" <=> Sn(OH)3- + 2e- as 0.90 and 0.88 V, respectively. From these data log Pi,3 values forreactions Sn2+ + 3H2O <=> Sn(OH)3" + 3H+ can be calculated (Table 4.14). These values arerather estimates than exact measurements due to possible errors introduced by assuming log Pvalues for the Sn2+/Sn(cr) redox reaction and a log Kw-

The selected hydrolysis constants describe the experimental data well. A closer look at thestepwise formation constants shows that they are unequally distributed (Sn2+ to SnOH+,Sn(OH)2°, and Sn(OH)3": 3.75, 3.96 and 9.83). The third stepwise hydrolysis constants is 6order of magnitude higher than the first two. This unequal distribution of the stepwisehydrolysis constants could be due to a number of reasons, e.g., not all (i.e. polynuclear)relevant species have been considered or due to experimental artefacts (detection limit).

49

JNC TN8400 9 9 - 0 1 1

4.5.4 Sn3(OH)42+

In solution containing more than 10"5 M Sn(II), the formation of polymeric Sn(II) species, canbe expected besides Sn2+ and SnOH+ under acidic conditions. [1958TOB, 1976GOB and1997SAL/FER] determined formation constants for Sn3(OH)4

2+ in 3 M NaC104 (Table 4.13).Extrapolation of the mean log p3>4 of-6.82 to 1=0 gives (assuming an As of -0.16 as observedfor Pb3(OH)4

2+, Section 6.1.7): '

3Sn2+ + 4H2O <=> Sn3(OH)42+ + 4H+ log ( 3 \ 4 =-6.51, Ae =-0.16

The value at 1=0 is rather tentative as no data at different I are available.

4.5.5 Sn2(OH)22+

[1928PRY2] and [1958TOB] propose the formation of Sn2(OH)22+ in concentrated Sn(II)

solutions. More recently, [1976GOB] and [1997SAL/FER] have presented potentiometric datawhich convincingly indicate the formation of SnOH+ and Sn3(OH)4

2+, but seem to exclude theformation of Sn2(OH)2

2+.

4.5.6 MSn(0H)3+

[1981 PET/MIL] determined an equilibrium constant for the ion pair formation Sn2+ + 3H2O +W+ & MSn(0H)3+ + 3H+, where MP+ is Mg2+ , Ca2+, or Sr2+ (Table 4.14). As this is theonly study of this kind, no log pi j 3 is calculated for 1=0.

4.5. 7 Additional equilibrium data compiled for Sn(II) hydrolysis

Table 4.14: Additional experimentally determined equilibrium data compiled for the hydrolysis of Sn2+,according to the equilibrium: mSn2+ + nH2O *=> Snm(OH)n

2m"n + nH*. These data were notchosen in the present report for the evaluation of recommended stability values. Reasons for notselecting these references are given in the text in Section 4.5 and Section 4.14. Method: sol =solubility measurements, sp = spectrophotometry, and pot = potentiometry.

log fWn

log Pij: •!

-1.70 '-2.06 >-1.70 ><-2.57-3.70-3.10 2

-3.10 2

-2.18 2

Reference

?n2+ + H2O <=> SnOH+

[1939GOR][1941GAR/HEI][1952VAN/RHO][1965MES/IRA][1974GOB][1981PET/MIL][1981PET/MIL][1995DJU/JEL1

Comments

+ H+

T= 298.15 K, 1=0.01-0.5T= 298.15 K, 1=0-0.04T= 298.15 K, 1=3T= 298.15 K, 1=1T= 298.15 K, 1=3T= 298.15 K, 1=0.5T= 298.15 K, 1=0.7T=298.15K, 1=3

KM)

0var313

0.50.73

Medium

NH4NO3/HC1O4

HCl/NaOHNaC104

NaC104

NaC104

NaClASW2

NaCl

Method

titsolpottitpotpotpotpot

50

JNC TN8400 9 9 - 0 1 1

Table 4.14. continued

log P,,2: Sn2+ + 2H2O <=> Sn(OH)2° + 2W

-7.06 '-8.20 2

-8.20 2

log P,,.

-16.61-16.59-15.63-16.88-17.8 2

-17.2 2

log P3,<

-6.81-2.70 »

[1941GAR/HEI][1981PET/MIL][1981PET/MIL]

T= 298.15 K, 1=0-0.04T= 298.15 K, 1=0.5T= 298.15 K, 1=0.7

,.- Sn2+ + 3H2O <=> Sn(OH)f + 3H+

1 [1941GAR/HEI]1 [1941GAR/HEI]3 [1973GAB/SRI]4 [1978DIC/LOT1]

[1981PET/MEL][1981PET/MIL1

T= 298.15 K, 1=0-0.4T= 298.15 K, 1=0.005T= 298.15 K, 1=0.25T= 298.15 K, 1=0.5-5T= 298.15 K, 1=0.5T= 298.15 K, 1=0.7

,: 3Sn2+ + 4H2O <=> Sn3(OH)42+ + 4H+

[1974GOB][1995DJU/JEL1

T= 298.15 K, 1=3T= 298.15 K, 1=3

var0.50.7

var0.0050.25

0.50.7

33

HCl/NaOHNaCl

ASW 2

HCl/NaOHNaOHNaOHNaOHNaCl

ASW2

NaClO4

NaCl

solpotpot

solsolpotpotpotpot

potpot

log 2Sn2+ + 2H2O <=> Sn2(OH)22+ + 2H+

-2.74 5

-2.70 5

-4.10 5-e

-4.45 5

log PiiU

-16.20 7

[1928PRY2][1928PRY2]

'< [1928PRY2][1958TOB1

,: Sn2+ + 3H2O + M2+ <=

[1981PET/MIL]

T=298.15K,I=dilT= 298.15 K, I=dilT= 298.15 K, 1=0.5T= 298.15 K, 1=3

> MSn(0H)3+ + 3H+

T= 298.15 K, 1=0.5

0.53

0.5

Sn(C104)2 •SnCl2

KC1NaC104

NaCl

tittittitpot

pot1 polynuclear complexes not considered2 formation of tin chloride complexes probable; ASW = Artificial seawater3 estimated from potentiometric data assuming for 1=0.25 a log Kw of -13.76, log P (Sn2+/Sn(cr)) of 5.11 (see Section

4.4: Redox reactions)4 extrapolated to 1=0 using Pitzer's approach by [1978DIC/LOT1]. Estimated in this report from potentiometric data

assuming a log Kw of -14.0, log K (Sn2+/Sn(cr)) of 4.63 (see Section 4.4: Redox reactions)5 this species does probably not exist, cf. Section 4.5.56 formation of Sn(II) chloride complexes7 M2+ = Mg2+, Ca2+, or Sr2+

Table 4.15: Thermodynamic data compiled for the hydrolysis of Sn2+ taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

Reference Comments I (M) Medium

log Sn2+ + H2O <=> SnOH+ + H+

-3.40-3.78-3.40

[1976BAE/MES][1976SMI/MAR]fl981BAE/MESl

T= 298.15 K,T= 298.15 K,T= 298.15 K,

1=01=3I=n/a

03

51

JNC TN8400 9 9 - 0 1 1


-1.67-3.70-3.88-2.10-1.54-3.67-3.41-3.40-3.78-3.64-1.55

log Pi.

-7.1-7.06-7.00-6.64-8.64-7.07-7.10-7.58-7.06

[1982WAG/EVA][1982SMI/MAR][1984HOU/KEL][1985BAB/MAT][1985GAL][1987BRO/WAN][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1992PEA/BER1

2: Sn2+ + 2H2O <=> Sn(OH)2°

[1963FEI/SCH][1976BAE/MES][1984HOU/KEL][1985BAB/MAT][1987BROAVAN][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1992PEA/BER]

T= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3T= 293.15 K, 1=0.5T= 298.15 K, 1=0

+ 2H+

T= 293 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 293.15 K, 1=0.5T= 298.15 K, 1=0

030000003

0.50

0000000

0.50

log fiu: Sn2+ + 3H2O <=> Sn(OH)y + 3H+

15.8815.8816.716.617.9618.7116.1616.0314.7816.6416.6017.9418.5416.61

[1952LAT][1955DEUZOU][1963FEI/SCH][1976BAE/MES][1982SMI/MAR][1984HOU/KEL][1985BAB/MAT][1985GAL][1987BRO/WAN][1988PHI/HAL][1989SM1/MAR][1989SMI/MAR][1989SMI/MAR][1992PEA/BER1

T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 293 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3

' T= 293.15 K, 1=0.5T= 298.15 K, 1=0

00030000003

0.50

log Pu: Sn2+ + 2H2O <=> SnOOH- + 3H+

-15.8-15.82-16.05

[1975KRA][1984HOU/KEL][1992PEA/BER]

T= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0

00

log p14; Sn2+ + 4H2O <=> Sn(OH)42- + 4H+

-22.05 [1987BROAVAN] T= 298.15 K, 1=0 0-21.98 [1988PHI/HAL] T= 298.15 K, 1=0 0_

52

JNC TN8400 9 9 - 0 1 1


log fh-4.58-4.77-4.46-4.46-4.39-4.79-4.80-4.46-4.77

- 2Sn2+ + 2H2O <=> Sn2(OH)22+ + 2H+

[1964LIN/TU] T=[1976BAE/MES] T=[1976SMI/MAR] T=[1982SMI/MAR] T=[1984HOU/KEL] T=[1988PHI/HAL] T=[1989SMI/MAR] T=[1989SMI/MAR] T=[1992PEA/BER] T=

298.15 K, 1=3298.15 K, 1=0298.15 K, 1=3298.15 K, 1=3298.15 K, 1=0298.15 K, 1=0298.15 K, 1=0298.15 K, 1=3298.15 K, 1=0

303300030

NaClO,

log (32J: 2Sn2+ + 3H2O

-6.66 [1964LIN/TU]-6.77 [1984HOU/KEL]

Sn2(OH)3+ + 3H+

T= 298.15 K, 1=3T= 298.15 K, 1=0

30

NaC104

log P26: 2Sn2+ + 3H2O <=>

-30.36 fl984HOU/KEL1

Sn2O32- + 6H+

T= 298.15 K, 1=0

log Ps,4- 3Sn2+ + 4H2O <=> Sn3(OH)42+ + 4H+

-6.85 [1964LIN/TU] T= 298.15 K, 1=3-6.88 [1976BAE/MES] T= 298.15 K, 1=0-6.79 [1976SMI/MAR] T= 298.15 K, 1=3-6.79 [1982SMI/MAR] T= 298.15 K, 1=3-6.65 [1984HOU/KEL] T= 298.15 K, 1=0-6.26 [1987BRO/WAN] T= 298.15 K, 1=0-6.93 [1988PHI/HAL] T= 298.15 K, 1=0-6.90 [1989SMI/MAR] T= 298.15 K, 1=0-4.47 [1989SMI/MAR] T= 298.15 K, 1=3-6.88 T1992PEA/BER1 T= 298.15 K, 1=0

3033000030

NaC104

log P3,5: 3Sn2+ + 5H2O <=> Sn3(OH)s+ + 5H+

-8.41 f!987BROAVANl T= 298.15 K, 1=0

log $4,4: 4Sn2+ + 4H2O <=* Sn4(OH)/+ + 4H+

-3.07 [1987BROAVAN1 T= 298.15 K, 1=0

log (l6i8: 6Sn2+ + 8H2O <=> Sn6(OH)84+ + 8H+

-7.81 [1987BROAVAN1 T= 298.15 K, 1=0

53

JNC TN8400 9 9 - 0 1 1

4.6 Solid tin(II) oxide/hydroxide

SnO(cr) is more soluble than SnO2(precip). Dissolved tin(II) species precipitate from solutionas Sn(OH)2(precip). The hydroxide is unstable with respect to the oxide, SnO(cr) and mayundergo progressive dehydration to SnO(cr) or oxidation to SnO2(precip) [1963SHA/DAV,1984HOU/KEL and references therein]. [1906GOL/ECK] observed that the whiteSn(OH)2(precip) precipitates changed to black SnO(cr) in aqueous solutions and that with thischange the amount of dissolved Sn(II) decreased. In contrast, [1955DEL/ZOU] stated thatdehydration of Sn(OH)2(precip) to SnO(cr) is complete only at T « 100 °C. [1955DEL/ZOU]give no experimental evidence for their statement.

Both, the crystalline SnO(cr), and precipitated Sn(OH)2(precip) are quite soluble (Table 4.16).Only few experimental measurements of Sn(II) solubility can be found in the literature. The dataused for the calculations of the formation constants of tin(II) hydroxide/oxide compounds arecompiled in Table 4.16. Additional data are given in Tables 4.17 and 4.18.

Table 4.16: Experimentally determined equilibrium data compiled for the formation of tin(II)hydroxide/oxide compounds. These data were chosen for the evaluation ofrecommended values in the present report. Additional information for the differentreferences see Section 4.14: 'Comments to selected references'. Method: pot =potentiometry, sol = solubility measurements.

log K so Reference Comments I (M) Medium Method

log K*so: Sn2+ + H2O <=> SnO(cr) + 2H+

-2.41 i [1941GAR/HEI] T= 298.15 K, I=dil-2.93 [1977MAR] T= 298.15 K, 1=3

03

waterNaClO.

solpot

log K*so: Sn2+ + 2H2O <=> Sn(OH)2(precip) + 2H+

-2.84 i [1906GOL/ECK] T= 298.15 K, I=dil 0-2.79 [1922PRY] T= 298.15 K, I=dil 0

waterwater

solsol

recalculated from solubility in water and log (3,2 = -7.71; see also text

54

JNC TN8400 99-011

4.6.1 SnO(cr) and Sn(OH)2(precip)

In water [1941GAR/HEI] determined the solubility of SnO(cr) to be 5 x 10-6 M S n ( T I) a n d

reported for Sn(OH)2(precip) a solubility of 1.35 x 10"5 M (determined by [1906GOL/ECK]).From these values and a log (31>2 of -7.71 (see Section 4.5.2) for the reaction Sn2+ + 2H2O <=>Sn(OH)2° + 2H+, the following solubility can be calculated (Figure 4.5)

Sn2+ + H2O SnO(cr) + 2H+ log K*°so = -2.41

For Sn(OH)2(precip) the solubility may be best defined by the mean of the measurements of[1922PRY2] and [1906GOL/ECK] (Table 4.16):

Sn2+ + 2H2O o Sn(OH)2(precip) + 2H+ log K*°so = -2.82

The values calculated by [1941GAR/HEI] themselves are different (Table 4.15 and 4.16) as[1941GAR/HEI] used a different log p u value.

Sn2+ + H2O *=> SnO(s) + 2H+

lm, molal

Figure 4.5: Plot of log K*so + 2 D vs. Im for the reaction : Sn2+ + H2O <=> SnO(s) + H+ at25 °C. The straight line shows the result of the 'linear regression': Ae = -0.01; logK*°so = -2.41. Calculated from data given in Table 4.16.

55

JNC TN8400 9 9 - 0 1 1

4.6.2 Additional equilibrium data compiled for tin(ll) hydroxide/oxide compounds

Table 4.17: Additional experimentally determined equilibrium data compiled for the formation of tin(II)hydroxide/oxide compounds. These data were not chosen in the present report for the evaluationof recommended stability values. Reasons for not selecting these references are given in Section4.14: 'Comments to selected references'. Method: sol = solubility measurements.


log K*s0: Sn2+ + H2O <=> SnO(s) + 2H+

-1.76 '-2.30

[1941GAR/HEI]H965MES/IRA1

T= 298.15 K, I=dilT= 298.15 K, 1=1

01

waterNaC104

sol

sol

log K*s0: Sn2+ + 2H2O <=> Sn(OH)2(precipitated) + 2H+

-2.19 ' [1941GAR/HEI]0.33 2 [1995DJU/JEL1

T= 298.15 K,I=dilT= 298.15 K, 1=3

waterNaCl pot

1 formation of polymeric species neglected;2 formation of tin chloride complexes probable

Table 4.18: Thermodynamic data compiled for the formation of tin(II) hydroxide/oxide compounds taken fromprevious compilations. As pointed out in Section 2 of this report only experimental data wereused for the present evaluation. The following table serves only for comparison.

log K*so

log K*so:

-1.12-1.08-1.17 '-1.07-1.80-1.17 l

-1.10 '-1.17 *-1.76-1.80-1.38-101.4 '-1.76-1.12 '-1.30-1.04-1.30-1.29-1.12 '-2.17-1.31

Reference

Sn2+ + H2O <=> SnO(s) •

[1937HOA][1952LAT][1954COU][1955DEL7ZOU][1963FEI/SCH][1963WIC/BLO][1964HIR][1971NAU/RYZ][1976BAE/MES],[1976SMI/MAR][1978COD][1979KUB/ALC][1981BAE/MES][1982PAN][1982WAG/EVA][1984HOU/KEL][1985BAB/MAT][1985GAL][1985JAC/HEL2][1986KOV/RYZ][1988PHL«AL1

Comments

f 2H+

T= 298.15 K,I=n/aT=298.15 K, I=n/aT=298.15K, I=n/aT= 298.15 K, 1=0T= 293 K, 1=0T= 298.15 K, I=n/aT=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT=298.15K, I=n/aT=298.15K, I=n/aT=298.15K, I=n/aT= 298.15 K, I=n/a

KM)

00

00

0

56

JNC TN8400 9 9 - 0 1 1


-2.25 [1989COX/WAG] T= 298.15 K, I=n/a-1.76 ri992PEA/BER] T= 298.15 K, 1=0

log K*so- Sn2+ + 2H2O <=> Sn(OH)2(precipitated) + 2H+

-1.51-1.5-2.21-1.43-1.73-1.67-2.42-1.75

[1952LAT][1955DEL/ZOU][1980BEN/TEA][1984HOU/KEL][1985BAB/MAT][1985GAL][1987BROAVAN][1988PHI/HAL]

T=298.15T= 298.15T= 298.15T= 298.15T= 298.15T=298.15T= 298.15T= 298.15

K, I=n/aK, 1=0K, I=n/aK, 1=0K, I=n/a

K, I=n/aK,I=0K, I=n7a

calculated with a AfG° of -26.42 kJ/mol for Sn2+ (Section 4.4.1)

57

JNC TN8400 9 9 - 0 1 1

4.7 Tin(II) chloride system

Tin(II) forms complexes and compounds with chloride. Also the existence of mixed tin(II)chloride hydroxide complexes and compounds is reported. Experimentally determined data aregiven in Table 4.19. Further thermodynamic data are compiled in Table 4.20 and 4.21.

Table 4.19: Experimentally determined equilibrium data compiled for the tin(II) chloridesystem. These data were chosen for the evaluation of recommended values in thepresent report. Additional information for the different references see Section4.14: 'Comments to selected references'. Method: pot = potentiometry, sol =solubility measurements.

log Pm,n

log Pif-1.111,06 11.14 2

1.451,091.021.181.341.801.080.73

Reference

Sn2+ + Cl- & SnCl+

[1951DUK/PIN][1961RAB/MOO][1961RAB/MOO][1962HAJ/ZOL][1975FED/BOL][1975FED/BOL][1975FED/BOL][1975FED/BOL][1975FED/BOL][1976SAM/LYA][1981PET/MEL]

Comments

T= 298.15 K, 1=2T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=6T= 298.15 K, 1=1T= 298.15 K, 1=1

KM:

223

40.5

134611

) Medium

HC1O4

cio;NaC104

H2SO4/HC1NaC104

NaC104

NaClO4

NaC104

NaC104

NaC104

NaNO,

Method

potpotpot

solpotpotpotpotpotpotpot

log Pi,2: Sn2+ + 2CI- <=> SnCl2°

1.72 11.70 2

2.351.361.131.782.123.041.851.08

[1961RAB/MOO][1961RAB/MOO][1962HAI/ZOL][1975FED/BOL][1975FED/BOL][1975FED/BOL][1975FED/BOL][1975FED/BOL][1976SAM/LYA][1981PET/MEL]

T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=6T= 298.15 K, 1=1T= 298.15 K, 1=1

234

0.5134611

cio;NaC104

H2SO4/HC1NaClO4

NaC104

NaClO4

NaClO4

NaC104

NaC104

NaNO,

potpotsol

pot

potpotpotpotpotpot

58

JNC TN8400 9 9 - 0 1 1


log PJi3: Sn2+ + 3CI- a SnClf

1.50 ]

1.66 2

2.461.652.123.30

log Pi,4:

2.31

log Pi,1,1

-2.77 3

-2.90

log K*so:

1.38 4

1.37 4

1.38 4

1.38 4

1.79 4

1.70 4

2.25 4

2.24 4

2.25 4

2.25 4

2.25 4

2.24 4

1.82 4

1.67 4

2.00 4

1.47 4

1.87 4

1.80 4

1.78 4

1.77 4

1.85 4

[1961RAB/MOO][1961RAB/MOO][1962HAI/ZOL][1975FED/BOL][1975FED/BOL][1975FED/BOL]

Sn2+ + 4CI- <=> SnCU

[1962HAI/ZOL]

: Sn2+ + H2O + Cl- 4=

[1952VAN/RHO][1981PET/MIL]

Sn2+ + H2O + Cl- <=t

[1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR]

T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=6

r7-

T= 298.15 K, 1=4

* SnOHCl°+ H+

T= 298.15 K, 1=3T= 298.15 K, 1=0.5

> SnOHCl(s) + H+

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/a

234346

4

30.5

0.1690.170.170.1790.0580.0680.0180.0150.0150.0140.0140.0140.0880.0820.0390.130.0450.0480.0480.0530.048

cio;NaC104

H2SO4/HC1NaC104

NaC104

NaClO4

H7SO4/HC1

NaC104

NaCl

Cl-Cl-Cl-Cl-Cl-Cl-

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-Cl-

potpotsolpotpotpot

sol

potpot

solsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsol

1 recalculation of measurements from [1950DUK/COU]2 recalculation of measurements from [1952VAN/RHO]3 calculated in this report from a log Ki i U of 1.04 for SnOH+ + CY <=> SnOHCl0 [1952VAN/RHO] and a log

(3U of -3.81 for Sn2+ + H2O <=> SnOH+ + H+ at I = 3 (see Table 4.13)4 recalculated in this report from measurements of [1930RAN7MUR]. Values have been corrected for the

formation of Sn chloride complexes. Uncorrected data are given in Table 4.20

59

JNC TN8400 9 9 - 0 1 1

4.7.1 SnCl+, SnCl2°, SnClf, and SnCl42-

A number of equilibrium constants for the formation of tin(II) chloride complexes can be foundin the literature. As early as 1928, [1928PRY] determined with potentiometric measurementsformation constants for tin(II) chloride complexes and corrected these values to 1=0.[1961RAB/MOO] recalculated the experiments of [1950DUK/COU] and [1952VANRHO] witha non-linear least square fit. From the data given in Table 4.19 the following formationconstants for the tin chloride complexes can be derived with the SIT equation (see also Figure4.6, 4.7 and 4.8):

Sn2+HSn2+HSn2+H

hCl-h2Cl-h3Cl"

<=» SnCl+

« SnCl2°«=> SnCl3-

log P°U

log P°i,2log P°i,3

= 1.65= 2.31= 2.09

These constants extrapolated with the SIT equation are comparable with the constantsdetermined by [1928PRY] for I = 0 (Table 4.20). No equilibrium constant is recommended inthis report for the reaction Sn2+ + 4C1" <=> SnCU2" as this constant was only determined at I = 4(Table 4.19). The Ae values calculated agree well with the Ae values calculated for Pb(II)chloride complexes (see Section 6.3).

c0

4.5

4

Q3.5

+ 3

c i 2 - 5

g) 21.5

1

0.5

n

Sn2* + CI-

0

y = 0.14x + 1

» SnCI+

.65

0 2 4 6 8

lm, molal10

Figure 4.6: Plot of log P u + 4 D vs. Im for the reaction : Sn2+ + Cl' <=* SnCl+ at 25 °C. Thestraight line shows the result of the linear regression: Ae = -0.14; log (3\ i = 1.65.Calculated from data given in Table 4.19.

60

JNC TN8400 99-011

QCO

+CM,-T

CO.

_g

c0

4.5

4

3.5

3

2.5

2

1.5

1

0.5 •

n •

Sn2+ + 2CI- «

O O^ô

y =

1 1

> SnCI2°

0.26X + 2.31

H 1

4 6lm> molal

8 10

Figure 4.7: Plot of log (3]>2 + 6 D vs. Im for the reaction : Sn2+ + 2C1- <=> SnCl2° at 25 °C.The straight line shows the result of the linear regression: Ae = -0.26; log (3°ii2 =2.31. Calculated from data given in Table 4.19.

Sn2+ + 3CI"

y = 0.29x + 2.09

4 6lm> molal

10

Figure 4.8: Plot of log (3li3 + 6 D vs. Im for the reaction : Sn2+ + 3C1" «=> SnCl3- at 25 °C.The straight line shows the result of the linear regression: AE = -0.29; log P \ 3 =2.09. Calculated from data given in Table 4.19.

61

J N C T N 8 4 0 0 99 - O i l

4.7.2 SnOHCl0

A mixed tin hydroxide chloride complex can be formed. [1952VAN/RHO] gives for the reactionSnOH+ + Cl- « SnOHCl0 a log K U ; i of 1.04 and [1981PET/MIL] gives a log K u , i of 1.14for the same reaction and a log J5ITI,I of -2.9 for the reaction Sn2+ + H2O + Cl" <=> SnOHCl0 +H+. Extrapolation to I = with the SIT equation in Figure 4.9 gives

Sn2+ + H2O + Cl- SnOHCl0 + H+ (3°i,i,i = - 2 . 2 7

Q

+

oa

0 -

-0.5 -

-1

-1.5

-2

-2.5

-3

-3.5

. 4 ,

-4.5 -

-5 1

Sn2 ++H2O+CI" <=>

— & - — — • " "

y = 0.12x

• 1 h

SnOHCI°+H+

-——-—"""^

-2.27

1

2mola!

Figure 4.9: Plot of log p u > 1 + 4 D vs. Im for the reaction : Sn2+ + H2O + Cl- <=> SnOHCl0

+ H+ at 25 °C. The straight line shows the result of the 'linear regression': As = -0.12; log (3O

1)U = -2.27. Calculated from data given in Table 4.19.

62

JNC TN8400 9 9 - 0 1 1

4.7.3 SnCl2(s)

Direct measurements of tin chloride solubility are not available. Based on the data compiled inTable 4.21 it can be concluded that SnCl2(s) is easily soluble. Under environmental conditions,SnOHCl(s) is thermodynamically more stable.

4.7.4 SnOHCl(s)

[1930RAN/MUR] determined the solubility of SnOHCl(s) in water (pH range form 1.1 to 2.2).Their experimental results are given in Table 4.19. [1930RAN/MUR] extrapolated their resultsto 1=0 with the Debye-Hiickel term, resulting in a log K*°So of 2.75 (Table 4.20). In this report,the values reported by [1930RAN/MUR] are corrected for the formation of Sn(II) chloridecomplexes. Extrapolation with the SIT equation to 1=0 gives (Figure 4.10):

Sn2++ H2O SnOHCl(s) + H+ log K*°so = 2.42

Sn2++H2O+CI- & SnOHCI(s)+H+

Q

+o0)

log

3 -I

2.5

2

1.5

1

0.5

0-

-0.5

-1

-1.5

C

y = -3.28x + 2.42

) 0.1 0.2 0.3 0.4lm, molal

0

Figure 4.10: Plot of log K*so + 4 D vs. Im for the reaction : Sn2+ + H2O + Cl" <=> SnOHCl(s)+ H+ at 25 °C. The straight line shows the result of the linear regression: Ae =3.28; log K*°so = 2.42. Calculated from the data given in Table 4.19.

63

JNC TN8400 9 9 - 0 1 1

4. 7.5 Additional equilibrium data compiled for the tin(II) chloride system

Table 4.20: Additional experimentally determined equilibrium data compiled for the tin(ll) chloride hydroxidesystem. These data were not chosen in the present report for the evaluation of recommendedstability values. Reasons for not selecting these references are given at the end of this table.Method: pot = potentiometry, sol = solubility measurements.

Reference Comments I (M) Medium Method

log Pu: Sn2+ + Cl- <=> SnCl"

1.51 »-2 [1928PRY]1.85 2 [1949RIC/POP]1.05 3 [1950DUK/COU]1.15 3 [1952VAN/RHO]0.704 [1962HAI/ZOL]1.87 2 [1975EED/BOL1

T= 298.15 K, 1=0.08-0.35 0T= 298.15 K, 1=0.1-4 0T= 298.15 K, 1=2 2T= 298.15 K, 1=3 3T= 298.15 K, 1=2-6 varT= 298.15 K, 1=0 0

HC1, KC1KC1

NaC104/HC104

NaC104

HC1NaC104

potpotpotpotsolpot

log P,t2- Sn2+ + 2CI <=> SnCl2°

2.25 I-2.31 2

1.76 3

1.70 3

0.78 4

2.38 2

[1928PRY][1949RIC/POP][1950DUK/COU][1952VAN/RHO][1962HAI/ZOL][1975FED/BOL]


HC1, KC1KC1

NaC104/HC104

NaC104

HC1NaC104

potpotpotpotsolpot

log Pu: Sn2+ + 3d <=P SnClf

2.03 1-2 [1928PRY]1.94 2 [1949RIC/POP]1.14 3 [1950DUK/COU]

[ 1952V AN/RHO][1962HAIÔL][1975FED/BOL1

1.68 3

0.38 4

1.93 2


HC1, KC1KC1

NaC104/HC104

NaC104

HC1NaCIO,

potpotpotpotsolpot

log Pl4:

1.50 >'2

2.00 2

1.14 3

2.31-0.24 4

log P,,u

-0.66 4

-1.35 5

Sn2+ + 4CI <=> SnCl42-

[1928PRY][1949RIC/POP][1950DUK/COU][1962HAI/ZOL][1962HAI/ZOL]

.• Sn2+ + H2O + Cl- <=>

[1952V AN/RHO][1986KOV/RYZ]

T= 298.15 K, 1=0.08-0.35T= 298.15 K, 1=0.1-4T= 298.15 K, 1=2T= 298.15 K, 1=4T=298.15K, 1=2-6

SnOHCl°+ H+

T=298.15K, 1=3T= 298.15 K, 1=0

0024

var

30

HC1, KC1KC1

NaC104/HC104

H2SO4/HC1HC1

NaC104

water

potpotpotsolsol

potsol

64

JNC TN8400 9 9 - 0 1 1


log K'so: i

1.29 7

1.27 7

1.28 7

1.28 7

1.76 7

1.66 7

2.24 7

2.23 7

2.23 7

2.24 7

2.24 7

2.23 7

1.777

1.62 7

1.97 7

1.39 7

1.84 7

1.77 7

1.75 7

1.74 7

1.82 7

2.75 2

5/2/+ + H2O + Ci <=>

[1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][193ORAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][1930RAN/MUR][193ORAN/MUR][1930RAN/MUR][1930RAN/MUR1

SnOHCl(s) + H+

T= 298.15 K,I=n/aT= 298.15 K.I^n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/a

0.1690.170.17

0.1790.0580.0680.0180.0150.0150.0140.0140.0140.0880.0820.0390.13

0.0450.0480.0480.0530.048

0

Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-CI-Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-Cl-

solsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsol

1 measured with Calomel electrode2 extrapolated to 1=0 by the respective authors with Debye-Huckel equation3 these values were recalculated by [1961RAB/MOO], see Table 4.194 I not constant5 calculated by [1952VAN/RHO] without considering polymeric tin(II) hydroxide species: Corrected value in6 extrapolated from measurement at 500 °C7 formation of Sn chloride complexes not considered. Constants corrected for the formation of Sn chloride

are given in Table 4.18

Table 4.19

complexes

Table 4.21: Thermodynamic data compiled for the tin(II) chloride hydroxide system taken from previouscompilations. As pointed out in Section 2 of this report only experimental data were used for thepresent evaluation. The following table serves only for comparison. Medium: Where data refer tospecific electrolyte solutions, this is indicated.


log 0U: Sn2+ + Ci SnCl+

1.141.511.081.171.451.641.001.011.061.161.06

[1967AHR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976V AS/KOK][1976V AS/KOK][1976V AS/KOK][1976V AS/KOK][1976 V AS/KOK][1979VAS/GLA]

T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T=298.15K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=2

302340

0.512

32

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

65

JNC TN8400 9 9 - 0 1 1


1.161.281.641.051.051.081.171.401.801.121.291.501.80 1

2.081.760.62

log /?/>2-

2.251.721.722.352.431.471.511.611.791.571.721.902.431.421.501.681.752.243.041.732.062.262.253.461.721.43

[ 1979V AS/GLA][1979VAS/GLA][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1984HOU/KEL][1985BAB/MAT][1985GAL][1987BRO/WAN],[1988PHI/HAL][1992PEA/BER]

Sn2+ + 2CI- & SnCl2°

[1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976VAS/KOK][1976V AS/KOK][1976VAS/KOK][1976V AS/KOK][1976 V AS/KOK][1979VAS/GLA][1979VAS/GLA][1979VAS/GLA][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982 WAG/EVA][1984HOU/KEL][1985BAB/MAT][1986KOV/RYZ][1987BRO/WAN],[1988PHI/HAL][1992PEA/BER]

T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T=298.15K, 1=6T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

T= 298.15 K, 1=0T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0.5T=298.15K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=41- 298.15 K, 1=6T=298.15K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T=298.15K, 1=0

340

0.5123463000000

02340

0.51232340

0.5123463000000

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

log Pu: Sn2+ + 3Cl- <=> SnClf

2.001.501.702.501.400.44

[1976SMI/MAR][1976SMI/MAR][1976SMIMAR][1976SMI/MAR][1976V AS/KOK][1976V AS/KOK]

T= 298.15 K, 1=0T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0.5

02340

0.5HC1O4

HC1O4

66

JNC TN8400 9 9 - 0 1 1


0.781.181.661.191.682.191.301.672.243.301.602.022.064.262.050.88

[1976VAS/KOK][ 1976V AS/KOK][ 1976V AS/KOK][1979VAS/GLA][1979VAS/GLA][ 1979V AS/GLA][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1984HOU/KEL][1985BAB/MAT][1987BROAVAN],[1988PHI/HAL][1992PEA/BER]

T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=6T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

1232342346300000

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

log Pl4: Sn2+ + 4CI- <=> SnCl42-

1.81 [1975KRA]1.50 [1976SMI/MAR]2.30 [1976SMI/MAR]1.70 [1984HOU/KEL]4.54 [1987BROAVAN],1.47 [1988PHI/HAL1

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

004000

log Pi

-0.62-0.41-0.64

Sn2+ + H2O + Ci <=> SnOHCl°+ H+

[1982WAG/EVA][1984HOU/KEL]ri988PHI/HAL1

T= 298.15 K, I=n/aT=298.15K, 1=0T= 298.15 K, I=n/a-

log Ks0: Sn2+ + 2CI <=> SnCl2(s)

2.31.9 >-0.1-0.12.12.22.22.2

[1952LAT] T=[1963WIC/BLO] T=[1977BAR/KNA] T=[1979KUB/ALC] T=[1984HOU/KEL] T=[1985BAB/MAT] T=[1985GAL] T=[1988PHI/HAL] T=

298.15 K,I=n/a298.15 K,I=n/a298.15 K,I=n/a298.15 K,I=n/a298.15 K, 1=0298.15 K, 1=0298.15 K, 1=0298.15 K, 1=0

0000

log K'so: Sn2+ + H2O + Ci <=> SnOHCl(s) + H+

2.28 [1980BEN/TEA]2.75 [1982WAG/EVA]3.00 [1984HOU/KEL]2.73 [1988PH1/HAL1

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/a

1 calculated with a A(G° of -26.42 kJ/mol for Sn2+ (Section 4.4.1)

67

JNC TN8400 9 9 - 0 1 1

4.8 Tin(II) fluoride system

Tin(II) forms complexes and compounds with fluoride. Experimentally determined data aregiven in Table 4.22. Further thermodynamic data are compiled in Table 4.23 and 4.24.

Table 4.22: Experimentally determined equilibrium data compiled for the tin(II) fluoridesystem. These data were chosen for the evaluation of recommended values in thepresent report. Method: pot = potentiometry.

log (3

log Pij:

4.14 14.05 l

4.004.48 2

3.60

logfr.2:6.858.18 27.04

Reference

Sn2+ + F- <=> SnF+

[1961CON/PAU][1961CON/PAU][1970BONATAY][1971BON1][1975NEL/AM]

Sn2+ + 2F- a SnF2°

[1970BON/TAY][1971BON1][1975NEL/AMI]

Comments

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

T= 298.15T= 298.15T= 298.15

K, 1=0.5K, 1=2K, 1=1K, 1=0.85K, 1=0.1

K, 1=1K, 1=0.85K, 1=0.1

KM)

0.521

0.850.1

10.850.1

Medium

HC1O4

HC1O4

NaC104

NaC104

NaF

NaC104

NaClO4

NaF

Method

potpotpotpotpot

potpotpot

log pli3: Sn2+ + 3F- <=> SnF3-

9.43 [1970BON/TAY] T= 298.15 K, 1=110.30 2 [1971BON1] T= 298.15 K, 1=0.859.00 [1975NEL/AMJJ T= 298.15 K, 1=0.1

10.850.1

NaC104

NaC104

NaF

potpotpot

1 mean of several determinations, no higher Sn(II) fluoride complexes considered. Calculated using a pK(HF/F")of 2.92 (1=0.5) and 3.06 (1=2).

2 recalculation of measurements of [ 1968HAL/SLA] by [ 1971 BON 1 ].

68

JNC TN8400 9 9 - 0 1 1

4.8.1 SnF+, SnF2°, and SnF3-

Equilibrium constants for the formation of tin(II) fluoride complexes found in the literature aregiven in Tables 4.22 and 4.23. From the data given in Table 4.22 the following formationconstants for the tin fluoride complexes can be derived with the SIT equation at 1=0 (see alsoFigure 4.11 to 4.13):

log .P° u =4.47log 3°i,2 = 7.74log (3°i,3 = 9.61

As the data at higher ionic strength (Figures 4.11 to 4.13) show a considerable spread, thegiven constants should be regarded as preliminary.

S n 2 + -

S n 2 + .Sn2+H

h F"h2Fh3F

<=> SnF+SnF2°SnF3-

Sn2+ SnF+

7.5 -•

7 ••

Q 6.5

+ 6 +

oa5.5

3.5

3

y = 0.32x + 4.47

0.5 1 1.5lm, molal

2.5

Figure 4.11: Plot of log (3U + 4 D vs. Im for the reaction : Sn2+ + F" <=> SnF+ at 25 °C. Thestraight line shows the result of the linear regression: Ae = -0.32; log (3°],] = 4.47.Calculated from data given in Table 4.22.

69

JNC TN8400 9 9 - 0 1 1

SnF,°

1 1.5L, molal

Figure 4.12: Plot of log (3U + 6 D vs. Im for the reaction : Sn2+ + 2F" <=> SnF2° at 25 °C. Thestraight line shows the result of the linear regression: Ae = -0.91; log 3°ii2 = 7.74.Calculated from data given in Table 4.22.

13Sn2+ + 3 P <=> SnF,

0.5 1 1.5lm> molal

2.5

Figure 4.13: Plot of log p u + 6 D vs. Im for the reaction : Sn2+ + 3P <=> SnF3- at 25 °C. Thestraight line shows the result of the linear regression: Ae = -1.40; log (3°i)3 = 9.61.Calculated from data given in Table 4.22.

70

JNC TN8400 9 9 - 0 1 1

4.8.2 SnF2(s)

Direct measurements of the solubility of SnF2(s) are not available. Based on the data compiledof [1963WIC/BLO] (Table 4.24) it can be expected that SnF2(s) is rather soluble.

4.8.3 Additional equilibrium data compiled for the tin(II) fluoride system

Table 4.23: Additional, experimentally determined equilibrium data compiled for tin(II) fluoride system. Thesedata were not chosen in the present report for the evaluation of recommended stability values.Method: pot = potentiometry.

log Pm,n,o Reference Comments I (M) Medium Method

log Pij: Sn2+ + F- <=> SnF+

6.26 ' ri968HAL/SLA1 T= 298.15 K, 1=0.85 0.85 NaC104 pot

log Pia: Sn2+ + 2F- <=> SnF2°

8.76 ' ri968HAL/SLAl T= 298.15 K, 1=0.85 0.85 NaC104 pot

log pu: Sn2+ + 3F- <=> SnFf

9.25 l [1968HAL/SLA1 T= 298.15 K, 1=0.85 0.85 NaCIO, pot

these values were recalculated by [1971BON1], see Table 4.22

Table 4.24: Thermodynamic data compiled for the formation of tin(II) fluoride system. As pointed out inSection 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison. Medium: Where data refer to specific electrolytesolutions, this is indicated.

log P Reference Comments I (M) Medium

log Pu: Sn2+ + F- <=> SnF+

4.08 [1976SMI/MAR]4.45 ' [1980BON/HEF]4.16 [1980BEN/TEA]4.62 [1982WAG/EVA]3.52 [1987BRO/WAN]6.9 [1992PEA/BER]

T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

110000

NaC104

log P,_2: Sn2+ + 2F- <=> SnF2°

6.68 [1976SMI/MAR]7.43 ' [1980BON/HEF]6.59 [1987BROAVAN]9.7 [1992PE A/BER]

T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0

1100

NaC104

71

JNC TN8400 9 9 - 0 1 1


log (3!3: Sn2+ + 3F- <=> SnFy

T= 298.15 K, 1=1 1NaC104

9.511.10 '9.3310.20

[1976SMI/MAR][1980BON/HEF][1987BROAVAN]fl992PEA/BERl

T= 298.15 K,T= 298.15 K,T= 298.15 K,T=298.15K,

1=11=11=01=0

1100

log pli4: Sn2+ + 4F- <=> SnF42-

9.81 [1987BROAVAN1 T== 298.15 K, 1=0

log K*so: Sn2+ + 2F- <=> SnF2(s)

5.6 2 ri963WIC/BL01 T= 298.15 K, I=n/a1 calculated using a pK (HF/F) of 2.94 (1=1).2 calculated with a Afi° of -26.42 kJ/mol for Sn2+ (Section 4.4.1)

72

JNC TN8400 9 9 - 0 1 1

4.9 Tin(II) carbonate system

No experimental data are available on complex formation between tin(H) and carbonate species.It is probable that Sn2+, similar to Pb2+ (cf. Section 6.6) forms complexes with carbonate andbicarbonate ions. Thermodynamic data calculated by [1987BROAVAN] are compiled in Table4.25.

Table 4.25: Thermodynamic data for the Sn(II) carbonate system, taken from [1987BRO/WAN]. Thefollowing table serves only for comparison.

log $i i log Pi 2 log Pi 3 log P, 4 References Comments

log pm,n: mSn2+ + nCO32- <=> Snm(CO3)n

2m-2n

9.72 17.85 24.65 30.21 [1987BRO/WAN] T= 298.15 K, 1=0

log pmn: mSn2+ + nHCOf <=> Snm(HCO3)n2">-"

4.57 7.66 9.51 10.24 fl987BROAVAN1 T= 298.15 K, 1=0

log K*so Reference Comments

log K*so: Sn2+ + CO3 <=> Sn(CO3)(s)

10.95 ri987BROAVANl T= 298.15 K, 1=0

73

JNC TN8400 9 9 - 0 1 1

4.10 Tin(II) nitrate system

Tin(II) forms weak complexes and compounds with nitrate. Experimental data were determinedby [1980AND/SAM] (Table 4.26). Further thermodynamic data are compiled in Table 4.27 and4.28.

Table 4.26: Experimentally determined equilibrium data compiled for the tin(II) nitrate system.These data were chosen for the evaluation of recommended values in the presentreport. Method: pot = potentiometry.

!„„ n Reference1U5 Pm,n,o

Comments I (M) Medium Method

log Pu: Sn2+ + NOf <=> SnNO3"

0.440.410.140.150.18

log Pi,:

0.450.05-0.060

[1980AND/SAM][1980AND/SAM][1980AND/SAM][1980AND/SAM][1980AND/SAM]

>; Sn2+ + 2NOf t=>

[1980AND/SAM][1980AND/SAM][1980AND/SAM][1980AND/SAM]

T= 298.T= 298.T= 298.T= 298.T= 298.

Sn(NO3)2°

T= 298.T= 298.T= 298.T= 298.

1515151515

15151515

K, 1=1K, 1=2K,I=3K, 1=4K, 1=6

K, 1=2K, 1=3K, 1=4K, 1=6

12346

2346

NaNO3NaNO3

NaNO3

NaNO3

NaNO3

NaNO3

NaNO3

NaNO3

NaNO3

potpotpotpotpot

potpotpotpot

log PU: Sn2+ + 3NO3- <=> Sn(NO3)f

-0.35 [1980AND/SAM] T= 298.15 K, 1=3-0.58 [1980AND/SAM] T= 298.15 K, 1=4-0.85 [1980AND/SAM] T= 298.15 K, 1=6

346

NaNO3NaNO3

NaNO3

potpotpot

log pli4: Sn2+ + 4NO3- & Sn(NO3)42-

-0.98 [1980AND/SAM] T= 298.15 K, 1=4-1.2 [1980AND/SAM] T= 298.15 K, 1=6

46

NaNO3

NaNO3

potpot

74

JNC TN8400 9 9 - 0 1 1

4.10.1 SnNO3+, Sn(NO3)2°, Sn(NO3)3- and Sn(NO3)42-

From the experimentally determined data by [1980AND/SAM] in 1, 2, 3, 4, and 6 M NaNO3

(Table 4.26) the following formation constants for the tin nitrate complexes can be derived (seealso Figures 4.14 to 4.17):

Sn2+ + NO3-Sn2+ + 2NO3-Sn2+ + 3NO3-Sn2+ + 4NO3-

SnNO3+Sn(NO3)2°Sn(NO3)3-Sn(NO3)4

2+

log P°i.i = 1-25log P°i,2= 1-74log P° l i 3 = 1.37log P° l i 2 = 0.30

Sn2+ + NO SnNO,

Figure 4.14: Plot of log | 3 U + 4 D vs. Im for the reaction : Sn2+ + NO3" <=> SnNO3+ at 25 °C.

The straight line shows the result of the linear regression: Ae = 0.02; log (3°ii =1.25. Calculated from data given in Table 4.26.

75

JNC TN8400 9 9 - 0 1 1

QCD

+

CO.

O)

.o

4.5 •

4

3.5

3

2.5

2

1.5

1

0.5

n

Sn2+ + 2NO 3 - <=>

y = -0.05x +

" • " o : " • ' ••"••'

° o

: J 1 _ _

Sn(NO3)2°

1.74

— — Q_

1

4

L molal

Figure 4.15: Plot of log p1>2 + 6 D vs. Im for the reaction : Sn2+ + 2NO3- o Sn(NO3)2° at 25°C. The straight line shows the result of the linear regression: Ae = 0.05; log (3°ii2= 1.74. Calculated from data given in Table 4.26.

Sn(NO3)3

Figure 4.16: Plot of log p1>3 + 6 D vs. Im for the reaction : Sn2+ + 3NO3" <=> Sn(NO3)3" at 25°C. The straight line shows the result of the linear regression: Ae = 0.12; log (3°i3

= 1.37. Calculated from data given in Table 4.26.

76

JNC TN8400 9 9 - 0 1 1

Sn2+ + 4NO3" <=> Sn(NO3)4;2-

lm, molal

Figure 4.17: Plot of log p1>4 + 4 D vs. Im for the reaction : Sn2+ + 4NO3" <=> Sn(NO3)42- at 25

°C. The straight line shows the tentative result of the 'linear regression': Ae =0.11; log p°1>4 « 0.30. Calculated from data given in Table 4.26.

4.10.2 Tin(II) nitrate compounds

No thermodynamic data about tin(II) nitrate compounds are available. It can be expected, thatthey are very easily soluble and will not limit for Srt(II) solubility.

4.10.3 Additional equilibrium data compiled for tin(II) nitrate system

Table 4.27: Additional, experimentally determined equilibrium data compiled for tin(II) nitrate system. Thesedata were not chosen in the present report for the evaluation of recommended stability values.Method: pot = potentiometry.

log (3m,n,0 Reference Comments I (M) Medium Method

log (5,.,: Sn2+ + NOy <=> SnNO3

1.1 [1980AND/SAM]0.43 ' [1980AND/SAM]

T= 298.15 K, 1=0T=298.15K, 1=8

NaNO3 potpot

log f3u: Sn2+ + 2NO3- <=> Sn(NO})20

0.57 ' [1980AND/SAM] T= 298.15 K, 1=8 pot

77

JNC TN8400 9 9 - 0 1 1


log p13: Sn2+ + 3NO3- <=> Sn(NO3)3-

0.46 ' [198QAND/SAM1 T= 298.15 K, 1=8 8 NaNCh pot

log P,,4: Sn2+ + 4NO3 <^> Sn(NO3)42-

-0.01 ' [1980AND/SAM1 T= 298.15 K, 1=8 8 NaNO^ pot1 the values determined at 1=8 are quite different from the others and are therefore not chosen for extrapolation to 1=0 in

this report.

Table 4.28: Thermodynamic data compiled for the formation of tin(II) nitrate system. . As pointed out inSection 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

l°g Pm.n,o Reference Comments I (M)

log PJJ: Sn2+ + NOf <=> SnNO3+

0.44 [1976SMI/MAR] T= 298.15 K, 1=1 11.90 ri987BROAVAN1 T= 298.15 K, 1=0 0_

log PL2: Sn2+ + 2NOy <=> Sn(NO3)2°

2.25 [1987BRO/WAN1 T= 298.15 K, 1=0

log Pu: Sn2+ + 3NOf <=> Sn(NO3)f

1.32 f!987BRO/WAN1 T= 298.15 K, 1=0

log P14: Sn2+ + 4NO3- <=> Sn(NO3)42-

-0.80 [1987BRO/WAN] T= 298.15 K, 1=0

78

JNC TN8400 9 9 - 0 1 1

4.11 Tin(H) phosphate system

[1969AWA/KAS] determined the redox potential of the reaction 3Sn(cr) + 2PO43" <=>

Sn3(PO4)2(s) + 6e" in phosphate solution. From the standard potential of -0.865 V given by[1969AWA/KAS], a log K*So of 73.93 for the reaction 3Sn2+ + 2PO4

3- <^ Sn3(PO4)2(s)(Table 4.29) can be calculated. The usage of this value is not recommended as the ionic strengthin the experiments of [1969AWA/KAS] is varied strongly.

[1969AWA/KAS] also proposed the formation of tin phosphate complexes without being ableto calculate formation constants. No other experimental data are available about complexformation between tin(II) and phosphate. Thermodynamic data given in previous compilationsare compiled in Table 4.30.

Table 4.29: Experimentally determined equilibrium data compiled for tin(II) phosphate system. These datawere not chosen in the present report for the evaluation of recommended stability values. Method:pot = potentiometry.


log K*so: 3Sn2+ + 2PO43- <=> Sn3(PO4)2(s)

73.84' [1969AWA/KASJ T = 303 K, 1=0.0005-2 var Na3PO4 pot1 calculated from E° = -0.865 V for 3Sn(cr) + 2PO4

3" <=> Sn3(PO4)2(s)+ 6e- and log K (Sn2+/Sn(cr)) of 4.63 (Section4.4.1)

Table 4.30: Thermodynamic data for the Sn(II) phosphate system, taken from previous compilations. Aspointed out in Section 2 of this report' only experimental data were used for the presentevaluation. The following table serves only for comparison.

log $! log P2 log P3 log p4 References Comments

log ft: Sn2+ + nPO/' + nH+ <=> Sn(HPO4)2-2n

18.62 36.40 53.60 70.31 [1987BROAVAN] T= 298.15 K, 1=0

log ft: Sn2+ + nPO/- + 2nH+ <=> Sn(H2PO4)2-n

21.71 42.52 62.70 82.32 [1987BRO/WAN] T= 298.15 K, 1=0


log K*s0: 3Sn2+ + 2PO43- <=> Sn3(PO4)2(s)

71.9 ' [1985GAL1 T= 298.15 K, I=n/a1 calculated with a A(G° of -26.42 kJ/mol for Sn2+ (Section 4.4.1)

79

JNC TN8400 9 9 - 0 1 1

4.12 Tin(II) sulfate system

Similar to Pb(II), Sn(II) forms complexes and compounds with sulfate. Experimental data forcomplex formation of Sn(II) with sulfate were determined by [1981PET/MIL] in 1 M NaNO3

(Table 4.31). Further data are compiled in Table 4.32.

Table 4.31: Experimentally determined equilibrium data compiled for the tin(EE) sulfate system.These data were chosen for the evaluation of recommended values in the presentreport. Method: pot = potentiometry.

R Reference Comments I (M) Medium MethodPm,n,o

log Pi,j: Sn2+ + SO42- & SnSO4°

1.29 [1981PET/MEL] T= 298.15 K, 1=1 1 NaNQ, pot

log pli2: Sn2+ + 2SO42- <^ Sn(SO4)2

2-

1.65 [1981PET/MIL] T= 298.15 K, 1=1 1 NaNO, pot

4.12.1 Tin(II) sulfate complexes

From the experimentally determined data by [1981PET/MTL] in 1 M NaNO3 (Table 4.31) thefollowing tentative formation constants for the tin sulfate complexes can be calculated:

Sn2+ + SO42- «=> SnSO4° log (3°u = 2.91, Ae = - 0.01

Sn2+ + 2SO42- <=> Sn(SO4)2

2- log p°i'i2 = 2.83, Ae = -0.43

The Ae values are assumed to be the same as calculated for Pb (Section 6.9.1: Lead sulfatecomplexes).

80

JNC TN8400 9 9 - 0 1 1

4.12.2 SnSO4(s)

SnSC>4(s) might be expected to be moderately soluble, similar to anglesite (PbSC>4(cr); Section6.9.2). The values compiled in Table 4.32, however, are unusually high, indicating a very lowsolubility and might be somewhat suspect. In these compilations no primary source ofinformation is given. [1987BRU] conducted a specific computer search in Chemical Abstractsfor solubility data about SnSO4(s) which did not turn up any useful information.

Therefore, no solubility product for SnSO4(s) is recommended. The solubility of SnSC>4(s) maybe much larger than indicated by the values compiled in Table 4.32.

Table 4.32: Thermodynamic data compiled for the formation of tin(II) sulfate system taken from previouscompilations. As pointed out in Section 2 of this report only experimental data were used for thepresent evaluation. The following table serves only for comparison.

log f3m,n,0 Reference Comments I (M)

log P,j: Sn2+ + SO42- <=> SnSO4°

3.12 [1987BRO/WAN] T= 298.15 K, 1=01.35 [1988PHI/HAL1 T= 298.15 K, I=n/a

log plX- Sn2+ + 2SO42- <=> Sn(SO4)2

2-

5.02 H987BRO/WAN1 T= 298.15 K, 1=0

log Pu: Sn2+ + 3SO42- <=> Sn(SO4)}

4-

5.94 ri987BRQ/WAN] T= 298.15 K, 1=0

log P!A:

5.99

log K'so

24.24 '24.24 '49.8649.8623.93

Sn2+ + 4SO42- <=> Sn(SO4)f-

[1987BRO/WAN]

.• Sn2+ + SO42- <=> SnSO

[1977BAR/KNA][1979KUB/ALC][1985GAL][1988PHKHAL][1992PEA/BER1

T= 298.15

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=0

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, 1=0

0

0calculated with a AfG° of -26.42 kJ/mol for Sn2+ (Section 4.4.1)

JNC TN8400 9 9 - 0 1 1

4.13 Tin(II) sulfide system

4.13.1 SnS(herzenbergite)

Chemically, SnS(cr), herzenbergite, is expected to be quite insoluble, as indicated by the valuescompiled in Table 4.33. No determination of solubility data could be found in the literature andin the compilations listed in Table 4.33 no primary source of information is given. Therefore,no solubility product for SnS(cr) is recommended in this report.

4.13.2 Sn2Ss(s) and SnjS^s)

Thermodynamic data for the formation of Sn2S3(s) and Sn3S4(s) have been proposed indifferent compilations (Table 4.33). However, no experimental determination of the solubilityof these solids could be found in the literature and in the compilations listed in Table 4.33 noprimary source of information is given. Thus, no solubility products for Sn2S3(s) andare recommended.

Table 4.33: Thermodynamic data compiled for the formation of tin(II) sulfide system taken from previouscompilations. As pointed out in Section 2 of this report only experimental data were used for thepresent evaluation. The following table serves only for comparison.

logK

logK

15.5514.7414.7816.1112.0015.8615.5715.7214.7416.3315.8612.0014.60

logK

41.7041.5841.58

logK

57.5457.5757.57

*so

•

so-'1

1

1

1

1

1

1

*so'J

1

1

*

so-'1

1

1

Reference Comments

Sn2+ + HS- <=> SnS(s, herzenbergite) + H+

[1964HIR][1968SUS/KHO][1971NAU/RYZ][1974MBL][1976SMI/MAR][1978ROB/HEM2][1979KUB/ALC][1980BEN/TEA][1982WAG/EVA][1985GAL][1985JAC/HEL2][1986MYE][1988PHI/HAL]

2Sn2+ + 3HS- <=> Sn2S3(s)

[1974MEL][1977BAR/KNA][1979KUB/ALC1

3Sn2+ + 4HS- <=> Sn3S4 (s)

[1974MIL][1977BAR/KNA][1979KUB/ALC1

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T=298.15K, I=n/aT=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

+ 3H++ 2e-

T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/a

+ 4//++ 2e-

T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, I=n/a

KM)

0

0

0

calculated with a AfG0 of -26.42 kJ/mol for Sn2+ (Section 4.4.1)

82

JNC TN8400 9 9 - 0 1 1

4.14 Comments on selected references

[1939GOR]: [1939GOR] measured Sn(II) solubility under acidic conditions, assuming onlythe presence of Sn2+ and SnOH+ and neglecting the formation of polymeric species.They extrapolated their results to 1=0 with the Debye-Hiickel term.

[1941GAR/HEI]: [1941GAR/HEI] measured Sn(II) solubility in 0 to 0.4 M HC1 and NaOHassuming only the presence of Sn2+, SnOH+, Sn(OH)2° and Sn(OH)3" at variable I.They neglected the formation of polymeric species. Under acidic conditions, theformation of tin(II) chloride complexes cannot be excluded. This work also suffersfrom the difficulty that the free acidity was not measured but was calculated indirectlyfrom assumed equilibria. In water [1941GAR/HEI] determined the solubility ofSnO(cr) to be 5 x 10"6 M and reported for Sn(OH)2(precip) a solubility of 1.35 x 10"5 M (determined by [1906GOL/ECK]). From these values and a log (3°i,2 of -7.71for the reaction Sn2+ + 2H2O <=> Sn(OH)2° + 2H+ (Section 4.5.2):

Sn2+ + H2O <=> SnO(s) + 2H+ log K*°so = -2.41

Sn2++2H2O<=} Sn(OH)2(precipitated) + 2H+ log K*°s0 =-2.84

The values given by [1941GAR/HEI] themselves are different (Table 4.17 and 4.18)as [1941GAR/HEI] used a different log p1>2 value.

[1942GOR/LEI]: [1942GOR/LEI] measured Sn(II) solubility in 0 - 0.1 M HC1O4 at variableI under acidic conditions. The free acidity was not measured but was calculatedindirectly from assumed equilibria.

[1952VAN/RHO]: Sn(II) concentration = 200 mM. Investigated hydrolysis of tin(II) andcomplex formation with chloride. Polymeric tin hydroxide species were not takeninto account by [1952VAN/RHO]. [1952VAN/RHO] give for the reaction SnOH+ +Cl- <=> SnOHCl0 a log K l i U of 1.04. With a log (3U of -3.81 for Sn2+ + H2O <=>SnOH+ + H+ at I = 3 (see Table 4.13) a log (3u , i of -2.77 for Sn2+ + H2O + Cl"<=> SnOHCl0 + H+ at I = 3 can be calculated.

[1963FEI/SCH]: Review of solubility of metal oxides and hydroxides in water. The data ofSn(OH)4, SnO2(precip) and SnO2(cassiterite) are from unpublished experimentalresults [1957EGG]. [1963FEI/SCH] gives a log K*so of 5.0 for Sn4+ + 4H2O ^>SnO2 + 4H+. They state that to them no data concerning the hydrolysis of Sn(IV) areavailable. Experimental details such as pH, ionic strength, ... are not given.However, considering that Sn(OH)4° dominates Sn(IV) speciation between pH 2 to8, it was assumed in this report that in the experiments cited by [1963FEI/SCH],Sn(OH)4°, and not Sn4+, is the dominant dissolved species.

83

JNC TN8400 9 9 - 0 1 1

[1964LIN/TU]: [1964LIN/TU] recalculated the results of [1958TOB] with least-squaresanalysis and questioned the existence of SnOH+. In place of this species theyproposed the existence of Sn2(OH)3+. However, the interpretation of [1958TOB]seems more probable. [1976GOB] confirmed in her work the existence ofSn3(OH)42+ and SnOH+ and rejected the formation of the species Sn2(OH)3+

proposed by [1964LIN/TU].

[1965MES/IRA]: The solubility measurements of [1965MES/IRA] in 1 M NaC104 and inpresence of 0.5-10 mM Sn allows one to assign a lower limit for the reactions:

Sn2+ + H2O <=> SnO(cr) + 2H+ log K*So < -2.29SnOH+ o SnO(cr) + H+ log K*Si = 0.28

As the formation of Sn(OH)2° and polynuclear Sn3(OH)42+ is not taken into account,this are only estimates.

[1970BAR/KLI]: [1970BAR/KLI] investigated the hydrolysis of tin(TV) in 0.2 - 2.5 MNaOH. In contrast to [1997AMA/CHI] who proposed the formation of Sn(OH)5- andSn(OH)62" under alkaline conditions, [1970BAR/KLI] found only evidence for thepresence of Sn(OH)5". From the experimental details reported in the paper of[1970BAR/KLI] it seems that pH was not measured directly but determined from theamount of NaOH initially added to the system, while the Sn concentration wasmeasured in the solutions after one month shaking in contact with atmospheric CO2which has a strong influence on both solution composition and resulting pH. Thusthe calculated pH is probably significantly larger than a measured pH would be.

[1970BAR/KLI] also measured the solubility of cassiterite in water, dilute HNO3 anddilute NaOH. They observed a constant, pH-independent (between pH 2 - 11)solubility of 4 x 10'7 M Sn(IV)/L. Again, the pH seems to be calculated and not to bemeasured directly. No detection limit for Sn(IV) is indicated which makes it difficultto decide whether the measured concentrations below pH 11 correspond to thedetection limit or whether they are real concentrations.

[1970KUR/BAR]: [1979KUR/BAR] determined constants for the hydrolysis of Sn(OH)4° at100 °C. In contrast to observations at 25 °C, protonation of the Sn(OH)4° ion wasobserved by [1979KUR/BAR] already at a pH of 7.

[1971NAZ/ANT]: [1971NAZ/ANT] studied the hydrolysis of tin(IV) under acidic conditionsin 1 M KNO3. [1971NAZ/ANT] used a spectrophotometric method in which thecompetition between the hydrolysis reaction and complexation with salicylfluoronewas measured. The total Sn(IV) concentration in these experiments was 10~5 M.Considering the possible interaction of Sn(IV) with the nitrate ions in solution, theselog p values can be considered as useful estimates, but not as exact values for the

84

JNC TN8400 9 9 - 0 1 1

stability constants of Sn(IV) hydrolysis at I = 1. Nevertheless, based on theobservation of [1971NAZ/ANT] and [1997AMA/CHI] it is clear that protonation ofSn(OH)4° will occur only at pH < 1.

[1973GAB/SRI]: [1973GAB/SRI] determined the redox potential of the reactions Sn(cr) +3OH- o Sn(OH)3- + 2e" and Sn(OH)3- + 3OH" <=> Sn(OH)6

2" + 2e- as 0.90 V and0.910 V, respectively. From their data log (3 values for the hydrolysis of Sn2+ andSn(OH)4 can be estimated (Table 4.2 and Table 4.14). Assumptions: log K\y = -13.76, log K (Sn2VSn(cr)) = 5.0 and log K (Sn(OHySn(cr)) of -0.77 (at 1=0.25).These values can be considered only as estimates due to possible error introduced byestimating the redox potential at 1=0.25.

[1973KLI/BAR]: The values given in [1973KLI/BAR] for the hydrolysis of Sn(OH)4° arecalculated from the data measured by [1971NAZ/ANT] in 1 M KNO3. The solubilityof synthetic cassiterite is determined at different temperatures. [1973KLI/BAR]determined in the temperature range 473 - 673 K a higher solubility of cassiterite than[1981DAD/SOR] and [1988BAR/SHA] (cf. Table 4.5). Also at 298 K, the solubilitymeasured by [1973KLI/BAR] of 4 10"7 M is much higher than observed in otherreferences [1926GRU/LIN, 1963FEI/SCH, 1997AMA/CHI] (Table 4.5). The reasonof this apparent higher solubility observed by [1973BAR/KLI] is not clear. Thispaper was not chosen in the present report.

[1974GOB]: same results as given in [1976GOB].

[1976GOB]: [1976GOB] studied with potentiometric titration the hydrolysis of the tin(II) ionin 0.02 -2.3 mM in the pH range 2.7 - 3.7. [1976GOB] showed in her work that thespecies Sn3(OH)42+, SnOH+ and Sn2* are dominant under the conditions of herexperiments, similar to [1958TOB]. With her data she could not confirm (norexclude) the presence of Sn2(OH)22+ which was also proposed by [1958TOB].

[1977MAR]: [1977MAR] studied the hydrolysis of tin(II) in 3 M NaC104 in alkaline solution.Sn(II) = 1 mM; log Kw = 14.18

[1978DIC/LOT1]: [1978DIC/LOT1] determined the redox potential of the reactions Sn(0) +3OH- <=> Sn(OH)3- + 2e\ They extrapolated their data to I = 0 with the Pitzer modeland obtained a redox potential of 0.88 V. From their data value for the reaction Sn2+

+ 3H2O <=> Sn(OH)3" + 3H+ a log (3°i;3 of -16.85 can be calculated, assuming a logK°w = -14.0, log K° (Sn2+/Sn(cr)) = 4.60 (Section 4.4.1).

[1979VAS/GLA]: The values given in [1979VAS/GLA] for the hydrolysis of Sn(OH)4° arecalculated from the data measured by [1971NAZ/ANT] in 1 M KN03.

[1981DAD/SOR]: [1981DAD/SOR] determined the solubility of synthetic cassiterite in thetemperature range 473 - 673 K. [1981DAD/SOR] determined an approximately 5

85

JNC TN8400 9 9 - 0 1 1

times lower solubility of cassiterite than [1973KLI/BAR]. Extrapolation of themeasurements of [1981DAD/SOR] to 298 K gives a log K*So of « 7.5 forcassiterite, a value which is in fair agreement with the a log K*so of 8.06 determinedrecently by [1997AMA/CHI] for cassiterite (SnO2(cr)).

[1981PET/MIL]: [1981PET/MIL] studied with anodic stripping voltammetry the hydrolysisof tin(II) and the complex formation of Sn(II) with CY, Br and SO42" in 0.1 - 1 MNaNO3, NaCl and artificial seawater with a Sn concentration of 10~7 M. Thehydrolysis constants measured in NaCl and artificial seawater also include theformation of tin chloride complexes.

[1988BAR/SHA]: [1988BAR/SHA] determined at 300 °C a much smaller solubility ofsynthetic cassiterite than [1973KLI/BAR]. They offer no explanation for thesefindings.

[1997AMA/CHI]: Based on solubility measurements, [1997AMA/CHI] proposed theformation of Sn(OH)s" and Sn(OH)(52" under alkaline conditions. They extrapolatedtheir measurements from I = 0.1 M NaC104 to I = 0 using the wrong sign for theDavies equation. For further comments see [1998ODA/AMA].

[1998ODA/AMA]: See also comments to [1997AMA/CHI]. [1998ODA/AMA] gives theuncorrected log P15 and (3ig values for I = 0.1 M NaC104 derived from theexperimental data from [1997AMA/CHI] and corrected these values to I = 0 using theDavies equation. The additional experiments of [ODA/AMA] in 0.1 M NaClC^confirmed the solubility data determined by [1997AMA/CHI]. [1998ODA/AMA] alsoshowed that the presence of chloride and sulfate did not influence the solubility ofSnC>2(precip).

86

J N C T N 8 4 0 0 99 - O i l

5 Antimony

Antimony exists in the oxidation states -III, 0, +III and +V. It occurs in nature primarily as thesulfide Sb2S3(stibnite) or as the oxide Sb2O3(valentinite) [1976BAE/MES, 1985PAS,1995WIB]. When Sb2C<3(cr) is heated in air, additional oxygen is taken up above 300 °C, andSb2O4 (a mixed compound of Sb(III) and Sb(V)) can be formed [1984BER/BRE]. In water,Sb(III) is stable under reducing conditions and Sb(V) under oxidizing conditions[1976BAE/MES, 1985PAS, 1995WIB].

5.1 Hydrolysis of antimony (III)

Solubility measurements in presence of Sb2C>3(valentinite), e.g. [1952GAY/GAR] anddiffusion experiments [1965JAN/HAR1, 1977KEP/TAL] show that Sb(III) exists below pH 2as Sb(OH)2+, and in basic solutions (pH > 11) as Sb(OH)4~. For a wide pH range the solubilityof Sb2O3 does not depend on pH, indicating the predominance of the uncharged Sb(OH)3°species. [1965JAN/HAR1] have shown by spectrophotometric measurements that no polymericspecies are formed in significant amounts in 0.1 M Sb(III) solutions. [1965JAN/HAR1]showed that under basic conditions (in 1 to 16 N NaOH), Sb(III) exists exclusively asSb(OH)4- (or SbO2-), while no other anionic species could be observed. [1970DAW/WIL]observed with spectrophotometric methods at Sb(III) concentrations > 0.1 mM a polymerizationof antimony(III) species in sulfate solutions. [1974AHR/BOV] and [1994AKI/ZOT] concludedfrom solubility measurements that Sb2(OH)42+ and Sb2(OH)6 exist in concentrated (Sb > 0.1mM) Sb(III) solutions.

In the following paragraphs, all log P values refer to Sb(OH)3° as master species as the constantfor the protonation of Sb(OH)3° to Sb3+ is not known with sufficient accuracy.

Table 5.1: Experimental equilibrium data used for the hydrolysis of antimony(III), accordingto the equilibrium: mSb(OH)3° + nH2O <=> Sbm(OH)3m+n "

n + nH+. These datawere chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 5.12: 'Commentson selected references'. Method: extr = solvent extraction, sol = solubility, sp =spectrophotometry. The notation log |3b refers to a reaction involving OH~.

] ° g Pm

log A,

0.43

Reference;, 3m+n

0: Sb(OH)3° + 3H+ <=> Sb

[1977ANT/NEV]

Comments

3+ + 3H2O

T= 298.15 K, 1=1

I (M)

1

Medium

NaClO4

Method

sp

87

JNC TN8400 9 9 - 0 1 1


2H2Olog pjj: Sb(OH)3° + 2H+ s=> SbOH2+

0.23 [1974AHR/BOV] T= 298.15 K, 1=51.04 [1977ANT/NEV] T= 298.15 K, 1=1

51

HNO3

NaCIO,solsp_

log fiu: Sb(OH)30 + H+ <=> Sb(OH)2" + H2O

1.37 *1.16 11.19 l

1.18 l1.431.401.421.091.091.231.231.05

log fP1A:

2.07 2

2.04 21.99 21.93 2

1.93 2

2.21 22.21 2

2.202

2.142

2.192

2.06 2

2.04 2

2.05 2

1.99 21.98 21.96 21.95 22.06 22.09

[1924SCH] T= 298.15 K, 1=0.2[1924SCH] T= 298.15 K, 1=0.5[1924SCH] T= 298.15 K, 1=0.9[1924SCH] T= 298.15 K, 1=1.1[1968MIS/GUP] T= 298.15 K, 1=0.02[1968MIS/GUP] T= 298.15 K, 1=0.05[1968MS/GUP] T= 298.15 K, 1=0.1[1974AHR/BOV] T= 298.15 K, 1=5[1974AHR/BOV] T= 298.15 K, 1=5[1974SHO/MAB] T= 298.15 K, 1=3[1974SHO/MAB] T= 298.15 K, 1=0.1[1977ANT/NEV] T= 298.15 K, 1=1

Sb(OH)3° + OH~ & Sb(OH)i

[1948TOU/MOU] T= 308 K, I=diluted[1948TOU/MOU] T= 308 K, I=diluted[1948TOU/MOU] T= 308 K, I=diluted[1948TOU/MOU] T= 308 K, I=diluted -[1948TOU/MOU] T= 308 K, I=diluted[1952GAY/GAR] T= 298.15 K, 1=0.1[1952GAY/GAR] T= 298.15 K, 1=0.075[1952GAY/GAR] T= 298.15 K, 1=0.04[1952GAY/GAR] T= 298.15 K, 1=0.04[1952GAY/GAR] T= 298.15 K, 1=0.02[1973VAS/SHO2] T= 298.15 K, 1=2.37[1973VAS/SHO2] T= 298.15 K, 1=2.03[1973VAS/SHO2] T= 298.15 K, 1=1.69[1973VAS/SHO2] T= 298.15 K, 1=1.36[1973VAS/SHO2] T= 298.15 K, 1=1.02[1973VAS/SHO2] T= 298.15 K, 1=0.68[1973VAS/SHO2] T= 298.15 K, 1=0.34[1973VAS/SHO2] T= 298.15 K, 1=2[1994AKI/ZOT] T=298, I=dil

0.20.50.91.1

0.020.050.1553

0.11

1.990.70.46

0.0920.042

0.10.0750.040.040.022.372.031.691.361.020.680.34

20

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HNO3HC1O4

HC1O4

HC1O4

NaClO4

NaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOH

NaClO4

self medium

sol

solsolsolspspspsolsolextrextrsp

solsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsol

JNC TN8400 9 9 - 0 1 1


log j324: 2Sb(OH)3° + 2H+ <=> Sb2(OH)42+ + 2H2O

2.86 [1974AHR/BOV] T= 298.15 K, 1=5 5 HC1O4 sol

log P2i6: 2Sb(OH)3o & Sb2(OH)6°

0.08 [1994AKI/ZOT] T=298, I=dil 0 self medium sol"T calculated with log K°*S3 = 8.72 (section 5.2.1)

2 recalculated in this report from experimental values of [1948TOU/MOU], [1952GAY/GAR] and[1973VAS/SHO].

5.1.1 Sb3+

[1965JAN/HAR1], [1965JAN/HAR2] concluded from diffusion measurements that onlymonomeric Sb(OH)2+, and possibly Sb3+, are present in perchloric acid solutions up to H+

concentration of 6 M and [Sb] < 1 mM. Spectrophotometric measurements of [1968MIS/GUP](Sb concentration = 0.3 mM) indicate the presence of Sb(OH)2+ in the pH range 1-2, and at pH< 0 (1 and 3 M HC1O4) the presence of a more protonated Sb complex (Sb3+ or SbOH2*).

From the data given by [1977ANT/NEV] (Table 5.1) a tentative log P \ o value can beextrapolated for the reaction Sb(OH)3° + 3H+ 0 Sb3+ + 3H2O assuming a Ae of 0.07 (frome(Am3+, CIO4-) = 0.49 and E(H+, CIO4-) = 0.14; [1995SIL/BID]):

Sb(OH)3° +3H+ o Sb3+ + 3H2O log (3°u = -0.73

The values calculated from the potentiometric measurements of [1970BON/WAU] and[1970BON] in 5 and 2 M HCIO4 (Table 5.2) agree well with the log p l i 0 values measured by[1977ANT/NEV]. It is not clear, however, if in the experiments of Bond and co-workers[1970BONAVAU, 1970BON] Sb3+ is the only species present. The potentiometricmeasurements of [1975HEI/SCH] in 1.5 M H F and H2SO4 give a log p of 0.36 for theprotonation of Sb(OH)3° to Sb3+. [1975HEI/SCH] assumed the formation of Sb(OH)2

+ (at apH < 0). We consider it as probable, considering the high H+ concentration, that the Sb3+

species is dominant in such solutions (cf. Table 5.2).

89

JNC TN8400 9 9 - 0 1 1

5.1.2 SbOH2+

For the reaction Sb(OH)3° + 2H+ <=> SbOH2+ + 2H2O, [1974AHR/BOV] determined withsolubility measurements a log P u of 0.23 in 5 M HNO3. [1977ANT/NEV] determined withsolubility and spectrophotometric measurements a log P u of 1.04 in 1 M HCIO4 (Table 5.1)From these measurements (Figure 5.1) a tentative value can be calculated:

Sb(OH)3°+2H+ SbOH2+ + 2H2O logP°u =0.83

The Ae = 0.19 obtained from only two datapoints (Figure 5.1) agrees satisfactorily with the Aeof 0.11 estimated from E(AmOH2+, CIO4-) = 0.39 and e(H+, CIO4-) = 0.14; [1995SIL/BID].

Sb(OH)3°+2H+ <4 SbOH2++2H2O

2.5

2

L m&lal

Figure 5.1: Plot of log 01,1 - 2 D vs. Im for the reaction Sb(OH)3° + 2H+ <=>2H2O at 25 °C. The straight line shows the result of the 'linear regression': Ae0.19; log P°i,i = 0.83. Calculated from data compiled in Table 5.1.

5.1.3 Sb(OH)2+

Many measurements of the hydrolysis of Sb(OH)3° in dilute acid exist. The formation constantgiven in [1952LAT] and [1957PIT/POU] for SbO+ (corresponding to Sb(OH)2

+) and SbO2~originate from the early work of [1883SCH]. The log p l i 2 value of 0.90 given by [1985PAS]for the reaction Sb(OH)3° + H+ <=> Sb(OH)2

+ + 2H2O is taken from the data measured by[1924SCH] in 0.2 - 1.1 M HCIO4.

90

JNC TN8400 9 9 - 0 1 1

[1924SCH] determined in 0.2 - 1.1 M HC1O4 both the solubility of Sb(III) in perchlorate acidand the electrode potential of the reaction Sb(cr) + 2H2O <=> Sb(OH)2

+ + 2H+ + 3e~. Bothmeasurements resulted in similar log (3ii2 values (Table 5.1 and 5.2). [1952GAY/GAR]measured in 0 - 0.1 M HC1 a mean log p l i 2 of 1.13 for the reaction Sb(OH)3° + H+ <=>Sb(OH)2+ + H2O (Table 5.2). As Sb(III) forms complexes with chloride, these measurementswere not used for extrapolation in this report. However, they agree well with othermeasurements reported in the literature. [1968MIS/GUP] determined spectrophotometrically amean log (3i)2 value of 1.42 in 0.02-0.1 M HCIO4. Based on potentiometric measurements,[1975HEI/SCH] calculated a log p1>2 of 0.53 in presence of I = 1.5 M (HF and H2SO4).Although they corrected their measurements for the interactions of Sb(III) with F~ and SO42-,their measurements differ from the other data reported. Considering the high H+ concentration,the presence of Sb3+ or SbOH24" in the experiments is probable.

Extrapolation of the data of [1924SCH (only solubility data), 1968MIS/GUP, 1974AHR/BOV,1974SHO/MAB, and 1977ANT/NEV (compiled in Table 5.1) gives at I = 0, as shown inFigure 5.2:

Sb(OH)3° Sb(OH)2+ + H2O log(3°li2 =1.30

The extrapolation of the experimental results with the SIT term shown in Figure 5.2 gives a Aevalue of 0.02 which is in good agreement with the expected Ae value of =0 for the isocoulombicreaction Sb(OH)3° + H+ <=> Sb(OH)2

+ + H2O.

Sb(OH)3°+H+o Sb(OH)2++H2O

Figure 5.2: Plot of log ( 3 U + 0 D vs. Im for the reaction Sb(OH)3° + H+ <=> Sb(OH)2+ +H2O at 25 °C. The straight line shows the result of the linear regression: Ae =0.04; log ( 3 \ 2 = 1.30. Calculated from data compiled in Table 5.1.

91

JNC TN8400 9 9 - 0 1 1

The potentiometric measurements of [1924SCH] and [1972VAS/SHO] (in 0.3-2.5 M HC1O4)refer all to the reaction Sb(cr) + 2H2O <=> Sb(OH)2+ + 2H+ + 3e~. For comparison the log ofthese values were corrected for the reaction Sb(cr) + 3H2O <=> Sb(OH)3 + 3H+ + 3e~ with afactor of 11.99 (cf. Section 5.7: Redox reactions). These values agree very well with the log(312 values measured with other techniques. If the these values were also included in theextrapolation to 1=0 a log (3° 1,2 value of 1.29 results, the nearly the same value as in absence ofthe potentiometric measurements.

5.1.4 Sb(OH)4-

Several authors determined formation constants for the reaction Sb(OH)3° + H2O <=> Sb(OH)4~+ H+. Extrapolation to I = 0 of the experimental data given by [1948TOU/MOU,1952GAY/GAR, 1973VAS/SHO2, 1994AKI/ZOT] compiled in Table 5.1 results in a log p b \ 4

of 2.07 for the reaction Sb(OH)3° + OH~ <=> Sb(OH)4- (Figure 5.3). From this log ( 3 \ 4 canbe obtained:

Sb(OH)3° + H2O <> Sb(OH)4- \A =-11.93

The extrapolation of experimental results with the SIT term shown in Figure 5.3 gives a Aevalue of 0.05 which is in good agreement with the expected Ae value of 0 for the isocoulombicreaction Sb(OH)3° + OH- <=> Sb(OH)4-

Sb(OH)3° + OH" « Sb(OH)4"

Figure 5.3: Plot of log pb1?4 + 0D vs. Im for the reaction Sb(OH)3° + OH~<=» Sb(OH)4- at 25

°C (and 35 °C). The straight line shows the result of the linear regression: Ae =0.05; log pb°i,4 = 2.07. For the reaction Sb(OH)3° + H2O <=> Sb(OH)4- + H+ alog P°ij4 of- 11.93 can be calculated. Calculated from data compiled in Table 5.1.

92

JNC TN8400 9 9 - 0 1 1

5.1.5 Sb2(OH)42+

For the reaction 2Sb(OH)3° + 2H+ <=> Sb2(OH)42+ + 2H2O, a tentative log p2,4 of 2.43 can be

extrapolated from the data given by [1974AHR/BOV], As this log (3 value is determined at highionic strength and only by one author, the usage of this value at low ionic strength may result ina considerable error. The species Sb2(OH)42+ will be important only in concentrated Sb(III)solutions (Sb>0.1 mM).

5.1.6 Sb2(OH)6°

For the reaction 2Sb(OH)3° <=> Sb2(OH)6 a log (32,6 of 0.08 is determined by [1994AKI/ZOT]from measurements at 25 - 300 °C. While this complex seems to be important at highertemperature, its concentration is small at 25 °C.

5.1.7 Additional equilibrium data compiled for the hydrolysis of antimony(III)

Table 5.2: Additional experimentally determined equilibrium data compiled for the hydrolysis ofantimony(IH), These data were not chosen in the present report for the evaluation of recommendedstability values. Method: pol = polarography, pot = potentiometry, sol = solubility.

log Pm, 3

log Pi,0:

0.89 >0.32 '0.53 2 ' '

log p12:

1.38 '1.18 '1.19 '1.19 '1.131.02 3

1.09 3

1.16 3

1.15 3

1.16 3

1.15 3

1.64 4 ' '1.61 ]

1.55 l

1.30 '1.34 '1.32 '0.53 5 ' !

log PK4:

-12.05 6

-11.82-11.83 6

-11.73 4

,ra+n Reference

Sb(OH)3° + 3W <=> Sb3* +

[1970BON/WAU][1970BON][1975HEI/SCH]

Sb(OH)3° + H+ <=> Sb(OH)2

[1924SCH][1924SCH][1924SCH][1924SCH][1952GAY/GAR][1952GAY/GAR][1952GAY/GAR][1952G AY/GAR][1952G AY/GAR][1952GAY/GAR][1952GAY/GAR][1972V AS/SHO][1972VAS/SHO][1972V AS/SHO][1972 V AS/SHO][1972VAS/SHO][1972V AS/SHO][1975HEI/SCH]

Comments

3H2O

T= 303 K, 1=5T= 303 K, 1=2T= 298.15 K, 1=1.5

,+ + H2O

T= 298.15 K, 1=0.2-1.1T= 298.15 K, 1=0.24.1T= 298.15 K, 1=0.2-1.1T= 298.15 K, 1=0.2-1.1T= 298.15 K, 1=0.01-0.1T= 298.15 K, 1=0.01T= 298.15 K, 1=0.02T= 298.15 K, 1=0.03T= 298.15 K, 1=0.05T= 298.15 K, 1=0.075T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=1.5

Sb(OH)3° + H2O <=> Sb(OH)i + H*

[1948TOU/MOU][1952GAY/GAR][1952GAY/GAR]

•! [ 1973 V AS/SHO 1]

T= 298.15 K, 1=0.04-2T= 298.15 K, 1=0.01-0.1T= 298.15 K, 1=0.01-0.1T= 298.15 K, 1=0.3-2.5

KM)

52

1.5

0.20.50.91.1var

0.010.020.030.050.080.10

0.360.711.782.142.491.5

0var00

Medium

HC1O4

Na/HClO4

HF, H,SO4

HC1O4

HC1O4

HC1O4

HC1O4

HC1, NaOHHC1HC1HC1HC1HC1HC1

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HF, H,SO4

NaOHHC1, NaOH

NaOHNaOH

Method

polpolpot

potpotpotpotsolsolsolsolsolsolsolpotpotpotpotpotpotpot

solsolsolpot

93

JNC TN8400 9 9 - 0 1 1


-12.10 6

-11.99 4-7

-11.91

[1973VAS/SHO2][1973VAS/SHO2][1994AKI/ZOT]

T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T=298, I=dil

000

NaOHNaOH

self medium

solsolsol

1 estimated from potentiometric data assuming a log K (Sb(OH)3°/Sb(cr)) = 11.99 (see Section 5.7: Redox reactions)2 we assumed, considering the high H+ concentration, that the Sb3+ species is dominant in such a solution3 calculated with log KO>

S3 = 8.48 [1952GAY/GAR] or = 8.30[1974AHR/BOV]4 linear extrapolation to 1=0 by [1972VAS/SHO]5 [1975HEI/SCH] assumed that the Sb(OH)2

+ species is dominant6 recalculated in this report from experimental values and extrapolated to 1=0 with SIT.7 calculated with log K°*S3 = 8.72 (section 5.2.1).

Table 5.3: Thermodynamic data for the antimony(III) hydroxide system taken from previous compilations.As pointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

logpm, 3m+n Reference Comments I (M) Medium

log /?,,„.• Sb(OH)3° + 3H+ <=>Sb3+ + 3H2O

0.35 ' H952LAT1 T= 298.15 K, I=n/a

log p,,: Sb(OH)3° + 2H+ <=> SbOHu + 2H2O

0.19 ri976SMI/MAR] T= 298.15 K, 1=5

log pu:

1.23 '0.871.201.411.47 >1.411.47 '1.170.901.181.181.42

log / V-10.63 >-10.99-11.80-11.82-11.50 '-11.82-11.50 l

-11.79-11.90-11.79-11.82

Sb(OH)3° + H* <=> Sb(OH)2+

[1952LAT][1957PIT/POU][1976SMI/MAR][1976BAE/MES][1980BEN/TEA][1981BAE/MES][1982WAG/EVA][1985BAB/MAT][1985PAS][1986ITA/NIS][1989SMI/MAR][1992PEA/BER]

Sb(OH)3° + H2O <=> Sb(OH)4

[1952LAT][1957PIT/POU][1976SMI/MAR][1976BAE/MES][1980BEN/TEA][1981BAE/MES][1982WAG/EVA][1985BAB/MAT][1985PAS][1986ITA/NIS][1992PEA/BER]

+ H2O

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0

- + H*

T= 298.15 K,I=n/aT= 298.15 K, 1=3-9T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T=298.15K, 1=0

00

0

00030

var00

0

0000

KOH

log (324: 2 Sb(OH)3° + 2H+ <=> Sb2(OH)42* + 2H2O

2.88 [1976SMI/MAR] T= 298.15 K, 1=5calculated in this report with a A,G0 of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1)

94

JNC T N 8 4 0 0 99 - O i l

5.2 Solid antimony(IH)-oxide/hydroxide

Antimony sesquioxide, Sb2C>3(cr) (often also the notation Sb4C>6(cr) is used) exists in theorthorhombic form (a-Sb2O3, valentinite) from room temperature to 846 K. Between 846 Kand the melting temperature of 929 K, Sb2C>3(cr) exists in the cubic modification ((3-Sb2C>3,senarmontite). While at room temperature valentinite (a-Sb2O3(cr)) precipitates readily fromsolutions ([1924SCH], [1939BLO], [1984BER/BRE]), the transformation of the cubicsenarmontite to the more stable orthorhombic valentinite does not take place. Where nodescription of the Sb2C>3(cr) is given, it can normally be assumed that the data given refer to a -Sb2O3(cr) (Table 5.6). Many measurements given in the literature are based on emf(electromotive force) measurements at high temperature extrapolated to 298 K, which mayreduce the reliability of the data (these data are given in Table 5.5). When Sb2C>3(cr) is heated inthe presence of oxygen above 300 °C, Sb2O4(cr) can be formed [1984BER/BRE],[1987PAN/SRE].

Table 5.4: Experimental equilibrium data used for the determination of precipitation ofantimony(UI) oxide/hydroxide. These data were chosen for the evaluation ofrecommended values in the present report. Additional information for the differentreferences see Section 5.12: 'Comments on selected references'. Method: pol =polarography, pot = potentiometry, sol = solubility, sp = spectrophotometry.

log K*S3

log K*S3:

8.57 i8.72 18.26 i8.709.298.488.72 i8.69 '8.77 i8.71 i8.69 >8.308.56

logK*S3:

8.558.968.96

Reference Comments

2 Sb(OH)3° <=> a-Sb2O3 (valentinite)+ 3H2O

[1924SCH][1924SCH][1924SCH][1948TOU/MOU][1948TOU/MOU][1952GAY/GAR][1972VAS/SHO][1972VAS/SHO][1972VAS/SHO][1972VAS/SHO][1972VAS/SHO][1974AHR/BOV][1990SHI/ZOT]

T= 298.15 K, 1=0.23T= 298.15 K, 1=0.49T= 298.15 K, 1=1.13T= 308 K, I=dilT= 308 K, I=dilT= 298.15 K, I=dilT= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=0.3-2.5T= 298.15 K, 1=5T=298, I=dil

2 Sb(OH)3° <=> P-Sb2O3(senarmontite) + 3H2O

[1974AHR/BOV][1990SHI/ZOT][1994AKI/ZOT]

T= 298.15 K, 1=5T=298, I=dilT=298, I=dil

KM,

0.230.491.13

000

0.360.711.782.142.49

50

500

Medium

HC1O4

HC1O4

HC1O4

self mediumNaOH

self mediumHC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

self medium

HC1O4


Method

potpotpotsolsolsolpotpotpotpotpotsolsol

solsolsol

calculated from potentiometric data with a log K (Sb(OH)3°/Sb(cr)) = 11.99 (see Section 5.7: Redox reactions)

95

JNC TN8400 9 9 - 0 1 1

5.2.1 a-Sb2O3 (valentinite)

The solubility of a-Sb2C>3(valentinite) in water was determined by [1939BLO,1948TOU/MOU and 1952GAY/GAR]. They determined log K*S3 values of 8.08 to 8.70 forthe reaction 2Sb(OH)3° <=> Sb2O3(valentinite) + 3H2O. The solutions used by [1939BLO]contained 0.01 M HCl. Since they did not measure pH and did not account for interaction withC1-, the solubility of Sb2O3(valentinite) is probably overestimated. A careful examination of thedata of [1948TOU/MOU] showed that the solubility a-Sb2O3 decreased slightly (minimum logK*s3 = 9.29) after the addition of small amounts of NaOH, indicating that in the measurementsin water, Sb2O3 was probably in equilibrium with Sb(OH)2+. The potentiometricmeasurements of [1924SCH and 1972VAS/SHO], as well as the measurements of[1973BEH/ROS, 1974AHR/BOV, and 1990SHI/ZOT] all resulted in log K*S3 values in therange of 8.3 - 8.8. Extrapolation of the data given in Table 5.4 results in:

2Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O log K*°S3 = 8.72

The extrapolation of the experimental results with the SIT term shown in Figure 5.4 gives a AEvalue of 0.02 which is in good agreement with the expected Ae value of 0 for the unchargedspecies involved in the reaction 2Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O.

10.5 -•

10 -

g 9.5+ 9oco

2Sb(OH)3° <=> Sb2O3+3H2O

valentinite

8.5° O O

j? 87.5

7 •

6.5

6

y = -0.02x + 8.72

4 6

i molal

Figure 5.4: Plot of log K*S3 + 0D vs. Im for the reaction 2Sb(OH)3° <=> a-Sb2O3 + 3H2O at25 °C (and 35 °C). The straight line shows the result of the linear regression: Ae =0.02; log K*°s3 = 8.72. Calculated from data given in Table 5.4.

96

JNC TN8400 9 9 - 0 1 1

5.2.2 fi-Sb2C>3 (senarmontite)

For the (theoretical) formation of senarmontite a log K*°s3 value of 8.96 is given by[1990SHI/ZOT] based on solubility experiments for the isocoulombic reaction 2Sb(OH)3° <=>(3-Sb2C>3(senarmontite) + 3H2O. This value agrees reasonably well with the value of 8.55determined by [1974AHR/BOV]. As the value determined by [1990SHI/ZOT] for valentiniteagrees very well with our calculation it seems reasonable to use for senarmontite the log K*°s3value of 8.96 given by [ 1990SHI/ZOT].

Table 5.5: Additional experimentally determined equilibrium data compiled for the precipitation ofantimony(III) oxide/hydroxide. These data were not chosen in the present report for the evaluationof recommended stability values. Method: pot = potentiometry, sol = solubility.

log K*S3 Reference Comments I (M) Medium Method

log K*S3: 2 Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O

8.08 '8.50 2

8.78 3'4

9.08 5>4

9.30 3>4

6.74 6-4

[1939BLO][1965FRI/VER][1972VAS/SHO][1973BEH/ROS][1973VAS/SHO1]fl986AZA/PAN]

T= 298.15 K,I=dilutedT= 298.15 K,I=n/aT= 298.15 K, 1=0.3-2.5T= 298.15 K, I=n/aT= 298.15 K, 1=0.3-2.5T= 298.15 K,I=n/a

0.010000

HC1n/a

HC1O4

n/aNaOH

n/a

soln/apotn/apotemf

log K*S3: 2 Sb(OH)3° <=> fi-Sb2O3(senarmontite) + 3H2O

8JJ3 f!939BLO] T=298.15 K, I=diluted 0.01 HC1 sol1 measured in presence of 0.01 M HC1, equilibrium with Sb(OH)2

+ probable2 as cited in [1968MIS/GUP]3 linear extrapolation to 1=0 by [1972VAS/SHO].4 calculated in this report with a Afi° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1)5 extrapolated from AH and AS values determined at 600-700 K6 extrapolated from = 800 K

Table 5.6: Thermodynamic data for the precipitation of antimony(III) oxide/hydroxide taken from previouscompilations. As pointed out in Section 2 of this report only experimental data were used for thepresent evaluation. The following table serves only for comparison.

logK* S3 Reference Comments

log K*S3: 2 Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O

6.386.368.548.309.10 '

[1957PIT/POU][1957PIT/POU][1976SMI/MAR][1976SMI/MAR][1980B EN/TEA]

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=5T= 298.15 K, I=n/a

05

97

JNC TN8400 99 - Oil


9.108.148.527.66

I

2

[1982WAG/EVA][1985PAS][1986ITA/NIS][1994AKI/ZOT]

T=298T=298T=298T=298,

.15 K,

.15 K,

.15 K,I=dil

I=n/aI=n/a1=0

log K*S3: 2 Sb(OH)3° <=> Sb2O3(c) + 3H2O)

8.54 •8.62 l

8.62 •10.23 '8.4810.78 '9.07 '10.01 '8.489.10 '7.648.48

[1952LAT][1954COU][1963WIC/BLO][1971NAU/RYZ][1976BAE/MES][1977BAR/KNA][1978ROB/HEM2][1979KUB/ALC][1981BAE/MES][1982PAN][1985BAB/MAT][1992PEA/BER]

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/a

T*55.- 2 Sb(OH)3° 3 <=> P-Sb2O3(senarmontite) + 3H2O

9.28 '7.847.839.28 '11.638.5410.4210.429.4612.77

i

l

i

3

[1954COU][1957PIT/POU][1957PIT/POU][1963WIC/BLO][1971NAU/RYZ][1976SMI/MAR][1980BEN/TEA][1982WAG/EVA][1985PAS][1986ITA/NIS1

T=T=X-T=

T=T=T=T=T=

298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK,I=5K, I=n/aK, I=n/aK, I=n/aK,I=0

log K*S3: Sb(OH)3° <=> Sb(OH)3 (s)

5.13 l [1971NAU/RYZ] T= 298.15 K, I=n/a7.40 ' [1980BEN/TEA] T= 298.15 K, I=n/a7.39 ' ri982WAG/EVA1 T= 298.15 K, I=n/a1 calculated in this report with a A[G° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1)2 probably copied wrongly by [ 1994AKI/ZOT] from literature3 copied wrongly by [198ITA/NIS] from [1982WAG/EVA]

98

JNC TN8400 9 9 - 0 1 1

5.3 Antimony(HI) chloride system

Sb(III) forms complexes and compounds with chloride. Also the existence of mixed Sb(III)chloride hydroxide complexes and compounds is reported. Experimentally determined data aregiven in Table 5.7 and 5.8. Thermodynamic data from earlier compilations are given in Table5.9.

Table 5.7: Experimentally determined equilibrium data compiled for the Sb(III) chloridesystem. These data were chosen for the evaluation of recommended values in thepresent report. Additional information for the different references see Section5.12. Method: pol = polarography, pot = potentiometry.

log Pm.n

log Pi, v

2.262.30

log 07,2:

3.494.003.39

Reference

Sb3+ + Or <=> SbCl2+

[1959PAN/DES][1970BON/WAU]

Sb3+ + 2CI- t=> SbCl2

[1959PAN/DES][1970BON/WAU][1975BIE/ZIE]

Comments

T= 298.15T= 303.15

+

T= 298.15T= 303.15T= 298.15

K,K,

K,K,K,

1=41=5

1=41=51=4

I (M) Medium

45

454

NaClO4, HC1O4

HC1O4

NaClO4, HC1O4

HC1O4

HC1O4

Method

polpol

polpolpot

log pU: Sb3+ + 3Ct <=> SbCl3°

4.18 [1959PAN/DES] T= 298.15 K, 1=4[1970BON/WAU] T= 303.15 K, 1=55.78

4.09 [1975BIE/ZIE] T= 298.15 K, 1=4

4 NaClO4, HC1O4 polHC1O4

HC1O.polpot

log PJ,4: Sb3+ + 4CI- <^ SbCl4-

4.72 [1959PAN/DES] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol

6.78

log Pi,s--

4.72

log PJ,6:

4.11

[1970BON/WAU] T=

Sb3+ + 5CI- <=> SbCl52-

[1959PAN/DES] T=

Sb3+ + 6Ct <=> SbCl63~

[1959PAN/DES] T=

303.

298.

298.

15

15

15

K,

K,

K,

1=5

1=4

1=4

5

4

4

HC1O4

NaClO4, HC1O4

NaC104, HC1O4

pol

pol

pol

99

JNC TN8400 9 9 - 0 1 1

5.3.1 Reactions ofSb3+: SbCl2+, SbCl2+, SbCl3

0, SbCl4-, SbCl52~ and SbCl6

3~

Antimony forms complexes with chloride in acid solutions. [1959PAN/DES, 1970BON/WAU,1975BIE/ZIE] determined in concentrated HCIO4 solutions the formation of Sb(III) chloridecomplexes from the reaction of Sb3+ with Ch (Table 5.7). From the experimental evidence it isclear that Sb(III) chloride complexes will be important only in very acid solution, under otherconditions, the hydrolyzed of complexes of Sb3+, i.e., SbOH2+ , Sb(OH)2

+ and Sb(OH)3°,will be dominant. [1974SHO/MAB] was able to show that the formation of Sb(III) chloridecomplexes will become important only at pH values < 2.

Antimony forms complexes with chloride. [1959PAN/DES, 1970BONAVAU, 1975BIE/ZIE]determined in concentrated HCIO4 solutions the formation of Sb(III) chloride complexes fromthe reaction of Sb3+ with Ch (Table 5.7). From these data equilibrium constants can becalculated (Figures 5.5 and 5.6). After conversion from Sb3+ to Sb(OH)3° (cf. Section 5.1.1)the following tentative values can be calculated:

Sb(OH)3° +3H+ + Cl-Sb(OH)3° +3H+ + 2C1-

SbCl2+ + 3H2OSbCl2

+ + 3H2Olog P ° u = 2.78log P°1>2 = 3.27

4.5 •

4 -

Q 3.5 -CD

+ 3 4-

D)2

1.5

1

0 . 5 ••

0

Sb3++CI"oSbCI2+

= 0.05x + 3.51

0 2 4 6 8lmi molal

Figure 5.5: Plot of log (3U + 6 D vs. Im for the reaction : Sb3+ + Cl~ <=» SbCl2+ at 25 °C.The straight line shows the result of the 'linear regression': Ae = -0.05; log $°\,\~ 3.51. Calculated from data given in Table 5.7.

100

JNC TN8400 9 9 - 0 1 1

Q

Figure 5.6: Plot of log (3li2 + 10 D vs. Im for the reaction : Sb3+ + 2C1- <=> SbCl2+ at 25 °C.

The straight line shows the result of the linear regression: Ae = -0.38; log P0]^ =4.00. Calculated from data given in Table 5.7.

No data are recommended in this report for the formation of SbCl30, SbCLr, SbCls2 , andSbCl6

3~ as the data reported for SbCl30 and for SbCl4- show quite a spread and for SbCl52-

and SbCl63- not enough data are available to extrapolate the constants to 1=0. The Ae values of

-0.05 and -0.38 obtained in Figures 5.5 and 5.6 can be compared to the As of -0.22 and -0.56estimated from the NEA americium data (£(AmCl2+, C1O4-) = 0.39, e(Am3+, CIO4-) = 0.49,and e(H+, C1-) = 0.12, e(Am(0H)2

+, CIO4-) = 0.17; [1995SIL/BID]).

5.3.2 SbCl4- and SbOHCl3~

A number of scientists, [1953HAI, 1967VAS/YUS, 1968NOR/KAZ, 1984VAS/SHO],measured the redox potential of the reaction Sb(cr) + 4Ch <=> SbCl4- +3e~. Assuming a log (3for Sb(OH)3°/Sb(cr) of 11.82 (see Section 5.7: Redox reactions), from these data a constant canbe estimated for the reaction Sb(OH)3 + 4C1" + 3H+ <=> SbCl4- + 3H2O (Table 5.8). Theseconstants are not very reliable, as none of the authors could show, that SbCl4~ was in fact thedominant species in their experiments or whether other antimony(III) chloride complexes werealso present (see Section 5.3.1: Reactions of Sb3+: SbCl2+, SbCl2

+, SbCl3°, SbCl4~, SbCl52-

and SbCl63-). Additionally, the solutions of [1953HAI] and [1967VAS/YUS] had a pH around

0 (H+ concentrations < 1.2 M). It is doubtful that really Sb3+ will dominate the antimony(III)speciation at this pH [1977ANT/NEV].

101

JNC TN8400 9 9 - 0 1 1

5.3.3 SbCl3(s) and SbOCl(s) (or Sb4O5Cl2(s))

Direct measurements of antimony(III) chloride and oxychloride solubility are not available.Based on the data compiled in Table 5.9, it can be concluded that both SbCl3(s) and SbOCl(s)are easily soluble. [1975HEN/LON, 1997KOL/HEN] showed that recrystallization of SbCl3(s)with water produces S b ^ s C ^ s ) and not SbOCl(s).

5.3.4 Additional equilibrium data compiled for the antimony(III) chloride system

Table 5.8: Experimentally determined equilibrium data compiled for the antimony(III) chloride hydroxidesystem. These data were not chosen in the present report for the evaluation of recommendedstability values (cf. Section 5.3.2). Method: pol = polarography, pot = potentiometry.

log P Reference Comments Medium Method

,: Sb(OH)3° + 4CI- + 3H+ » SbCl4- + 3H2O

2.914.184.383.372.663.904.715.696.076.326.977.53

,2

log Pi.i,3

1.84 '

[1953HATJ[1953HAI][1967VAS/YUS][1968NOR/KAZ][1984VAS/SHO][1984VAS/SHO][1984VAS/SHO][1984VAS/SHO][1984VAS/SHO][1984VAS/SHO][1984VAS/SHO]f 1984V AS/SHO]

: Sb(OH)3° + 3Ct + 2H+

ri967VAS/YUSl

T= 303 K, 1=4T= 303 K, 1=1-6T= 298.15 K, 1=0.8-1.2T= 303.15 K, 1=5T= 298.15 K, 1=4-7.5T= 298.15 K, 1=4.0T= 298.15 K, 1=4.7T= 298.15 K, 1=5.5T= 298.15 K, 1=5.9T= 298.15 K, 1=6.3T= 298.15 K, 1=6.9T= 298.15 K, 1=7.4

i=> SbOHClf + 2H2O

T= 298.15 K, 1=0.6

4varvar504

4.75.55.96.36.97.4

0.6

HCIHCIHCI

HC1O4

HCIHCIHCIHCIHCIHCIHCIHCI

HCI

polpolpolpotpotpotpotpotpotpotpotpot

pol1 estimated from potentiometric data assuming a log K (Sb(OH)3°/Sb(cr)) = 11.99 (see Section 5.7: Redox reactions)2 extrapolated to 1=0 by [1984VAS/SHO] with extended Debye-Huckel term.

Table 5.9: Thermodynamic data for the formation of antimony(III) chloride hydroxide system taken fromprevious compilations. As pointed out in Section 2 of this report only experimental data wereused for the present evaluation. The following table serves only for comparison. Medium: Wheredata refer to specific electrolyte solutions, this is indicated.


log Pu:Sb3+ + Ct <=> SbCl2*

2.30 [1976SMI/MAR]2.30 [1989SMI/MAR1

T= 298.15 K, 1=4T= 303.15 K, 1=5

HC1O,

102

JNC TN8400 99-011


log pli2: Sb3+ + 2CI- <=> SbCl2"

3.50 [1976SMI/MAR]4.00 R989SMI/MAR1

T= 298.15 K, 1=4T= 303.15 K, 1=5

45

HC1O.

log pu: Sb3+ + 3Cl- <=> SbCl3°

4.20 [1976SMI/MAR]5.80 P989SMI/MAR1

T= 298.15 K, 1=4T= 303.15 K, 1=5

HC1O,

log p1A: Sb3+ + 4CI- <=> SbCl4-

4.70 [1976SMI/MAR]6.80 [1989SMI/MAR1

T= 298.15 K, 1=4T= 303.15 K, 1=5

45

HC1O4

log P,i5: Sb3+ + 5Cl~ <=> Sb

4.70 [1976SMI/MAR] T= 298.15 K, 1=4 HC1O.

log pL6: Sb3+ + 6CI- <=> SbCl63~

4.10 [1976SMI/MAR1 T= 298.15 K, 1=4 HC1O.

log K*S3: Sb(OH)3° + 3Cl~ + 3H+ <=> SbCl3(s) + 3H2O

4.3 >4.5 '4.2 '4.1 l

4.2 '4.2 ]

3.9

[1952LAT][1963WIC/BLO][1971NAU/RYZ][1979KUB/ALC][1980BEN/TEA][1982WAG/EVA][1985PAS]

T=T=T=T=T=T=T=

298.15298.15298.15298.15298.15298.15298.15

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/a

log K*S3: Sb(OH)3° + Cl- + H+ <=> SbOCl(s) + 2H2O

1.8 ' [1977BAR/KNA]1.8 ' ri979KUB/ALCl

T= 298.15 K,I=n/aT= 298.15 K, I=n/a

calculated in this report with a AfG° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1)

103

JNC TN8400 9 9 - 0 1 1

5.4 Antimony(IH) fluoride system

Antimony(III) forms complexes and compounds with fluoride. Experimentally determined dataused in this report are given in Table 5.10. Further constants are compiled in Tables 5.11 and5.12.

Table 5.10: Experimentally determined equilibrium data compiled for the antimony(III)fluoride system. These data were chosen for the evaluation of recommendedvalues in the present report. Additional information for the different references seeSection 5.12. Method: pol = polarography, pot = potentiometry.

Pm,n,oReference Comments I (M) Medium Method

log Ku: Sb3+ + HF& SbF2+ + H+

3.00 [1970BON] T= 303 K, 1=2 2 Na, HC1O4 pol

log Klt2: Sb3+ + 2HF <=> SbF2+ + 2H+

5.70 [1970BON] T= 303 K, 1=2 2 Na, HC104 pol

log KJJ: Sb3+ + 3HF <=> SbF3° + 3H+

8.30 [1970BON] T= 303 K, 1=2 2 Na, HC1O4 pol

log Kh4: SbF3°

2.65 [1970BON]

SbF4~ + H+

T= 303 K, 1=2 2 Na, HC1O4 pol

5.4.1 Reactions ofSb3+: SbF2+, SbF2+, SbF30, and SbF4~

Antimony(in) forms weak complexes with fluoride. From the data determined by [1970BON]at 1=2 (Table 5.10) the tentative formation constants for Sb(III) fluoride complexes can bederived with the SIT equation. The Ae values of 0.04 and -0.04 are estimated from the NEAamericium data (£(Am3+, C1O4-) = 0.49, e(AmF2+, CIO4-) = 0.39, £(AmF2

+, CIO4-) = 0.17,and e(H+, CIO) = 0.14; [1995SIL/BID]).

Sb3+ + HFSb3+ + 2HFSb3+ + 3HF

SbF2+

SbF30

log K°u = 4.03, Ae = 0.04log K°i,2 = 7.02, Ae = -0.04log K°i,3 = 9.55, Ae =-0.07

104

JNC TN8400 9 9 - 0 1 1

Correction of these values with a log Ka of 3.18 for F" + H+ o HF [1992GRE/FUG] and forthe reaction Sb(OH)3° + 3H+ <=> Sb3+ + 3H2O (Section 5.1.1) gives the following tentativevalues:

Sb(OH)3°+3H++ F- <=> SbF2+ + 3H2O log p ° u = 6.48Sb(OH)3o +3H+ + 2F- <=> SbF2+ + 3H2O log ( 3 \ 2 = 12.65Sb(OH)3° +3H+ + 3F- «• SbF3° + 3H2O log p \ 3 = 18.36

No data is recommended in this report for the formation of SbF^.

5.4.2 Reactions ofSbF3°: SbF4~, and SbF3OH-

Recently, [1993DEL/MIL] determined constants for the formation of SbF^ from SbF3° and forthe hydrolysis of SbF3° to SbF3OH~ (Table 5.11). Unfortunately, their experiments werecarried out at varying ionic strength. Extrapolation to 1=0 is therefore not possible.

5.4.3 SbOF°orSb(OH)2F°

In Table 5.12 data are compiled for the formation of SbOF0 or Sb(OH)2F°. Direct experimentalmeasurements of these constants, however, are not available.

5.4.4 SbF3(s)

Direct measurements of the solubility of SbF3(s) are not available. Based on the data compiledin Table 5.12 it can be concluded that SbF3(s) is quite soluble.

Table 5.11: Additional, experimentally determined equilibrium data compiled for Sb(III) fluoride system.These data were not selected in the present report for the evaluation of recommended stabilityvalues. Method: NMR = 19F NMR, tit = titration (pH).

log Pm,n,0 Reference

log K, ,: Sb3+ + HF <=> SbF2+ +

3.11 ' [1975HEI/SCH]

Comments

H+

T= 298.15 K, 1=1.5

I (M) Medium

HF, H,SO4

Method

pot

log K12: Sb3+ + 2HF <=> SbF2* + 2H+

5.11 ' [1975HEI/SCH] T= 298.15 K, 1=1.5 HF, H7SO4 p_ot_

log Ku: Sb3+ + 3HF <=> SbFf + 3H+

6.23 ' [1975HEI/SCH1 T= 298.15 K, 1=1.5 HF, H,SO4 pot^

105

JNC TN8400 9 9 - 0 1 1


log K14: SbF3° + HF <=> SbFf + H+

1.72 2 [1993DEL/MIL] T= 298 K, 1=0.03-1 var self medium NMR

log KUJ: SbF3° + H2O <=> SbF3OH~ + H+

3.92 2 [1993DEL/MIL1 T= 298 K, 1=0.03-0.1 var self medium tit' exact I not known.2 in this report the presence of HF (instead of F~ as by [1993DEL/MIL]) is assumed, as measurements are carried out at

pH 1.6 - 2.8). I not constant.

Table 5.12: Thermodynamic data for the Sb(III) fluoride system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log Pm.n.o Reference

log K,j: Sb3+ + HF <=> SbF2+ +

3.00 [1980BON/HEF1

log Ku: Sb3+ + 2HF <=> SbF2+ •

5.70 [1980BON/HEF1

log Ku: Sb3+ + 3HF <=> SbF3° -

8.30 [1980BON/HEF]

log KK4: SbF3° + HF<=> SbF4- -

2.70 [1980BON/HEF1

Comments

H+

T= 298.15 K, 1=2

+• 2H+

T= 298.15 K, 1=2

f 3H+

T= 298.15 K, 1=2

¥ H+

T= 298.15 K, 1=2

KM;

2

2

2

2

i Medium

Na, HC1O,

Na, HC1O,

Na, HC1O4

Na, HC1O4

Sb(OH)3° H+

6.51 '6.19

[1982WAG/EVA][1985PAS1

SbOF0 + 2//2O

T= 298.15 K,I=n/aT= 298.15 K,I=n/a

log jij

6.52 '6.52 '6.21

: Sb(OH)3° H+ <=> Sb(OH)2F° + H2O

[1980BEN/TEA][1982WAG/EVA][1985PAS1

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

log K'S3: Sb(OH)3° + 3F'+ 3H+ <=> SbF3(s) + 3H2O

10.48 '12.96 '12.78 '12.78 '10.17

[1952LAT][1963WIC/BLO][1977BAR/KNA][1979KUB/ALC][1985PAS1

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/a

calculated in this report with a Afi" of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1)

106

JNC TN8400 9 9 - 0 1 1

5.5 Sb(III) sulfate system

Antimony(III) is reported to form complexes and compounds with sulfate. Experimental datafor complex formation of Sb(III) with sulfate were determined by [1970DAWAVTL] inconcentrated H2SO4 solutions (Table 5.13). Additional constants are compiled in Table 5.14.

5.5.1 SbOSO4~

[1970DAWAVTL] proposed the formation of SbOSO4- in 2 - 6 M H2SO4 solutions and ofSb(SO4)2~ in 14 - 16 M H2SO4 solutions. For the calculation of formation constants (Table5.13) he assumed in 0.2 - 12 M H2SO4 solutions the presence of SbO+, which seems quitedoubtful in the light of the results discussed in Section 5.1.1 and 5.1.2 of this report. Noformation constants for the Sb(III) sulfate complexes are proposed in this report. Suchcomplexes, however, will be important only in concentrated sulfate solutions[1970DAW/WIL].

5.5.2 Sb2(SO4)3(s)

Thermodynamic values for Sb2(SC>4)3(s) were compiled by [1977BAR/KNA] and[1979KUB/ALC]. Solubility products calculated in this report from these thermodynamic dataare shown in Table 5.14. From these data, Sb2(SC>4)3(s) is expected to be very soluble. Nodirect experimental determination of Sb2(SC>4)3(s) solubility is available.

Table 5.13: Experimentally determined equilibrium data compiled for Sb(III) sulfate system. These data werenot chosen in the present report for the evaluation of recommended stability values. Reasons fornot selecting these references see text. Method: cat = cation exchange.

l°g Pm.n Reference Comments I (M) Medium Method

log pij: SbO+ + SO42~ <=> SbOSOf

0.30 ri970DAW/W]L1 T= 298.15 K, 1=1-4 var H,SO^ cal_

Table 5.14: Thermodynamic data for the Sb(III) sulfate system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

log P m n o Reference Comments

log K'S3: 2Sb(OH)3° + 3SO42- + 6H+ <=> Sb2(SO4)3(s)

-5.0 ' [1977BAR/KNA] T= 298.15 K, I=n/a-5.0 ' [1979KUB/ALC1 T= 298.15 K, I=n/a1 calculated in this report with a A(G° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1)

107

JNC TN8400 9 9 - 0 1 1

5.6 Antimony(III) sulfide system

Antimony(III) forms complexes with sulfide and also forms a stable salt, stibnite (Sb2S3(cr)).[1988KRU] and [1994AKI/ZOT] studied complex formation with sulfide from measurementsat 25 - 300 °C. [1994AKI/ZOT] included also the results determined by [1988KRU] and put upa consistent thermodynamic data set for the Sb(III)-S(II)-OH system. Equilibrium constants arecompiled in Table 5.15.

Table 5.15: Experimentally determined equilibrium data compiled for the Sb(III) sulfidesystem. These data were chosen for the evaluation of recommended values in thepresent report. Additional information for the different references see Section5.12: 'Comments on selected references'. Method: sol = solubility measurements.

log Pm,n,o

log Po,2J:

42.55 142.51 2

Reference

2 Sb(OH)3° +4HS-

[1988KRU][1994AKI/ZOT]

Comments

+ 2H+

T=298,T=298,

<=> Sb2S<

I=dilI=dil

t2- + 6H2O

I(M)

00

Medium

H2SH,S

Method

solsol

log Pi,2,4.- 2 Sb(OH)3° + 4HS- + 3H+ <=> HSbtfr + 6H2O

52.07 i52.29 2

log /32,2,4:

56.98 ]

57.02 2

log K*S3:

55.14 2


2 Sb(OH)3° + 4HS-


2Sb(OH)3° + 3HS- +

[1994AKI/ZOT]

T=298, I=dilT=298, I=dil

00

+ 4H+ <=> H2Sb2S4° + 6H2O

T=298, I=dilT=298, I=dil

00

3H+ <=> Sb2S3(stibnite) + 6H2O

T=298, I=dil 0

H2SH,S

H2SH,S

H,S

solsol

solsol

sol

1 extrapolated to 1=0 with Debye-Hiickel by [1988KRU]; calculated with a log K*so for stibnite of 55.14.2 extrapolated to 1=0 with Debye-Hiickel by [1994AKI/ZOT]; calculated from Afi° given by [1994AKI/ZOT].

5.6.1 Sb2S42', HSb2S4-, and H2Sb2S4°

Sb(III) forms complexes with sulfide (Table 5.15 and 5.16). As the determination of the Sb(III)sulfide complex formation constants at 1=0 by [1988KRU] and [1994AKI/ZOT] are the only

108

JNC TN8400 9 9 - 0 1 1

experimental value reported in the literature. The mean of these logrecommended in this report as tentative values:

2Sb(OH)3 + 4HS-+ 2H+2Sb(OH)3 + 4HS- + 3H+2Sb(OH)3 + 4HS-+ 4H+

<=> Sb2S42-+6H2O

HSb2S4-+6H2OH2Sb2S4° + 6H2O

values (Table 5.15) are

log p°0,2,4 = 4 2 - 5 3

log P°i'2 '4= 52.18log p° 2 ^ 4 = 57.00

5.6.2 Sb2S3(stibnite)

Stibnite, Sb2S3(cr), is quite insoluble, as indicated by the values compiled in Table 5.15 and5.16. The log K*°s3 value determined by [1994AKI/ZOT] is recommended as tentative value:

2Sb(OH)3° + 3HS-+ 3H+ Sb2S3(stibnite) + 6H2O logK*°S3 =55.14

Table 5.16: Thermodynamic data for the Sb(II) sulfide system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

log r s o

log Po,2,4-50.0 '50.0 >49.249.4

lOg Po,l,2-25.8 '

log K*S3:

66.05 '60.1960.2556.68 '

log K*S3:

3.07 2-66.0560.8066.0560.8055.0260.82 '

log K'S3:

53.9 '

Reference

2Sb(OH)3° +4HS-[1980BEN/TEA][1982WAG/EVA][1985PAS][1986ITA/NIS]

Sb(OH)3° + 2HS- +

[1952LAT1

2Sb(OH)3° + 3HS- +

[1974MIL][1985PAS][1986ITA/NIS]ri992SEAl

2Sb(OH)3° + 3HS- +

[1971NAU/RYZ][1977BAR/KNA][1978ROB/HEM2][1979KUB/ALC][1980BEN/TEA][1981JOH/PAP][1982WAG/EVA1

2Sb(OH)3° + 3HS- +

[1952LAT1

+ 2W

H+ <^

3H+

3H+

3H+

Comments

<=> Sb^/- + 6H2O

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

> SbS2- + 3H2O -

T= 298.15 K, 1=0

« Sb2S3(stibnite) + 6H2O

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298, I=n/a

<=> Sb2S3(s) +6H2O

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/a

i=> Sb2S3(am, orange) + 6H2O

T= 298.15 K, I=n/a

KM)

0000

0

0

1 calculated in this report with a AfG° of-643.0 kJ/mol for Sb(OH)3° (section 5.7.1)2 difference to others values probably due to missing minus sign in the Afi° of 156.1 kJ/mol [1971NAU/RYZ].

109

JNC TN8400 9 9 - 0 1 1

5. 7 Redox reactions

Antimony exists in the oxidation states -III, 0, +III and +V [1976BAE/MES, 1985PAS,1995WIB]. The -III oxidation state is thermodynamically not stable [1952LAT, 1985PAS] inpresence of aqueous solutions. In water, aqueous species of Sb(III) and Sb(V) are stable.

Experimentally determined equilibrium data for the redox potential of antimony chosen in thisreport for the calculation of equilibrium constants at 1=0 are given in Table 5.17. Additional dataare compiled in Table 5.18 and 5.19.

Table 5.17: Experimentally determined equilibrium data compiled for the redox potential ofantimony. These data were chosen for the evaluation of recommended values inthe present report. Additional information for the different references see Section5.12: 'Comments on selected references'. Method: pot = potentiometry, sol =solubility measurements.

logK Reference Comments I (M) Medium Method

log K: Sb(OH)2+ + 2H+ + 3e- ^Sb(cr) + 2H2O

10.3810.4410.6810.6410.67

[1972VAS/SHO][1972VAS/SHO][1972VAS/SHO][1972VAS/SHO][1972VAS/SHO]

T=T=rp

T=T=

298.15298.15298.15298.15298.15

K,K,K,K,K,

1=0.361=0.711=1.81=2.11=2.5

0.360.711.782.142.49

HCIO4HC1O4

HCIO4HCIO4HCIO4

potpotpotpotpot

log K: Sb(OH)3° + 3e~ + 3H+ 4=> Sb(cr) + 3H2O

11.88 [1990SHI/ZOT] T=298, I=dil11.83 [1994AKI/ZOT] T=298, I=dil

00


solsol

log K: Sb(OH)6- + 2e~ <^Sb(OH)4~+ 20H~

-16.36-17.44-18.22-18.73-18.97-19.20-19.54-19.91

[1923GRU/SCH][1923GRU/SCH][1923GRU/SCH][1923GRU/SCH][1923GRU/SCH][1923GRU/SCH][1923GRU/SCH][1923GRU/SCH]

T=T=T=

T=T=

T=

293293293293293293293293

K, 1=4K, 1=5K, 1=6K, 1=7K, 1=7.5K, 1=8K, 1=9K,1=10

4567

7.58910

KOHKOHKOHKOHKOHKOHKOHKOH

potpotpotpotpotpotpotpot

110

JNC TN8400 9 9 - 0 1 1


-19.54-17.68-16.36-14.98

[1953HAI][1953HAI][1953HAI][1953HAI]

T=300K, 1=10 MT= 300 K, 1=6 MT= 300 K, 1=3 MT= 300 K, 1=1 M

10631

KOHKOHKOHKOH

potpotpotpot

5.7.1 Sb(cr)/Sb(OH)3°

The redox couple Sb(cr)/Sb(OH)3° is not well investigated. Sb(cr) has a hexagonalrhombohedral structure and has a gray color [1995WIB]. [1924SCH] showed that in 0.2 to 1.1M HC1O4 solutions Sb(OH)2+ dominates the speciation. Also the hydrolysis constantscalculated in Section 5.1 indicate that in this pH region Sb(OH)2+ dominates the speciation.[1972VAS/SHO] determined with potentiometric measurements under acidic conditions (Table5.17) constants for the reaction Sb(OH)2

+ + 2H+ + 3e- <=> Sb(cr) + 2H2O.

The values of [1972VAS/SHO] were extrapolated to I = 0 in Figure 5.7 giving a log K° of10.83. Using the log (3°li2 of 1.30 (Section 5.1.3: Sb(OH)2

+), a log K° of 12.13 for thereaction Sb(OH)3° + 3H+ + 3e- <=> Sb(cr) + 3H2O is obtained.

Sb(OH)2++3e+2H+<^>Sb(cr)+2H2O

13

lm, molal

Figure 5.7: Plot of log K + 3 D vs. Im for the reaction : Sb(OH)2+ + 2H+ + 3e- «=> Sb(cr) +

2H2O at 25 °C. The straight line shows the result of the linear regression: Ae =-0.21; log K° = 10.83. Calculated from data given in Table 5.17.

I l l

JNC TN8400 9 9 - 0 1 1

[1990SHI/ZOT and 1994AKI/ZOT] determined constants for the reaction Sb(OH)3° + 3H+ +3e- o Sb(cr) + 3H2O with solubility measurements (Table 5.17) and extrapolated these valuesto 1=0 with the Debye-Hiickel equation. This value shows a good agreement with thepotentiometric data of [1972VAS/SHO]. The mean of the measurements of [1972VAS/SHOand 1990SHI/ZOT, 1994AKI/ZOT] is:

Sb(OH)3° + 3H+ + 3e- Sb(cr) + 3H2O logK° = 11.990.236 V

resulting in a AfG° of -643.0 kJ/mol for Sb(OH)3°.

5.7.2 Sb(III)/Sb(V)

The equilibrium between Sb(V) and Sb(III) has been studied electrochemically in KOHsolutions by [1923GRU/SCH] and [1953HAI]. The log K values calculated in this report fromthe experimentally determined redox potentials are listed in Table 5.17 (log K = 2xEmeaSured[V]/0.05916). Extrapolation to 1=0 is shown in Figure 5.8 and gives a log Kb° value of -15.36 forthe reaction Sb(OH)6- + 2e- <=> Sb(OH)4~ + 2OH~. This value is similar to the mean log Kb

value of -15.74 (Table 5.19) calculated by [1957PIT/POU] based on the data measured by[1923GRU/SCH]. No other data for the redox equilibria between Sb(III) and Sb(V) have beenfound in the literature.

Sb(OH)6-+2e-oSb(OH)4-+20H-

g> -18 +

- 1 9 ••

-20 -m

-215 10

lmi molal15

Figure 5.8: Plot of log Kb - 2 D vs. Im for the reaction: Sb(OH)6- + 2e~ <=> Sb(OH)4- +2OH~ at 25 °C. The straight line shows the result of the linear regression: Ae =0.38; log Kb° = -15.36. Calculated from data given in Table 5.17.

112

JNC TN8400 9 9 - 0 1 1

Using the log Kb° of-15.36 (Figure 5.10), a log Kw of- 14.00, a log p°1>4 of -11.93 (for thereaction Sb(OH)3° + H2O <=> Sb(OH)4- + H+; see Section 5.1.4) and a log p° l t 6 of 2.72 (forthe reaction Sb(OH)5° + H2O <=> Sb(OH)6- + H+; see Section 5.8), the following constant forthe redox equilibria between Sb(OH)3° and Sb(OH)5° can be calculated:

Sb(OH)5° + 2H++ 2e- Sb(OH)3° + 2H2O logK° = 21.84E> = 0.646 V

5.7.3 Sb(cr)/Sb(OH)5°

From the data obtained in Section 5.7.1 and 5.7.2 also a log K° value for the theoreticalequilibrium between Sb(cr)/Sb(OH)5° can be calculated:

Sb(OH)5°+5 H++5e- <=> Sb(cr) + 5H2O logK° =33.83F = 0.400 V

corresponding to a AfG° of -992.62 kJ/mol for Sb(OH)5°.

5. 7.4 Additional data compiled for the redox potential of antimony

Table 5.18: Additional, experimentally determined equilibrium data compiled for the redox potential ofantimony. These data were not chosen in the present report for the evaluation of recommendedstability values. Method: pot = potentiometry.


log K: Sb(OH)2+ + 2H+ + 3e- <=> Sb(cr) + 2H2O

10.60 >10.80 '10.80 '10.80 >10.69 2

10.82 2

11.03 2

10.34 3

[1924SCH][1924SCH][1924SCH][1924SCH][1924SCH][1924SCH][1924SCH][1972VAS/SHO1

T= 298.15 K, 1=0.2T= 298.15 K, 1=0.5T= 298.15 K, 1=0.9T= 298.15 K, 1=1.1T= 298.15 K, 1=0.2T= 298.15 K, 1=0.5T= 298.15 K, 1=1.1T= 298.15 K, 1=0.3-2.5

0.20.50.91.10.20.51.10

HC1O4

HCIO4HCIO4HC1O4

HCIO4HCIO4HC1O4

HC1O,

potpotpotpotpotpotpotpot

13

JNC TN8400 9 9 - 0 1 1


log K: Sb(OH)6- + 2e~ <=> Sb(OH)f + 2OH~

-14.47 4

-15.19 5

-15.31 5

[1923GRU/SCH][1953HAI][1953HAI]

T= 293 K, 1=3T= 300 K, 1=1-10 MT= 303 K, 1=1-10 M

300

KOHKOHKOH

potpotpot

1 values corrected for H+ activity by [1924SCH]2 calculated in this report from E° values for the reaction Sb/Sb2O3(cr) as given by [1924SCH] and Sb(III) and H+

concentrations measured by [1924SCH] in presence of Sb2O3(cr).3 linear extrapolation to 1=0 by [1972VAS/SHO]4 this value is very different from the other values and it is not clear if the redox potential is -0.428 or -0.488 V. This

value is therefore not chosen in this report5 as given by [1953HAI]; from data at variable ionic strength

Table 5.19: Thermodynamic data for the redox potential of antimony taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions this is indicated.

logP

log K:

11.2611.6711.70

log K:

-15.72-13.49-14.74-15.38-15.79-15.86-15.86-15.86-15.69

Reference

Sb(OH)3° + 3e~ + 3H+ <=>

[1985B AS/MAT][1985PAS]ri986ITA/NIS]

Sb(OH)6-+ 2e~ <=>Sb(OH)1 [1957PIT/POU]2-3 [1957PIT/POU]2'3 [1957PIT/POU]2 [1957PIT/POU]2 [1957PIT/POU]2 [1957PIT/POU]2 [1957PIT/POU]2 [1957PIT/POU]2 [1957PIT/POU1

Comments

Sb(cr) + 3H2O

T=298.15,I=0T=298.15,I=0T=298.15,I=0

f + 2OH-

T= 293 K, 1=3-9 M,T= 293 K, 1=3-9 MT= 293 K, 1=3-9 MT= 293 K, 1=3-9 MT= 293 K, 1=3-9 MT= 293 K, 1=3-9 MT= 293 K, 1=3-9 MT= 293 K, 1=3-9 MT= 293 K, 1=3-9 M

KM)

000

034567

7.589

Medium

KOHKOHKOHKOHKOHKOHKOHKOHKOH

1 as calculated by [1957PIT/POU]; mean of data after correction for activity of OH~2 data given by [1957PIT/POU], corrected for activity of OH~ (original data from [1923GRU/SCH])3 these data were not chosen by [1957PIT/POU] for the calculation of a mean value

114

J N C T N 8 4 0 0 9 9 - O i l

5.8 Hydrolysis of antimony(V)

[1963LEF/MAR] studied the hydrolysis of Sb(V) in 0.5 M (CH3)4NH4C1 solutions and Sb(V)concentrations from 0.0013 to 0.346 M. They interpreted their data in terms of a sequence ofdodecamers Hi2_y[Sb(OH)6]i2y- (Table 5.21). These data were recalculated later by[1976BAE/MES] (Table 5.20).

[1948TOU/MOU] determined the solubility of Sb2O5(precip) in dilute acid. Their results indiluted acids are consistent with the existence of S b ^ O H ^ 4 " (or Sbi2(OH)633-) as proposedby [1963LEF/MAR] and [1976BAE/MES]. In more acidic solutions (pH < - 0.25) Sb(OH)2

3+

(or a polynuclear compound having the same Sb(V)/H+ ratio) seems to dominate the solution([1948TOU/MOU]). Also the observation of [1975ALY/ABD] indicate further hydrolysis ofSb(V) at pH < - 0.3.

The data determined by [1963LEF/MAR] are the only experimental data found in the literature.These data were recalculated by [1976BAE/MES] and its use is recommended in this report:

Sb(OH)5° + H12Sb(OH)5° +12Sb(OH)5° +12Sb(OH)5° +12Sb(OH)5° +

2 04H2O5H2O6H2O7H2O

«- Sb(OH)6- + H+

<=> Sbi2(OH)644-+« Sb12(OH)65

5-+« Sb12(OH)666-+<=> Sb12(OH)677-+

4H+

5H+

6H+

7H+

log P°,,6

log B°1 2 6 4

log P ° l M S

log P°12,66

log P°, , 6 7

= -2.72= 20.34= 16.72= 11.89= 6.07

Table 5.20: Equilibrium data used for the hydrolysis of antimony(V). These data were chosenas recommended values in the present report. Method: cal = calculated by[1976BAE/MES] based on experimental data reported by [1963LEF/MAR]

log R& r'm,5m+

Reference

log P1A: Sb(OH)5° +H2O <=>

-2.47 i-2.72 i

log Pl2,64

23.06 '20.34 1

[1976BAE/MES][1976BAE/MES]

: 12Sb(OH)5° + 4H2

[1976BAE/MES][1976BAE/MES]

Comments

Sb(OH)6- + H+

T= 298.15 K, 1=0.5T= 298.15 K, 1=0

,0 <=> Sb12(OH)6/- +

T= 298.15 K, 1=0.5T= 298.15 K, 1=0

I(M)

0.50

4H+

0.50

Medium

(CH3)4NH4C1(CH,)4NH4C1

(CH3)4NH4C1(CH,)4NH4C1

Method

calcal

calcal

log pi2,65- 12Sb(OH)5° + 5H2O <^> Sb!2(OH)655- +

16.72' [1976BAE/MES] T= 298.15 K, 1=0 0 (CHQ4NH4C1 cal

115

JNC TN8400 9 9 - 0 1 1


log Pj2,66- 12Sb(OH)5° + 6H2O <=> SbirfOH)^6- + 6H+

11.89 l [1976BAE/MES] T= 298.15 K, 1=0 0 (CH,)4NH4C1 cal

log Pi2,67- 12Sb(OH)5° + 7H2O & Sb}2(OH)677-

6.07 l [1976BAE/MES] T= 298.15 K, 1=0

7H+

0 (CH,)4NH4C1 calrecalculated by [1976BAE/MES] from data of [1963LEF/MAR]

Table 5.21: Experimentally determined equilibrium data compiled for the hydrolysis of antimony(V). Thesedata were not chosen in the present report for the evaluation of recommended stability values.Reasons for not selecting these references see text. Method: tit = pH titration.

log Pm,5m+n Reference Comments I (M) Medium Method

log j5L6: Sb(OH)5° + H2O <=> Sb(OH)6~ + H*

-2.55 ' ri963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CHQ.NH.C1 tit

log P12,63: 12Sb(OH)5° + 3H2O O Sb,2(OH)633~ + 3H+

24.45 • [1963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CH,).NHdCl tit

log j3!2M: 12Sb(OH)5° + 4H2O <=> Sb12(OH)6/- + 4H*

22.90 ' fl963LEF/MARl T= 298.15 K, 1=0.5 0.5 tit

log Pn,65-- 12Sb(OH)s° + 5H2O <=> Sb12(OH)655- + 5H+

19.95 ' [1963LEF/MAR1 T= 298.15 K, 1=0.5 0.5 (CH,).NH4C1 tit

log pil66: 12Sb(OH)5° + 6H2O « Sb12(OH)666~ + 6H+

15.80 • [1963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CH,).NH4C1 tit

log p,l67: 12Sb(OH)5° + 7H2O

9.85 ' [1963LEF/MAR1

Sb12(OH)677~ + 7H+

T= 298.15 K, 1=0.5 0.5 (CHQ.NH.Cl tit

'OS P 12.68- 12Sb(OH)5° + 8H2O <=> Sb12(OH)68s- + 8H+

^/70> [1963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CH,)aNHaCl titSb = 1.3 - 346 mM; later recalculated by [1976BAE/MES], see Table 5.20

116

JNC TN8400 9 9 - 0 1 1

Table 5.22: Thermodynamic data for hydrolysis of antimony(V) taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions this is indicated.

log Pm,5m+n Reference Comments I (M) Medium

log KIA: Sb(OH)6- + 2H+ <=> Sb(OH)4+ + 2H2O

-0.54 !-0.55 2

[1957PIT/POU]ri986ITA7NISl

T= 308 K, 1=0.05-5T= 298.15 K, I=n/a

HC1

log fi,,6: Sb(OH)5° + H2O <=> Sb(OH)6~ + H+

-2.47-2.72

log P)2.6

23.0620.34

log P,2,6

16.72 3

log P12.6

11.89 3

6.07 3

[1976SMI/MAR][1976SMI/MAR]

4: 12Sb(OH)5° + 4H2O <=

[1976SMI/MAR][1976SMI/MAR1

5: 12Sb(OH)5° + 5H2O <z

[1976SMI/MAR]

6: 12Sb(OH)5° + 6H2O <=

[1976SMI/MAR1

7: 12Sb(OH)5° + 7H2O <=

[1976SMI/MAR]

T= 298.15 K, 1=0.5T= 298.15 K, 1=0

* Sb12(OH)644- + 4H+

T= 298.15 K, 1=0.5T= 298.15 K, 1=0

* Sb12(OH)65s- + 5H+

T= 298.15 K, 1=0

* Sb!2(OH)666' + 67T

T= 298.15 K, 1=0

* Sb12(OH)677~ + 7H+

T= 298.15 K, 1=0

0.50

0.50

0

0

0

(CH3)4NH4C1

(CH3)4NH4C1

1 calculated by [1957P1T/POU] from the data of [1948TOU/MOU]2 data from [1957PIT/POU]3 these values are cited wrongly in the NEA database

117

JNC TN8400 9 9 - 0 1 1

5.9 Sb2O5(precip)

The only reported solubility measurements for Sb2O5(precip) is that of [1948TOU/MOU] at 35°C in water and dilute acid. [1948TOU/MOU] reported a solubility of 2.71xlO"4 M Sb(V) inwater. [1976BAE/MES] calculated later from these data (Table 5.23):

2Sb(OH)5° <=> Sb2O5(precip) + 5H2O log K*°so =7.40

[1976BAE/MES] state that in view of the ease with which gels are prepared from antimonic acidsolutions, it is not likely that a pure phase was present in the experiments of [1948TOU/MOU].[1984BER/BRE] find that the precipitation product obtained from the hydrolysis of Sb(V) indilute HC1 is amorphous to X-ray detection up to 650 °C.

Table 5.23: Equilibrium data determined for Sb2Os(precip). These data were chosen asrecommended values in the present report. Method: cal = calculated by[1976BAE/MES] based on solubility data reported by [1948TOU/MOU] (Table5.24).

a Reference Comments I (M) Medium MethodPm.5i,5m+n

log K*so: 2Sb(OH)5° <^> Sb2O5(precip) + 5H2O

7.40 l [1976BAE/MES] T= 308 K, 1=0 0 water cal_1 recalculated by [1976BAE/MES] from data of [1948TOU/MOU]

Table 5.24: Experimentally determined equilibrium data compiled for Sb2O5(precip). These data were notchosen in the present report for the evaluation of recommended stability values. Method: sol =solubility.

l°g Pm,5m+n Reference Comments I (M) Medium Method

log K*so: 2Sb(OH)5° <=> Sb2O5(precip) + 5H2O

7.13 ' [ 1948TOU/MOU1 T= 308 K, I=diluted self medium sol' approximate value; equilibrium with Sb(OH)5° assumed; calculated in this report based on the solubility of Sb in water

= 2.71X10"4 M

118

JNC TN8400 9 9 - 0 1 1

Table 5.25: Thermodynamic data for Sb2O5(s) taken from previous compilations. As pointed out in Section2 of this report only experimental data were used for the present evaluation. The following tableserves only for comparison.

log pm i5m + n Reference Comments I (M) Medium

log K*so: 2Sb(OH)5° <=> Sb2Os(s) + 5H2O

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,L=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/a

log K*S4: 2Sb(OH)4+ <^> Sb2O5(precip) + 4H2O + 2W

8.32 2 [1957PIT/POU] T= 308 K, 1=0.05-5 HC11 all these values were calculated assuming a log K (Sb(OH)5°/Sb(cr)) of-33.83 (Section 5.7.3)

6.88 '3.22 '3.22 '11.41 '11.43 '11.42 '5.20 '5.19 '5.20 '5.20 '5.21 '5.21 '

[1952LAT][1954COU][1963WIC/BLO][1971NAU/RYZ][1977BAR/BCNA][1979KUB/ALC][1980BEN/TEA][1982PAN][1982WAG/EVA][1985BAB/MAT][1985PAS][1986ITA/NIS]

2 calculated by [1957PIT/POU] from the data of [1948TOU/MOU]

119

JNC TN8400 9 9 - 0 1 1

5.10 Sb2O4(cr) and Sb6O13(cr)

When Sb2C>3(cr) is heated in the presence of oxygen above 300 °C, orthogonal Sb2C>4(cr) isformed [1971HEG/BAK, 1984BER/BRE, 1987PAN/SRE, 1997KOL/HEN]. OrthorhombicSb2O4(cr) is generally assumed to be a mixed compound of Sb2C>3(cr) and Sb2C>5(s)[1971HEG/BAK, 1976BAE/MES, 1995WIB]. [1984BER/BRE] observed in nitric acid theformation of orthorhombic oc-Sb2C>3(cr) at room temperature and the formation of Sb2C>4(cr)above 360 °C in air from powdered Sb(cr). [1984BER/BRE] also showed that the precipitationproduct obtained from Sb(III) and Sb(V) solutions under acidic conditions is different,indicating a kinetic barrier for the equilibria between Sb(III) and Sb(V) under acidic conditions.

[1984BER/BRE] found that the precipitation product obtained from the hydrolysis of Sb(V) indilute HC1 was amorphous to X-ray detection up to 650 °C. After calcination at 735 °C theyfound that SbgO^cr) was formed.

Based on these observations it is not expected that Sb2C>4(cr) or Sb6Oi3(cr) are formed at roomtemperature and no thermodynamic values for these solids are recommended in this report.

Solubility products (with relation to Sb(OH)5°), converted in this report from AfG° values givenin the literature, are compiled in Tables 5.26 and 5.27.

Table 5.26: Experimentally determined equilibrium data compiled for Sb2O4(s). These data were not chosenin the present report for the evaluation of recommended stability values. Method: emf =electromotive force measurements at high temperatures.

log K*so Reference Comments - I (M) Medium Method

log K*so: 2Sb(OH)5°+ 2H+ + 2er <=> Sb2O4(s) + 6H2O

35 .63 ' - 2 [1987PAN/SRE] T= 298.15 K, I=n/a n/a emf1 this value was calculated assuming a log K (Sb(OH)5°/Sb(cr)) of -33.83 (Section 5.7.3)2 extrapolated from = 800 K; originally calculated with for A,G° of a-Sb2O3(cr) (-613.03 kJ/mol) (1986AZA/PAN]; in

this report corrected for AfG° of a-Sb2O3(cr) (-624.32 kJ/rnol) derived in this report

120

JNC TN8400 9 9 - 0 1 1

Table 5.27: Thermodynamic data for Sb2O4(s) and Sb6Ol3(s) taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison. All these values were calculated assuming a log K(Sb(OH)5°/Sb(cr)) of-33.83 (Section 5.7.3)

logPm

log K\

23.0635.2335.2340.9240.9240.8840.9140.8740.8940.89

log r

77.71

5m+n Reference

0: 2Sb(OH)5°+ 2H+ + 2e~ <=

[1952LAT][1954COU][1963WIC/BLO][1977BAR/KNA][1979KUB/ALC][1980BEN/TEA][1982PAN][1982WAG/EVA][1985PAS][1986ITAyNISl

0: 6Sb(OH)5° + 4H+ + 4e~ <=

fl952LATl

Comments

? Sb2O4(s) +

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

* Sb6OI3(s) -

T= 298.15

6H2O

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/a

I- 77//2O

K, I=n/a1 all these values were calculated assuming a log K (Sb(OH)5°/Sb(cr)) of-33.83 (Section 5.7.3)

121

JNC TN8400 9 9 - 0 1 1

5.11 Other antimony(V) complexes and compounds

For other antimony(V) species and compounds the amount of experimentally determined dataavailable in the literature is rather limited.

[1975ALY/ABD] determined stability constants of Sb(V) chlorocomplexes by cation exchange.They give no indication of the speciation of Sb(V) present in 0.1 - 2 M HC1. However, in thepH range -0.3 to 1, the presence of Sb(OH)5° (or of is probable (see Section5.8). The formation constants given by [1975ALY/ABD] for the formation of Sb(V)chlorocomplexes are compiled in Table 5.28. [1969BRY/IOF] state that SbClg- dominates inconcentrated hydrochloric acid, SbCl5OH- in 9 M HC1 and SbCl4OH2- in 6 M HC1.

[1974BLA/BUR] observed a solubility of 0.0033 M NaSb(OH)6(s) in water. As pH was notmeasured it is not possible to calculate a solubility product. [1996KAS/MUK] give standardenthalpies and entropies for alkaline-earth metal antimonates produced in the temperature rangeof 200-600 °C.

From the available data it can be expected that Sb(V) forms weak complexes with chloride (andfluoride) and that easily soluble solids with alkali and earth alkali metal ions can be formed.However, there are not sufficient data available to recommend thermodynamic data for thesecomplexes and solids. Data for other inorganic ligands (e.g., sulfate, nitrate, fluoride, ..) havenot reported in the literature.

Table 5.28: Experimentally determined equilibrium data compiled for the antimony(V) system. These datawere not chosen in the present report for the evaluation of recommended stability values.Reasons for not selecting these references are given in Section 5.12: 'Comments on selectedreferences'. Method: ex = cation exchange.


log KUJ: "Sb(V) + Cl- <=> Sb(V)Cl-

0.340.320.200.461.321.82

[1975ALY/ABD][1975ALY/ABD][1975ALY/ABD][1975ALY/ABD][1975ALY/ABD][1975ALY/ABD1

T= 298.15 K, I=.lT= 298.15 K, I=.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=4T= 298.15 K, 1=6

0.10.5

1246

HC1HC1HC1HC1HC1HC1

exexexexexex

1 Sb = 1 mM; [1975ALY/ABD] give no indication of the speciation of Sb(V) present in 0.1 - 2 M HC1. However, at pH 1to -0 .3 the presence of Sb(OH)5° (or of Sb ]2(OH)M

4-) is probable (see Section 5.8). [1975ALY/ABD] observed in 4 and 6M HC1 solutions an increase of the measured log K values, which is consistent of the hydrolysis as observed by[1948TOU/MOU] (Section 5.8).

122

JNC TN8400 9 9 - 0 1 1


[1923GRU/SCH]: [1923GRU/SCH] determined the redox potential of the reactions Sb(cr) +4OH- <=> Sb(OH)4- + 3e~ and Sb(OH)4- + 2OH~ & Sb(OH)6- + 2e~ as -0.675 Vand -0.589 V, respectively, in 10 N KOH. The redox potential of Sb(OH)4- + 2OH~<=> Sb(OH)<5~ + 2e~ was also determined more diluted solutions (Table 5.17).

[1924SCH]: [1924SCH] determined in 0.2 - 1.1 M HC1O4 both the solubility of Sb(III) inperchlorate acid and the electrode potential of the reaction Sb(cr) + H2O <=> SbO+ +2H+ + 3e~. Both measurements resulted in similar log (31 2 values (Table 5.1 and5.2). For comparison the log of these values were corrected for the reaction Sb(cr)+ 3H2O « Sb(OH)3 + 3H+ + 3e~ with a factor of 11.99 (see Section 5.7: Redoxreactions). These values agree very well with the log pi 2 values for the reactionSb(OH)3° + H+ <=> Sb(OH)2

+ + H2O measured with other techniques (Tables 5.1and 5.2).

[1948TOU/MOU]: [1948TOU/MOU] determined the solubility of valentinite (a-Sb2O3) inwater and diluted NaOH at 35 °C. From their experimental results, constants for thereactions V2 a-Sb2O3 + 1.5H2O <=> Sb(OH)3° and V2 a-Sb2O3 + 2.5H2O <=>Sb(OH)4- + H+ and Sb(OH)3° + H2O «• Sb(OH)4- + H+ can be calculated. Theirvalues agree well with data measured at 25 °C (see Tables 5.1 and 5.4) and wereincluded in our calculations.

[1948TOU/MOU] determined the solubility of Sb2O5(precip) in water and indiluted HC1 at 35 °C. They reported a solubility of 2.71x10^ M Sb(V) in water.Their results in diluted acids are consistent with the existence of Sbi2(OH)64

4~ (orSb12(OH)63

3-) as proposed by [1963LEF/MAR] and [1976BAE/MES]. In moreacidic solutions (pH < - 0.25) Sb(OH)2

3+ (or a polynuclear compound having thesame Sb(V)/H+ ratio) seems to dominate the solution.

[1952GAY/GAR]: [1952GAY/GAR] measured in 0 - 0.1 M HC1 and 0 - 0.1 M NaOH Sbhydrolysis and determined a mean log P12 of 1.13 for the reaction Sb(OH)3° + H+

<=> SbO+ + 2H2O. As Sb is known to form weak complexes with chloride,measurements in HC1 medium were not used for extrapolation. However, thevalues in acidic medium agree very well with other measurements reported in theliterature (Tables 5.1 and 5.2). In alkaline medium, constants can be calculated as afunction of I from their experimental results for the reactions V2 a-Sb2O3 +1.5H2O <=> Sb(OH)3° and V2a-Sb203 + 2.5H2O <=> Sb(OH)4- + H+ (respectivelySb(OH)3° + H2O <=> Sb(OH)4- + H+). The constants calculated from themeasurements agree well with other measurements given in the literature (Table5.1) and were included in our calculations.

[1957PIT/POU]: [1957PIT/POU] calculated the Sb2O3 solubility from data from differentsources.

123

JNC TN8400 9 9 - 0 1 1

[1968MIS/GUP]: [1968MIS/GUP] determined spectrophotometrically a mean log p \ 2 valueof 1.42 in 0.02-0.1 M HC1O4 for the reaction Sb(OH)3° + H+ <=> SbO+ + 2H2O.From their measurements it can be concluded that from pH 2 to 1 (0.02-0.1 MHCIO4) no further hydrolysis of Sb(III) takes place (predominance of SbO+, orSb(OH)2

+), while at pH < 0 (1 and 3 M HC1O4) extinction coefficients are differentindicating the presence of SbOH2+ or Sb3+.

[1970BON/WAU]: [1970BON/WAU] and [1970BON] determined polarographically in 5and 2 M HCIO4 values for the reaction Sb(cr) <=> Sb3+ + 3e~. From the redoxpotential given by [1970BONAVAU] and [1970BON] (corrected in this report forthe potential of the AgCl/Ag electrode; 0.1988 V in saturated media,[1996STU/MOR]), log p°i)0 values for the reaction Sb(OH)3° + 3H+ <=> Sb3+ +3H2O can be calculated (using a log K° of 11.99 for the reaction Sb(cr) + 3H2O <=>Sb(OH)3 + 3H+ + 3e~ Section 5.7.1). The values are given in Table 5.2 and agreequite well with the log Pi.o values measured by [1977ANT/NEV]. It is not clear,however, if in the experiments of Bond and co-workers [1970BONAVAU,1970BON], Sb3+ is the only species present.

[1970BON]: see [1970BON/WAU]

[1973BEH/ROS]: [1973BEH/ROS] calculated AfG° values at 298 K for a-Sb2O3 based ondetermination of AH and AS values at 600-700 K.

[1973VAS/SHO2]: [1973VAS/SHO2] determined the solubility of valentinite in diluteNaOH at 25 °C. From their experimental results, constants for the reactionSb(OH)3° + H2O o Sb(OH)4~ + H+ can be calculated as a function of I (Table5.1).

[1974SHO/MAB]: [1974SHO/MAB] studied the hydrolysis of Sb(III) in 0.1 and 3 MNaC104 and at a Sb concentration of < 10~8 M. Their results show that no Sb-ClO4

complexes are formed, and that the formation of Sb(III) chloride complexes willbecome important only at pH values < 2.

[1975ALY/ABD]: [1975ALY/ABD] determined stability constants of Sb(V)chlorocomplexes by cation exchange. They give no indication of the speciation ofSb(V) present in 0.1 - 2 M HC1. However, at pH 1 to -0.3 the presence ofSb(OH)5° (or of Sb12(OH)64

4-) is probable (see Section 5.8). [1975ALY/ABD]observed in 4 and 6 M HC1 (pH < -0.5 ) solutions an increase of the measured logK values, which is indicates further protonation of Sb(V) and is consistent with thehydrolysis as observed by [1948TOU/MOU] (Section 5.8).

[1975HEI/SCH]: [1975HEI/SCH] determined potentiometrically in 1.5 M H2SO4 and HFthe redox potential of the Sb(OH)2

+/Sb(cr) couple. From this value an approximatelog p l i 2 of 0.36 for the reaction Sb(OH)3° + H+ £=> Sb(OH)2+ + H2O (Table 5.2)

124

JNC TN8400 9 9 - 0 1 1

was calculated, using log K of 11.99 for the reaction Sb(cr) + 3H2O <=> Sb(OH)3

+ 3H+ + 3e~ (see Section 5.7: Redox reactions). Antimony forms complexes withthe fluoride and sulfate present in the electrolyte. Although [1975HEI/SCH]corrected their measurements for the interactions of Sb(III) with F~ and SO42-,their measurements differ from the other data reported. Considering the high H+concentration, the presence of Sb3+ or SbOH2+ is quite probable. Assuming thepresence of Sb3+ gives a log p ] ) 0 of 0.53 for the reaction Sb(OH)3° + 3H+ <=>Sb3+ + 3H2O which agrees well with other values found in literature (Table 5.2).From the indications given by [1975HEI/SCH] the exact value of ionic strength ofthe electrolyte present is not clear (1.5 - 4.5 M).

[1977ANT/NEV]: [1977ANT/NEV] studied the hydrolysis of Sb(III) under acidicconditions in 1 M NaC104 containing 2 x 10"5 M Sb(ffl). [1977ANT/NEV] used aspectrophotometric method in which the competition between the hydrolysisreaction and complexation with gallein was measured.

[1994AKI/ZOT]: [1994AKI/ZOT] determined Sb(III) hydrolysis and complex formationwith sulfide from measurements at 25 - 300 °C and an Sb(III) concentration of 40mM. They put up a consistent thermodynamic data set. They also included resultsdetermined in other studies in their calculations. The complex formation constantsfor the Sb(III) sulfide system are based on solubility measurements of stibnite madeby [1994AKI/ZOT] and by [1988KRU] at different pH values, temperatures andH2S concentrations.

125

JNC TN8400 9 9 - 0 1 1

6 Lead

The oxidation states 0, +11, +IV are found in naturally occurring lead compounds. The mostcommon oxidation state is +11. The hydrolysis behavior of Pb(II) has been the subject ofextensive work. There are at least eight species which exist under widely varying conditions.The plumbous ion, Pb2+, is stable in acid solutions and forms complexes with most negativelycharged ions. At Pb concentrations > 10~5 M, polynuclear Pb species play a dominant role; themost prominent polynuclear species are Pb4(OH)4

4+ and Pb6(OH)84+ [1976BAE/MES]. The

stable solid phase in water under most conditions is red PbO(s). The compounds Pb3O4(s) andPbC>2(s) exist in very oxidizing environments. It is known that PbC>2(s) is very insoluble. Pb4+

also hydrolyzes extensively but its hydrolysis products are not well known [1952LAT,1976BAE/MES, 1985GAL, 1995WIB].

Besides complex formation with organic ligands [1976SMI/MAR], thermodynamic data areavailable for the formation of Pb2+ complexes or compounds with the following inorganicligands: hydroxide, chloride, fluoride, nitrate, phosphate, sulfate and sulfide. Equilibriumconstants for the hydrolysis of Pb2+ and the complex formation with chloride, fluoride, nitrate,phosphate, sulfate and sulfide are given in the Sections 6.1, 6.3 - 6.10. Also the solubilityproduct of PbO(red), PbO(yellow), Pb(OH)2(precip), PbCl2(s), and PbOHCl(s), PbF2(s),PbClF(s), cerrusite, hydrocerrusite, Plumbonacrite, PbOHNO3(s), different lead phosphatesolids, anglesite and galena are calculated from experimental data given in literature (Sections6.2-6.10).

The redox reactions of Pb are discussed in Section 6.11 and the hydrolysis of Pb(IV) in Section6.12.

6.1 Hydrolysis of lead

The hydrolysis of Pb(II) is complex. At least eight different Pb(II) complexes are proposed tobe present in aqueous solutions. The most prominent polynuclear species in concentrated leadsolutions (Pb > 10"5 M) are Pb4(OH)4

4+ and Pb6(OH)84+ [1976BAE/MES]. The Pb2+ ion will

dominate lead speciation in acidic solutions; in solutions with pH 12 or higher the anionPb(OH)3" will be the most important species. Experimental data of the hydrolysis of lead usedin this report for extrapolation to I = 0 are given in Table 6.1. Further experimental resultsreported in the literature are shown in Table 6.2 and log [3 values proposed by authors ofprevious reviews are collected in Table 6.3.

Pedersen [1945PED] has published one of the early works investigating the hydrolysis ofPb(II). Unfortunately, [1945PED and 1954FAU] measured lead hydrolysis in nitrate media. Aslead forms complexes with nitrate (see Section 6.6), these data cannot be used directly for theevaluation of the hydrolysis constants of lead. In the 1960s Pb(II) hydrolysis has beendetermined in 0.3 and 3 M perchlorate media by Olin and co-workers [1960OLI1, 1960OLI2,1960CAR/OLI, 1961OLI, and 1962PAJ/OLI] in concentrated (1 - 1500 mM Pb) lead solutions.

126

JNC TN8400 9 9 - 0 1 1

These data have been critically reviewed and extrapolated to I = 0 by [1976BAE/MES]. Severalother authors determined lead hydrolysis in 2-3 M perchlorate media [1964HUG, 1976LEE2,1980KAW/ISH, and 1981KOG/OKA], see Table 6.1. The log p values reported by theseauthors are all quite similar.

Comparison with the data reported by [1965HUG] for 2 M NaNO3 (Table 6.2) shows a stronginfluence of nitrate on the measured log P values for the hydrolysis of Pb(II). Similarly, thedata of [1967SCH/ING] (Table 6.2) illustrate that lead chloride complexes are formed.Consequently, also the data measured by [1973BIL/STU, 1976BIL/HUS, and 1980SYL/BRO]in 0.1 M KNO3 were not used in the present evaluation of formation constants. In the case ofnitrate media, the complex formation constants between Pb2+ and nitrate are well defined (seeSection 6.6), and theoretically the formation constants can be corrected for the effect of thecomplex formation with nitrate. However, a closer examination of the data showed a distinctshift of the log P values for the polynuclear Pb(II) species depending on the nitrateconcentration which indicates that not only Pb2+ but also Pb2OH3+, Pb4(OH)4

4+, Pb3(OH)42+,

Pb3(OH)5+, and Pb6(OH)s4+ interact with nitrate. At low nitrate concentrations, the interactionof Pb(II) with nitrate becomes very small. At NO3- = 0.05 M less than 5% of total Pb(II) arepresent as PbNCV and less than 0.2% are present as Pb(NO3)2°. Other lead nitrate complexeswill be probably even less important. Thus, data measured at NO3" < 0.05 M were included inthe calculations (Table 6.1).

The data used for the calculation of the formation constants of lead(II) hydroxide complexesvalid at I = 0 are listed in Table 6.1 and shown in Figure 6.1 to 6.8. Additional data for thelead(II) hydroxide system are compiled in Table 6.2 and 6.3. These data were not chosen forthe calculation of log (3° values in this report.

Table 6.1: Experimentally determined equilibrium data compiled for the lead hydroxidesystem, according to the equilibrium* mPb2+ + nH2O <=> Pbm(OH)n

2m-n. Thesedata were chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 6.13: Comments onselected references. Method: sol = solubility measurements, tit = titration (pH),and pot = potentiometry.

log Pm

logfr-1.11 1-7.88 i-7.86 !

-7.90-7.80-7.93

Reference

,:Pb2+ + H20<=^Pb0H+

[1945PED][1945PED][1945PED][1960OLI1][1960OLI1][1964HUG]

Comments

+ H+

T= 291 K, 1=0.06T= 291 K, 1=0.03T= 291 K, 1=0.015T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=2

KM,

0.060.03

0.0153

0.32

Medium

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

NaClO4

NaClO4

NaClO4

Method

tittittit

potpotpot

127

JNC TN8400 9 9 - 0 1 1


-8.0 2

-7.32 3

-7.36 3

-7.80

log fax-17.08 4

-17.46-17.18-17.02 2

-17.01 2

log fax-28.04 4

-28.88-27.99-28.103

-29.24 5

log fa,i:

-6.93 i-7.05 i-7.13 i-6.45-6.45-6.30-6.39-6.24-6.49-6.57-6.45-6.30 6

log faA--2025 ]

-20.45 1-20.54 i-19.25-19.90-19.25

[1976LEE2][1978LIN][1978LIN][1993CRU/VAN]

Pb2+ + 2H2OaPb(O

[1939GAR/VEL][1960CAR7OLI][1960CAR/OLI][1978LIN][1978LIN]

T=T=T=

T=

H)2°

T=T=T—

T=T=

Pb2+ + 3H2O 4=>Pb(OH)3-

[1939GAR/VEL][1960CAR/OLI][1960CAR/OLI][1978LIN][1987FER/GRE]

2Pb2+ + H20t=>Pb20.

[1945PED][1945PED][1945PED][1960OLI2][1960OLI2][1962PAJ/OLI][1962PAJ/OLI][1962PAJ/OLI][1962PAJ/OLI][1962PAJ/OLI][1962PAJ/OLI][1962PAJ/OLI]

4Pb2+ + 4H2O&Pb4(

[1945PED][1945PED][1945PED][1960OLI1][1960OLI1][1960OLI2]

•y

T=T=T=T=

T=T=T=T=T=T=T=T=T=T=T=T=

OH)

T=T=T=T=

T=T=

298.15 K, 1=3298.15 K, 1=0.01298.15 K, 1=0.01298.15 K, 1=1

+ 2H+

298.15 K, 1=0-0.1298.15 K, 1=3298.15 K, 1=0.3298.15 K, 1=0.01298.15 K, 1=0.01

+ 3H+

298.15 K, 1=0-0.1298.15 K, 1=3298.15 K, 1=0.3298.15 K, 1=0.01298.15 K, 1=3

+ H+

291 K, 1=0.06291 K, 1=0.03291 K, 1=0.015298.15 K, 1=3.5298.15 K, 1=4298.15 K, 1=4.5298.15 K, 1=4.5298.15 K, 1=4.5298.15 K, 1=4.5298.15 K, 1=4.5298.15 K, 1=4.5298.15 K, 1=4.5

,44+ + 4tf+

291 K, 1=0.06291 K, 1=0.03291 K, 1=0.015298.15 K, 1=3298.15 K, 1=0.3298.15 K, 1=3.5

30.010.01

1

03

0.30.010.01

03

0.30.01

3

0.060.030.015

3.54

4.54.54.54.54.54.54.5

0.060.030.015

30.33.5

NaClO4

NaC104

NaC104

NaClO4

NaOHNaC104

NaC104

NaClO4

NaC104

NaOHNaC104

NaC104

NaC104

NaClO4

Ba(NO,)2

Ba(NO02

Ba(NO,)2

NaClO4/Pb(ClO4)2

NaC104/Pb(C104)2

Ba(C104)2

Ba(C104)2

Ba(C104)2

Mg(C104)2

Mg(C104)2

Mg(C104)2

Pb(C104),

Ba(NO,)2

Ba(NO,)2

Ba(NO3)2

NaClO4

NaC104

NaC104/Pb(C104),

potpolpol

tit

solpotpotpolpol

solpotpotpolsol

tittittit

potpotpotpotpot

potpotpotpot

tit

tittit

potpotpot

128

JNC TN8400 9 9 - 0 1 1

Table 6.1: continued:

-19.23 [1960OLI2] T=-19.16 [1962PAJ/OLI] T=-19.12 [1962PAJ/OLI] T=-19.11 [1962PAJ/OLI] T=-19.12 [1962PAJ/OLI] T=-18.9 [1962PAJ/OLI] T=-18.95 [1962PAJ/OLI] T=-18.98 [1962PAJ/OLI] T=-19.05 [1962PAJ/OLI] T=-19.19 6 [1962PAJ/OLI] T=-19.35 [1964HUG] T=-19.1 2 [1976LEE2] T=-19.42 [1980KAW/ISH] T=-18.90 7 [1981KOG/OKA] T=-19.58 [1993CRU/VAN] T=

298.15 K, 1=4 4 NaC104/Pb(C104)2 pot298.15 K, 1=4.5 4.5 Ba(C104)2 pot298.15 K, 1=4.5 4.5 Ba(ClO4)2 pot298.15 K, 1=4.5 4.5 Ba(C104)2 pot298.15 K, 1=4.5 4.5 Ba(ClO4)2 pot298.15 K, 1=4.5 4.5 Mg(ClO4)2 pot298.15 K, 1=4.5 4.5 Mg(C104)2 pot298.15 K, 1=4.5 4.5 Mg(C104)2 pot298.15 K, 1=4.5 4.5 Mg(C104)2 pot298.15 K, 1=4.5 4.5 Pb(C104)2 pot298.15 K, 1=2 2 NaC104 pot298.15 K, 1=3 3 NaClO4 pot298.15 K, 1=3 3 LiC104 pot298.15 K, 1=3 3 LiC104 pot298.15 K, 1=1 1 NaClO4 tit

log foA:-22.87-23.35-22.6 2

-22.78-23.03 7

-22.69

log p3,5-'

-31.62 8-32.27-30.80

log p6,8-'

-42.14-42.66-42.1 2-42.33-41.68 7

-42.43

3Pb2+ + 4H2O^Pb3(

[1960OLI1][1960OLI1][1976LEE2][1980KAW/ISH][1981KOG/OKA][1993CRU/VAN]

OH)42++ 4H+

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=1

3Pb2+ + 5H2O <=> Pb3(OH)5+ + 5H+ .

[ 1980S YL/BRO][ 1980S YL/BRO][1993CRU/VAN]

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=1

6Pb2+ + 8H2O <^Pb6(OH)84+ + 8H+

[1960OLI1][1960OLI1][1976LEE2][1980KAW/ISH][1981KOG/OKA][1993CRU/VAN]

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=1

30.33331

30.3

1

30.33331

NaClO4

NaClO4

NaC104

LiC104

LiClO4

NaC104

NaC104

NaClO4

NaC104

NaC104

NaClO4

NaC104

LiC104

LiC104

NaClO4

potpotpotpotpottit

potpottit

potpotpotpotpottit

corrected in this report for the effect of the formation of PbNO3+ and Pb(NO3)2° complexes using the constants

derived in Section 6.7.1. Uncorrected values are given in Table 6.2.Pb concentration = 0.3-1.2 mM.calculated with a pKw of 13.91 at 1=0.01.extrapolated to 1 = 0 with SIT in this report, see Section 6.13: Comments on selected references.calculated with a pKw of 14.18 at 1=3.recalculation of data of [1960OLI2].Pb concentration = 5-40 mMrecalculation of data of [ 1960OLI1 ].

129

JNC TN8400 9 9 - 0 1 1

6.1.1 PbOH+

In diluted solutions, Pb2+ hydrolyzes to PbOH+, and higher-hydrolyzed species (see alsoSections 6.1.2, 6.1.3, and 6.1.4). To calculate the formation constant log p \ i at I = 0, themeasurements of Olin and co-workers and the data measured in perchlorate medium by[1964HUG, 1976LEE2, 1978LIN, and 1993CRU/VAN] (see Table 6.1) were used. The datameasured in nitrate media by [1945PED] in 0.01, 0.02 and 0.04 M nitrate media were correctedfor the formation of lead nitrate complexes using the constants calculated in Section 6.7.1 (cf.Table 6.1). Extrapolation of these measurements to I = 0 gave the following formation constantlog P ° u for PbOH+, as illustrated in Figure 6.1.

Pb2+ +H2O PbOH+ + H+ log p° u =-7 .51

Pb2++H,0 «=> PbOH++H+

QCM

+CO.

-b -I

-5.5

-6

-6.5 •

- 7 •

t-7.5 !

-8.5

-9

-9.5 --m -

C

-jjlllllfli

)

i - ••• : . ^ . ; j . . ; I , , . -U^ . i . : . . : ,.j:,,:,L—•;.= ;.i>|= , , ; . . . . , ..:..-.,•... , , ^.-ji:^;;,;;.:;;; : , ;> • . j

•••>".;--t i::;jjjiij;-*'i!?l'=viîiO'Of1'"'1''^' ^ 7 ^ ' C * 1 ' ' " - • ' ^^ ijî.fifii'ri":';;":

2 4 C

L, molal

Figure 6.1: Plot of log (3U + 2D vs. Im for the reaction Pb2+ + H2O <=> PbOH+ + H+ at 25°C. The straight line shows the result of the linear regression: Ae = - 0.02; logP°ij = - 7.51. Calculated from data compiled in Table 6.1.

The experimental values given by [1980KAW/ISH] and [1989DOR/MAR] (see Table 6.2) werenot used. [1980KAW/ISH] themselves classified their value as doubtful due to the low PbOH+concentration in their experiments. [1989DOR/MAR] did not report experimental details, suchas concentration and type of electrolyte present, and no Pb hydroxide species other than the 1:1complex was included in their calculations.

130

JNC TN8400 9 9 - 0 1 1

6.1.2 Pb(OH)2°

Measurements of lead hydrolysis at higher pH values were carried out by [1939GAR/VEL] and[1960CAR/OLI]. [1939GAR/VEL] measured Pb hydrolysis in 0.002 - 1.2 M NaOH. From themeasurements of [1939GAR/VEL] a log p \ 2 value at I = 0 of -17.08 for the reaction Pb2+ +2H2O <=> Pb(OH)2° + 2H+ could be calculated in this report (experimental data of[1939GAR/VEL], extrapolation to 1=0 and conversion to log p \ 2 value see [1939GAR/VEL] inSection 6.13: Comments on selected references). The extrapolation of the data of[1939GAR/VEL, 1960CAR/OLI and 1978LIN] gives (see also Figure 6.2):

Pb2+ +2H2O Pb(OH)2° + 2H+ log ( 3 \ 2 = -16.95

Pb2++2H2O <=> Pb(OH)2°+2H+

Q

»-

D)O

-15.5 -

-16 -

-16.5 -

-17.5 -

-18

- 1 8 . 5 •

-19

-19.5

-20

•;.;': Ji'';::"^--iJ:"i?; : \ ! T::':';':: "" !;-;i;:i:;iiJi;-;:i:;;;";i-.iiV:;i;i-:i-=J:

- '.'•'• -Vvl'' ?.\>"-\}'~z~Js'..l'.? il'^W

•••••. '::i£^^WMB^k

=! = . - : : - - : . r . . : ' \ ' . ; . V - ^ . ; : . ; : . ~ ~ ' : . . ^ : : - v . : : . : ' > . . i ! . V - . : • ' - • '•. , . : ; / " - - . - V " • . - • • - • ;

lilS;iillti;ilii?:s iSIf ;il" S:iiIiiii:i:iliiiiiiil®^l:'S?-#ffia«i:'

•:::;;ii;^;:!^::;n::^ :;;{:y;^:^::;i>;^:s^s;i^,:::;::;::^^:::^:) r;';;; p;; \1 v ; ; .

2 - 4

lm, molal

Figure 6.2: Plot of log pi,2 + 2D vs. Im for the reaction Pb2+ + 2H2O <=> Pb(OH)2° + 2H+ at25 °C. The straight line shows the result of the linear regression: Ae = -0.0, log(3°i2 = - 16.95. Calculated from data compiled in Table 6.1.

Again, the data measured by [1973BHVSTU] and [1976BIL/HUS] in 0.1 M KN03 and by[1967SCH/ING] in 3 M NaCl were not chosen for the extrapolation to I = 0. Also the valuereported by [1980KAW/ISH] was not used as the authors themselves classified this value asdoubtful due to the low Pb(OH)2 concentration in their experiments.

131

JNC TN8400 9 9 - 0 1 1

6.1.3 Pb(OH)3-

[1939GAR/VEL] measured Pb hydrolysis in dilute, alkaline solutions. Extrapolation of thesemeasurements to 1=0 with the SIT model gave a log p ] 3 value of -28.04 for the reaction Pb2+

+ 3H2O <=> Pb(OH)3- + 3H+ (experimental data of [1939GAR/VEL], extrapolation to I = 0 andconversion to log (313 value see: [1939GAR/VEL] in Section 6.13: Comments on selectedreferences). Further data were determined by [1960CAR/OLI, 1978LIN and 1987FER/GRE] insodium perchlorate media (see Table 6.1). The extrapolation of the data of [1939GAR/VEL,1960CAR/OLI, 1978LIN and 1987FER/GRE] to I = 0 gives:

Pb2+ +3H2O Pb(OH)3- log = - 28.02

The data measured by [1967SCH/ING] in 3 M NaCl and by [1980WAL/SIN] in wastewater(given in Table 6.2) were not chosen for the extrapolation to I = 0.

Pb2++3H2O « Pb(OH)3-+3H+

Qo+CO

CO.

-<io -

-25.5

-26

-26.5

-27

-27.5 •

-28 (

-28.5

-29

-29.5 -

-sn -

llilII|B.,••••;. ' > , • . • , " * " • • £ Ji-iMVil-Bi-,

1!-;:.::-.-;

1'!';;^.-:.;..!!;,

1.

; ; : - i ;r' ;«:£;£5ir;;?=:;=L?; ' ' V . ^ A ^ - ' ; r i s ^ ii:ii;

pî^lil[îi;!biHi;iîi—^"••-il;; i:v;:-:î:JI'î(ii*

-;.:::•; i./;^:H^";'v;:?;i::;ii?*î^^;-^;,::

•:.;•••/. *..'T~,1T:.'.?;•"~V' :>'• •«".*V:i!""'iV.'=H"':="4S"•:jf:"7-^

;!-i"i:-Vin.-f'-i!'iii-i;"i::":i-^-'; :''-.- ::'i

"L'i.iiiiiiriSn'lisli!!^: '•,'?:'?'ll'.rj^'.-';

iiliiiiil•||S||jK-J: . . " . 1 : ' L ' - ^ , n : i ; J ^ : : : ' : - ' J : ' ^ ' : " - • " ' • : ' - - •::.;. •

iliSilliililS;|:i!i;r:'>-™-5'-ii;:-'ii"i!'-":':ii"'"i=i:=i";-:;

"••f-'i, ;'~;i:^-:::?i^::i-:;"":Vv.-^

;i:;

:!:;

liijllllll•:l:-!;';:v'l~", \>--, ^T:[•".?-!" : : i p ^ S -

. . . . ^ ; : , : . . ; ; : . : : • : ^ . . ^ • _ • .

iylllillll;.

Illlt=::::ii!J".:'::i:;i!;'::':i-;:î:-:ii1!::.

Sli'ilil.pii-':'

" :: •; ' y ^ ' " ; ., : L " . • - : " l l -

2 4

lm, molal

Figure 6.3: Plot of log (3i>3 + 0D vs. Im for the reaction Pb2+ + 3H2O <> Pb(OH)3" + 3H+ at25 °C. The straight line shows the result of the linear regression: Ae = 0.29, logP°i)3 = - 28.02. Calculated from data compiled in Table 6.1.

6.1.4 Pb(OH)42-

[1976HEM] and [1987BROAVAN] proposed a log pli4of-39.7 and -37.2 for the formation ofthe Pb(OH)42" species (compiled in Table 6.3). Direct experimental data, however, are notavailable and the experiments of [1939GAR/VEL], [1960CAR7OLI] and [1987FER/GRE]show that below a pH of 14, the species Pb(OH)3" dominates lead speciation under alkalineconditions.

132

JNC TN8400 9 9 - 0 1 1

6.1.5 Pb2OH3+

In solutions containing more than approximately 10"5 M Pb, polynuclear lead complexes areformed. Around pH 6 to 7, the Pb2OH3+ species can occur in very concentrated lead solutions(Pb > 0.1 M). Evaluation of a formation constant for Pb2OH3+ is somewhat difficult, as thePb2OH3+ species exists only in concentrated lead solutions (Pb > 0.1 M) in detectable amounts.[1980SYL/BRO] considered their data determined in presence of 0.1 - 2 mM Pb for thePb2OH3+ complex as unreliable, due to the low Pb2OH3+ concentration in their experiments.Reliable data are measured by Olin and co-workers [1960OLI2, 1962PAJ/OLI] in I = 3 - 4.5perchlorate media (Table 6.1). [1945PED] determined log $2,\ values at different ionic strengthin nitrate media. The data measured in 0.01, 0.02 and 0.04 M nitrate media were corrected forthe formation of lead nitrate complexes using the constants calculated in Section 6.7.1 (cf. Table6.1). Extrapolation to I = 0 is shown in Figure 6.4 and results in:

2Pb2+ + H2O <=> Pb2OH3+ + H+ log p°2,i =-7.18

Extrapolation of the data measured by [1945PED] in nitrate medium gives a similar result (seeTable 6.2), as at low nitrate concentrations, complex formation of Pb(II) with nitrate becomesnegligible.

j3+ , u+

-2D

CM"EQ.

D)

-5

-5.5 -

-6 -

-6.5

-7 (-7.5

-8 •

-8.5 -

-9

-9.5

-m -

*•?•"• :•••':] : i:

: ^1

^ r:

Ul: F ^ i

; :^ . H ^ =

1' ; ^ r : ^ r i - ^ r ^ - y ' - i \

1; ^ i

:^ > ^ ' ^ i - ^ i

8! ^

1: ^ ^ ; ' . ^ ! ! * ! : ; - ^ !

1:

1! " ;:•:••!':<>;"/• :'":!:

1;: -:r

:: '•'••

; ^ : i . i. i'-: ; - - :" : ! :" ; i , ' !"" ' 'i;-i :'.\7:', '*.,'' '.'• ''•:' '[ l-_\ '\ '= : ' • : : : .^ '.• '. ' i ; ";=!!i\ ',' }'• '• '*•'' \'-: • ••','. V i:-1-:^1-'^^ :! :? ; '.!:.- ..":•":"

'" '.'-- i v .'•'." : ' , " - ' - ' ' , r ' " . - ' " - r - ' r ' \ '"'" • ' . : - ' ' ' - - ' " 1 '••'•••'• • • ; : " > " : : : . : ' ' • . " • ' > ' : • ' . ; : : . . - = ! :; , ' • :". •; • : : . ' i " : : ; :. . • . • : . : - . " - • . • . ; '

'•''-. ••. •• ' : '•:-'"•'•'. : : ? • ' . ' • • : = . ' . : . , ' • , - : ; ; - : ; C . . ; i : - : ! = ^ : - ' = : ' " : - - 1 - - : - : : ' / l - - ' - ' - • - £ , " • - ' . ' : ' • ' • : . • : ! ' : ' - . L - ; : • • : : - - v ^ • • ' • J i : : : . . ' = ' - • : ' . "

• _ : . • • • • • " . " ; : : : : - ; ; : . I . . : 1 " ; . ' . : ' ; : • i ^ ; l . : : i • . • • - • . • | • ^ ; ~ " " . : ; : ^ • : " . - : • • • : ] r - ' " " • . • • \ y . \ r , •.-.••..•• y I L ! " - . 1 " 1 " . : . ; • • ; • : • • = - .

0 2 4 6

lm, molal

Figure 6.4: Plot of log P2,i - 2D vs. Im for the reaction 2Pb2+ + H2O <> Pb2OH3+ + H+ at 25°C. The straight line shows the result of the linear regression: Ae = - 0.04; logP°2,i = - 7.18. Calculated from data compiled in Table 6.1.

133

J N C T N 8 4 0 0 99 - O i l

6.1.6 Pb4(OH)44+

Many independent measurements of the formation constants for Pb4(OH)44+ exist [1945PED,1960OLI2, 1962PAJ/OLI, 1964HUG, 1976LEE2, 1980KAW/ISH, 1981KOG/OKA,1991CRU/VAN] (cf. Table 6.1). The data measured in 0.01,0.02 and 0.04 M nitrate media by[1945PED] were corrected for the formation of lead nitrate complexes (cf. Table 6.1).Extrapolation of these measurements to 1=0 (see also Figure 6.5) gives:

4Pb2+ + 4H2O <=> Pb4(OH)44+ + 4H+ log (3°4,4 = - 20.63

4Pb2++4H2O <=> Pb4(OH)44++4H+

Q

1

so.D)O

- 10

-18.5 i

-19 -

-19.5 -

-20 -

-20.5 jV

-21

-21.5

-22

-22.5 -

jo Q ° ^

: 1

-

- , . • .

o a

-:—S^"""°—^

7x : 20.63

2 4

L, molal

Figure 6.5: Plot of log p4;4 - 4D vs. Im for the reaction 4Pb2+ + 4H2O <=> Pb4(OH)44+ + 4H+

at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.07;log P°4>4 = - 20.63. Calculated from data compiled in Table 6.1.

134

JNC TN8400 9 9 - 0 1 1

6.1.7 Pb3(OH)42+

The formation constant for Pb3(OH)42+ was calculated based on the measurements by

[1960OLI, 1976LEE2, 1980KAW/ISH, 1981KOG/OKA, 1993CRU/VAN] as given in Table6.1. Extrapolation to 1=0 gives (Figure 6.6):

3Pb2+ + 4H2O Pb3(OH)42+ + 4H+ log (3°3)4 = - 22.48

3Pb2++4H2O <=> Pb3(OH)42++4H+

Q

aL

-20.5 •• I;

- 2 1 ••

-21.5 - - i ;

-22 -;|;

-22.5 -M-23 -W

-23.5 -§,

-24 -§ |

-24.5 --I

.:•:::.. ?LZ> ,; ^ l ; , ; ; \A : .^ ; . : O i ' T b X - — ^ ^ ] i 4 o : ; i - ••• :^ : / ; : J '• ;••• ^ :

;il;ft6iillli|gl8lil|IS|:itil^li5|K^V:;;;:; \Sy •

,fl3ia::i;.;:;=5:-:3:ffljil':.i: i±*JI Ji:yH;3!ft£!p:;Si:-.;-S#i3r.!i:':.£:•:%r&3=l', . '. '

: ' ' " ' ''}~::':":-1 ' " B ^ t i f :g :° = : j ' : : ; g N ' i i M ^ - : ; ; -•••alii j j ' w -•• • <•• ;?•«•;-:> -: " ^ ; : •• •

0 2 4 6

lm, molal

Figure 6.6: Plot of log p3,4 + 4D vs. Im for the reaction 3Pb2+ + 4H2O & Pb3(OH)42+ + 4H+

at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.16;log P°3,4= - 22.48. Calculated from data compiled in Table 6.1.

135

JNC TN8400 9 9 - 0 1 1

6.1.8 Pb3(OH)33+ and Pb3(OH)5+

[1980SYL/BRO and 1980KAW/ISH] proposed the presence of additional polymeric species:Pb3(OH)3

3+ and Pb3(OH)5+. [1980SYL/BRO] showed that the consideration of the additionalpresence of Pb3(OH)5+ had hardly an influence on the other log |3 values, as recalculated fromthe experiments of Olin and co-workers by [1980SYL/BRO] (see Table 6.1 and 6.2). Theextrapolation of the formation constant of Pb3(OH)5+ based on the values calculated by[1980SYL/BRO] for 1=0.3 and 3 M NaC104 and the values determined by [1993CRU/VAN]gives (Figure 6.7):

3Pb2+ +

QCD

CO

CO.CD

5H2O

-28 -

-28.5 •

-29 -

-29.5 •

-30-

-30.5 -

-31 •

-31.5

-32 -

-32.5 -

-33 -

o Pb3(OH)5+ + 5H+ log P°

3Pb2 ++5H2O <=> Pb3 (OH)5++5H+

^::'jJ:\ij$BMi:!il!fi:^

.' l,~i-'^ ;:'.'-. ;J::.:--L.--:=:'-: . i i ' v ' : :•= : . .: ' y V 1 ' ; - j ^ ' . ! / S K ! i H i ' f j : w -'j .•• ' Q / ^ ' ^ n . - > ; • ' : :B: I:: ',;| '

1::-;J:• : ,:.;;1i..::-,-f::i::.^|!.::;:;::-,-i.1.: :"^;;';". . V ; • " ^ • . \ / ^ / 1 i C / . / V ^ ' ^ V O V / • / £f''-'--'- ••••-1'-"- '' !- -: •••.:••• v\-:t'i;.i:!^::;;]^:.^:;i:::?,.iij:;1;>.;.i^:r::-:»-:;'""^"•/'."•:•<!.':i-;;7;.!:;i!"!---::-^ " • = - i :" ; i>. ; i : ' i : ' " . . • . - :•.••: •••

':'.'BMiM{g§§MMS^

3,5 = - 30.72

n, moral

Figure 6.7: Plot of log p \ 5 + 6D vs. Im for the reaction 3Pb2+ + 5H2O <=> Pb3(OH)5+ + 5H+at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.19,log P°3i5 = - 30.72. Calculated from data compiled in Table 6.1.

The only experimentally determined formation constant for Pb3(OH)33+ has been determined by[1980KAW/ISH] in 3 M LiClO4. This value is not extrapolated to 1=0 and no log (3 value forthis complex is recommended in this report. In any case, such a complex is not important insolutions containing less than 1 mM Pb.

136

J N C T N 8 4 0 0 99 - O i l

6.1.9 Pb6O(OH)64+

Several independent measurements exist for the formation constant of PbeO(OH)64+, log p6,8([1960OLI, 1976LEE2, 1980KAW/ISH, 1981KOG/OKA, 1993CRU/VAN]; Table 6.1).Extrapolation to 1=0 gives:

6Pb2+ + 8H2O Pb6(OH)84+ + 8H+ log {3°6i8 = - 42.68

[1981ISH/OHT] proposed, based on calorimetric measurements, that the complex in fact isPb6O(OH)64+, which corresponds to X-ray diffraction data indicating a central oxygen atomsurrounded by four lead atom [1968JOH/OLI].

6Pb2++8H2O «• Pb6(OH)84++8H+

-40

-40.5 -

-41

Q -41.5 ••O+ -42 +

CO.

O

-42.5 ;

-43 -

-43.5

-44

-44.5 -f

-45

y = 0.08x - 42.68

0 2 4 6

lm, molal

Figure 6.8: Plot of log |36,8 + 0D vs. Im for the reaction 6Pb2+ + 8H2O <=> Pb6(OH)84+ + 8H+

at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.08,log p°6,8 = - 42.68. Calculated from data compiled in Table 6.1.

137

JNC TN8400 9 9 - 0 1 1

6.1.10 Additional equilibrium data compiled for the lead hydroxide system

Table 6.2: Additional experimentally determined equilibrium data compiled for the lead hydroxide system,according to the equilibrium: mPb2+ + nH2O <=> Pbm(OH)n

2"n + nH+. These data were notchosen in the present report for the evaluation of recommended stability values. Other reasons fornot selecting these references are given at the end of this table or in Section 6.13: 'Comments onselected references'. As pointed out in Section 6.1, studies carried out in nitrate, chloride andsulfate media were not chosen as these ions form complexes with Pb(II). Method: pot =potentiometry, sol = solubility measurements, tit = pH titration.

log Reference Comments Medium Method

log pu: Pb2+ + H2O <=> PbOH* + H*

-6.18 '-8.74 2

-8.35 2

-8.18 2

-8.07 2

-7.85 2 '3

-7.94 2 '3

-7.89 2-3

-7.73 5

-8.66 5

-8.37 5

-8.84 6

-6.96-6.96-8.31 7

-7.86-8.17-6.84-8.72-7.94 10

[1939GAR/VEL][1945PED][1945PED][1945PED][1945PED][1945PED][1945PED][1945PED][1945PED][1954FAU][1954FAU][1965HUG1][1973BDL/STU][1976BEL/HUS][1980KAW/ISH][1980SYL/BRO][1980WAL/SIN][1989DOR/MAR][1993CRU/VAN][1993CRU/VAN1

T= 298.15 K, 1=0-0.1T=291 K, 1=1.2T=291 K, 1=0.6T= 291 K, 1=0.3T= 291 K, 1=0.15T=291 K, 1=0.06T=291 K, 1=0.03T= 291 K, 1=0.015T=291 K, 1=0.0T= 293 K, 1=0.6T= 293 K, 1=0.06T= 298.15 K, 1=2T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=3T= 298.15 K, 1=0.1T= 298 K, I=dilT=298.15K,I=dilT= 298.15 K, 1=1T= 298.15 K, 1=1

1.20.60.3

0.150.060.030.015

00.6

0.062

0.10.1

30.1

11

NaOHBa(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

NaNO3

KNO3

KNO3LiC104

KNO3

H2SO4

Pb(NO3)2, NaHCOjKNO3

KNO,

soltittittittittittittittittittit

potpotpotpotpotsoltittittit

log 2H2O <=> Pb(OH)2° + 2H+

-20.58 "-16.42-16.42-16.37 7

-17.85

[1967SCH/ING][1973BDL/STU][1976BIL/HUS][1980KAW/ISH][1980WAL/SIN1

T= 298.15 K, 1=3T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=3T= 298 K, I=dil

30.10.13

NaClKNO3

KNO3

LiC104

H,SO4

potpotpotpotsol

log [)u: Pb2+ + 3H2O <=> Pb(OH)/ + 3H+

-32.58 n

-28.3

log p2J: 2Pb

-7A22

-7.05 2

-7.05 2

-7.05 2

[1967SCH/ING][198OWAL/SIN1

2+ + H20 <=> Pb2OH3* -

[1945PED][1945PED][1945PED][1945PED]

T= 298.15 K, 1=3T= 298 K, I=dil

T=291 K, 1=1.2T=291 K, 1=0.6T=291 K, 1=0.3T=291 K, 1=0.15

3

1.20.60.3

0.15

NaClH,SOd

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

potsol

tittittittit

138

JNC TN8400 9 9 - 0 1 1


-7.10-7.15 2

-7.19 2--7.28 4

-6.34 8

-7.11 6

-6.26-6.79

2-3[1945PED][1945PED][1945PED][1945PED][1960OLI2][1965HUG1][ 1980S YL/BRO][1993CRU/VAN1

T=291 K, 1=0.06T=291 K, 1=0.03T=291 K, 1=0.015T= 291 K, 1=0.0T= 298.15 K, 1=4.5T= 298.15 K, 1=2T= 298.15 K, 1=0.1T= 298.15 K, 1=1

log fty 3Pb2+ + 3H2O <=> Pb3(OH)33+ + 3H+

-8.31 7 [1980KAW/ISH] T= 298.15 K, 1=3

0.060.030.015

04.52

0.11

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Pb(C104)2

NaNO3

KNO3

KNO,

tittittittitpotpotpottit

log fty 4P&+

-20.97 2

-20.80 2

-20.65 2

-20.64 2

-20.59 2-3

-20.66 2-3

-20.66 2.3

-20.96 4

-18.75 s

-18.05 5

-19.28 8

-21.72 6

-20.4 •

-21.01-19.01 10

+ 4H2O <=> Pb4(OH)44+

[1945PED][1945PED][1945PED][1945PED][1945PED][1945PED][1945PED][1945PED][1954FAU][1954FAU][1960OLI2][1965HUG1][1980S YL/BRO][1993CRU/VAN][1993CRU/VAN1

+ 4H+

T=291T=291T=291T=291T=291T=291T=291T=291T=293T=293T=298T=298T=298T='298T=298

K, 1=1.2K, 1=0.6K, 1=0.3K, 1=0.15K, 1=0.06K, 1=0.03K, 1=0.015K, 1=0.0K, 1=0.6K, 1=0.06.15 K, 1=4.5.15 K, 1=2.15 K, 1=0.1.15 K, 1=1.15 K, 1=1

1.20.60.3

0.150.060.030.015

0.60.064.52

0.111

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Ba(NO3)2

Pb(ClO4)2

NaNO3

KNO3

KNO3

KNO,

tittittittittittittittittittitpotpotpottittit

LiCIO, pot

log 3Pb2+

-23.91-24.33-22.83 10

4H2O Pb3(OH)42+ + 4H+

[1980SYL/BRO][1993CRU/VAN]ri993CRU/VAN]

T= 298.15 K, 1=0.1T= 298.15 K, 1=1T= 298.15 K, 1=1

0.111

KNO3

KNO3

KNO,

pottittit

log P3,5: 3Pb2+ + 5H2O <=> Pb3(OH)5+ + 5FT

-31.75 f!980SYL/BRO1 T= 298.15 K, 1=0.1 0.1 KNO, pot

log p6S: 6Pb2+ + 8H2O <=> Pb6(OH)84+ + 8H+

-43.38-44.91-41.55 10

[1980SYL/BRO][1993CRU/VAN]ri993CRUAÂNl

T= 298.15 K, 1=0.1T= 298.15 K, 1=1T= 298.15 K, 1=1

0.111

KNO3

KNO3

KNO,

pottittit

91011

polynuclear species not considered, I not constant, calculated with a log K*so of -12.68 for PbO (Section 6.2.1)formation of PbNO3' complexes, only small concentration of Pb2OH3* present in his experiments. Pb concentration= 5-400 mMvalues corrected for the effect of the formation of Pb-nitrate complexes are given in Table 6.1.after extrapolation to 1=0 with SIT (this report)[1954FAU] assumed Pb4(OH)4

4+ to be only species present at Pb > 10 mM. Pb concentration = 0.25-200 mM.Pb concentration =10-200 mM, formation of lead nitrate complexes not considered[1980KAW/ISH] themselves classified their value as doubtful due to the low PbOH+ and Pb(OH)2 concentration intheir experiments.corrected value in [1962PAJ/OLI]see Section 6.1.8values corrected for the effect of the formation of Pb-nitrate complexes by [1993CRU/VAN]. Pb = 1-50 mM.Pb concentration =10-500 mM, formation of lead chloride complexes not considered; pKw used = 14.18

139

JNC TN8400 9 9 - 0 1 1

Table 6.3: Thermodynamic data for the hydrolysis of lead taken from previous compilations. As pointed outin Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison. Medium: Where data refer to specific electrolytesolutions, this is indicated.

log Pm ,n

log (3U: Pb2+

-7.90 '-7.80 •-7.81 2

-7.84 2

-7.71-6.57-7.70-7.8-7.9-7.22-7.89 2

-7.79 2

-7.71-6.17-7.71-7.70-7.8-7.9-6.18-7.71-7.70-6.17-7.71-8.00-7.64-7.60-7.8-7.8-8.0-7.9

log $l2: Pb2+

-17.46 '-17.18 '-17.1-17.12-15.82-17.1-17.2-17.5-16.91-17.12-17.12-17.1

Reference

+ H2O <=> PbOH* + H+

[1961OLI][1961OLI][1974VAD][1974 V AD][1976BAE/MES][1976HEM][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980SCH][1980SYL/BRO][1980SYL/BRO][1981BAEMES][1981STU/MOR][1981TUR/WHI][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1983LAN][1984TAY/LOP][1985BAB/MAT][1987BROAVAN][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR1

+ 2H2O <=> Pb(OH)2° +

[1961OLI][1961OLI][1963FEI/SCH][1976BAE/MES][1976HEM][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980SCH][1981BAE/MES][1981TUR/WHI][1982SMI/MAR]

Comments

T=298.15K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, I=n/a.T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T = 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T=291 K, 1=0.7T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=2T= 298.15 K, 1=3

2H*

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/a.T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T = 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

KM)

30.3330

00.3303

0.300

- 00

0.3300

00

0.700

0.10.323

30.300

00.330000

Medium

NaClO4

NaC104

NaC104

NaC104

NaClO4

NaClO4

NaClO4

NaC104

140

JNC TN8400 9 9 - 0 1 1


-17.2-17.5-17.12-17.75-17.16-16.55-17.41-17.06-17.1-17.2-17.5

[1982SMI/MAR][1982SMI/MAR][1983LAN][1984TAY/LOP][1985BAB/MAT][1987BROAVAN][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR]

T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T= 291 K, 1=0.7T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3

0.330

00

0.700

0.33

sw

log Pu: Pb2+ + 3H2O <=> Pb(OH);

-27.98-28.89 '-27.99 '-28.1-28.06-28.08-28.1-28.0-28.8-54.20-28.02-28.07-28.07-28.06-28.05-28.06-28.1-28.0-28.8-28.08-28.07-28.09-28.07-27.98-26.39-28.01-28.1-28.0-28.8

[1952LAT][1961OLI][1961OLI][1963FEI/SCH][1976BAE/MES][1976HEM][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1977PAU][1979PAT/OBR][1980SCH][1980BEN/TEA][1981BAE/MES][1981STU/MOR][1981TUR/WHI][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1983LAN][1984TAY/LOP][1985BAB/MAT][1985GAL][1987BROAVAN][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR1

T= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, I=n/a.T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=0T = 298.15 K, 1=0T = 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3"T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3

0.0153

0.300

00.33

0000000

0.3300

00000

0.33

NaC104

NaC104

log P,4: Pb2+ + 4H2O <=> Pb(OH)/- + 4H+

-39.703

-37.19[1976HEM][1987BRO/WAN]

T= 298.15 K, I=n/aT= 298.15 K, 1=0

141

JNC TN8400 9 9 - 0 1 1


log P21:

-6.40 '-6.36-6.40-6.30-6.36-6.40-6.30-6.36-6.40-6.35-6.34-6.40-6.40

2Pb2+ H2O Pb2OH3+ + H*

[1961OLI][1976BAE/MES][1976SMI/MAR][1976SMI/MAR][1980SCH][1982SMI/MAR][1982SMI/MAR][1983LAN][1984TAY/LOP][1985BAB/MAT][1987BRO/WAN][1989SMI/MAR][1989SMI/MAR]

T= 298.15 K, 1=4.5T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3T = 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3

4.50030030

0003

Pb(C104)2

log P4,4: 4Pb2+

-19.25 '-19.90 '-19.27 2

-19.26 2

-20.88-20.9-19.9-19.2-20.87-19.26 2

-19.94 2

-19.27-20.9-19.9-19.1-19.27-20.88-20.89-21.02-18.98-20-19.6-19.9-19.8-19.3-19.2

+ 4H2O <=> Pb4(OH)/

[1961OLI][1961OLI][1974VAD][1974V AD][1976BAE/MES][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980SCH][1980SYL/BRO][1980SYL/BRO][1980BEN/TEA][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1983LAN][1984TAY/LOP][1987BROAVAN][1988PH17HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR1

+ + 4H+

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T = 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=0.5T= 298.15 K, 1=2T= 298.15 K, 1=3

30.33300

0.3303

0.300

0.3300

000

0.10.30.523

NaC104

NaC104

NaC104

NaCIO4

NaC104

NaC104

log ft,: 3Pb2+

-22.87 >-23.35 '-22.90 2

-22.92 2

-23.88-23.9-23.3

4H2O Pb3(OH)47+ + 4H+

[1961OLI][1961OLI][1974 V AD][1974V AD][1976BAE/MES][1976SMI/MAR][1976SMI/MAR]

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K,I=3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3

30.33300

0.3

NaC104

NaC104

NaC104

NaCIO,

142

JNC TN8400 9 9 - 0 1 1


-22.9-23.86-23.35-23.01 2

-23.18 2

-23.9-23.3-22.7-23.34-23.88-23.89-23.77-23.65-23.9-23.3-22.7

[1976SMI/MAR][1980SCH][198OBEN/TEA][ 1980S YL/BRO][ 1980S YL/BRO][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1983LAN][1984TAY/LOP][1987BRO/WAN][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR]

T= 298.15 K, 1=3T = 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3

3003

0.30

0.3300

000

0.33

NaC104

NaC104

log p3J: 3Pb2+ + 5H2O <=> Pb3(OH)5+ + 5H+

-30.29-7.86

log PM: 6Pb2+

-42.14-42.66-42.13 2

-42.12 2

-43.61-43.6-42.7-42.1-43.61-42.13 2

-42.76 2

-42.66-43.6-42.7-41.9-42.66-43.61-43.59-43.6-43.35-43.6-42.7-42.1

[1987BRO/WAN]|"1988PHI/HAL1

+ 8H2O <=> Pb6(OH)84+

[1961OLI][1961OLI][1974 V AD][1974VAD][1976BAE/MES][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980SCH][1980SYL/BRO][1980SYL/BRO][1980BEN/TEA][1982SMI/MAR][1982SMI/MAR)[1982SMI/MAR][1982WAG/EVA][1983LAN][1984TAY/LOP][1987BROAVAN][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR]

T= 298.15 K, 1=0T= 298.15 K, 1=0

+ 8H+

T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3-T= 298.15 K, 1=3T = 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0.3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.3T= 298.15 K, 1=3

00

30.33300

0.3303

0.300

0.33000000

0.33

NaC104

NaC104

NaC104

NaC104

NaC104

NaC104

1 same values as given in [1960CAR/OL1], [1960OLI1], or [1960OL12].2 recalculation of the data of [1960OLI1].3 cited by J1976HEM] from a reference in Russian.

143

JNC TN8400 9 9 - 0 1 1

6.2 Solid lead-oxide/hydroxide phases

PbO(s) occurs in two crystalline forms: litharge (red), the stable form under ambient conditions,and massicot (yellow). The existence of a hydrated form, freshly precipitated form is reportedby [1928RAN/SPE].

6.2.1 PbO(litharge) and PbO(massicot)

The solubility of litharge and massicot has been determined by [1928RAN/SPE] and[1939GAR/VEL], who measured lead solubility in alkaline medium (see Table 6.5). Bothshowed that in alkaline solutions, the red PbO is more stable than the yellow form. Thesolubility constants for litharge and massicot can more accurately be determined from thepotentiometric data given by [1922APP/REI] and [1923SMI/WOO] (Table 6.4). From theirdata:

Pb2+ + H2O <=> PbO(s, red) + 2H+ log K*°so= - 12.68

Pb2+ + H2O <=> PbO(s, yellow) + 2H+ log K*°so = - 12.96

These values also are in good agreement with the solubility measurements listed in Table 6.5 formassicot and litharge and with the log K*°so of -12.65 and -12.71 (Table 6.6) proposed forlitharge (red PbO(cr)) by [1979PAT/OBR] and [1984TAY/LOP]. While litharge isthermodynamically favored at room temperature, at temperature above 488 °C the yellowmassicot is more stable. However, the precipitation of massicot can also been observed at roomtemperature [1984TAY/LOP, 1995WD3] as the conversion rate between the two phases issmall.

Table 6.4: Experimentally determined equilibrium data compiled for the for the precipitationof lead hydroxide/oxide, according to the equilibrium: mPb2+ + nHÔ <=>Pbm(OH)n

2"n + nH+. These data were chosen for the evaluation of recommendedvalues in the present report. Additional information for the different references seeSection 6.13: 'Comments on selected references'. Method: sol = solubilitymeasurements, and pot = potentiometry.

log K* Reference Comments I (M) Medium Method

log K*so: Pb2+ + H2O a PbO (red, litharge) + 2H+

-12.67 > [1922APP/REI] T= 298.15 K, I=n/a 0 NaOH pot-12.68 2 [1923SMI/WOO] T= 298.15 K, I=n/a 0 no pot

144

JNC T N 8 4 0 0 99 - O i l


log K*so: Pb2+ + H2O t=> PbO(yellow, massicot) + 2H+

-12.96 1 [1922APP/REI] T= 298.15 K, I=n/a 0 NaOH pot

log K*so: Pb2+ + 2H2O <^Pb(OH)2 (freshly precipitated) + 2H+

-13.05 [1980SCH] T = 298.15 K, 1=0 0 sol1 calculated by [1922APP/REI] from potentiometric measurements2 calculated in this report with a log K of -4.25 for the reaction Pb2++ 2e" <=> Pb(cr) (Section 6.11.1)

6.2.2 Precipitated lead hydroxide

Based on the calculations for litharge and massicot and the observations of [1928RAN/SPE](Table 6.5) and [1980SCH] (Table 6.4) a log K*°So of -13.05 is proposed for freshlyprecipitated lead hydroxide. The log K*°So of -13.15 determined by [1980WAL/SIN] forprecipitated lead hydroxide or oxide is slightly larger due to the presence of anionic species inthe wastewater examined.

6.2.3 Pb(OH)2(s)

In many compilations, a log K*°so of approximately -8.1 is indicated for the reaction Pb2+ +2H2O <=> Pb(OH)2(s) + 2H+. This value refers to crystalline Pb(OH)2(cr) which is formed at300 °C [1995MAR/MAC]. Literature data [1980SCH, 1995MAR/MAC] indicate that this soliddoes not precipitate directly from solutions at room temperature. The absence of Pb(OH)2(cr) inprecipitates is consistent with [1956ROB/THE] who pointed out the difficulty of precipitatingPb(OH)2(cr) directly. Their method of preparing Pb(OH)2(cr) involved heating overnight to280-300 °C.

6.2.4 Additional data for lead hydroxide/oxide compounds

Table 6.5: Additional experimentally determined equilibrium data compiled for the formation of leadhydroxide/oxide compounds. These data were not chosen in the present report for the evaluationof recommended stability values. Reasons for not selecting these references are given in the textin Section 6.2 and in Section 6.13: 'Comments on selected references'. Method: emf =electromotive force measurements at high temperature, sol = solubility measurements.

log K*s Reference Comments Medium Method

log K'so-

-12.68 '-12.63 2

-12.55 2

Pb2+ + H2O <=> PbO (red) + 2H+

[1922APP/REI][1928RAN/SPE][1939GAR/VEL1

T= 298.15 K, 1=1T= 298.15 K, I=n/aT= 298.15 K, 1=0-0.

NaOHNaOHNaOH

solsolsol

145

JNC TN8400 9 9 - 0 1 1


log K*so: Pb2+ + H2O <=> PbO (yellow) + 2H+

-12.90 '-12.75 2

-12.95 2

[1922APP/REI][1928RAN/SPE][1939GAR/VEL]

T= 298.15 K, 1=1T= 298.15 K,I=n/aT= 298.15 K, 1=0-0.1

100

NaOHNaOHNaOH

solsolsol

log K*so: Pb2+ + 2H2O <=> Pb(OH)2 (fresh precipitated) + 2H+

-12.91 2 [1928RAN/SPE] T= 298.15 K, I=n/a NaOH sol-13.15 [ 1980WAL/SIN1 T= 298 K, I=dil 0 H,SO,. sol

log K*so: Pb2+ + 2H2O <=> Pb(OH)2 (crystalline, formed atT = 300 °C) + 2H+

-7.51 [1972NRI] T= 298.15 K, 1=0.1 0 J KOH sol

log K*so: Pb2+ + H2O <=> PbO (crystallinity not defined) + 2H+

-13.18 3 ' 4 [1968CHA/FLE] T= 298.15 K, I=n/a n/a emf-12.55 3 [1984BAN1 T= 298.15 K, I=n/a n/a emf1 calculated with a log p13 value of -28.32 (1=1; Section 6.1.3) and a log Kw of -13.79 (1=1)2 the solubility constants for litharge and massicot can more accurately be determined from the potentiometric data

given in Table 6.43 extrapolated from measurements at 600 - 1000 K4 criticized by [1984BAN]

Table 6.6: Thermodynamic data for the fonnation of lead hydroxide/oxide compounds taken from previouscompilations. As pointed out in Section 2 of this report only experimental data were used for thepresent evaluation. The following table serves only for comparison. Medium: Where data refer tospecific electrolyte solutions, this is indicated.

log K*so Reference Comments I (M) Medium

log K*so: Pb2+ + H20 <=> PbO(red) + 2H+

-12.64 [1952LAT] T= 298.15 K, I=n/a-12.62 [1954COU] T= 298.15 K, I=n/a-12.64 [1960NAS/MER] T= 298.15 K, I=n/a-12.70 [1963FEI/SCH] T= 298.15 K, 1=0-12.74 [1963WIC/BLO] T= 298.15 K, I=n/a-12.69 [1971NAU/RYZ] T= 298.15 K, I=n/a-12.72 [1976BAE/MES] T= 298.15 K, I=n/a-12.70 [1976SMI/MAR] T= 298.15 K, I=n/a

27.93 [1977PAU] T= 298.15 K, I=n/a-12.67 [1978ROB/HEM2] T= 298.15 K, I=n/a-12.63 [1979KUB/ALC] T= 298.15 K, I=n/a-12.65 ' [1979PAT/OBR] T = 298.15 K, 1=0-12.68 [1980BEN/TEA] T= 298.15 K, I=n/a-12.69 [1982PAN] T= 298.15 K, I=n/a-12.64 [1982PAU] T= 298.15 K, I=n/a-12.73 [1982WAG/EVA] T= 298.15 K, I=n/a-12.68 [1983LAN1 T= 298.15 K, I=n/a

146

JNC TN8400 99 - O i l


-12.71 [1984TAY/LOP] T= 295 K, I=dil 0 HCO3"-12.72 [1984VEE/TAR] T= 298.15 K, I=n/a-12.66 [1985BAB/MAT] T= 298.15 K, I=n/a-12.65 [1985GAL] T= 298.15 K, I=n/a-12.61 fl988PHI/HAL] T= 298.15 K, I=n/a

log K*so: Pb2+ + H2O <=> PbO(yellow) + 2H+

-12.79-12.77-12.9-12.86-12.88-12.90-12.79-12.78-12.79-12.87-12.79-12.91-12.79-12.91-12.80-12.79-12.72

log K'so: Pb2+

-13.63-13.63-13.1-27.99-12.99-13.63-12.98-13.64-13.63-12.91

[1952LAT][1954COU][1963FEI/SCH][1963WIC/BLO][1971NAU/RYZ][1976SMI/MAR][1977PAU][1978ROB/HEM2][1980BEN/TEA][1982PAN][1982PAU][1982WAG/EVA][1983LAN][1983SAN/BAR][1985BAB/MAT][1985GAL][1988PHI/HAL1

+ 2H2O<=>Pb(OH):

[1952LAT][1960NAS/MER][1963FEI/SCH][1977PAU][1980BEN/TEA][1982PAU][1982WAG/EVA][1985BAB/MAT][1985GAL][1988PHI/HAL]

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT=298.15K, 1=0T=298.15K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT=298.15K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

>(fresh precipitated) + 1

' T= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0 ,T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

log K*so: Pb2+ + 3H2O <=> PbOPb(OH)2 + 4H+

-27.10 [1976SMI/MAR] T= 298.15 K, I=n/a

• K*so: Pb2+ + 2H2O <=> Pb(OH)2 (crystalline, formed atT' = 300 °C) + 2H+

-8.11-7.74-8.14-8.13-8.14-8.15-8.08

[1971NAU/RYZ][1976NRI][1980BEN/TEA][198OSCH][1981STU/MOR][1982WAG/EVA][1988PHI/HAL1

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT = 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/a

0000000

n/an/an/an/an/an/an/a

147

JNC TN8400 9 9 - 0 1 1


log K*so: Pb2+ + H2O t=>PbO (crystallinity not defined) + 2H+

-12.88-12.71-12.72-12.90-12.63

[1964HIR][1973BAR/KNA][1981BAE/MES][1981STU/MOR][1985CHA/DAV]

T=T=T=T=T=

298.15298.15298.15298.15298.15

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/a

-8.93 [1987BRQAVAN1 T= 298.15 K, I=n/a1 [1979PAT/OBR] selected thermodynamic data for lead and then tested their dataset with tested with experimental data

from different sources.

148

JNC TN8400 99-011

6.3 Lead chloride system

6.3.1 Lead chloride complexes

The existing literature on lead chloride complexes is extensive. Comparison of the resultsreported by different workers is complicated by the fact that the experimentally determinedvalues depend on the composition of the respective medium [1984BYR/MIL, 1984MHVBYR].For the present report, experimental values determined in chloride or perchlorate media (Table6.7) were extrapolated with the SIT equation to I = 0 (Figures 6.9, 6.10, 6.11 and 6.12):

Pb2+ + Cr <=> PbCl+ log p°,,i = 1.55Pb2+ + 2Cr <=} PbCl2° log p°i)2 = 2.00Pb2+ + 3Cr <=> PbCl3- log P°i>3 = 2.01Pb2+ + 4CT <=> PbCl4

2- log P°i,4 = 1.35

A number of investigators extrapolated the formation constants of lead chloride complexes to1=0. Some of these data have also been included in the calculation of the formation constant toexpand the data basis (see Table 6.7). The included data have been either determined at lowionic strength [1930RIG/DAV, 1955PIG/PAR, 1955NAN, 1992LOZ/SCH] where the methodof extrapolation to 1=0 is expected to have only a small influence, or are based on a largenumber of experiments which have been carried out at different ionic strength [1984BYR/MEL,1980MEL/BYR, 1984SEW, 1991MAG/FUE]. These results have been extrapolated by therespective authors to 1=0 using the Pitzer equation. These data were also included in ourcalculations assuming that ideally at 1=0, the same formation constant is obtained with the Pitzercorrection as with the SIT approach. The omission of all constants reported at 1=0 would resultinlogP°i,i, p°i)2, P°i,3, and P°i i4 values of 1.59, 2.10, 2.07 and 1.26, respectively.

The log P values for the formation of Pb chloride'complexes determined in LiClO4 medium aregenerally larger than formation constants determined in NaQC>4 medium. Thus, the log P valuesare plotted in the SIT plot using:

a) the log P values measured in all medium to determine log P°and Ae.b) only the log P values measured in LiClO4 to determine a Ac^cio value for the

IJCIO4 medium.c) only the log P values measured in NaClO4 to determine a A e ^ o o value for the

NaClO4 medium.

149

JNC TN8400 9 9 - 0 1 1

Table 6.7: Experimentally determined equilibrium data compiled for the lead chloride system,according to the equilibrium: mPb2+ + nCl" <=> Pbm(Cl)n

2m"n. These data werechosen for the evaluation of recommended values in the present report. Additionalinformation for the different references see Section 6.13: 'Comments on selectedreferences'. Method: con = conductivity measurements, el = electrophoresis, kin =kinetic measurements, pol = polarography, pot = potentiometry, sol = solubilitymeasurements, and sp = spectrophotometry.

log Pm,

b« A.1.52 i1.57 2,1.593

1.23 4

1.110.900.851.041.151.341.28 50.94 61.231.080.93 70.83 80.69 8

0.90 80.61 80.70 81.59 30.910.840.920.890.980.820.830.850.921.051.15

Referencen

r-pb2+ +cr<=>Pbcr

[1930RIG/DAV]3 [1955BIG/PAR]

[1955NAN][1964MER/KUL][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972VBE][1973BON/HEF][1973HUT/HIG][1973HUT/HIG][1977SIP/VAL][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980PRA/PRA][1981BYR/YOU][1981BYR/YOU][1981BYR/YOU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU]

Comments

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T - 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T - 298.15T= 298.15T= 298.15T= 298.15T= 298.15T - 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=0.001-0.02K, 1=0.0.02-0.1K, 1=0.001K, 1=3K, 1=0.1K, 1=0.5K, 1=1K, 1=2K, 1=3K, 1=4K, 1=4K, 1=1K, 1=0.1K, 1=1K, 1=6.7K, 1=0.7K, 1=0.7K, 1=0.7K, 1=3K, 1=3K, 1=0.01-0.03K, 1=1K, 1=1K, 1=1K, 1=0.5K, 1=1K, 1=2K, 1=0.5K, 1=1K, 1=2K, 1=3K, 1=4

.(M>

0003

0.10.5

12344

10.1

10.70.70.70.7330111

0.51

20.5

1234

Medium

NaClPb(C104)2

NaClO4

LiClO4

LiClO4

LiClO4

LiC104

LiC104

LiC104

LiClO4

NaClO4

NaClO4

NaClO4

NaClO4

NaClNaClO4

NaClO4

NaClO4

NaC104

NaClO4

HC1HC1/HC1O4

MgCl2HC1

NaClO4

NaClO4

NaC104

NaClO4

NaC104

NaC104

NaClO4

NaC104

Method

consp

consolpotpotpotpotpotpotpotpotkinkinpolpolpolpolpolpolpotspspspspspsppotpotpotpotpot

150

JNC TN8400 9 9 - 0 1 1


1.001.33 91.54 91.37 91.38 91.141.111.151.48 1°1.41 n1.60 i2

1.60 3

1.87 4

1.561.301.261.401.702.061.57 5

1.08 61.35 ^1.19 8

1.108

1.26 81.25 8

1.38 81.211.061.231.131.301.331.341.241.271.531.941.041.769

2.08 9

[1982ROH][1984BYR/MIL][1984BYR/MIL][1984BYR/MIL][1984BYR/MIL][1984LOZ/SCH][1984LOZ/SCH][1984LOZ/SCH][1984ML/BYR][1984SEW][1991MGA/FUE][1992LOZ/SCH]

b2++ 2Cr t=> PbCl2°

[1964MIR/KUL][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972 VIE][1973BON/HEF][1977SIP/VAL][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1981BYR/YOU][1981BYR/YOU][1981BYR/YOU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982ROH][1984BYR/MIL][1984BYR/MIL]

T= 296 K, 1=0.7T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=0.5T= 298.15 K, 1=0.2T= 298.15 K, 1=0T= 298.15 K, 1=0

T= 298.15 K, 1=0-6T= 298.15 K, I=dil

T= 298.15 K, 1=3T= 298.15 K, 1=0.1T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=4T= 298.15 K, 1=1T= 298.15 K, 1=0.7T= 298.15 K, 1=0.7T= 298.15 K, 1=0.7T= 298.15 K, 1=0.7T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 296 K, 1=0.7T= 298.15 K, 1=0T= 298.15 K, 1=0

0.7

00001

0.50.2

0000

3

0.10.5

1234

41

0.70.70.70.733111

0.512

0.5

1234

0.700

NaC104

HC1/HC1O4

NaCl/NaC104

MgCl2

CaCl2

HC1O4

HC1O4

HC1O4

differentNaCl, HC1

NaClPbCl,

LiC104

LiC104

LiC104

LiC104

LiC104

LiC104

LiC104

NaC104

NaC104

NaClNaC104

NaC104

NaClO4

NaC104

NaC104

HC1/HC1O4

MgCl2

HC1NaC104

NaClO4

NaClO4

NaC104

NaC104

NaC104

NaClO4

NaClO4

NaClO4

HC1/HC1O4

NaCl/NaClO4

elspspspspsolsol

sol

spsp

solcon

sol

potpotpotpotpotpot

potpotpolpolpolpolpolpol

spspspspspsppot

potpotpotpotelsp

sp

151

JNC TN8400 9 9 - 0 1 1


1.77 9 [1984BYR/MIL]1.81 9 [1984BYR/MIL]1.01 [1984LOZ/SCH]0.92 [1984LOZ/SCH]0.97 [1984LOZ/SCH]2.03 10 [1984MDDL/BYR]1.97 n [1984SEW]1.66 12 [1991MGA/FUE]1.74 3 [1992LOZ/SCH]

T= 298.15 K, 1=0 0T= 298.15 K, 1=0 0T= 298.15 K, 1=1 1T= 298.15 K, 1=0.5 0.5T= 298.15 K, 1=0.2 0.2T= 298.15 K, 1=0 0T= 298.15 K, 1=0 0T= 298.15 K, 1=0-6 0T= 298.15 K,I=dil 0

MgCl2 spCaCl2 spHC1O4 solHC1O4 solHC1O4 sol

different spNaCl, HC1 sp

NaCl solPbCL con

log /3!3: PI

1.98 4

1.201.401.952.402.28 5

1.72 60.86 8

0.75 8

0.74 8

1.71 8

1.79 81.160.921.181.171.231.091.431.831.971.251.72 92.58 91.729

1.75 9

1.521.571.791.86 1°1.66 ii2.27 12

* +set etna;[1964MIR/KUL][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972VIE][1973BON/HEF][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1980LOV/BRA][1981BYR/YOU][1981BYR/YOU][1981BYR/YOU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982BEN/MEU][1982ROH][1984BYR/MIL][1984BYR/MIL][1984BYR/MIL][1984BYR/MIL][1984LOZ/SCH][1984LOZ/SCH][1984LOZ/SCH][1984MIL/BYR][1984SEW][1991MGA/FUE]

T=T=T=T=

T=T=T=T=T=T=T=T=

T=T=T=T=T=T=T=T=T=T=T=T=T=T=nn

T=T=T=

298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15296 K,298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15

K, 1=3K, 1=1K, 1=2K, 1=3K, 1=4K, 1=4K, 1=1K, 1=0.7K, 1=0.7K, 1=0.7K, 1=3K, 1=3K, 1=1K, 1=1K, 1=1K, 1=1K, 1=2K, 1=1K, 1=2K, 1=3K, 1=41=0.7K, 1=0K, 1=0K, 1=0K, 1=0K, 1=1K, 1=0.5K, 1=0.2K, 1=0K, 1=0K,1=0-6

3123441

0.70.70.733111121234

0.700001

0.50.2000

LiC104

LiClO4

LiClO4

LiClO4

LiClO4

NaC104

NaC104

NaC104

NaClO4

NaClO4

NaC104

NaClO4

HC1/HC1O4

MgCl2

HC1NaC104

NaC104

NaClO4

NaC104

NaClO4

NaClO4

NaC104

HC1/HC1O4

NaCl/NaC104

MgCl2

CaCl2

HC1O4

HC1O4

HC1O4

differentNaCl, HC1

NaCl

solpotpotpotpotpotpotpolpolpolpolpolspspspspsppotpotpotpot

elspspspspsolsolsolspsp

sol

152

JNC TN8400 9 9 - 0 1 1


log A/1.72 4

0.851.181.901.43 5

1.46 ii1.27 12

log KSo:

5.00 4

4.78 4

4.114.014.072.80 I2

Pb2+ + 4CI <=> PbCl/'

[1964MIR/KUL][1972FED/SHI][1972FED/SHI][1972FED/SHI][1972VIE][1984SEW][1991MGA/FUE]

Pb2+ + 2Ct & PbCl2(s)

[1964MIR7KUL][1964MIR/KUL][1984LOZ/SCH][1984LOZ/SCH][1984LOZ/SCH][1991MGA/FUE]

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=3K, 1=2K, 1=3K, 1=4K, 1=4K, 1=0K, 1=0-6

K, 1=3K, 1=3K, 1=1K, 1=0.5K, 1=0.2K, 1=0-6

3234400

331

0.50.20

LiC104

LiC104

LiC104

LiClO4

NaC104

NaCl, HC1NaCl

LiC104

LiClHC1O4

HC1O4

HC1O4

NaCl

solpotpotpotpotspsol

potpotsolsolsolsol

log Kso: Pb2+ + H2O + Cl t=>PbOHCl(cr) + Ct

-1.27 13-1.33 13

-1.33 13

-1.28 13

-0.82 13

-0.88 14

-1.25 14

-1.31 14

-1.26 14

-0.62 14

[1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976NAS/LIN][1976N AS/LIN]

T= 298.T= 298.T= 298.T= 298.T= 298.T= 298.T= 298.T= 298.T= 298.T= 298.

15 K, 1=15 K, 1=15 K, 1=15 K, 1=15 K, 1=15 K, 1=15 K, 1=15 K, 1=15 K, 1=15 K, 1=

:0.052=0.554=1.06=1.08=2.078:0.0158=0.029=0.53=1.0332.031

0.050.551.061.082.08

0.0160.0290.531.0332.031

NaC104

NaC104

NaC104

NaClO4

NaC104

NaClO4

NaC104

NaC104

NaC104

NaCIO,

solsolsolsolsolsolsolsolsolsol

corrected to I = 0 by [193ORIG/DAV] with a simplified Debye-Huckel equation50-600 mM Pbcorrected to 1=0 with the Davies equation by the respective authorsonly results used where < 10% LiCl presentthe reported constants are dependent on the Cl concentrationPb concentration = 0.19 mMPb concentration = 0.001 mMPb concentration = 2E-8 Mcorrected to 1=0 with extended Debye-Huckel equation with parameters were fitted by [1984BYR/MIL]

10 corrected to 1=0 with Pitzer's equation by [1984MIL/BYR], same experimental data as [1984BYR/MIL].11 Pb = 0.1 mM, extrapolated to 1=0 with Pitzer equation by [1984SEW]. Also measurements at higher

temperature, AH, AS, ACp given.12 corrected to 1=0 with Pitzer's equation by [1991MGA/FUE].13 measured in presence of lead nitrate14 measured in presence of lead chloride

153

JNC TN8400 9 9 - 0 1 1

Pb2+ + Cl- <=> PbCI+

Figure 6.9: Plot of log p\ i + 4 D vs. Im for the reaction : Pb2+ + Cl- <=> PbCl+ at 25 °C. Thestraight line shows the result of the linear regression: Ae = - 0.12; log f3\i =1.55; AeLjci04 = -0.16 (in LiC104), AeNaCiO4 = -0.10 (in NaC104). Calculatedfrom data compiled in Table 6.7.

QCD

+

CO.

log

5 -

4.5 -

4

3.5

3

2.5

2 |

1.5 -

1

0.5

0 -

Pb2+ + 2Ch <=> PbCI2°

"'•.'''"•'~".'"".\''"]. '-'•• ;"* '

sT-;-!! :

1;:?D

:' - : ; '

:' '

:' * : : : i i

;i , ! !r ' j ; ' ;

;; . '?l ' : : '^ i^.H-i i ' i is i^^.^'v.-vi i^î/ :^: ' i ; ;" : :

1' : ; - .•• •

t-":-::-it.:j"T''^^•;••;"-;

:•

'.'.'. ;••••;;>.;•':-.»"i •i/vv:-:!: j . r ' ; :;:;iv.:l;î';^F:::-\"i^:tj,:'.-L:^^1?:™ : ' i : ;':!i:i - ; i ' : ?;••:;: : i : : - . • .:•••.!.:..»•;•:•.: " .::-\- •••.: ;••:;

' :: ;. 'L:. .-"• •' '^'"ii.-t)-;;^' ^'1"'.-vj:.';;il' i!7" ^ / ^ ^ ^ * : i i i : : i : ' : -|=. ":i "-'-ri T" r r'""-l- H"| |=: f.; -Lf.r i :r- '- •;•;! : '.'..\-:-'- ':•.'•:.' •. .'. '•;;".•

T;i?|i^fsj..::..i:v.;:.1;::::i'|^W;i,•,,;•;

iii^-v':fh^,i'i5t.;:V:;iK, - « * « - • ; . : • !• ' ' i f f i . : : - . : ; - : 1 : > : , ; ; : . r . v ' n > : ; s ; v - ^ ^ p y . : ^ ^ : , , - • ;;.,.,-•• . : • , : - . ; . : • • . : • : .

•I'-.?.:.-; ;•••;;•.-,• :.i:;:;2«': "i.::.,: :a:?KW' !pra|iftt?«K"' ^;;:- s - p ' . * : ; : :^ :« «i «•«;- :s.::

2 4lm, molal

Figure 6.10: Plot of log (5 l i2 + 6 D vs. Im for the reaction : Pb2 + + 2C1- <=> PbCl2° at 25 °C.The straight line shows the result of the linear regression: Ae = - 0.27; log P°i2 =2.00; AeLiCiO4 = -0.32 (in LiC104), AeNaciO4 = -0.24 (in NaC104). Calculatedfrom data compiled in Table 6.7.

154

JNC TN8400 9 9 - 0 1 1

Pb2+ + 3CI- & PbCI3

4.5 -•

4 - LiCIO4

y = 0.35x + 2.01

NaCIO4

= 0.29x + 2.01

y = 0.31x+2.01

2 4lm> molal

Figure 6.11: Plot of log p i i 3 + 6 D vs. Im for the reaction : Pb2 + + 3C1- <=> PbCl3- at 25 °C.The straight line shows the result of the linear regression: Ae = - 0.31; log P°i 3 =2.01; AeLiciO4 = -0.35 (in L1CIO4), AeNaCiO4 = -0.29 (in NaC104). Calculatedfrom data compiled in Table 6.7.

Pb2+ + 4CI- <^ PbCI42"

lm, molal

Figure 6.12: Plot of log p1>4 + 4D vs. Im for the reaction : Pb2 + + 4C1" <=> PbCl42- at 25 °C.

The straight line shows the result of the linear regression: Ae = - 0.21; log (3 \4 =1.35; AeLiCiO4 = -0.24 (in LiC104), AeNaCiO4 = -0.15 (in NaClO4). Calculatedfrom data compiled in Table 6.7.

155

JNC TN8400 9 9 - 0 1 1

6.3.2 PbCl2(s)

Lead also forms salts with chloride. Extrapolation of the values given by [1964MIR/KUL,1984LOZ/SCH and 1991MGA/FUE] (see Table 6.7) is shown in Figure 6.13:

Pb2+ + 2C1" PbCl2(s) log K*°so = 4.81

This value agrees well with the value of 4.77 proposed by [1980CLE/JOH], who made acareful and extensive review of the solubility of lead salts.

Pb2++2CI" « PbCI2(s)

jij4p|ilMga|g||a||p;S;ii|S|];|S|;:

Figure 6.13: Plot of log K*So + 6D vs. Im for the reaction : Pb2+ + 2C1- <=> PbCl2(s) at 25 °C.The straight line shows the result of the linear regression: Ae = - 0.44; log K*°So =4.81. Calculated from data compiled in Table 6.7.

6.3.3 Mendipite: Pb2(OH)3Cl(cr) or Pb4(OH)6Cl2(cr)

Based on data given in [1980BEN/TEA] and [1980MAN], log K*So for mendipite(Pb2(OH)3Cl(cr) or Pb4(OH)6Cl2(cr)) is approximately -8. Primary experimental data howeverare not available. Thus, no solubility product for mendipite is recommended in this report.

156

JNC TN8400 9 9 - 0 1 1

6.3.4 Laurionite andparalaurionite

For laurionite (PbOHCl(cr)) a log K*so of 15.36 can be calculated from the thermodynamic datagiven in [1971NAU/RYZ] and [1988PHI/HAL]. [1987BRU], however, writes that this valueis probably based on an erroneous copying (AjG° -480.3 instead of -408.3) of a value whichalready originally was too low. Experimental results of the formation of PbOHCl(cr) are givenin the careful work of [1976NAS/LIN] (Table 6.7). [1976NAS/LIN] do not state if they uselaurionite or paralaurionite in their experiments. [1987BRU], however, states that the differencein solubility between laurionite (PbCl2-Pb(OH)2(cr)) and paralaurionite (PbCl2-PbO-H2O(cr))is too small to make any apparent difference. Extrapolation of the measurements of[1976NAS/LIN] to 1=0 gives (Figure 6.14):

Pb2+ + Cl- + H2O PbOHCl(cr) + logK *o_so = - 0.62

This value is similar to the value of -0.29 given in different compilations [1982WAG/EVA,1983LAN, 1983SAN/BAR] for laurionite or paralaurionite.

The formation of lead complexes with other halides (bromide, iodide) has also been observed indifferent studies, but is not addressed in the present report.

Pb2++C|-+H2O<^>PbOHCI(cr)+H+

1 2 3lmi molal

Figure 6.14: Plot of log K*s0 + 4D vs. Im for the reaction : Pb2+ + Cl- + H2O <=> PbOHCl(cr)+ H+ at 25 °C. The straight line shows the result of the linear regression: Ae =-0.11; log K*°so = -0.62. Calculated from data compiled in Table 6.7.

157

JNC TN8400 9 9 - 0 1 1

6.3.5 Additional equilibrium data compiled for the lead chloride system

Table 6.8: Additional experimentally determined data for the lead chloride system, according to theequilibrium: mPb2+ + nCl" <=> Pbm(Cl)n

2m'n. These data were not chosen in the present report forthe evaluation of recommended stability values. Reasons for not selecting these references areindicated at the end of the table. Method: con = conductivity measurements, el = electrophoresis,pol = polarography, pot = potentiometry, sol = solubility measurements, and sp =spectrophotometry.

log ?„ Reference Comments Medium Method

log Pb2+ + Cl <=> PbCl*

1.57 '1.16 2

1.16 >1.52 3

1.48 4

0.89 5

1.61

[1955BIG/PAR][1971BON1][1970FED/SAM][1972FED/SHI][1972FED/SHI][1982BEN/MEU]fl984LOZ/SCH]

T= 291.15 K, 1=0T= 298.15 K, 1=0.2-1.8T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

0

30000

n/aCl-

LiClO4

LiC104

LiClO4

NaC104

HC1O,

n/apoln/apotpotpotsol

log pu: Pb2+

1.26 2

2.19 3

2.08 4

1.32 5

1.67 6

h 2CI & PbCl2°

[1971BON1][1972FED/SHI][1972FED/SHI][1982BEN/MEU][1984LOZ/SCH]

T= 298.15 K, 1=0.2-1.8T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

0000

Cl-LiClO4

LiClO4

NaC104

HC1O4

polpotpotpotsol

log /5U: Pb2+ + 3CV <=> PbClj

1.94 x [1960FRI/SAR]1.45 2 [1971BON1]1.97 3 [1972FED/SHI]1.81 4 [1972FED/SHI]0.83 5 [1982BEN/MEU]2.62 6 [1984LOZ/SCH1

T= 298.15 K, 1=5T= 298.15 K, 1=0.2-1.T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

0000

NaC104Cl-

LiC104

LiC104

NaC104

HC1O,

n/apolpotpotpotsol

log p,4: Pb2+ + 4Ci <=>

0.77 3

0.85 4[1972FED/SHI][1972FED/SHI]

T= 298.15 K, 1=0T= 298.15 K, 1=0

00

LiClO4

LiClO,potpot

log Kso: Pb2+ + 2CI

4.907

4.77 6

PbCl2(s)

[1972VIE]H984LOZ/SCH1

T= 298.15 K, 1=4T=298.15K, 1=0

40

NaC104

HC1O4

solsol

158

JNC TN8400 99 - O i l

Table 6.8 : continued

log Kso: Pb2+ + H20 + Cl & PbOHCl(s) + Cl

-0.62 [1976NAS/LIN] T= 298.15 K, 1=0 NaCIO, sol1 no experimental details are known2 I not constant-* extrapolated to 1=0 with SIT model4 extrapolated to 1=0 with Vasilev equation (= Debye-Hiickel)5 linear? extrapolated to 1=0 by [1982BEN/MEU]6 corrected to 1=0 with SIT by [1984LOZ/SCH]7 the reported solubility constant is dependent on Cl concentration8 extrapolated to 1=0 with Debye-Hiickel

Table 6.9: Thermodynamic data for the lead chloride system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison. Medium: Where data refer to specific electrolytesolutions, this is indicated.

Reference Comments KM) Medium

log pa: Pb2+ + Ct <=> PbCl+

1.601.621.590.900.901.021.171.291.641.581.641.641.321.53 i0.94 '1.17 l

0.861.371.551.080.870.860.901.001.121.23

[1969HEL][1976HEM][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980BEN/TEA][1981TUR/WHI][1982WAG/EVA][1983LAN][1987BRO/WAN][1987BRU][1987BRU][1987BRU][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI7MAR]

T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3 ^T= 298.15 K, 1=4T=298.15K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=3T= 291 K, 1=0.7T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=0.5T= 298.15 K, 1=0.7T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4

0

00.5

123400000013

0.7 sea water00

0.10.50.7

1234

159

JNC TN8400 9 9 - 0 1 1


log pu: Pb2+ + 2CV <=> PbCl2°

1.782.441.801.301.301.401.702.001.821.841.841.571.93 '1.16 '1.72 !1.161.932.201.401.201.201.201.301.601.80

[1969HEL][1976HEM][1976SM1/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1981TURAVHI][1982WAG/EVA][1983LAN][1987BRO/WAN][1987BRU][1987BRU][1987BRU][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR]

T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0T=298.15K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=3T=291 K, 1=0.7T= 298.15 K, 1=0T=298.15K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=0.5T= 298.15 K, 1=0.7T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4

0

00.5

12340000013

0.7 sea water00

0.10.50.7

1234

log Pu: Pb2+ + 3CI <=> PbCli

1.682.041.701.401.501.902.301.471.711.471.760.881.84 '1.19 >1.95 '1.061.601.801.301.101.101.101.401.902.20

[1969HEL][1976HEM][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980BEN/TEA][1981TUR/WHI][1982WAG/EVA][1983LAN][1987BRO/WAN][1987BRU][1987BRU][1987BRU][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR]

T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=3T=291 K, 1=0.7T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=0.5T= 298.15 K, 1=0.7T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4

0

0123400000013

0.7 sea water00

0.10.50.7

1234

160

JNC TN8400 9 9 - 0 1 1


log PM: Pb2+ + 4CI <=> PbCl/-

1.38 [1969HEL] T= 298.15 K, 1=0 01.40 [1976SMI/MAR] T= 298.15 K, 1=0 00.70 [1976SMI/MAR] T= 298.15 K, 1=2 21.20 [1976SMI/MAR] T= 298.15 K, 1=3 31.70 [1976SMI/MAR] T= 298.15 K, 1=4 41.40 [1981TUR/WHI] T= 298.15 K, 1=0 01.46 [1983LAN] T= 298.15 K, 1=0 0-0.71 [1987BRO/WAN] T= 298.15 K, 1=0 01.24 : [1987BRU] T= 298.15 K, 1=3 31.38 [1988PHI/HAL] T= 298.15 K, 1=0 01.10 [I989SMI/MAR] T= 298.15 K, 1=0 00.70 [1989SMI/MAR] T= 298.15 K, 1=2 21.00 [1989SMI/MAR] T= 298.15 K, 1=3 31.40 f 1989SM1/MAR1 T= 298.15 K, 1=4 4_

log Kso: Pb2+ + 2Ct « PbCl2(s)

4.78 [1952LAT] T= 298.15 K, I=n/a4.80 [1963W1C/BLO] T= 298.15 K, I=n/a5.02 [1968ROB/WAL] T= 298.15 K, I=n/a4.80 [1971NAU/RYZ] T= 298.15 K, I=n/a5.04 [1973BAR/KNA] T= 298.15 K, I=n/a4.78 [1976SMI/MAR] T= 298.15 K, 1=0 05.00 [1976SMI/MAR] T= 298.15 K, 1=3 34.78 [1977PAU] T= 298.15 K, I=n/a4.79 [1978ROB/HEM2] T= 298.15 K, I=n/a4.81 [1979KUB/ALC] T= 298.15 K, I=n/a4.79 [1980BEN/TEA] T= 298.15 K, I=n/a4.77 [1980CLE/JOH] T=298.15 K, I=n/a4.67 [1980MAN/DEU] T=298.15 K, 1=0 , 04.78 [1982PAU] T= 298.15 K, I=n/a4.81 [1982WAG/EVA] T= 298.15 K, I=n/a4.82 [1983LAN] T= 298.15 K, I=n/a4.82 [1983SAN/BAR] T= 298.15 K, I=n/a4.81 [1985CHA/DAV] T= 298.15 K, I=n/a4.78 [1985GAL] T= 298.15 K, I=n/a

_4J?0 ri988PHKHAL] T= 298.15 K, I=n/a

log Kso: Pb2+ + H2O + Cl- <=> PbOHCl(s) + H+

-1.00-0.30-0.28-0.29

log Kso: Pb2+

15.37-0.2915.36

[1973BAR/KNA][1980CLE/JOH][1982PAU][1982WAG/EVA]

T= 298.15 K, I=n/aT=298.15K, I=n/aT=298.15K, I=n/aT= 298.15 K, I=n/a

+ H2O + Cl <=> PbOHCl(laurionite) + Cl

[1971NAU/RYZ][1983SAN/BAR][1988PHI/HAL1

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/a

161

JNC TN8400 9 9 - 0 1 1


log Km: 2Pb2+ + 3H2O + Ci <=> Pb2(OH)3Cl(s), mendipite

-8.90 2 [1980BEN/TEA] T= 298.15 K, I=n/a-8.00 3 [1980MAN/DEU] T=298.15 K, 1=0

log Ks0: 4Pb2+ + 6H2O + 2Ct <=> (Pb(OH)2)3PbCl2(s)

-17.59 [1982WAG/EVA1 T= 298.15 K, I=n/a1 mean calculated by [1987BRU] from data from different sources.2 mendipite (Pb2(OH)3Cl(cr)) solubility as given in [1987BRU],3 data selected from literature, corrected to 1=0 by [1980MAN/DEU].

162

JNC TN8400 9 9 - 0 1 1

6.4 Lead fluoride system

In the presence of fluoride, lead forms different complexes. The complex formation constants inbinary lead fluoride system are well documented (Table 6.10).

Table 6.10: Experimentally determined equilibrium data compiled for the lead fluoride system,according to the equilibrium: mPb2+ + nF- <=> Pbm(F)n

2m"n. These data werechosen for the evaluation of recommended values in the present report. Additionalinformation for the different references see Section 6.13: 'Comments on selectedreferences'. Method: fe = fluoride selective electrode, pol = polarography, pot =potentiometry, sol = solubility measurements.

log Kn

log fa,:

1.261.48 11.53 21.62 21.73 2

1.40 3

1.40 4

1.46

log fa,:

2.552.59 2

2.52 32.52

log Kso:

7.43 5

6.607.44 6

7.67 7

6.60 7

Reference

Pb2+ + F- <=>PbF+

[1963MES/HUM][1965BOT/CIA][1970BON/HEF][1970BON/HEF][1970BON/HEF][1971BON2][1972HEF][1973BON/HEF]

Pb2+ + 2F- <=> PbF2°

[1963MES/HUM][1970BON/HEF][1971BON2][1973BON/HEF]

Pb2+ + 2F- <^>PbF2(s)

[1963MES/HUM][1963MES/HUM][1963MES/HUM][1981 CIA][1981 CIA]

Comments

T=291 K, 1=2T= 298 K, 1=1T= 288 K, 1=1T=288K,I=1T= 288 K, 1=0.1T= 298 K, 1=1T= 298 K, 1=1T= 298 K, 1=1

-*

T= 291 K, 1=2T= 288 K, 1=1T= 298 K, 1=1T= 298 K, 1=1

T= 298 K, 1=0T=291 K, 1=2T=291 K, 1=0T= 298 K, 1=0T= 298 K, 1=1.05

KM,

2111

0.1111

2111

0200

1.05

Medium

NaClO4

NaC104

NaC104

NaC104

NaC104

NaClO4

NaC104

NaClO4

NaC104

NaClO4

NaClO4

NaClO4

n/aNaC104

NaClO4

NaClO4

NaClO4

Method

polpolpolfefefefefe

polpolfefe

potsolsolsolsol

Pb < 10 mMPrecision lower than in other studies of Bond and Hefter [1971BON2, 1972HEF, 1973BON/HEF]. Pbconcentration = 10 mM.Pb concentration = 0.4 mM.Pb concentration = 31 mM.recalculation of a value given in [1947BRO/DEV]extrapolation of own measurements to 1=0no experimental details given

163

JNC TN8400 9 9 - 0 1 1

6.4.1 PbF+andPbF2°

The formation constants for Pb fluoride complexes reported in literature [1963MES/HUM,1965BOT/CIA, 1970BON/HEF, 1971BON, 1972HEF, and 1973BON/HEF] (Table 6.10) areextrapolated to 1=0 ( Figures 6.15 and 6.16):

pb2+ ,

P b 2 + -

HF

h 2 F

« PbF+

<=> P b F 2 °

log p°i

log P°i

,1 =

,2 =

2.27

3.01

4 . 5 -

4 -•

Q 3.5

+ 3 +

og

1.5 -

1

0.5

0

o

0

Pb2+ + F- a PbF+

y ==-0.02x_+2.27

0.5 1 1.5lm, molal ^

2.5

Figure 6.15: Plot of log p u + 4 D vs. Im for the reaction : Pb2 + + F- o PbF+ at 25 °C. Thestraight line shows the result of the linear regression: Ae = 0.02; log P°i i = 2.27.Calculated from data compiled in Table 6.10.

164

JNC TN8400 9 9 - 0 1 1

5 -

4.5 •

4

Q 3.5CD

+ 3CM

-2.5CO.

1.5

1

0.5 •

n

Pb2+ + 2F- <=> PbF2°

.y'ljjllllplilil^filllllllllllilll| -:;.il|||J ||||||f||||||::

0.5 1 1.5lm, molal

2.5

Figure 6.16: Plot of log (3U + 6 D vs. Im for the reaction : Pb2+ + 2F- <=>-PbF2° at 25 °C. Thestraight line shows the result of the linear regression: Ae = - 0.04; log (3°i 2 =3.01. Calculated from data compiled in Table 6.10.

6.4.2 PbF2(s)

Extrapolation of log K*So values for the reaction: Pb2+ + 2F- <=> PbF2(s) from[1962MES/HUM] and [1981CIA] results in (Figure 6.17):

Pb2+ + 2F PbF2(s) log K*°so = 7.52

This value is in very good agreement with the value of 7.48 proposed by [1980CLE/JOH], whomade a careful and extensive review of the solubility of lead salts.

165

JNC TN8400 9 9 - 0 1 1

QCD

+o

log

1 0 •

9.5 -

9

8.5 -

8 -

7.5 <

7

6.5 -

6 -

5.5-

F. -

Pb2+

•-I K

• - - • ' \

+ 2F- <=> PbF2 (s)

'r-;V''" „• •«" * V

J* * " ' , — — - ^

_j;= 0:1^ + 7,52 "

« • r~^* • ^^* ?\ ' t ' ** ^ '

Z , 1 1 , ( •

0.5 1 1.5m, molal

2.5

PbF2(s) at 25 °G.Figure 6.17: Plot of log K*So + 6 D vs. Im for the reaction : Pb2+ + 2F"The straight line shows the result of the linear regression: Ae = - 0.19; log K*°so =7.52. Calculated from data compiled in Table 6.10.

Table 6.11: Thermodynamic data for the lead fluoride system taken from previous compilations. As pointedout in Section 2 of this report only experimental^ data were used for the present evaluation. Thefollowing table serves only for comparison.

Reference Comments KM)

log pu:Pb2+ +F <=>PbF+

1.441.261.442.061.441.260.771.012.091.08

[1976SMI/MAR][1976SMI/MAR][1980BON/HEF][1981TUR/WHI][1982SMI/MAR][1982SMI/MAR][1982WAG/EVA][1985BAB/MAT][1987BROAVAN][1988PHI/HAL]

T= 298.15 K, 1=1T= 298.15 K, 1=2T=298K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

1210120000

166

JNC TN8400 9 9 - 0 1 1


log pu: Pb2+ + 2F- <=> PbF2°

2.54 [1976SMI/MAR] T= 298.15 K, 1=1 12.55 [1976SMI/MAR] T= 298.15 K, 1=2 22.53 [1980BON/HEF] T= 298 K, 1=1 13.42 [1981TUR/WHI] T= 298.15 K, 1=0 02.53 [1982SMI/MAR] T= 298.15 K, 1=1 12.55 [1982SMI/MAR] T= 298.15 K, 1=2 21.60 [1982WAG/EVA] T= 298.15 K, 1=0 01.75 [1985BAB/MAT] T= 298.15 K, 1=0 03.85 [1987BRO/WAN] T= 298.15 K, 1=0 0-0.89 [1988PHI/HAL] T= 298.15 K, 1=0 0_

log Pu: Pb2+ + 3P «• PbFf

5.39 [1987BROAVAN] T= 298.15 K, 1=0 02.71 ri988PHI/HAL1 T= 298.15 K, 1=0 0_

log p!4: Pb2+ + 4F- <=> PbF42-

6.78 ri987BROAVANl T= 298.15 K, 1=0

log Kso: Pb2+ + 2F- <=> PbF2(s)

0120

7.647.446.266.607.487.68

[1973BAR/KNA][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980CLE/JOH][1988PHI/HAL]

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, I=n/aK, 1=0K, 1=1K,I=2K, 1=0K, I=n/a

log Kso: Pb2+ + 2F- <=> a-PbF2

5.675.647.067.547.607.547.645.64

[1952LAT][1963WIC/BLO][1971NAU/RYZ][1979KUB/ALC][1982WAG/EVA][1985CHA/DAV][1985CHA/DAV][1985GAL]

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n7a

167

JNC TN8400 9 9 - 0 1 1

6.5 Mixed lead fluoride chloride system

Independently, [1971BON2] and [1972HEF] determined the formation constant of mixedPbCIF complexes in 1 M NaClC>4 medium and [1968AND] determined the solubility ofPbClF(cr) (Table 6.12).

Table 6.12: Experimentally determined equilibrium data compiled for the lead chloride fluoridesystem, according to the equilibrium: mPb2+ + nQ~ + oF~ <=> PbmCl nF0

2m-n"°.These data were chosen for the evaluation of recommended values in the presentreport. Additional information for the different references see Section 6.13:'Comments on selected references'. Method: fe = fluoride selective electrode, pol= polarography,.

log Pm,n Reference Comments Medium Method

log Pl2: Pb2+ + CI-+ F- <=$ PbCIF0

2.72 ] [1971BON2] T= 298 K, 1=12.78 2 [1972HEF] T= 298 K, 1=12.87 2 [1972HEF] T= 298 K, 1=1

111

NaC104

NaClO4

NaClO.

polfefe

log KSO: Pb2+ + CI-+ F- ^PbClF(cr)

8.008.158.028.108.228.408.548.64

[1968 AND][1968AND][1968AND][1968 AND][1968 AND][1968 AND][1968 AND][1968 AND]

T= 293 K, 1=0.6T= 293 K, 1=0.2T= 293 K, 1=0-0.2T= 293 K, 1=0-0.2T= 293 K, 1=0-0.2T= 293 K, 1=0-0.2T= 293 K, 1=0-0.2T= 293 K, 1=0-0.2

0.60.2

0.250.20.1

0.0250.010.005

NaClO4

NaClO4

NaClO4

NaClO4

NaClO4

NaC104

NaClO4

NaC104

solsolsolsolsolsolsolsol

1 Pb concentration = 0.4 mM2 Pb concentration = 1.8 mM

168

JNC TN8400 9 9 - 0 1 1

6.5.1 PbFCP

Correction with the SIT term, assuming a Ae of -0.43 (as observed for PbFCl(cr) in thefollowing section), yields a value of:

PbFCl0 log = 3.55

Pb2+ + Cl- +F- PbCIF0

0.5 1lm, molal

1.5

Figure 6.18: Plot of log p i , u + 6 D vs. Im for the reaction : Pb2+ + Cl- + P <=> PbFCl0 at 25°C. Ae = - 0.43 (see text); log $°vj,\ = 3.94. Calculated from data compiled inTable 6.12.

6.5.2 Matlockite (PbClF(cr))

Based on data given in different compilations (Table 6.14), a log K*so for matlockite(PbClF(cr)) of approximately 8.6 can be calculated. The solubility has been determinedexperimentally by [1968AND] (Table 6.12). Extrapolation to 1=0 is shown in Figure 6.19.

+ Cl- + F- <=> PbClF(matlockite) log K*°so= 8.82

169

JNC TN8400 9 9 - 0 1 1

Pb;2+ PbCIF(cr)10

5.5 +• ,

5

y = 0.43x + 8.82

o 0.2 0.4 0.6lm, molal

0.8

Figure 6.19: Plot of log K*so + 6 D vs. Im for the reaction : Pb2+ + Cl- + F" oPbClF(matlockite) at 25 °C. The straight line shows the result of the linearregression: Ae = -0.43; log K*°so= 8.82. Calculated from data compiled in Table6.12.

Table 6.13: Additional experimentally determined data for the lead chloride fluoride system. These data werenot chosen in the present report for the evaluation of recommended stability values. Method: sol= solubility measurements.

log K*so

log K*so:8.82 '

Pb2+

Reference

+ CI-+ F- <=>T1968AND1

Comments

PbClF(cr)T= 293 K, 1=0-0.2

KM)

0

Medium

NaClOi

Method

solextrapolated to 1=0 as a function of I°-5/(l+I0-5)

Table 6.14: Thermodynamic data for the lead chloride fluoride system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison.


log K*so: Pb2+ + Cl- + F- <=> PbClF(s)

8.378.948.668.41

[1971NAU/RYZ][1982WAG/EVA][1983LAN]fl988PHI/HAL1

T= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

170

JNC TN8400 9 9 - 0 1 1

6.6 Lead carbonate system

Under environmental conditions carbonate is an important ligand for lead and has a significantinfluence on lead solubility under neutral to alkaline conditions. The formation constants forlead carbonate complexes and in particular, the solubility products of lead carbonates (cerrusite,hydrocerrusite) reported in the literature are controversial. The experimental data of[1975ERN/ALL, 1976BIL/HUS, 1977SIP/VAL, 1982BIL/SCH, 1987FER/GRE and1989DOR/MAR] were used to obtain formation constants valid at 1=0. Some of the constants[1973BIL/STU, 1976BIL/HUS] are determined in 0.1 M KNO3 medium. These values werealso included in the calculations as the complex formation between lead and carbonate is muchstronger than between lead and nitrate and as the constants can be corrected for the interactionbetween Pb and nitrate using the constants derived in Section 6.7.1.

The data measured by Byrne and co-workers [1981BYR, 1988BYR/KUM] in seawater(interaction with Cl) and [1980WAL/SIN] measured in wastewater were not included in theevaluation. The experimentally determined data used to obtain the formation constants valid at1=0 are given in Table 6.15.

Table 6.15: Experimentally determined equilibrium data compiled for the lead carbonatesystem, according to the equilibrium: mPb2+ + nCO3

2-<=> Pbm(CO3)n(OH)o2m-2n.

These data were chosen for the evaluation of recommended values in the presentreport. Additional information for the different references see Section 6.13:'Comments on selected references'. Method: sol = solubility measurements, sp =spectropho tome try, and pot = potentiometry, pol = polarography.

log (3

log (5

6.376.476.576.275.625.375.345.407.40

m,n

1.2

1, 2

1

1

3

3

Reference

>2+ + CO32- <=> PbCO3°

[1975ERN/ALL][1975ERN/ALL][1976BEL/HUS][1976BIL/HUS][1976BIL/HUS][1977SIP/VAL][1977SIP/VAL][1982BIL/SCH][1989DOR/MAR]

Comments

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=0.1K, 1=0.1K, 1=0.1K, 1=0.1K, 1=0.7K, 1=0.7K, 1=0.7K, 1=0.3K, I=dil

KM)

0.10.10.10.10.70.70.70.30

Medium

KNO3

KNO3

KNO3KNO3

NaC104

NaClO4

NaClO4

NaClO4

no

Method

potpotpotpotpolpolpolsolsol

171

JNC TN8400 9 9 - 0 1 1


log (3]2: Pb2+ + 2CO32- <=>Pb(CO3)2

2-

9.979.278.2 4

8.378.738.868.90

l

l

3

4

[1976BIL/HUS][1976BIL/HUS][1976BIL/HUS][1977SIP/VAL][1977SIP/VAL][1982BIL/SCH][1987FER/GRE]

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=0.1K, 1=0.1K, 1=1K, 1=0.7K, 1=0.7K, 1=0.3K, 1=3

0.10.1

10.70.70.33

KNO3

KNO3NaClO4

NaC104

NaC104

NaC104

NaC104

potpotpotpolpolsolsol

log Kso: Pb2+ + CO32- <=> PbCOrfcerrusite)

1312121313

.13 6

.71 i

.15

.20

.18

[1961NAS/MER][1973BIL/STU][1982BEL/SCH][1984TAY/LOP][1989DOR/MAR]

T= 298.15T= 298.15T= 298.15T= 295 K,T= 298.15

K, 1=0.2-2K, 1=0.1K, 1=0.3I=dilK, I=dil

00.10.300

KNO3

KNO3NaClO4

HCO3no

solpotsolsolsol

log Kso: 3Pb2+ + 2CO32- + 2H2O t=> Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+

17.64 [1984TAY/LOP] T= 295 K, I=dil 0 HCCy sol

log Kso: 10Pb2+ + 6CO32- + 7H2O <^PbI0(CO3)6(OH)6O(plumbonacrite) + 8H+

41.21 [1984TAY/LOP] T= 295 K, I=dil 0 HCO,- solcorrected in this report for the effect of the formation of PbNO3

+ and Pb(NO3)2° complexes using the constantsderived in Section 6.7.1. Uncorrected values are given in Table 6.16.Pb concentration = 0.0025 mM.no PbHCO3" complexes are formed. Pb concentration = 0.001 mM.recalculation of the experimental data of [1965BAR/BAR] by [1976BIL/HUS]pKw used for calculations in this report (3 M NaClO4) = 14.18extrapolated to 1=0 by [1961NAS/MER] as a function of I with a poiynom.

172

JNC TN8400 9 9 - 0 1 1

6.6.1 PbCO3°

The formation constant for the PbCC>30 complex at 1=0 was derived from measurements by[1975ERN/ALL, 1976BIL/HUS, 1982BIL/SCH, 1977SIP/VAL, 1989DOR/MAR] (Figure6.20). The resulting log pV i of 7.30 is in good agreement with the value of 7.10 and 7.14used by [1980SCH] and [1984TAY/LOP], respectively.

Pb2 + + CO32 PbCO3° log P ° u = 7.30

Pb2+ + CO,2" <=> PbCO,

lmi molal

Figure 6.20: Plot of log (3U + 8 D vs. Im for the reaction : Pb2+ +CO3- <=> PbCO3° at 25 °C.The straight line shows the result of the linear regression: Ae = 0.56; log (3°u =7.30. Calculated from data compiled in Table 6.15.

173

JNC TN8400 9 9 - 0 1 1

6.6.2 Pb(CO3)22-

Experimental values reported in the literature show some spread (see Table 6.15 and Figure6.21). The largest difference, however, is observed for two values determined by the[1976BIL/HUS] in 0.1 M KNO3 medium with anodic pulse voltammetry and differential pulsepolarography.

Extrapolation of the results of [1976BIL/HUS, 1977SIP/VAL, 1982BIL/SCH, and1987FER/GRE] to 1=0 is shown in Figure 6.21 and gives:

Pb(CO3)22- log p°i,2 = 10.13

The resulting log P°i2 of 10.13 is comparable with the value of 10.33 and 10.62 used in thereviews of [1980SCH] and [1984TAY/LOP].

Q00

+CO.

log

Pb2+ + 2CO32" *

1 O

12.5

12

11.5

11 -

10.5

10

9.5

9 -

8.5 -

8 -

• ' * * ^ • " % *

O : ; !> - :"*:~- } '-, - ••'

"Illll-—-^-^° ° v"' ' :,

" • •

I- +-

* Pb(CO3)22"

3x> 16.13 , _.

• v ' O .

I

1 2 3lm, molal

Figure 6.21: Plot of log p u + 8 D vs. Im for the reaction : Pb2+ + 2CO32- <=> Pb(CO3)22" at 25

°C. The straight line shows the result of the linear regression: Ae = - 0.13; logP°i,2 = 10.13. Calculated from data compiled in Table 6.15.

174

JNC TN8400 9 9 - 0 1 1

6.6.3 Pb(CO3)34- and Pb(CO3)4

6-

[1986BROAVAN] calculated formation constants for Pb(CO3)34- and Pb(CO3)4

6\ Theseformation constants are theoretical values and the existence of these species is not proven.

6.6.4 PbCO3OH-

[1987FER/GRE] determined experimentally a log (5 value for the reaction: Pb2+ + CO32" + H2O

<=> PbCO3OH- + H+ of -3.28 at 1=3 (see Table 6.16). Correction with the SIT term of a singlevalue measured at I = 3 to I = 0 is not possible without making uncertain assumptions. Thus,no log P value for the formation of PbCCÔH" is recommended in this report.

6.6.5 PbHCOf, Pb(HCO3)2° and Pb(HCO3)f

Several lead bicarbonate complexes have been reported. For the PbHCO3+ complex, the onlyactual measurement is that of [1989DOR/MAR] (Table 6.16), who give a log p ° u of 2.63 forthe reaction Pb2+ + HCO3- <=> PbHCO3+ or, using a log p value of 10.33 for the reaction CO3-+ H+ <=> HCO3- [1995SIL/BID], a log p ° U i l = 12.96 for the reaction Pb2+ + CO3- + H+ <=>PbHCC>3+. As this is only formation constant determined for PbHCC>3+ and as in theexperiments of [1989DOR/MAR] less than 10% of the total dissolved Pb are present asPbHCC>3+, the usage of this value is not recommended in this report.

[1976BIL/HUS] showed that the results of [1965BAR/BAR and 1967BAR], who proposed theexistence of Pb(HCO3)2° and Pb(HCC>3)3~ (given In Table 6.16), can be explained satisfactorilywith the formation of a Pb(CC>3)22" complex [1976BIL/HUS], which makes the formation ofhigher bicarbonate complexes not probable.

175

JNC TN8400 9 9 - 0 1 1

6.6.6 PbCO3(cerrusite)

A number of lead carbonate solids are reported in the literature. From the data of[1961NAS/MER], [1973BIL/STU], [1982BIL/SCH], [1984TAY/LOP], [1989DOR/MAR](Table 6.15) a log K*°So for cerrusite (PbCO3(cr)) can be derived (Figure 6.22).

PbCO3(cerrusite) log K*°so = 13.23

This value can be compared with the log K*°So of 13.10 selected and verified with experimentalresults by [1995MAR/MAC] (Table 6.16).

Pb2++CO32-

15PbCO3(cerrusite)

Figure 6.22: Plot of log K*so + 8 D vs. Im for the reaction : Pb2+ +CO3" <=> PbCO3(cerrusite)at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.76;log K*°so = 13.23. Calculated from data compiled in Table 6.15.

176

JNC TN8400 9 9 - 0 1 1

6.6.7 Pb3(CO3)2(OH)2(hydrocerrusite)

In the case of hydrocerrusite (Pb3(CO3)2(OH)2(cr)), two distinct groups of log K*So values canbe seen in Figure 6.23, Tables 6.15 and 6.16. From the data of [1928RAN/SPE,1973BIL/STU, 1982BIL/SCH] (Table 6.16) one extrapolates a log K*So of 19.31 forhydrocerrusite at 1=0. However, more recent measurements by [1983SCH/GAR] (Table 6.16)and [1980PAT] indicate that this log K*so for hydrocerrusite is too large, i.e., underestimatinglead solubility. [1984TAY/LOP] determined with X-ray analysis (see also comments in Section6.13) the following solubility product at 1=0 for hydrocerrusite:

3Pb2+ + 2CO32- + 2H2O <=> Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+ log K*°so = 17.64

This value is in accordance with the values proposed by [1983SCH/GAR] and[1995MAR/MAC] (cf. Table 6.16).

The difference between the two groups of data can not be explained satisfactorily at themoment. We propose to use a log K*°so of 17.64 for hydrocerrusite.

3Pb2+ + 2CO32" + 2H2O21

Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+

0.2 0.4 0.6lm, molal

0.8

Figure 6.23: Plot of log K*So + 18 D vs. Im for the reaction: 3Pb2+ + 2CO32" + 2H2O <=>

Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+ at 25 °C. log K*°so = 17.64.Calculated from data compiled in Table 6.15.

177

JNC TN8400 9 9 - 0 1 1

6.6.8 Plwnbonacrite

[1984T AY/LOP] showed with X-ray analysis the existence of plumbonacrite(Pbio(C03)6(OH)60(cr)) under certain conditions in alkaline medium (see also comments inSection 6.13). They determined a log K*°So of 41.21 at 1=0 (Table 6.15), which differsstrongly from the value given by [1981HAA/WIL] (Table 6.16), who concluded thatplumbonacrite is metastable with respect to hydrocerrusite and litharge. [1981HAA/WTL] didnot report any experimental details or details of the calculations, while the value given by[1984TAY/LOP] is based on extensive observation of interconversion reactions ofplumbonacrite. In this report the value determined by [1984TAY/LOP] is proposed:

10Pb2+ + 6CO32- + 7H2O 4=> Pb10(CO3)6(OH)6O(plumbonacrite) + 8H+ log K*°s0 = 41.21

6.6.9 PbCO3PbO(s) and PbCO3(PbO)2(s)

Based on data given in different thermodynamic compilations (Table 6.17), log K*so values forPbCO3PbO(s) and PbCO3(PbO)2(s) can be calculated Primary experimental data, however, arenot available and the conditions under which these minerals are formed do not seem to beknown.

6.6.10 Phosgen ite

A log K*so for phosgenite (Pb2CO3Cl2(cr)) of approximately 19.9 can be calculated from thedata compiled in previous compilations (Table 6.17). Primary experimental data are notavailable and the conditions which these minerals are formed do not seem to be known.

178

JNC TN8400 9 9 - 0 1 1

6.6.11 Additional data for the lead carbonate system

Table 6.16: Additional experimentally determined data for the lead carbonate system, according to theequilibrium: mPb 2 + + nCO3

2" + oH2O <=> Pbm(CO3)n(OH)o2m-2n-° + H+. These data were not

chosen in the present report for the evaluation of recommended stability values. Reasons for notselecting these references are given in the text in Section 6.6 and in Section 6.13: 'Comments onselected references'. Method: pol = polarography, pot = potentiometry, sol = solubilitymeasurements, and sp = spectrophotometry.

Reference Comments Medium Method

log pu: Pb2+ + CO32- & PbCO3°

7 ' [1969BAR]6.4 2 [1973BBL/STU]6.2 3 - 4 [1975ERN/ALL]6.3 3- 4 [1975ERN/ALL]6.40 3 [1976BIL/HUS]6.10 3 [1976BIL/HUS]6.34 [1980WAL/SIN]7.10 [1980SCH]4.00 5 [1981BYR]6.00 6 [1987FER/GRE]

T= 298.15 K, 1=1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T = 298.15 K,I=dilT = 298.15 K, 1=0T= 298.15 K, 1=0.7T= 298.15 K, 1=3

10.10.10.10.10.100

0.73

NaC104

KNO3

KNO3

KNO3KNO3KNO3H2SO4

n/aSW

NaCIO,

solpotpotpotpotpotsolsolspsol

log J5U: Pb2+ + 2CO32- <=> Pb(CO3)2

2-

8.2 7

9 19.8 2

9.8 3

9.1 3

10.8 3

10.3310.40 8

[1959FAU/BON][1969BAR][1973BEL/STU][1976BIL/HUS][1976BIL/HUS][1976BIL/HUS][1980SCH][1987FER/GRE1

T=291 K, 1=1.7T= 298.15 K, 1=1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0T = 298.15 K, 1=0T= 298.15 K, 1=0

1.71

0.10.10.1000

KNO3NaC104

KNO3KNOjKNO3KNO3

n/aNaClO4

polsolpotpotpotpotsolsol

,,u: Pb2+ + CO32- + H2O <=> Pb(CO3)OH-+ H+

-3.28 9 [1987FER/GRE1 T= 298.15 K, 1=3 NaC104 sol

log Pu: Pb2+ + HCO3<=> PbHCO3+

2.63 10 [1989DOR/MAR] T= 298.15 K,I=dil no sol

log pl2: Pb2+ + 2HCO3<=> Pb(HCO3)2°

4.77 ' ' [1965BAR/BAR] T= 293 K, 1=14.78 "• 12. ri967BAR] T= 298 K, 1=1

NaHCOjNaHCO,

solsol

179

J N C T N 8 4 0 0 9 9 - O i l


log Pu: 3HCO3<=> Pb(HCO3)3

5.19 u [1965BAR/BAR]5.20 "• 12 [1967BAR1

T=293K,I=1T= 298 K, 1=1

NaHCO3

NaHCO,solsol


log Kso: Pb2

12.54 '12.60 l 3

13.10

f + C O / " <=> PbCO3(cerrusite)

[1973BIL/STU][1980WAL/SIN]fl995MAR/MACl

T= 298.15 K, 1=0.1T= 298 K, I=dilT = 298.15 K, 1=0

0.100

KNO3

H2SO4

no

potsolsol

log Kso: 31

19.04 ! 4

17.11 3

17.63 15

16.5616.2516.88 15

17.40 9

17.8017.50

>b2+ + 2CO32- + 2H2O

[1928RAN/SPE][1973BEL/STU][1973BIL/STU][1982BIL/SCH][1982BBL/SCH][1982BIL/SCH][1982BIL/SCH][1983SCH/GAR][1995MAR/MAC]

«• Pb3(CO3)2Pb(OH)2(hydrocerrusite) -

T= 298.15 K, I=n/aT= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=0.3T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T = 298.15 K, 1=0.0075T = 298.15 K, 1=0

\-2H+

0

00.10.10.30.30.10.1.0075

0

NaOHKNO3

KNO3NaC104

NaC104

KNO3KNO3

CaCO3

no

solpotpotsolsolsolsolsolsol

log Kso: 10Pb2+ + 6CO32- + 7H2O <=> Pbl0(CO3)6(OH)6O(plumbonacrite) + 8W

8.76 16 [1981HAA/WIL1 T= 298.15 K,I=n/a n/a solapproximate values, calculated by [1969BAR] with a log Kso*° of 12.83 for cerrusite.same values also reported in [1976BIL/HUS]values corrected for the effect of the formation of Pb-nitrate complexes are given in Table 6.15.Pb concentration = 0.0025 mMSW = seawater. Formation of lead chloride complexes probablelog P,, for the reaction Pb2+ + CO3

2" <=> PbCO3° is only an approximate value. pK^ in 3 M NaClO4 = 14.18.measured in 1.7 M nitrate medium, formation of Pb nitrate complexes probable. Pb concentration = 0.1 mM.extrapolated to 1=0 with SIT term by [1987FER/GRE].pKw used for calculations in this report (3 M NaClO4) = 14.18only value reported in the literature, relatively small amount of PbHCO3

- present in the experiments of[1989DOR/MAR]recalculated by [1976BIL/HUS] assuming the formation of Pb carbonate complexes .given in [1969BAR].determined in wastewater, see also comments in Section 6.13.

14 calculated with a log p ! 3 value of -28.02 (1=0; Section 6.1.3)15 corrected in this report for the effect of the formation of PbNO3

+ and Pb(NO3)2° complexes using the constantsderived in Section 6.7.1. Uncorrected values are given in Table 6.16.extrapolated to 1=0, no experimental details reported.16

180

JNC TN8400 9 9 - 0 1 1

Table 6.17: Thermodynamic data for the lead carbonate system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison. Medium: Where data refer to specific electrolytesolutions, this is indicated.

log Pm,n

log PlA: Pb2+

7.57.247.507.006.387.146.816.135.075.106.535.40

Reference

+ CO32- <=> PbCO3°

[1972ZIR/YAM][1976HEM][1980MAN/DEU][1981TURAVHI][1983LAN][1984TAY/LOP][1987BROAVAN][1987BRU][1988BYR/KUM][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR]

Comments

T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 295 K, I=dilT= 298.15 K, 1=0T= 298.15 K, 1=0- 1T= 298 K, 1=0.7T= 298 K, 1=0.7T= 298.15 K, 1=0T= 298 K, 1=0.3

I(M)

0

000000

0.70.70

0.3

Medium

SWSW

log Pu: Pb2+ + CO/- <=> Pb(CO3)22

10.648.208.6210.299.6610.6212.299.027.2610.088.86

[1976HEM][1976SMI/MAR][1980MAN/DEU][1981TUR/WHI][1983LAN][1984TAY/LOP][1987BRO/WAN][1987BRU][1988BYR/KUM][1988PHI/HAL][1989SMI/MAR]

T= 298.15 K,I=n/a 0T= 291.15 K, 1=1.7 1.7T=298.15K,I=0 0T= 298.15 K, 1=0 0T= 298.15 K, 1=0 0T= 295 K, I=dil 0T= 298.15 K, 1=0 0T= 298.15 K, 1=0-1 0T= 298 K, 1=0.7 0.7 SWT= 298.15 K, 1=0 0T= 298 K, 1=0.3 03

log pu: Pb2+ + 3CO32- <=> Pb(CO3)/-

16.70 IT987BRO/WAN1 T= 298.15 K, 1=0

log p, „: Pb2+ + 4CO32- <=> Pb(CO3)4

6-

20.14 [1987BROAVAN1 T= 298.15 K, 1=0

log ft,,: Pb2+ + HCOf<=> PbHCO3+

2.92.302.785.48

[1972ZIR/YAM][1983SCH/GAR][1987BRO/WAN][1988PHLHAL1

T= 298.15 K, 1=0T = 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0

0

00

181

JNC TN8400 9 9 - 0 1 1


log Pu

4.438.76

,: Pb2+ + 2HCO3<=> Pb(HCO3)2°

[1987BROAVAN]ri988PHI/HALl

T= 298.T= 298.

15 K, 1=015 K,1=0

00

log pu: Pb2+ + 3HCOf<=> Pb(HCO3)f

5.20 [1987BRQ/WAN1 T= 298.15 K, 1=0

log P1A: Pb2+ + 4HCOf<=> Pb(HCO3)42-

T= 298.15 K, 1=0

log Kso:

12.8213.2413.4612.8313.1313.1311.0112.9712.8012.9112.8513.4513.1313.1313.1112.8312.9712.8212.7812.8112.8010.91 '12.9512.8912.8213.1312.1513.6412.8613.1312.0012.1511.01

Pb2+ + CO/- <=> PbCO3(cerrusite)

[1952LAT][1959UGG][1971NAU/RYZ][1973BAR/KNA][1976HEM][1976SMI/MAR][1976SMI/MAR][1977PAU][1978ROB/HEM2[1979KUB/ALC][1979PAT/OBR][1980BEN/TEA][1980CLE/JOH][1980MAN/DEU][1980SCH][1981 STUMOR][1982PAU][1982WAG/EVA][1983LAN][1983SAN/BAR][1984NRI2][1984SVE][1985BABMAT][1985GAL][1985MUL][1986FLE/JOH][1986FLE/JOH][1987BROAVAN][1988PHI/HAL][1989SMI/MAR][1989SMI/MAR][1989SMI/MAR][1989SMIMAR1

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT = 298.15 K, 1=0T= 298.15 K,I=n/aT=298.15 K, I=n/aT=298.15K,I=0T = 298.15 K, 1=0,T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT = 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=1

001

0

00

0

00.10.3

1

182

JNC TN8400 9 9 - 0 1 1


log Kso:

19.0016.7718.9319.04 2

19.04 2

16.7218.817.4618.7816.7219.1316.9719.0016.0819.0516.24

3Pb2+ + 2CO32- + 2H2O

[1952LAT][1965GAR7CHR][1971NAU/RYZ][1976BAE/MES][1976SMI/MAR][1977PAU][1979PAT/OBR][1980MAN/DEU][198OSCH][1982PAU][1983LAN][1983SAN/BAR][1985GAL][1986FLE/JOH][1988PHI/HAL][1989SMI/MAR1

<=> Pb3(C03)2Pb(OH)2(hydrocerrusite) -

T= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T = 298.15 K, 1=0T = 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K,I=n/a

-. T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0.3

I-2//+

00

000

0.3


log Kso: 10Pb2+ + 6CO32- + 7H2O <=> Pbw(CO3)6(0H)60(plumbonacrite) + 8H+

8.76 [1989SMI/MAR1 T= 298.15 K, 1=0 0_

log Kso: 2Pb2+ + CO32- + H2O t

0.81-17.280.81-17.280.480.810.490.780.81

[1952LAT][1977BAR/KNA][1977PAU][1979KUB/ALC][1980BEN/TEA][1982PAU][1982WAG/EVA][1985BAB/MAT][1985GAL]

* PbCO3PbO(s)

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,l=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/a

log Kso: 3Pb2+ + CO32' + 2H2O <=> PbCO3(PbO)2(s) + 4H+

-10.91-10.91-10.99-11.04-11.09

log Kso:

19.8719.8019.8019.8019.95

[1952LAT][1977PAU][1982PAU][1985BAB/MAT][1985GAL1

2Pb2+ + CO32- + 2CI- c

[1971NAU/RYZ][1980BEN/TEA][1980MAN/DEU][1982WAG/EVA][1988PHI/HAL]

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT=298.15K, I=n/aT= 298.15 K,I=n/a

=> Pb2CO3Cl2(phosgenite)

T= 298.15 K,I=n/aT= 298.15 K, I=n/aT=298.15K, 1=0T= 298.15 K,I=n/aT= 298.15 K, I=n/a

1 calculated with linear free energy relation from the Afi° value of Pb2+ by [1984SVE]2 calculated with a log P u of - 28.02 (Pb2+ + 3 H2O <=> Pb(OH)3- + 3H+)

183

JNC TN8400 9 9 - 0 1 1

6. 7 Lead nitrate system

Lead forms weak complexes with nitrate. Several measurements of the formation constants canbe found in the literature. Experimental data for the formation of lead nitrate complexes used inthis report for extrapolation to I = 0 are given in Table 6.18. Further experimental resultsreported in the literature are shown in Table 6.19 and log (3 values proposed by authors ofprevious reviews are collected in Table 6.20.

Table 6.18: Experimentally determined equilibrium data compiled for the lead nitrate system,according to the equilibrium: mPb2+ + nNO3" <=> Pbm(NO3)n

2 m-n . These datawere chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 6.13: 'Commentson selected references'. Method: kin = kinetic measurements, pol = polarography,pot = potentiometry, sol = solubility measurements, and sp = spectrophotometry.

log Pm.n

log pu:

0.450.520.200.25 1. 2

0.31 1.2

0.36 1.2

0.82 3

0.77 3

0.70 3O.7O3

0.33 3

0.23 3

0.23 3

0.193

0.15 4

0.26 4

0.530.360.430.400.450.52

Reference

Pb2++ NOf d PbNO3

[1953HER/SMI][1953HER/SMI][1953HER/SMI]

[1955BIG/PAR][1955BIG/PAR]

[1955BIG/PAR][1956BAL/DAV][1956BAL/DAV][1956BAL/DAV][1956BAL/DAV][1956BAL/DAV][1956BAL/DAV]

[1956BAL/DAV][1956BAL/DAV][1965HUG2][1965HUG3][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB]

Comments

+

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, 1=2K, 1=2K, 1=2K, 1=0.5K, 1=1

K, 1=2K, 1=0.051K, 1=0.066K, 1=0.079K, 1=0.092K, 1=0.624K, 1=0.943

K, 1=1.168K, 1=1.455

K, 1=2K, 1=2K, 1=0.5K, 1=0.75K, 1=1K, 1=2K, 1=3K, 1=3

KM)

222

0.51

20.0510.0660.0790.0920.6240.943

1.168

1.45522

0.50.75

1233

Medium

NaClO4

NaC104

NaC104

HC1O4

HC1O4

HC1O4

NaC104

NaClO4

NaClO4

NaC104

NaC104

NaClO4

NaClO4

NaClO4

NaClO4

NaClO4

LiClO4

LiClO4

LiClO4

LiClO4

LiClO4

LiClO4

Method

polpotspspsp

spspspspspspsp

sp

sppotpolpolpolpolpolpolpol

184

JNC TN8400 9 9 - 0 1 1


0.53 5

0.57 5

0.46 5

0.62

log P12:

0.39 4

0.38 4

0.430.480.430.230.410.450.48 5

0.48 5

0.30 5

log p]3:

-0.050.080.260.30 5

0.18 5

[1972FED/ROB][1972FED/ROB][1972FED/ROB][1973HUT/HIG]

T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K,I=3T= 298.15 K, 1=1

Pb2+ + 2NO3-<^> Pb(NO3)2°

[1965HUG2][1965HUG3][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB]

T= 298.15 K, 1=2T= 298.15 K, 1=2T= 298.15 K, 1=0.5T= 298.15 K, 1=0.75T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3

Pb2+ + 3NOf <=> Pb(NO3)f

[1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB][1972FED/ROB]

T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3

3331

22

0.50.75

1233333

23333

LiClO4

LiC104

LiC104

NaClO4

NaC104

NaC104

LiC104

LiClO4

LiC104

LiC104

LiC104

LiC104

LiC104

LiClO4

LiC104

LiClO4

LiC104

LiClO4

LiC104

LiC104

potspsolkin

potpolpolpolpolpolpolpolpotspsol

polpolpolpotsol

log K*°so: Pb2+ + NO3 + H2O <=> PbOHNO3(s) +

-3.55 6 [1945PED] T= 291 K, 1=0.3 0.3 Ba(NO,)7 sol50-600 mM Pbcorrected to 1=0 with the Davies equation by the respective authorsoriginal values corrected with Daviesby [1956BAL/DAV]; corrected back in this report0.1-100 m M P bput together by [1972FED/ROB] from earlier measurements of Fedorov and Coworkers.5-400 mM Pb

185

JNC TN8400 99-011

6.7.1 Lead nitrate complexes

Extrapolation of the measurements compiled in Table 6.18 is shown in Figures 6.24, 6.25 and6.26:

3NO3

PbNO3+

Pb(NO3)2°Pb(NO3)3_0

log P ° u = 1.06log P°i.2= 1-48log P°i,3 = 0.76

For the reaction Pb2+ + 4NO3" O Pb(NO3)42" no log §°\£ is proposed in this report as the

correction with the SIT term of values measured at I = 3 to I = 0 is not possible without makinguncertain assumptions. The complex Pb(NO3)4

2" is important only in very concentrated nitratesolutions, which makes the extrapolation to 1=0 difficult and the occurrence of Pb(NO3)4

2" notvery likely under environmental conditions.

Pb2+ + NO PbNCV

PbNO3+ at 25 °C.Figure 6.24: Plot of log (3U + 4 D vs. Im for the reaction : Pb2+ + NO3-

The straight line shows the result of the linear regression: Ae = - 0.09; log1.06. Calculated from data compiled in Table 6.18.

186

JNC TN8400 9 9 - 0 1 1

4.5 -

4-Q 3 . 5 -

+ 3CM

^2.5

g) 21 . 5 •

1

0.5

0 -

Pb2+ + 2NO3" <:; - • • . - . : . • . • .-:•„'«.:;y., { " . - - • ! • ; • • < • ; : • . - . ; V •'•—..i.-.v . : • : • . : ! • : > ! . : • • . : : : • : :

• . . • ; . • . : " : : • • . • - ; ' • : : ; : ;:

\~:..^,: '^r-i'~ - - ; • • ; • [>

-'':'::'.i:"t:-[::\' •'-.

'.' -.'.'•'".'''•1"!^':;.;:'[h1'::'!::!îi;i:i!:i;:!i*iii;ii;Diii-!':-"^'-;--:r:"^;i^î:'i

" • • • • . : : . " - . • ; ; ;

/ • • . • : " ; • . I ' -1

' : " ! ; ' ; - ' :

! = " ::

. : ; ; i " i:

- : - ; * - v ': • • " ^ • •

:"

: • " • • : - . • : : i . - i v

••'• . . - • : • • ; • ; . : ' • : • - •" ' "-: : - ' • . - : ^ ; ^ ^ = - : j : " : v : : i t ?:.•;•?•«::<:j••!-.<;. ' /• :•": ;

•?: -f S'L? S ? ?:•;"; :3- . :?R; '«S; ' : : .TP:H|

. : : ; - • ' • • • • " • . . • : : • : . ] • : : • ^ ":

- " ' ":

":

' i1

' ^ ^ ^ ; : V : . • : . ; • : ' : • : . : : ' . . : : ;

v'T!~^.1":::i''\/;;'LASf;;--:;i::!r«î*i:r:A|::ii:"|:t!;i;S-. :

:' . •••• • • : • ; ; : • ! ; • ! « ; = : i r v i " , - - . : - : - - '

1.

1; - . : / ; - ' ' : -

1! - - : ?

l:-^-:1.p1:S:;i;S'll:B|:;;l;:s;!w

* Pb(NO3)2°

l:fl'-:Si;V;y:--:.t': ::i:-;fi:r7.;:.•:•,;:•:«*.;; :..:.».:A::::-.':-:V;:;:^:::;;,;P.r

1 2lm, molal

Figure 6.25: Plot of log p1>2 + 6 D vs. Im for the reaction : Pb2+ + 2NO3- o Pb(NO3)2° at 25°C. The straight line shows the result of the linear regression: Ae = -0.09; logP°i,2 = 1-48. Calculated from data compiled in Table 6.18.

Pb2+ + 3NO3- <=> Pb(NO3)3-

Figure 6.26: Plot of log p]>3 + 6 D vs. Im for the reaction : Pb2+ + 3NO3" <=> Pb(NO3)3" at 25°C. The straight line shows the result of the linear regression: Ae = - 0.22; logP°j 3 = 0.76. Calculated from data compiled in Table 6.18.

187

JNC TN8400 9 9 - 0 1 1

6. 7.2 Pb(NO3)2(s) and PbOHNO3(cr)

The high solubility of Pb(NO3)2(s) in water is well established and is given in the review of[1980CLE/JOH] as 1.80 mol/kg H2O (in agreement with e.g. [1935AKEATUR] and[1960KAZ]). Assuming the predominance of Pb(NC>3)20 and 1=0, this value can be correctedfor the formation of Pb nitrate complexes, which gives a tentative log K*so of -0.28. In solutionwith pH > 3, however, [1945PED] has shown the precipitation of PbOHNO3(cr) with X-raydiffraction. The only experimentally determined value is by [1945PED] in 0.3 M nitratemedium. Extrapolation to 1=0 with the SIT model (assuming Ae ~ 0) results in:

NO3-+H2O PbOHNO3(cr) + H+ log K*°so= - 2.94

6. 7.3 Additional data compiled for the lead nitrate system

Table 6.19: Additional experimentally determined data for the lead nitrate system, according to theequilibrium: mPb2+ + nNO3" <=> Pbm(NO3)n

2m'n. These data were not chosen in the present reportfor the evaluation of recommended stability values. Reasons for not selecting these references aregiven in Section 6.13: 'Comments on selected references'. Method: con = conductivitymeasurements, pol = polarography, pot - potentiometry, sol = solubility measurements, and sp= spectrophotometry.

log Pm,n

log pu: Pb

1.19 l

0.52 2

1.19 3

1.15 3

1.11 4

Reference

'2+ + NO3 <=> PbNO3+

[1930RIG/DAV][1953HER/SMI][1955NAN][1956BAL/DAV]F1972FED/ROB1

Comments

T= 298.15 K, 1=0.001-0,02T= 298.15 K, 1=6T= 298.15 K, 1=0.001T= 298.15 K, 1=0.9-1.5T= 298.15 K, 1=0

KM)

06000

Medium

NaClNaC104

NaC104

NaClO4

LiC104

Method

consp

consppol

log j5u: Pb2+ + 2N0f <=> Pb(NO3)2°

1.40 4 [1972FED/ROB1 T= 298.15 K, 1=0 LiClO4 pol

log pu: Pb2+ + 3NOf <=> Pb(NO3)f

-2.30 5 [1965HUG31 T= 298.15 K, 1=2 NaCIO, pol

log j3l4: Pb2+ + 4NOf <=> Pb(NO3)42-

-0.30 [1972FED/ROB]0.11 6 [1972FED/ROB]-0.52 6 [1972FED/ROB]

T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=3

LiClO4

LiClO4

LiClO.

polpotsol

JNC TN8400 9 9 - 0 1 1


log K*°so: Pb2+ + 2NO3 <=> Pb(NO3)2(s)

-0.28 7 [1935AKE/TUR] T=298.15 K, I=dil 0 water sol-0.28 7 [1960KAZ] T=298.15 K, I=dil 0 water sol^1 corrected to 1=0 with a simplified Debye-Hiickel equation2 not used because of high ionic strength3 corrected to 1=0 with Davies equation.4 corrected to 1=0 with Debye-Huckel equation5 tentative value6 put together by [1972FED/ROB] from earlier measurements of Fedorov and Coworkers.7 after correction for the formation of Pb nitrate complexes using the constants calculated in Section 6.7.1.

Table 6.20: Thermodynamic data for the lead nitrate system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

log Pm n Reference Comments I (M)

log hi

1.170.250.330.400.511.181.151.20

: Pb2* + NOf <=> PbNO3*

[1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1980BEN/TEA][1982WAG/EVA][1987BROAVAN]

y

T=T=T=T=T=T=T=

298.15298.15298.15298.15298.15298.15298.15298.15

K, 1=0K, 1=0.5K, 1=1K, 1=2K,I=3K, 1=0K, 1=0K, 1=0

00.5

123000

log (lu: Pb2+ + 2NO3- <=> Pb(NO3)2°

00.5

123

0.83 [1987BROAVAN1 T= 298.15 K, 1=0 0

log pu: Pb2* + 3NOf <=> Pb(NO3)f

0.10 [1976SMI/MAR] T= 298.15 K, 1=2 20.20 [1976SMI/MAR] T= 298.15 K, 1=3 3-0.86 ri987BRO/WAN] T= 298.15 K, 1=0 0_

log $lA: Pb2* + 4NOf <=> Pb(NO3)42-

-0.30 [1976SMI/MAR] T= 298.15 K, 1=3 3-3.76 [1987BROAVAN] T= 298.15 K, 1=0 0

1.400.400.400.400.40

[1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR]

T=T=T=T=T=

298.15298.15298.15298.15298.15

K,I=0K, 1=0.5K,I=1K, 1=2K, 1=3

189

JNC TN8400 9 9 - 0 1 1


log Ktoso: Pb2* + 2NO/ <=> Pb(NO3)2(s)

1.09 [1952LAT] T= 298.15 K, I=n/a1.09 [1985GAL] T= 298.15 K, I=n/a-0.28 ' [1980CLE/JOH1 T= 298.15 K, I=n/a

log K*°so: Pb2+ + NO/ + H2O <=> PbOHNO3(s) + H+

-12.01 [1952LAT] T= 298.15 K, I=n/a-12.03 fl985GAL1 T= 298.15 K, I=n/aI after correction for the formation of Pb nitrate complexes using the constants calculated in Section 6.7.

190

JNC TN8400 9 9 - 0 1 1

6.8 Lead phosphate system

Lead forms complexes with phosphate. Several measurements of the formation constants can befound in the literature. Experimental data for the formation of lead phosphate complexes andsolids used in this report are given in Table 6.21. Further experimental results reported in theliterature are shown in Table 6.22 and thermodynamic values proposed by authors of previousreviews are compiled in Table 6.23.

Table 6.21: Experimentally determined equilibrium data compiled for the lead phosphatesystem, according to the equilibrium: mPb2+ + nPO4

3" <=> Pbm(PO4)n2m-3n. These

data were chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 6.13: 'Commentson selected references'. Method: sol = solubility measurements.

Pm,nReference Comments I (M) Medium Method

log plu: Pb2+ + PO43- + H+ <=>PbHPO4°

15.45 1 [1972NRI] T= 298.15 K, I=dil 0 H,PO4 sol

log fiI21: Pb2+ + PO43' + 2H+

21.05 ] [1972NRI] T= 298.15-K, I=dil 0 H,PO4 sol

log K*so-Pb2+ + PO43- + H+ <=>PbHPO4(secondary lead phosphate)

23.78 1 [1972NRI] T= 298.15 K, 1=0.1 0 KOH sol

log K*so: 3Pb2+ + 2PO43' <=> Pb3(PO4)2(tertiary lead phosphate)

44.40 ! [ 1972NRI] T= 298.15 K, 1=0.1 0 KOH sol

log K*so: Pb2+ + 2PO43- + 4H+ <=> Pb(H2PO4)2(primary lead orthophosphate)

48.94 J [1973NRI2] T= 298.15 K, 1=0.1 () NaCl sol

log K*S0:4Pb2+ + 2PO43- + H2O <^ Pb4(PO4)2O(tetraplumbite orthophosphate) + 2H+

37.09 ' [1972NRI] T= 298.15 K, 1=0.1 0 KOH sol

191

JNC TN8400 9 9 - 0 1 1


log K*S0:5Pb2+ + 3PO43- + H2O <=> Pb5(PO4hOH(hydroxypyromorphite)+ H+

62.80 * [1972NRI] T= 298.15 K, 1 = 0 . 1 0 K O H s o l

log K*so: 5Pb2+ + 3PO43- + Cl- <=> Pb5(PO4)3Cl(s) (chloropyromorphite)+ H+

84.4 ! [1973NRI2] T= 298.15 K, 1=0.1 0 NaCl sol

log K*S0:5Pb2+ + 3PO43- + F- <=>Pb5(PO4)3F(s) (fluoropyromorphite)+ H+

71.60 * [1973NRI1] T= 298.15 K, 1=0.1 0 NaF sol1 extrapolated to 1=0 with Davies equation by (1972NRI], [1973NRI1] and [1973NRI2].

6.8.1 PbH2PO4+ and PbHPO4°

Nriagu [1972NRI] made measurements in dilute phosphoric acid solution and determined theconstants for the formation of the aqueous species PbH2PO4+ and PbHPO4°. These values arethe only experimentally determined values available and a correction to 1=0 with the Daviesequation has been made by [1972NRI].

Pb2+ + PO43" + H+ <=> PbHPO4° log p \ u = 15.45

Pb2+ + PO43- + 2H+ <=» PbH2PO4+ log p°i'2,i = 21.05

6.8.2 PbHPO4(cr) and Pb3(PO4)2(s)

The log K*so for PbHPO4(secondary lead phosphate) and for Pb3(PO4)2(tertiary leadphosphate) given in [1972NRI] (Table 6.20) agree very well with the values determined at 310Kby [1932JOW/PRI] and by [1969AWA/ELH] in 1 M phosphate solution (compiled in Table6.21). As only Nriagu determined log K*so values for several different lead phosphates, it isrecommended to use the consistent dataset given by Nriagu:

Pb2+ + PO43- + H+ <=> PbHPO4(secondary lead phosphate) log K*°So = 23.78

3Pb2+ + 2PO43- 4=> Pb3(PO4)2(tertiary lead phosphate) log K*°So = 44.40

[1972NRI] states that secondary and tertiary lead phosphate, however, are not stable in thepresence of water at ambient temperatures but will transform to hydroxypyromorphite or otherlead phosphates.

192

JNC TN8400 9 9 - 0 1 1

6.8.3 Pb(H2PO4)2(s), Pb4(PO4)2O(s) and plumbogummite(PbAl3(PO4)2(OH)5(cr))

The following values have been determined by [1972NRI] and [1973NRI2]:

4Pb2++ 2PO43-+ H2O & Pb4(PO4)2O(s) + 2H+ log K*°so = 37.09

tetraplumbite orthophosphate

Pb2 ++ 2PO43" + 4H+ <=> Pb(H2PO4)2(s) log K*°so = 48.94

primary lead orthophosphate

Again it is recommended to use the internally consistent dataset given by Nriagu. The log K*°soof 99.3 given by [198NRI] (Table 6.23) for plumbogummite (PbAl3(PO4)2(OH)5(cr)),however, is only an estimate.

6.8.4 Pyromorphites: Pb5(PO4)3Cl(s), Pb5(PO4)3F(s), and Pb5(PO4)3OH(s)

Again, not many authors have determined log K*so values for pyromorphites and it isrecommended to use the internally consistent dataset given by Nriagu (Table 6.21):

5Pb2++ 3PO43-+ H2O <=> Pb5(PO4)3OH(cr) + H+ log K*°so = 62.8

hydroxypyromorphite

5Pb2++ 3PO43-+ Cl- <=> Pb5(PO4)3Cl(cr) log K*°s0 = 84.4

chloropyromorphite

5Pb2++ 3PO43" + F- « Pb5(PO4)3F(cr) log K*°so = 71.6

fluoropyromorphite

The log K*so of 84.4 for chloropyromorphite (Pb5(PO4)3Cl(cr)) determined in [1973NRI2]after an aging time of 48 h agrees well with the value determined at 310 °K by [1932JOW/PRI]after a few hours (Table 6.22). [1964BAK] determined after an equilibration time of severalmonths a much lower log K*so of 34.5 for chloropyromorphite (Table 6.22). An examinationof his data leads to the conclusion that [1964BAK] assumed all phosphate to be present asPO4

3" (in a Na2HPO4 solution). Thus, the value given by [1964BAK] was not used in thisreport.

6.8.5 Pb10(PO4)6(OH)2(hydroxylapatite)

For lead hydroxylapatite (Pbi0(PO4)6(OH)2(cr)), [1976RAO] reported a slightly highersolubility than for calcium hydroxylapatite which leads to the conclusion that the precipitation oflead hydroxylapatite can not be expected in the presence of calcium. However, the formation ofmixed lead-calcium hydroxylapatite has also been observed [1976RAO].

193

JNC TN8400 9 9 - 0 1 1

6.8.6 Additional data compiled for the lead phosphate system

Table 6.22: Additional experimentally determined data for the lead phosphate system, according to theequilibrium: mPb2+ + nPO4

3" <=> Pbm(PC>4)n2m"3n. These data were not chosen in the present

report for the evaluation of recommended stability values. Reasons for not selecting thesereferences are given in the text in Section 6.8. Method: pot = potentiometry, sol = solubilitymeasurements.

Reference Comments KM) Medium Method

log K*S0: Pb2+ + P04}- + //+<=> PbHPO4(s)

22.28 ' [1929MIL/JOW]22.20 ' [1929MIL/JOW]23.71 2 [1932JOW/PRI]23.75 [1969AWA/ELH1

T=310K, I=dilT= 298.15 K, I=dilT= 310 K, 1=0.02T= 303 K, I=dil

0000

PbCl2> Na2HPO4

PbCl2, Na2HPO4

NaClH,PO4, NaH,PO,

solsolsolpot

log K*S0:3Pb2+ + 2PO43- <=> Pb3(PO4)2(s

42.0 ' [1929MIL/JOW]42.1 ' [1929MIL/JOW]43.53 2 [1932JOW/PRI]42.03 ri969AWA/ELH1

T=31OK,I=dilT= 298.15 K,I=dilT= 310 K, 1=0.02T= 303 K, I=dil

0000

PbCl2, Na2HPO4

PbCl2, Na2HPO4

NaClH,PO4, NaH,P04

solsolsolpot

log K*S0:5Pb2+ + 3PO43' + Cl- <=> Pb5(PO4)3Cl(s) (chloropyromorphite)+ H+

79.12 2 [1932JOW/PRI]34.50 3 H964BAK]

T= 310 K, 1=0.16T= 298.15 K, 1=0.008-0.028

00

NaCl solPbCl,, Na,HP04 sol

1 extrapolated to 1=0 with Debye-Huckel equation; later criticized by [1932JOW/PRI]2 extrapolated to 1=0 with Debye-Hiickel equation3 given by [1964BAK] assuming equilibria with PO4

3' in a Na2HPO4J solution

194

JNC TN8400 9 9 - 0 1 1

Table 6.23: Thermodynamic data for the lead phosphate system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison.

l og |3m n Reference Comments I (M)

log P,.u: Pb2

15.4515.2415.4523.2816.59

H+ <=> PbHPO4°

[1976SMI/MAR][1983LAN][1984VIE/TAR][1985GAL][1987BRO/WAN1

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K, 1=0

000

log Pi.2.2- Pb2+ + 2P0J- + 2H+ <=> Pb(HPO4)22

32.52 ri987BROAVANl T= 298.15 K, 1=0

log PUJ: Pb2+ + 3PO43- + 3H+ <=> Pb(HPO4)/-

48.02 [1987BRO/WAN] T= 298.15 K, 1=0

log P,,4.4: Pb2+ + 4PO/- + 4H+ & Pb(HPO4)46'

63.21 n987BRO/WANl T= 298.15 K, 1=0

log Pu.i: Pb2+ + PO43- + 2H+ PbH2PO4

+

21.0519.9021.0520.89

[1976SMI/MAR][1983LAN][1984VIE/TAR]ri987BROAVANl

T=

T=

298.298.298.298.

15151515

K,K,K,K,

1=01=01=01=0

0000

log P,.4.2: Pb2+ + 2PO/- + 4H+ <=> Pb (H2PO4)2°

41.09 ri987BRO/WANl T= 298.15 K, 1=0

log P,.6j- Pb2+ + 3PO/- + 6H+ <=> Pb (H2PO4)3-

60.85 f!987BROAVAN1 T= 298.15 K, 1=0

log P,,8.<: pb2+ + 4PO4

3' + 8H+ <=> Pb (H2PO4)42-

80.27 [1987BROAVAN1 T= 298.15 K, 1=0

195

JNC TN8400 9 9 - 0 1 1


log K*S0: Pb2+ + PO43- + H+ <=> PbHPO>4( secondary lead phosphate)

23.4126.8223.7823.7826.7823.7823.28

[1952LAT][1971NAU/RYZ][1976SMI/MAR][1980CLE/JOH][1982WAG/EVA][1984VIE/TAR][1985GAL1

T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

0

0

log K*S0:3Pb2+ + 2PO43' <=> Pb3(PO4)2(tertiary lead phosphate)

5443.5342.0844.444.4043.4142.18

[1952LAT][1976SMI/MAR][1977TAR/VIE][1980CLE/JOH][1984VIE/TAR][1985GAL][1988PHI/HAL1

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/a

00

log K*so.- Pb2+ + 3Al3+ + 2PO43'

99.30 ' [1984NRI2]

5H2O <^> PbAl3(PO4)2(OH)5(plumbogummite) + 5H+

T= 298.15 K, I=n/a

log K*S0:5Pb2+ + 3PO/- + H2O « Pb5(PO4)3OH(hydroxypyromorphite) + H+

62.80 [1980CLE/JOH]62.80 [1984VIE/TAR1

T= 298.15 K, I=n/aT= 298.15 K, I=n/a

log K*sa-5Pb2+ + 3PO43- + Ci <=> Pb5(P04)3Cl(chloropyromor,phite)+ H+

9.1179.1284.4584.484.40 '84.40

[1971NAU/RYZ][1972NRI][1976HEM][1980CLE/JOH][1984NRI2][1984VIE/TAR1

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/a

log K*S0:5Pb

71.6

2+ + 3PO43- + F- <=> Pb5(P04)3F(fluoropyromorphite)+ H+

ri980CLE/JOH] T=298.15 K, I=n/a[1984NRI2] claims that chloropyromorphites and plumbogummite or mixtures of them precipitate in soils

196

JNC TN8400 9 9 - 0 1 1

6.9 Lead sulfate system

Lead(II) forms complexes with sulfates in aqueous solutions and several authors determinedstability constants (Table 6.24). Also, a number of solid lead sulfate compounds are known toexist. The data used in this report for the calculations of the formation constants of lead(II)sulfate complexes are compiled in Table 6.24 and shown in Figure 6.27 and Figure 6.28.Additional data for the lead(II) sulfate system that were not selected for the calculation of log (3°values in this report are compiled in Table 6.25 and 6.26.

Table 6.24: Experimentally determined equilibrium data compiled for the lead sulfate system,according to the equilibrium: mPb2+ + nSO4

2" <=> Pbm(SO4)n2m-2n. These data

were chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 6.13: 'Commentson selected references'. Method: con = conductivity, el = electrophoresis, pol =polarography, pot = potentiometry, sol = solubility measurements

log Pm,n

log fa ,:

2.072.70 12.75 l

0.700.741.051.20 2

log Pj 2:

1.982.001.18

Reference Comments

Pb2+ + SO42-& PbSO4°

[1969DRY/TVA][1970GAR/NAN][1970GAR/NAN][1971BON1][1972BON][1982ROH][1989NYHAVIK]

Pb2+ + 2SO42-<^Pb(SO4)2

[1971BON1][1972BON][1982ROH]

T=298T=298T=298T=298T=298T=296T=298

2-

T=298T=298T=296

K,K,K,K,K,K,

.15

K,K,K,

1=0.2I=dilI=dil1=31=31=0.7K, 1=1

1=31=31=0.7

KM)

0.20033

0.71

33

0.7

Medium

NaClO4/Na2SO,NaC104

noNaC104

NaC104

NaC104

NaClO4

NaC104

NaClO4

NaClO4

Method

! SOl

potconpolpolel

pol

polpolel

197

JNC TN8400 9 9 - 0 1 1

log Kso: Pb2+ + SO42-<=> PbSO4(anglesite)

7.647.70 3

6.207.037.76 4

[1942KOL/PER][1959JAG][1961RAM/STE][1969DRY/IVA][1970LIT/NAN]

T= 298 K, I=dilT= 298 K, I=dilT= 298 K, 1=1T= 298 K, 1=0.2T= 298 K, I=dil

5E-0401

0.20

noHNO3

NaC104

NaClO4/Na2SO4

no

solsolsolsolpot

1 extrapolated to 1=0 with Davies equation, measured at 1=0.001 M2 0.0001-0.1 raM Pb; mean of 6 measurements3 measured at 1=0.001 M, extrapolated to 1=0 by [ 1959JAG] as a function of 1°5

4 measured at 1=0.001 M, extrapolated to 1=0 with Davies equation

6.9.1 Lead sulfate complexes

Several values of stability constants for PbSC>40 are reported in literature (Table 6.24), thesedata show a good agreement. Extrapolation of these measurements to 1=0 with the SIT methodis shown in Figure 6.27:

2SO42-

PbSO4°Pb(SO4)2

2-log (3°u= 2.82log p°lt2= 2.37

For Pb(SO4)2" the log (3 2 value of 2.37 (Figure 6.28) is calculated from only two independent

measurements ([1971BON1, 1972BON], 1=3 and [1982ROH], 1=0.7).

Q00

COL

D)JO

4.5 -

4 -

3.5 •

3 •

2.5 •

2 •

1.5 •

1 -

0.5

n -

-

o

) .

Pb2 +

5 _\-Y

0, - o; -

+ SO42" <:

'\__7V.-.- „->,»

* PbSO4°

;:V; •;. V*;/. :

^——t ^ -

1 2lm> molal

Figure 6.27: Plot of log p1(1 + 8 D vs. Im for the reaction : Pb2+ + SO42- <=> PbSO4° at 25 °C.

The straight line shows the result of the linear regression: Ae = - 0.01; log (3° \t\ =2.82. Calculated from data compiled in Table 6.24.

198

JNC TN8400 9 9 - 0 1 1

Pb2+ + 2SO42" <=> Pb(SO4) :

2"a •

4 5 •

4 •

+ 3 •

j? 21.5

1

0.5 •

0

: - [ z "1 • Tb:;! J;=i I! i > • : I • H~f:; L : Ti - - i'; M [J: =;-1" ;1 J n - MTr 1:1! I :fii :=^* 1 ;^ :^ lÊ ' ; ! : M i IWiii I=! 1 U1 Ij 1= ! i^; J ! i ' : i; - -f! / i \-^V-CfW^^V<^\jAf^' •-•

'•"•: • '''•'•'":::'r':;:':-:;*::p:}r'• •' : j : ' : :"' : | : 1 ; - ! j ' " : - w : i f 'H• • ' ' - ' : • v t ! : V : ^ T ' l ^ y S f f •>•-:-r ' :-

lm, molal

Figure 6.28: Plot of log pi,2 + 8 D vs. Im for the reaction : Pb2+ + 2SO42" <=> Pb(SO4)2

2- at 25°C. The straight line shows the result of the linear regression: Ae = -0.43; logP°i>2 = 2.37. Calculated from data compiled in Table 6.24.

6.9.2 PbSO4(anglesite)

A number of solid lead sulfate species are reported in different compilations (Table 6.26). Mostof these solids form only at higher temperatures, however [1989KEL, 1995MAR/MAC].[1995MAR/MAC] showed that only anglesite (PbSO^cr)) precipitates from aqueous solutionsat room temperature. From several solubility measurements, a log K*so of 7.81 can becalculated (Figure 6.29), a value which is in excellent agreement with the value of 7.80 givenby [1995MAR/MAC].

PbSO4(s) log K*°so= 7.81

199

JNC TN8400 9 9 - 0 1 1

oo

+

i nI U

9.5

9

8.5

8t

7 •

6.5

6

5.5

fi

Pb2 +

., o -"

+ SO42- «

, • • -

y = 0;09x

\ — •

^ PbS04(s)

^ . •• v • "

+ 7.81 ' J;'

1 _ -

0.5 1lmi molal

1.5

Figure 6.29: Plot of log K*So + 8 D vs. Im for the reaction : Pb2+ + 2SO42- <=> PbSO4(s) at 25

°C. The straight line shows the result of the linear regression: Ae = - 0.09; logK*°so = 7.81. Calculated from data compiled in Table 6.24.

6.9.3 Hinsdalite: PbAl3PO4SO4(OH)6(cr)

The values given in different compilations for hinsdalite (PbAl3PO4SO4(OH)6(cr)) (compiled inTable 6.26), spread over a wide range and are based either on an estimate [1984NRI2] or on themeasurements by [1964BAK]. [1964BAK] (Table 6.25) determined a log K*So of 2.11 forhinsdalite using an equilibration time of several months. However, a closer examination of hisdata led to the conclusion that [1964BAK] assumed all phosphate to be present as PO43- (at apH of 3). This value is therefore not used in this report and no log K*so for hinsdalite can beproposed in this report.

200

JNC TN8400 9 9 - 0 1 1

Table 6.25: Additional experimentally determined data for the lead sulfate system, according to theequilibrium: mPb2+ + nSO4

2" <=> Pbm(SO4)n2m"2n. These data were not chosen in the present

report for the evaluation of recommended stability values. Method: sol = solubilitymeasurements, SC>2= pSC>2 measurements at high temperatures.

log Pm,n Reference Comments Medium Method

log Pu: Pb2+ + SO42-<=> PbSO4°

2.62 ' [1969DRY/TVA] T= 298 K, 1=0 0 NaClO4/Na?SO4 sol

log Kso: Pb2+ + SO42-<=> PbSO4(s, anglesite)

7.82 '7.2415.40 2

7.80 3

7.80 4

[1969DRY/TVA][198OWAL/SIN][1989KEL][1993MAC/PAG][1995MAR/MAC1

T= 298 K, 1=0T= 298 K, I=dilT= 298.15 K,I=n/aT = 298.15 K, 1=0T = 298.15 K, 1=0

0

00

NaC104/Na2S04H2SO4

n/awaterwater

solsolSO,solsol

log Kso: 2Pb2+ + SO42-+ H2O <=> Pb2OSO4(s, lanarkite) +2H+

5.83 2 [1989KEL] T= 298.15 K, I=n/a n/a SO,

log Kso: 3Pb2+ + SO/-+ 2H2O <=> Pb3O2SO4(s) + 4H+

-7.22 2 [1989KEL] T= 298.15 K, I=n/a n/a SO,

log Kso: 5Pb2+ + SO42-+ 4H2O <=> Pb5O4SO4(s) + 8H+

-29.61 2 [1989KEL] T= 298.15 K, I=n/a' n/a SO,

log Kso: Pb2+ + PO43- + SO4

2-

2.11 5 H964BAK]

3Al3+ + 6H2O <=> PbAl3PO4SO4(OH)6(s, hinsdalite) + 6H+.

T= 298 K, 1=0.008-0.028 PbCl,, Na,HPO4 sol' not indicated how extrapolation to 1=0 was made2 measurements at high temperatures; this value for PbSO4 may refer to a high temperature modification3 [1993MAC/PAG] compare lead solubility calculated with thermodynamic data with experimental lead solubility4 [1995MAR/MAC] compare lead solubility calculated with thermodynamic data with experimental lead solubility5 calculated by [1964BAK] assuming equilibria with PO4

3" at a pH of 3

201

JNC TN8400 9 9 - 0 1 1

Table 6.26: Thermodynamic data for the lead sulfate system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison. Medium: Where data refer to specific electrolytesolutions, this is indicated.

log Pm,n Reference Comments Medium

log Pu: Pb2+ PbSO4°

2.72.622.750.742.752.862.582.281.292.65

2.692.070.74

[1972ZIR/YAM][1976HEM][1976SMI/MAR][1976SMI/MAR][1981TURAVHI][1983LAN][1985BAB/MAT][1987BRO/WAN][1988BYRyKUM][1988PHI/HAL1

[1989SMI/MAR][1989SMI/MAR][1989SMI/MAR1

T= 298.15 K, 1=0T= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T=291 K, 1=0.7T= 298.15 K, 1=0

T=298.15K, 1=0T= 298.15 K, 1=0.1T= 298.15 K, 1=3

0

030000

0.70

00.13

sea water

log P12: Pb2+ + 2SO42-<=> Pb(SO4)2

2-

3.473.471.994.513.682.48

log Pi,.

4.45

,: Pb2+

[1972ZIR/YAM][1976HEM][1976SMI/MAR][1981TURAVHI][1987BRO/WAN]fl988BYR/KUMl

+ 3SO/'<=> Pb(SO4)/-

[1987BRO/WAN]

T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=3T= 298.15 K, 1=0 T=298.15K, 1=0 'T=291 K, 1=0.7

T= 298.15 K, 1=0

0

300

0.7

0

sea water

log p1A: Pb2+ + 4SO42-& Pb(SO4)/-

4.70 ri987BROAVAN1 T= 298.15 K, 1=0

log Ks0: F

7.507.877.757.757.877.796.20

V + + SO/-<=> PbSO4(anglesite)

[1952LAT][1968ROBAVAL][1969HEL][1971NAU/RYZ][1973BAR/KNA][1976SMI/MAR][1976SMI/MAR1

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=1

0

01

202

JNC TN8400 9 9 - 0 1 1


7.50 [1977PAU] T= 298.15 K, I=n/a7.88 [1978COD] T= 298.15 K, I=n/a7.82 [1978ROB/HEM2] T= 298.15 K, I=n/a7.86 [1979KUB/ALC] T= 298.15 K, I=n/a7.60 [1980CLE/JOH] T= 298.15 K, 1=0 07.83 [1980BEN/TEA] T= 298.15 K, I=n/a7.85 [1981STU/MOR] T= 298.15 K, I=n/a7.50 [1982PAU] T= 298.15 K, I=n/a7.83 [1982WAG/EVA] T= 298.15 K, I=n/a7.87 [1983LAN] T= 298.15 K, I=n/a7.70 [1984NRI2] T= 298.15 K, I=n/a7.50 [1985BAB/MAT] T= 298.15 K, I=n/a8.02 [1985GAL] T= 298.15 K, I=n/a7.85 [1988PHI/HAL] T= 298.15 K, I=n/a7.85 [1989COX/WAG] T= 298.15 K, I=n/a7.79 [1989SMI/MAR] T= 298.15 K, 1=0 07.03 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.16.20 [1989SMI/MAR1 T= 298.15 K, 1=1 1_

log Kso: 2Pb2+ + SO/-+ H2O <=> Pb2OSO4(lanarkite) + 2H+

9.34 [1952LAT] T= 298.15 K, I=n/a0.44 [1971NAU/RYZ] T= 298.15 K, I=n/a3.22 [1973BAR/KNA] T= 298.15 K, I=n/a0.42 [1977BAR/KNA] T= 298.15 K, I=n/a0.37 [1980BEN/TEA] T= 298.15 K, I=n/a0.37 [1983LAN] T= 298.15 K, I=n/a9.34 [1985GAL] T= 298.15 K, I=n/a0.50 [1988PHI/HAL] T= 298.15 K, I=n/a

log Kso: 3Pb2+ + SO42-+ 2H2O <=> Pb3O2SO4(s) + 4H+ '

-1.72 [1973BAR/KNA] T= 298.15 K, I=n/a-10.70 [1977BAR/KNA] T= 298.15 K, I=n/a-10.79 [1980BEN/TEA] T= 298.15 K, I=n/a-10.79 [1982WAG/EVA] T= 298.15 K, I=n/a

log Kso: 4Pb2+ + SO42'+ 3H2O «• Pb4O3SO4(s) + 6H+

-21.91 [1977BAR/KNA] T= 298.15 K, I=n/a-22.02 [1980BENATEA] T= 298.15 K, I=n/a-22.01 [1982WAG/EVA1 T= 298.15 K, I=n/a

lag Kso: 5Pb2+ + SO/ + 4H2O <=> Pb5O4SO4(s) + 8H+

-38.24 [1973BAR/KNA] T= 298.15 K, I=n/a

log Kso: Pb2+ + PO43- + SO4

2' + 3Al3+ + 6H2O <=> PbAl3PO4SO4(OH)6(hinsdalite) + 6H+

1.63 [1971NAU/RYZ] T= 298.15 K, I=n/a15.10 [1984NRI2] T= 298.15 K, I=n/a

203

JNC TN8400 9 9 - 0 1 1

6.10 Lead sulfide system

Lead forms complexes with sulfide in aqueous solutions, see Section 6.10.1. PbS(s), galena,has a very low solubility and often controls lead solubility in reducing environments. The dataused for the calculations of the formation constants of lead(II) sulfide complexes are compiledin Table 6.27. Additional data for the lead(II) sulfide system that are not chosen for thecalculation of log (3° values in this report are compiled in Table 6.28 and 6.29.

Table 6.27: Experimentally determined equilibrium data compiled for the lead sulfide system.These data were chosen for the evaluation of recommended values in the presentreport. Additional information for the different references see Section 6.13:'Comments on selected references'. Method: col = colorimetric analysis, sol =solubility measurements.

cr R Reference Comments I (M) Medium Method6 Fm,n

log Kia: PbS(galena) + HS' + H+ <=> Pb(HS)2°

0.13 l [1953HEM] T= 298.15 K, I=dil 0 n/a col0.21 ! [1979GIO/BAR] T= 303 K, 1=0.5-3 0 n/a sol_

log KJi3: PbS(galena)+ 2HS' + H+ <=>Pb(HS)3-

1.43 i [1953HEM] T= 298.15 K, I=dil 0 n/a col1.41 l [1979GIO/BAR] T= 303 K, 1=0.5-3 0 n/a sol_

log K*so: Pb2+ + HS-<=>PbS(s, galena) + H+

111111

.79

.95

.91

[1984UHL/HEL][1984UHL/HEL][1984UHL/HEL]

T=T=T=

298.298.298.

151515

K,K,K,

1=0.1=0.1=0.

160203

0.0.0.

160203

HHH

2S2S

?s

solsolsol

extrapolated by [the respective authors to 1=0 with extended Debye-Huckel equation

204

JNC TN8400 9 9 - 0 1 1

6.10.1 Lead sulfide complexes

Determination of stability constants for Pb(HS)2° and Pb(HS)3~ (with regard to PbS(s)) arereported by [1953HEM] and [1979GIO/BAR]. [1984UHL/HEL] recalculated the values of[1953HEM] and [1979GIO/BAR]. However, from the comments given by [1984UHL/HEL],their calculations are not traceable and are therefore not used in this report.

[1953HEM] and [1979GIO/BAR] corrected their results with the Debye-Hiickel equation to 1=0(Table 6.27). As the values are similar, the mean may be used: log K°]2 = 0.17 for the reactionPbS(galena) + HS" + H+ <^ Pb(HS)2°, and log K ° u =1.42 for PbS(galena) + 2HS" + H+ <=>Pb(HS)3-). Using the log K*°So of 12.17 calculated in Section 6.10.2, the following constantscan be calculated:

Pb2+ + 2HS- « Pb(HS)2° log (3O1)2 = 12.34

Pb2+ + 3HS- o Pb(HS)3- log p ° u = 13.59

6.10.2 Galena (PbS)

Data reported in the literature for the equilibrium Pb2+ + 2HS" <=> PbS(galena) + 2H+ are oftenbased either on calculations from AS and AH values measured at high temperature or on theselection made by [1953HEM] (Table 6.28 and 6.29). Experimentally determined values arescarce. The solubility product of galena (PbS) with respect to Pb(HS)2° and Pb(HS)3- can bedetermined quite precisely [1953HEM, 1956KIV/RIN]). The complex formation between leadand sulfide, however, is so strong, that direct determination of the equilibria Pb2+ + 2HS' <=>PbS(galena) + 2H+ is difficult. [1956KIV/RIN] tried to determine the Pb2+ concentrationdirectly via competitive complex formation with chloride ions. This method, however, is proneto errors as the complex formation with chloride is much weaker than the complex formation ofPb with HS\ The log K*so of galena was determined more precise by [1984UHL/HEL] thewith help of competitive complex formation of Pb with EDTA and TRIS. Extrapolation of thevalues [1984UHL/HEL] to I = 0 with the SIT equation is shown in Figure 6.30:

Pb2+ + HS- <=> PbS(galena) + H+ log K*°so= 12.17

205

JNC TN8400 99 - O i l

Pb2+ + HS- <=> PbS(galena) + H+

15

14.5

14

13.5 •;

0.2 0.4 0.6

L molal

Figure 6.30: Plot of log K*So + 4 D vs. Im for the reaction : Pb2+ + HS" <=> PbS + H+ at 25 °C.The straight line shows the result of the linear regression: As = - 0.82; log K*°so =12.17. Calculated from data compiled in Table 6.27.

6.10.3 Additional data compiled for the lead sulfide system

Table 6.28: Additional experimentally determined data for the.lead sulfide system. These data were not chosenin the present report for the evaluation of recommended stability values. Method: emf = emfmeasurements, pol = polarography, sol = solubility measurements.

log Pm.n

log K'SQ:

14.94 1

15.66 2

15.66 2

12.42 3

11.78 3

Reference

Pb2+ + HS'<=> PbS(galena)

[1956KTV/RIN][1977SHA][1981SHA/MIS][1984UHL/HEL][1984UHIVHEL1

Comments

+ H+

T= 298.15 K, 1=0.8-1.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.7

KM)

0000

0.7

Medium

HC1n/an/aH2SH,S

Method

polemfemfsolsol

1 extrapolated to 1=0 by [1956KIV/RIN] with Davies equation; calculated in this report using a log K(H2S/HS)= 6.99, log KH(H2S(g)/H2S(aq)) = 1.02measured at 633-698 K, AH, AS also determined.extrapolated by [1984UHL/HEL] to 1=0 and 1=0.7 with the Davies equation

206

JNC TN8400 9 9 - 0 1 1

Table 6.29: Thermodynamic data for the lead sulfide system taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

log K Reference Comments I (M)

log Klt2: PbS + HS- + H+ <=> Pb(HS)2°

-0.590.3 '-0.05 2

[1983LAN][1984UHL/HEL][1984UHL/HEL1

T= 298.15 K,T= 298.15 K,T= 298.15 K,

I=n/a1=01=0

000

log Ku: PbS + 2HS- + H+<=> Pb(HS)f

1.42 [1983LAN] T= 298.15 K, I=n/a1.55 2 ri984UHL/HELl T= 298.15 K, 1=0

log K*so: Pb2+ + HS-<=> PbS(galena) + H+

14.1214.5814.0614.6715.1915.1914.8213.60 3

14.8214.7014.8414.8115.70 4

15.1615.1614.8115.1614.8415.1914.7313.51 5

14.77

[1952LAT][1953HEM][1964HIR][1969HEL][1971NAU/RYZ][1973BAR/KNA][1974MIL][1976SMI/MAR][1977BAR/KNA][1978ROB/HEM2][1979KUB/ALC][1980BEN/TEA][1980CLE/JOH][1981STU/MOR][1982WAG/EVA][1983LAN][1983SAN/BAR][1985CHA/DAV][1985GAL][1985MUL][1986MYE][1988PHI/HAL1

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T=298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

K, I=n/aK, I=dilK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, 1=0 0K, I=n/aK, I=n/aK, I=n/aK,J=n/a


1 recalculation of measurements by [1953HEM]2 recalculation of measurements by [1979GIO/BAR]; the constants used by [1984UHL/HEL] do not agree in all

cases with the original references3 calculated with pK(HS7S2>13.9 as given in [1976SMI/MAR]4 corrections made are not traceable5 calculations of [1986MYE] are based on the values from [1976SMI/MAR] and new values determined by

[1986MYE] for the protonation of S2" to HS".

207

JNC TN8400 9 9 - 0 1 1

6.11 Redox equilibria

6.11.1 Pb2+/Pb(cr)

Based on a critical review of available experimentally determined data the COD AT A[1989COX/WAG] Key Values recommended for Pb2+ (aq) an AfG° value of -24.20 kJ/molresulting in:

Pb2++2e- <=> Pb(cr) log K° = -4.25E° =-0.126 V

6.11.2 Pb2+/Pb4+

Solid and dissolved Pb(IV) species exist only under very oxidizing conditions [1976BAE/MES,1985GAL]. From the data given by [1952LAT and 1985GAL] a tentative log K value of 57.23can be calculated for the redox equilibrium Pb2+ <=> Pb4+ + 2e" (Table 6.31). However, nodirect measurements are available.

6.11:3 PbO2andPb3O4

[1985GAL] reviewed available experimental data for PbO2 and Pb3C>4. He gives for theformation of PbO2 from Pb(cr) a E° of 1.690 V and states that PbO2 is thermodynamicallyunstable under acidic conditions, where this E° value is more positive than that of the oxidationof H2O to O2. From the data given in [1985GAL] the following tentative formation constantscan be calculated:

3Pb2+ + 4H2O o Pb3O4(s) + 8H++2e- log K*°so = - 70.98

Pb2+ + 2H2O <=> PbO2(s) + 4H++2e- log K*°so = - 48.98.

208

JNC TN8400 9 9 - 0 1 1

6.11.4 Data compiled for the lead redox system

Table 6.30: Experimentally determined data for the lead redox system. These data were not chosen in thepresent report for the evaluation of recommended stability values. Reasons for not selecting thesereferences are given in Section 6.13: 'Comments on selected references'. Method: pot =potentiometry.

log K Reference Comments I (M) Medium Method

log K: Pb2+ + 2e <=> Pb(0)+

-4.27-14.20 '-4.18 2

[1929CAR][1945LIN][1971VAS/GLA]

T= 298.15 K, 1=0.001-0.15T=298.15,1=0.5T= 298.15 K, 1=0

00.50

waterKC1n/a

potpotpot

1 probably formation of PbCl complexes, hydrolysis not controlled2 no experimental details reported

Table 6.31: Thermodynamic data for the lead redox system taken from previous compilations. As pointed outin Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

logK

log K: Pb2+ H

-4.26-4.28-4.26-4.20-4.27-4.28-4.21-4.27-4.26-4.28-4.28-4.28-4.27-4.28-4.26-4.24-4.21-4.25

log K: Pb2+ <

57.23-48.7657.23

Reference

h 2e- « Pb(0)

[1952LAT][1969HEL][1977PAU][1978COD][1978ROB/HEM2][198OBEN/TEA][1981HEL/KIR][1981STU/MOR][1982PAU][1982WAG/EVA][1983LAN][1983SAN/BAR][1984VIE/TAR][1985BAB/MAT][1985GAL][1985RAI/RYA][1988PHI/HAL][1989COXAVAG1

^ Pb4+ + 2e-

[1952LAT][1983LAN][1985GAL]

Comments

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

I CM)

000000000000000

00

000

209

JNC TN8400 9 9 - 0 1 1


log K*so: 3Pb2+ + 4H2O <=> Pb3O4(s) + 8H++ 2c

-66.38-70.73-70.81-70.84-70.99-73.52-73.57-73.52-70.60-73.58-73.58-73.59-73.53-73.59-73.60-73.60-70.73-73.52-70.98-73.58

log K*so: Pb2+

-49.04-49.14-49.01-49.01-49.23-50.12-49.60-49.62-49.60-49.02-49.62-49.24-49.25-49.09-49.26-49.26-49.26-48.98-49.60-48.98-49.62

log K*so: Pb2+

35.13

log K*so: 2Pb2

-61.00

log K*so: Pb2+

-32.47

[1929MIL][1952LAT][1954COU][1963WIC/BLO][1973BAR/KNA] .[1977BAR/KNA][1978ROB/HEM2][1979KUB/ALC][1979PAT/OBR][198OBEN/TEA][1980SCH][1981STU/MOR][1982PAN][1982WAG/EVA][1983LAN][1983S AN/BAR][1985BAB/MAT][1985CHA/DAV][1985GAL][1988PHI/HAL1

+ 2H2O <=> PbO2(s) + 4H++

[1929MEL][1952LAT][1954COU][1963WIC/BLO][1971NAU/RYZ][1973BAR/KNA][1977BAR/KNA][1978ROB/HEM2][1979KUB/ALC][1979PAT/OBR][1980BEN/TEA][1980SCH][1981STU/MOR][1982PAN][1982WAG/EVA][1983LAN][1983S AN/BAR][1985BAB/MAT][1985CHA/DAV][1985GAL]ri988PHI/HALl

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT = 298.15 K, 1=0T= 298.15 K,I=n/aT = 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/a

• 2c

T= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a'T = 298.15 K, 1=0T= 298.15 K,I=n/aT = 298.15 K, 1=0T= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/a

+ 4H2O <=> Pb(OH)4(s) + 4H++ 2c

[1976TAR/GAR]

+ + 3H2O <=> Pb2O3(s) + 6H

[1985GAL1

+ 1.57H2O <=> PbO!57(s) +

H985GAL1

T= 298.15 K, I=n/a

+ + 2c

T= 298.15 K, I=n/a

3.14H++ 1.14c

T= 298.15 K, I=n/a

0

0

0

0

210

JNC TN8400 9 9 - 0 1 1

6.12 Lead(IV)

The hydrolysis of Pb(OH)4° in solution is not known with reliability: Based on differentreports, [1976BAE/MES] give the following estimates:

Pb(OH)4° + 4H+ <=> Pb4+ +4H2O log p° l i 0 = -4.0Pb(OH)4° + H+ <=> Pb(OH)3+ + H2O log ( 3 \ 3 = -0.8Pb(OH)4°+2H2O o Pb(OH)6

2- + 2H+ log [ 3 \ 6 =-28.3

Pb(OH)4° o PbO2 + 2H2O log K*°so = 4.0

No newer experimental data have been found and no values for the hydrolysis of Pb(IV) arerecommended here.

Table 6.32: Experimentally determined data for the hydrolysis of lead(IV). These data were not chosen in thepresent report for the evaluation of recommended stability values. Method: sol = solubilitymeasurements.


logpI:6.Pb(OH)4°+2H2O <=> Pb(OH)62' + 2H+

-28.06 [1969CHA] T= 298.15 K, I=n/a NaOH sol

log K*S0Pb(OH)40 <=> PbO2 + 2H2

<-4.00 [1969CHA1 T= 298.15 K, I=n/a NaOH sol

Table 6.33: Thermodynamic data for the hydrolysis of lead(IV) taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

log K Reference Comments I (M)

log p,i0.- Pb(OH)4° + 4H + » Pb4+ + 4H2O

-4.00-4.00 '-4.26

[1963FEI/SCH][1976BAE/MES][1979PAT/OBR]

T= 298.T= 298.T = 298

15 K,15 K,.15 K

1=0I=n/a,1=0

0

0

log Pu: Pb(OH)4° + H+ <=> Pb(OH)3+ + H2O

-0.80 ' [1976BAE/MES] T= 298.15 K, I=n/a

211

JNC TN8400 99-011


bgP,,6:Pb(OH)40+2H2O <=> Pb(OH)6

2- + 2H+

-28.5-28.30 '-27.32 2

[1963FEI/SCH][1976BAE/MES][1979PAT/OBR1

T= 298.T= 298.T = 298

15 K,1=015 K, I=n/a.15 K, 1=0

0

0

log K*S0Pb(OH)4° <=> PbO2 + 2H2

-4.00 ' [1976BAEMES1 T= 298.15 K, I=n/a1 reported by [1976BAE/MES] from different sources2 data originally from Pourbaix

212

JNC TN8400 99 - Oil


[1922APP/REI]: [1922APP/REI] determined the solubility of litharge and massicot(PbO(cr)) in 1 M NaOH. From these measurements a log K*so can be calculatedusing a log |31>3 for Pb(OH)3- of-28.32 (1=1; Section 6.1.3) and a log Kw of -13.79(1=1). [1922APP/REI] also determined the redox potential of the electrodePb/PbO(cr) from which [1922APP/REI] calculated the solubility of litharge andmassicot.

[1928RAN/SPE]: [1928RAN/SPE] determined the solubility of litharge and massicot(PbO(cr)) in NaOH. From these measurements a log K*so can be calculated using alog p l i 3 for Pb(OH)3- of-21.91 (1=0; Section 6.1.3) and a log Kw of -14.0 (1=0).

[1939GAR/VEL]: [1939GAR/VEL] measured Pb(II) solubility in dilute alkaline solutions(Table 6.34). Extrapolation of their measurements to 1=0 with the SIT model gave alog Kb

S3 of -1.36 (see Figure 6.31) for the reaction PbO(red) + H2O + OH" <=>Pb(OH)3-, instead of -1.34 as calculated by [1939GAR/VEL]. [1939GAR/VEL]calculated also a log K3 of 10.96 value for the reaction Pb(OH)3" + H+ <=> Pb(OH)2°+ H2O.

The values calculated from [1939GAR/VEL] can be converted to log (3° 12 and logP°ii3 values using a log K°So value of -12.68 for PbO(s, red) (cf. Chapter 6.2.1:PbO(litharge)):

Pb2+ +2H2OPb2+ +3H2O

Pb(OH)2° + 2H+ log p°1>2 = -17.08Pb(OH)3- + 3H+ log J 3 \ 3 = -28.04

Further formation constants for Pb(OH)2° and Pb(OH)3- are discussed in Section6.1.2 and 6.1.3 of this report.

Table 6.34: Experimentally determined equilibrium data compiled for the for the reactionPbO(red) + H2O + OH' <=> Pb(OH)3\ Method: sol = solubility measurements.

log Kbs3 Reference Comments KM) Medium Method

log KbS3: PbO(red) + H2O + OH- <=> Pb(OH)f

-1.31 [1939GAR/VEL] T= 298.15 K, 1=0.0-1.42 [1939GAR/VEL] T= 298.15 K, 1=0.0-1.25 [1939GAR/VEL] T= 298.15 K, 1=0.01-1.32 [1939GAR/VEL] T= 298.15 K, 1=0.01

0.00180.0050.0060.007

NaOHNaOHNaOHNaOH

solsolsolsol

213

JNC TN8400 9 9 - 0 1 1


-1.35-1.36-1.33-1.28-1.32-1.42-1.47-1.40-1.40-1.37-1.40-1.33-1.35-1.40-1.41-1.31-1.29-1.30-1.28-1.33

[1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL][1939GAR/VEL]

T=T=T=T—

T=TT—

T=T=T=T=T=T=T=T=T=T=T=T=T=T=

298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15298.15

K, 1=0.01K, 1=0.01K, 1=0.01K, 1=0.02K, 1=0.02K, 1=0.03K, 1=0.03K, 1=0.04K, 1=0.05K, 1=0.06K, 1=0.1K, 1=0.11K, 1=0.15K, 1=0.17K, 1=0.2K, 1=0.3K, 1=0.4K, 1=0.45K, 1=0.65K, 1=1.2

0.010.0110.0130.0160.02

0.0270.03

0.0350.050.060.1

0.110.150.170.20.30.40.450.651.2

NaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOHNaOH

solsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsolsol

red PbO(s)+H2O+OH" o Pb(OH)3-

0 5 Lmolal 1 1.5

Figure 6.31: Plot of log KbS3 + 0D vs. Im for the reaction PbO(s) + H2O + OH' <=> Pb(OH)3-

at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.05,log Kb°S3= - 1.36. Calculated from data compiled in Table 6.34.

214

JNC TN8400 9 9 - 0 1 1

[1952LAT]: Careful review of the early literature. Solubility of lead chloride calculated frompolarographic measurements by [1922GER].

[1960CAR/OLI]: cf. [1960OLI1]; 0.25-10 mM Pb; pKw = 14.18 (1=3) and 13.76 (1=0.3).

[1960OLI1]: Olin and co-workers [1960OLI1, 1960OLI2, 1960CAR/OLI, 1961OLI, and1962PAJ/OLI] determined lead hydrolysis in 0.3 and 3 M perchlorate medium. Theircareful experimental work included large range of pH values and lead concentrations.They put up a consistent system of hydrolyzed lead species including Pb2OH3+,Pb4(OH)4

4+, Pb3(OH)42+, and Pb6(OH)8

4+ which was later slightly enlarged byPb3(OH)5+ [ 1980SYL/BRO]. Pb concentration = 1.25- 80 mM.

[1960OLI2]: cf. [1960OLI1]; the value determined at I = 4.5 is corrected in [1962PAJ/OLI].Pb concentration = 500 - 1490 mM.

[1961OLI]: cf. [1960OLI1]; same values as reported by [1960CAR/OLI], [1960OLI1], and[1960OLI2]; pKw = 14.18 (1=3) and 13.76 (1=0.3). Pb concentration = 1 -80 mM.

[1962PAJ/OLI]: cf. [1960OLI1]; Pb concentration = 5 - 1450 mM.

[1964BAK]: [1964BAK] determined a log K*so for chloropyromorphite and hinsdalite usinga equilibration time of several months. A closer examination of his data leads to theconclusion that he probably assumed all phosphate to be present as PO43- (in aNa2HPO4 solution in the case of chloropyromorphite and at a pH of 3 in the case ofhinsdalite). These values are therefore not used in this report.

[1971NAU/RYZ]: small solubility of laurionite probably due to a transposing of digits (AfG°given as -480.3 instead -408.3) of value which already originally was too low[1987BRU].

[1972NRI]: Nriagu [1972NRI] made measurements in dilute phosphoric acid solution anddetermined the constants for the formation of soluble PbH2PO4+ and PbHPO4° aswell as the solubility of different solid lead phosphates. In 0.1 M solutions he alsodetermined the stability of chloro-, fluoro- and bromopyromorphite [1973NRI1,1973NRI2]. All constants were corrected to 1=0 with the Davies equation. In[1984NRI] a log K*so for plumbogummite and hinsdalite is estimated.

[1972ZIR/YAM]: calculated log P values from electronegativity.

[1973BIL/STU]: same values as [1976BIL/HUS].

[1973NRI1]: cf. [1972NRI].

215

JNC TN8400 99-011

[1973NRI2]: cf. [1972NRI].

[1976BAE/MES]: Careful and extensive review of existing literature concerning hydrolysisof lead. [1976BAE/MES] extrapolated the selected values to I = 0.

[1976BIL/HUS]: Bilinski and co-workers [1976BIL/HUS, 1973BIL/STU] determined leadcarbonate complex formation constants in 0.1 M KNO3 medium (pK , used by[1976BIL/HUS] is 13.96) in presence of 0.001-0.01 mM Pb potentiometrically andmeasured also the solubility of lead carbonate complexes. They also recalculated thedata of [1965BAR/BAR] in 1 M NaC104. The log (3 value given in 0.7 NaClO4 isbased on a private communication from M. Sipos (310"9 M Pb).

The first and second lead hydrolysis constants determined by [1976BIL/HUS] werenot chosen in this report as [1976BIL/HUS] neglected the formation of thepolynuclear species Pb3(OH)4 at a Pb concentration of 10~5 M.

[1976SMI/MAR]: Calculated mean of values at a given ionic strength: The values areselected after a critical review, however, no comments to the individual paperselected or not selected are made by Smith and Martell. Several compilations havebeen published: [1982SMI/MAR] and [1989SMI/MAR].

[1979PAT/OBR]: [1979PAT/OBR] selected thermodynamic data for lead carbonate andhydroxide complexes and solids and then tested their dataset with experimental datafrom different sources.

[1980CLE/JOH]: [1980CLE/JOH] made a careful and extensive review of data concerningthe solubility of sparingly soluble lead salts. Based on this review, log K*so formany salts are recommended. In cases were no newer data could be found, the datarecommended by [1980CLE/JOH] are used.

[1980KAW/ISH]: [1980KAW/ISH] determined with a potentiometric method the hydrolysisof Pb(II) in rather concentrated solutions (Pb(II) = 1 - 8 0 mM). [1980KAW/ISH]themselves classified their log P values for PbOH+ and Pb(OH)2 as doubtful, as onlyminor concentration of PbOH+ and Pb(OH)2 are present under the experimentalconditions . p ^ in 3 M LiClO4, as determined in [1980KAW/ISH] = 13.87.

[1980MAN/DEU]: selected data (from the literature) for Pb oxide/hydroxide, chloride,sulfate and carbonate complexes and solids have been corrected to 1=0 by[1980MAN/DEU].

[1980SCH]: [1980SCH] selected formation constants from the literature and verified hisdataset with experimental data from different sources in systems containingconsiderable amounts of carbonate; pK , used by [1980SCH] is 13.99.

216

JNC TN8400 9 9 - 0 1 1

[1980SYL/BRO]: Their determination of lead hydrolysis was carried out in 0.1 M KNO3 at0.1-2 raM Pb concentration. As these values were determined in nitrate solutionsthey were not chosen in this report for extrapolation to I = 0. [1980SYL/BRO]introduced an additional lead species (Pb3(OH)5+) and recalculated the valuesdetermined by Olin and co-workers at I = 0.3 and 3, including the new speciesPb3(OH)5+. [1980SYL/BRO] also showed that the additional consideration ofPb3(OH)5+ had hardly an influence on the other log P values.

[1980WAL/SIN]: [1980WAL/SIN] determined the solubility of cerrusite, anglesite and leadhydroxide/oxide as well as constants for lead hydrolysis in a wastewater containingpredominantly H2SO4. The values given are the mean calculated from measurementsin wastewater. As also additional anionic components were present in the wastewaterused which were not considered in their calculations, the calculated log K*so valuesare too large and the complex formation constants too small. Values from[1980WAL/SIN] were not chosen for the calculation. Nevertheless, they can be usedfor comparison.

[1982BIL/SCH]: cf. [1976BIL/HUS]; pK^ used = 13.96 (for 1=0.1) and 13.76 (for 1=0.3).

[1982SMI/MAR]: Calculated mean of values at a given ionic strength. See also[1976SMI/MAR] and [1989SMI/MAR].

[1983SCH/GAR]: Thermodynamic data selected in a review, dataset tested withexperimental data from different sources. The constant for PbHCO3+ could not beverified under the conditions of the experiments. The value given for hydrocerrusiteis an approximate value. [1983SCH/GAR] criticizes the older value forhydrocerrusite as too high compared to solubility measurements.

[1984NRI]: cf. [1972NRI].

[1984TAY/LOP]: [1984TAY/LOP] determined with X-ray analysis the conditions underwhich massicot (or litharge), plumbonacrite, hydrocerrusite and cerrusite werepredominantly formed in aqueous carbonate solutions (I, pCO2, pH were varied) andextrapolated their results to 1=0. They could observe the formation of all thesesubstances at room temperature. Calculations carried out by [1984TAY/LOP]demonstrated that at a total C concentration of 10"4 M, the solubilities ofplumbonacrite, cerrusite and hydrocerrusite are similar, especially near neutral pHvalues.

[1984UHL/HEL]: [1984UHL/HEL] determined the log K*so of galena (PbS(cr)) with helpof competitive complex formation with EDTA and TRIS. Extrapolation of theirvalues to I =0 with SIT yields a value of 12.17. [1984UHL/HEL] recalculated alsothe values given by [1953HEM] and [1979GIO/BAR] for Pb(HS)2° and Pb(HS)3-.However, from the information given in [1984UHL/HEL], their calculations are not

217

J N C T N 8 4 0 0 99 - O i l

traceable and the constants used by [1984UHL/HEL] did not agree in all cases withthe original references.

[1987BRU]: in the case of lead carbonate complexes the influence of I is not taken intoaccount.

[1988BYR/KUM]: extrapolated to 1=0.7 for seawater. The values for the lead carbonatecomplexes are a recalculation of [1981BYR]. These values are probably less precisethan measurements e.g. by [1977SIP/VAL].

[1989DOR/MAR]: These titration experiments are unfortunately not well documented. Theoriginal I (probably the measurements were carried out in diluted suspensions) wasnot reported. However, the data were extrapolated to 1=0 by the authors. Their valuedetermined for the first lead hydrolysis constant is not chosen in our calculations asthe formation of polynuclear lead hydroxide species can not be excluded at a Pbconcentration of 0.02-0.2 mM [1989DOR/MAR]. Their values given for theformation of PbCO3° and PbHCCV complexes and for the solubility of cerrusiteseem to be more accurate.

[1989SMI/MAR]: Calculated mean of values at a given ionic strength. See also[1976SMI/MAR] and [1982SMI/MAR].

[1993MAC/PAG]: [1993MAC/PAG] compare lead solubility calculated with thermodynamicdata with experimental lead solubility.

[1995MAR/MAC]: [1995MAR/MAC] tried to verify the MINTEQA2 database withsolubility measurements under different conditions and identified the precipitatedsolids with X-ray diffraction. They observed the precipitation of cerrusite,hydrocerrusite and anglesite at room temperature and showed that neither lanarkite(Pb2OSO4) nor crystalline Pb(OH)2 precipitates from solutions. In absence oflanarkite and Pb(OH)2 their calculations showed a reasonable agreement with theexperimental observations.

218

JNC TN8400 9 9 - 0 1 1

7 Bismuth

The most stable oxidation state of bismuth is +III, but the oxidation states -III and -I can beprepared in liquid nitrogen or in the gas phase. Bismuth compounds of the +V state are strongoxidizing agents and are able to oxidize water to oxygen [1985LOV/MEK, 1995WIB]. Inaqueous environments, Bi(III) is a strong acid and Bi3+ starts to hydrolyze at a pH of about 1.Bismuth has a strong tendency to form polynuclear complexes [1986BAE/MES]. For thepresent evaluation, thermodynamic data from literature are compiled for the formation of Bicomplexes and solids with hydroxide, chloride, fluoride, carbonate, nitrate, phosphate, andsulfate.

Based on experimental data reported in the literature, thermodynamic data are recommended forthe complex formation of bismuth with hydroxide, chloride, nitrate and for the redox equilibriaBi3+ + 3e- <=> Bi(cr). Thermodynamic data are also selected for the formation of the solidsa-Bi2O3(cr), BiOCl(s), (BiO)4(OH)2CO3(cr), (BiO)2CO3(cr), and BiONO3(s).

For other potentially relevant solid Bi phases, like BiPO^s) and Bi2(SO4)3(s), noexperimentally determined solubility data are available.

7.1 Hydrolysis of bismuth

The Bi3+ ion starts to hydrolyze at a pH of about 1 and has the tendency to form polynuclearcomplexes at Bi concentrations > 1(H M under acidic and neutral conditions [1976BAE/MES].The solubility of Bi2O3 does not depend on pH in the pH range of ~ 8 - 12, indicating thepresence of Bi(OH)3°. Only in alkaline solutions (pH > 12), bismuth solubility increases again.The data used for the evaluation of equilibrium constants for the hydrolysis of bismuth(III) aregiven in Table 7.1 and in Figure 7.1 to Figure 7.8. Additional equilibrium data that were notselected for the calculation of the bismuth(III) hydrolysis are compiled in Tables 7.2 and 7.3.

Olin [1957OLI, 1959OLI, 1961OLI, 1975OLI] was one of the first investigators whodetermined the hydrolysis of Bi(III). Based on potentiometric measurements he reported for theformation of mononuclear BiOH2+ a log P u value of -1.58 and for the polynuclearBi6(OH)i26+ a log 06,12 value of 0.33 in 3 M NaClO4. His careful experimental workencompassed a large range of pH values and bismuth concentrations and he developed aconsistent dataset for the hydrolysis of bismuth. [1976BAE/MES] critically reviewed thesemeasurements and extrapolated the measurements of Olin, [1960TOB, 197 IBID,1972DRA/NIM2] to 1=0. These measurements, as well as the studies of bismuth hydrolysis by[1972DRA/NIM1, 1982SUG/SHI, 1985SED/SIM] in perchlorate medium were used in thisreport for the evaluation of bismuth hydrolysis at 1=0. The measurements of [1975ANT/NEV,1975HEI/SCH, 1982HAT/SUG, 1987SUG/SHI, 1993KRA/DEC] in nitrate, chloride orsulfate medium were not chosen in this report for the calculations of the hydrolysis constants,because bismuth forms complexes with these anions. Also the data of [1960TOB] were

219

JNC TN8400 9 9 - 0 1 1

excluded because he neglected the possible formation of other Bi hydroxide complexes(BiOH2+, Bi9(OH)2o

7+) in his data interpretation.

Table 7.1: Experimentally determined equilibrium data compiled for the bismuth hydroxidesystem, according to the equilibrium: mBi3+ + nH2O <=> Bim(OH)n

3"n + nH+.These data were chosen for the evaluation of recommended values in the presentreport. Additional information for the different references see Section 7.1:2:'Comments on selected references'. Method: extr = solvent extraction, mig =migration, pot = potentiometry, sol = solubility measurements, sp =spectrophotometry,

-"m,nReference Comments Medium Method

log pu: Bi3+ + H2O

-1.43 [1971BID] T= 298.15 K, 1=0.1-1.55 [1972DRA/NIM1] T= 298.15 K, 1=1-1.59 [1975OLI] T= 298.15 K, 1=3-1.84 [1982HAT/SUG] T= 298.15 K, 1=1-1.46 [1985SED/SIM] T= 293.15 K, 1=0.4-1.40 [1987MIL/ROE] T= 298 K, 1=0.25

0.1131

0.40.25

NaC104

H, NaClO4

NaC104

NaC104

NaC104

NaC104

extrsppotextrextrmig

log pli2: Bi3+ + 2H2O <=> Bi(0H)2+ + 2H+

-2.82 [1972DRA/NIM1]T= 298.15 K, 1=1-4.74 [1982HAT/SUG] T= 298.15 K, 1=1-3.36 [1985SED/SIM] T= 293.15 K, 1=0.4-3.57 [1987MIL/ROE] T= 298 K, 1=0.25

11

0.40.25

H, NaC104

NaClO4

NaClO4

NaClO4

spextrextrmig

log p]f3: Bi3+ + 3H2O & Bi(0H)3° + 3H+

-7.5-6.41

log Kjt4:

-12.84-13.07

[1982HAT/SUG][1987MTL/ROE]

Bi(OH)3° + H2O&.

[1971BID][1987ME7ROE]

T=298T=298

Bi(OH)4-

T=298T=298

.15K,

+

.15K,

K, 1=11=0.25

K, 1=11=0.25

0

0

1.25

1.25

NaClO4

NaClO4

extrmig

NaC104

NaClO4solmig

220

JNC TN8400 9 9 - 0 1 1


log P6,J2-- 6Bi3+ + 12H2O a Bi6(OH)126+ + 12H+

0.26 [1972DRA/NIM1] T= 298.15 K, 1=10.33 [1975OLI] T= 298.15 K, 1=3

13

H, NaC104

NaC104

sppot

log K9i20:1.5 Bi6(OH)]26+ + 2 H2O <^> Bi9(OH)20

7+ •

[1959OLI] T= 298.15 K, 1=0.1[1972DRA/NIM2] T= 298.15 K, 1=0.1

2H+

-3.5-3.9

0.10.1

NaClO4

NH4C1O4

potsp

log K9i2i: Bi9(OH)2o7+ + H2O <=> Bi9(OH)2i

6+ + H+

-3.2 [1959OLI] T= 298.15 K, 1=0.1 0.1-3.2 [1972DRA7NIM2] T= 298.15 K, 1=0.1 0.1

NaClO4

NH4C104

potsp

log K9i22: Bi9(OH)2j6+ + H2O <=> Bi9(OH)22

5+ + H+

-2 .6 [1959OLI] T= 298.15 K, 1=0.1 0.1-2 .8 [1972DRA/NIM2] T= 298.15 K, 1=0.1 0.1

NaC104

NH4C1O4

pot

sp

log p3A: 3Bi3+ + 4H2O <=> Bi3(OH)45+ + 4H+

-0.80 [1975OLI] T= 298.15 K, 1=3 NaCIO, pot

221

J N C T N 8 4 0 0 99 - O i l

7.1.1

Based on potentiometric measurements, [1957OLI] determined a log BJJ value of-1.58 for theformation of mononuclear BiOH2+ according to the reaction Bi3+ + H2O o BiOH2+ + H+ (SeeTable 7.2). Bidleman [1971BID] measured in 0.1 M NH4CIO4 a log B u of -1.43.[1976BAE/MES] extrapolated in his careful review from these two values a log B 0 ^ of -1.09for 1=0 (Table 7.3). In the present report, the data determined experimentally by [1971BID,1972DRA/NIM1, 1975OLI, 1982HAT/SUG, 1985SED/SIM, 1987MH7ROE] (Table 7.1)were used for evaluation of log B°ii. Extrapolation to I = 0 with the SIT term is shown in

Figure 7.1.

Bi3+ + H2O <=> BiOH2+ log = -0.92, Ae =-0.09

The data measured by [1975ANT/NEV] and [1982HAT/SUG] in nitrate medium and[1987SUG/SHI] in chloride medium (Table 7.2) are not used for the determination of B°1;1

values because both, nitrate and chloride, tend to form complexes with bismuth.

Bi3++H?O<^BiOH2++H+

1 2 3

L molal

Figure 7.1: Plot of log Bij + 4 D vs. Im for the reaction : Bi3+ + H2O <=> BiOH2+ + H+ at 25°C. The straight line shows the result of the linear regression: Ae = - 0.09; logB°! i = - 0.92. Calculated from data compiled in Table 7.1.

222

J N C T N 8 4 0 0 9 9 - O i l

7.1.2 Bi(OH)2+

The determination of the precise value of log (3 j is difficult as in most experimental systemcompiled in this report, Bi(0H)2+ is only a minor species and BiOH2+ or polynuclearBig(OH)i26+ dominate the speciation in the solutions (see also [1976BAE/MES]). Thus, themeasured values given in Table 7.1 [1972DRA/NIM1, 1982HAT/SUG, 1985SED/SIM,1987MIL/ROE] show a considerable spread. [1976BAE/MES] estimated in their work (due tolack of experimental data) a value of -4.0 for log (3° 1,2- Extrapolation of the data given in Table7.1 to 1=0 gives:

Bi3+ + 2H2O Bi(OH)2+

logP°i>2 = -2.56

This value is similar to the value calculated by [1982LAP/KOL] from measurement at 75 -300 °C and the determinations by [1975ANT/NEV] (Table 7.2). The uncertainty of log pi,2 ismuch larger than for the other bismuth hydroxide species. However, even at bismuthconcentrations of 10"11 M, the dominating species in acidic solutions are Bi3+, BiOH2+ andBi(OH)3° and not the Bi(OH)2

+ complex [1987SUG/SHI].

It should be mentioned that the formation constant given in [1952LAT] for BiO+ (correspondingto Bi(OH)2

+) (Table 7.3) originates from the early works of [1923SWI, 1923SMI,1947GRA/SIL]. [1957OLI] could not establish the presence of Bi(OH)2

+ but demonstrated thatthe polynuclear Bi6(OH)i26+ complex is stable compared to Bi(OH)2+. [1957OLI] pointed outthat as polynuclear Bi6(OH)i26+ complex is so stable that the equilibrium concentration ofBi(OH)2+ is too low to be detected in his experiments.

Bi3++ 2H2O & Bi(OH)2+ + 2H+

Figure 7.2: Plot of log p1>2 + 6 D vs. Im for the reaction : Bi3+ + 2H2O <=> Bi(OH)2+ + 2H+ at

25 °C. The straight line shows the result of the linear regression: Ae = —0.03; logP°i,2 = -2.56. Calculated from data compiled in Table 7.1.

223

JNC TN8400 9 9 - 0 1 1

7.1.3 Bi(OH)3°

[1971BID] measured in 0.1 and 1 M perchlorate solution the bismuth hydrolysis with anorganic extractant (dithizone). He was able to show that no (or only little) polynuclear specieswere present in 0.001-0.01 mM bismuth solutions in the pH range 9.9 - 11.3. [1971BID]determined in 0.1 M NaC104 a log p l i 3 of-9.43 for the formation of Bi(OH)3 (Table 7.2).[1976BAE/MES] extrapolated from this a log [313 value of -8.86 for 1=0. However, newermeasurements in perchlorate medium by [1982HAT/SUG, 1985SED/SIM, 1987MIL/ROE]indicate a less negative log Pi 3 value. [1982HAT/SUG] assumed that the difference betweentheir results and the result of [1972BID] may be due to the formation of polynuclear species inthe experiments of [1972BID] at Bi(III) concentrations of 0.01 mM. Also the measurements of[1975ANT/NEV] in KN03 indicate a log p ] i3 value near -5.5 (Table 7.2). Extrapolation of themeasurements of [1982HAT/SUG] and [1987MEL/ROE] to 1=0 results in:

Bi3+ + 3H2O Bi(OH)3° log p° l i3 = -5.31

1 2lm, molal

Figure 7.3: Plot of log p1>3 + 6 D vs. Im for the reaction : Bi3+ + 3H2O «• Bi(OH)3° + 3H+ at25 °C. The straight line shows the result of the 'linear regression': Ae = 0.91; logP°i 3 = - 5.31. Calculated from data compiled in Table 7.1.

224

JNC TN8400 9 9 - 0 1 1

7.1.4 Bi(OH)4-

For Bi(OH)4-, [1973BID] measured in 0.1 M NH4CIO4 a log p ]>4 of -22.26 (Table 7.2). Thedata reported in the literature for log (314 are quite different (see Table 7.2), while for thereaction Bi(OH)3° + H2O <=> Bi(OH)4- + H+ the variation between the different experimentalpapers is smaller (Table 7.1), as these values do not depend on the value of log Pi 3 (cf. Section7.1.3). Extrapolation of the log Kj>4 values determined by [1971BID] and [1987MH7ROE]gives a log Ki>4 of -13.40 for the reaction Bi(OH)3° + H2O <=> Bi(OH)4- + H+. Using a logP°i,3 of-5.31 (Section 7.13) one obtains:

Bi3+ + 4H2O Bi(OH)4- logp°1 |4 = -18.71

O

Bi(OH)3°+ H20 o Bi(OH)4"+H+

- 1 0 T

- 1 0 . 5 ••

-11

CM - 1 1 . 5

I -12 +

s? -12.5 •

-13

-13.5 -:

-14-

-14 .5 ••

-150

y = 0.14x-13.40

0.5 1

lmi molal

1.5

Figure 7.4: Plot of log Ki,4 + 4 D vs. Im for the reaction : Bi(0H)3° + H2O «=> Bi(0H)4- +H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae = -0.14; log K ° M = -13.40. Using a log p° l i3 of -5.31 one obtains a log p ° M of -18.71 for the reaction Bi3+ + 4H2O « Bi(0H)4- + 4H+. Calculated from datacompiled in Table 7.1.

225

JNC TN8400 9 9 - 0 1 1

7.7.5 Bi6(OH)]26+

Polynuclear Big(OH)i26+ dominates the speciation below pH 3 in solution containing more than0.01 mM Bi. Experimental values are compiled in Table 7.1. Dragulescu and co-workers[1972DRA/NIM1, 1972DRA/NIM2] determined bismuth hydrolysis in 0.1 and 1 M perchloratemedium. Their results are comparable to the results of Olin. Extrapolation of the valuesdetermined by [1972DRA/NIM1] and [1975OLI] to 1=0 is shown in Figure 7.5 and results in:

6Bi3+ + 12H2O Bi6(OH)126+ + 12H+ log (3°6>i2 = 1-34

The value determined by [1960TOB] was not chosen for the extrapolation, as [1960TOB] didnot include the species BiOH2+ and Bi(OH)2

+ in his calculations.

QCD

+CM

CO*

CO.

log

A*r

3.5

3

2.5

2

1.5 •

1

0.5

0

-0.5

.1

6Bi3++12H2O<=>Bi6(OH)126++12H+

•

y - 6.0"4x 4-1.34

1 1 1

1 2lm, molal

Figure 7.5: Plot of log p6,i2 + 6 D vs. Im for the reaction : 6Bi3+ + 12H2O <=> Bi6(OH)i26+ +

12H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae = -0.04; log p°6,i2 = 1.34. Calculated from data compiled in Table 7.1.

226

JNC TN8400 9 9 - 0 1 1

7.1.6 Bi9(OH)207+, Bi9(OH)2i

6+, and Bi9(OH)225+

Based on potentiometric measurements, [1959OLI] reported values for the formation of thepolynuclear Bi9(OH)20

7+, Bi9(OH)2i6+, and Bi9(OH)225+ complexes from Bi6(OH)i2

6+ abovea pH of ~ 3 and at Bi concentrations of 0.25 - 4 mM. [1960TOB] proposed the existence of aBi6(OH)i53+ complex. Spectrophotometric measurements made by [1972DRA/NIM1] areconsistent with the formation of polynuclear Big(OH)2o

7+, Bi9(OH)2i6+, and Bi9(OH)225+

complexes as proposed by [1959OLI]. The consecutive formation constants given by[1959OLI] and [1972DRA/NEVI1] at 1=0.1 are listed in Table 7.1. The mean of these valueswas corrected in this report to 1=0 using the SIT approximation assuming a Ae of 0:

1.5Bi6(OH)i26++2H2O «• Bi9(OH)20

7++ 2H+ log K°9>20 =-3.37Bi9(OH)20

7+ + H2O a Bi9(OH)216+ + H+ log K \ 2 ! = -1.89

Bi9(OH)2i6++ H2O <=> Bi9(OH)225++H+ log K \ 2 2 = -1.61

These values were then converted to log (3 values that refer to Bi3+, using a log P°6,j2 of 1.34(see Figure 7.5):

9 Bi3+ + 20 H2O <=> Bi9(OH)207+ + 20H+ log (39,2o = -1-36

9 Bi3++ 21 H2O <=> Bi9(OH)2i6++ 21H+ log (39>21 = -3.25

9 Bi3+ + 22 H2O <=* Bi9(OH)225+ + 22H+ log fj9>22 = -4.86

7.1.7 Bi3(OH)45+

[1975OLI] calculated a tentative log (33i4 value of -0.80 in 3 M NaClO4 solution for thereaction 3Bi3+ + 4H2O <=> Bi3(OH)4

5+ + 4H+. As this is the only determination of thisconstant it is extrapolated to 1=0 with SIT, log p \ 4 - 2D = P°3?4 - Ae Im, assuming a Ae of0.18 (from Ae(Al3(OH)4

5+) = 1.30, Ae(Fe3+) = 0.56 and Ae(H+) = 0.14; [1992GRE/FUG]):

3 Bi3+ + 4 H2O <=> Bi3(OH)45+ + 4H+ log (3°3>4 = -0.80

7.1.8 Bi6(OH)!53+

[1960TOB] proposed, based on titration experiments, the existence of a Bi6(OH)]53+ complexinstead of Bi9(OH)20

7+, Bi9(OH)2i6+, and Bi9(OH)225+ as proposed by [1959OLI] (see Table

7.2). Spectrophotometric measurements made by [1972DRA/NIM1], however, are consistentwith the formation of polynuclear Bi9(OH)2o

7+, Bi9(OH)2j6+, and Bi9(OH)225+ complexes as

proposed by [1959OLI]. Thus, the existence of a Bi6(OH)i53+ complex instead ofBi9(OH)2o

7+, Bi9(OH)2]6+, and Bi9(OH)225+ complexes seems rather unlikely.

227

JNC TN8400 9 9 - 0 1 1

7.1.9 Additional equilibrium data compiled for the bismuth hydroxide system

Table 7.2: Additional experimentally determined equilibrium data compiled for the bismuth hydroxidesystem, according to the equilibrium: mBi3+ + nH2O <=> Bim(OH)n

3-n + nH+. These data werenot chosen in the present report for the evaluation of recommended stability values. Reasons fornot selecting these references see text and Section 7.12. Method: extr = solvent extraction, pot =potentiometry, mig = migration, sol = solubility measurements, sp = spectrophotometry, tit =titration (pH).

log Pm ,n

log pu: Bi3+ •

-1.58 '-1.58 '-1.552

-1.50 2

-1.41 2-1.342

-1.67 2

-2.72 2

log j3u: Bi3+ -

-3.52 2

-3.45 2

-3.26 2

-3.10 2

-0.15 2-4.51 2

-2.29 3

-9.72

-4.00 2.4-4.10 4

Reference

f //26> <=> fi/Otf24" +

[1957OLI][1961OLI][1975ANT/NEV][1975ANT/NEV][1975ANT/NEV][1975ANT/NEV][1982HAT/SUG]n987SUG/SHIl

v 2H2O « Bi(0H)2+

[1975ANT/NEV][1975ANT/NEV][1975ANT/NEV][1975ANT/NEV][1975HEI/SCH][1982HAT/SUG][1982LAP/KOL][1987SUG/SHI][1993KRA/DEC][1993KRA/DEC]

Comments

H+

T= 298.15 K, 1=3T= 298.15 K, 1=0.1,3T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K, 1=1

+ 2H+

T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=1.5T= 298.15 K, 1=1T=298.15K, I=dilT= 298.15 K, 1=1T= 29610.5 K, 1=1T= 29610.5 K, 1=1

KM,

33

0.10.30.5111

0.10.30.5

11.51

111

Medium

NaClO4

NaC104

KNO3KNO3KNO3KNO3

NaNO3

NaCl

KNO3KNO3KNO3KNO3H2SO4

NaNO3

NaOH, HC1O4

NaClNO3-

cio4-

Method

potpotspspspsp

extrextr

spspspsppotextrsolextrsolsol

log Pu: Bi3+ + 3H2O <=> Bi(OH)3° + 3H+

-9.43-5.94 2

-5.93 2

-5.58 2

-5.29 2

-7.6 2

-9.02 3

-5.59 5

-10.7 2

-9.90 4

-10 .00 2.4

[1971BID][1975ANT/NEV][1975ANT/NEV][1975ANT/NEV][1975ANT/NEV][1982HAT/SUG][1982LAP/KOL][1985SED/SIM][1987SUG/SHI][1993KRA/DEC][1993KRA/DEC]

T= 298.15 K, 1=0.1T= 298.15 K, 1=0.1T= 298.15 K, 1=0.3T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K,I=dilT= 293.15 K, 1=0.4T= 298.15 K, 1=1T= 29610.5 K, 1=1T= 29610.5 K, 1=1

0.10.10.30.5

11

0.4111

NH4CIO4KNO3KNO3KNO3KNO3NaNO3

NaOH, HC1O4

NaC104

NaCl

cio4-NO,"

solspspspsp

extrsolextrextrsolsol

228

JNC TN8400 9 9 - 0 1 1


log PL4: Bi3+ + 4H2O <=> Bi(OH)4- + 4H+

-19.48 6

-22.26 6

-21.50 4

-21.50 2-

log K1A:

-11.60 4

-11.70 2

[1987MDL/ROE][197 IBID][1993KRA/DEC]

4 [1993KRA/DEC1

Bi(OH)3°+ H2O<=>Bi(OH)

[1993KRA/DEC]|"1993KRA/DEC1

T= 298 K, 1=0.25T= 298.15 K, 1=0.1T= 296±0.5 K, 1=1T= 296+0.5 K, 1=1

4-+ H+

T= 298.15 K, 1=1T= 298.15 K, 1=1

0.250.1

11

11

NaClO4

NH4CIO4

C1O4"NCV

cio4-NO,~

migsolsolsol

solsol

2-- 6Bi3+ + 12H2O <=> Bi6(OH)]26+ + 12H+

0.33 '0.33 '-0.53

[1957OLI][1961OLI][1960TOB1

T= 298.15 K, 1=3T= 298.15 K, 1=3T= 298.15 K, 1=1

331

NaClO4

NaClO4

NaClO4

potpottit

log K9,20:1.5 Bi6(OH),26+ + 2 H2O <=> Bi9(OH)20

7+ + 2H+

-3.5 ' [1961OLI] T= 298.15 K, 1=0.1 0.1 NaClO4 pot

log K9i21: Bi9(OH)207+ + H2O <=> Bi9(OH)21

6+ + H+

-3.2 ' [1961OLI] T= 298.15 K, 1=0.1 0.1 NaClO4 pot

log K9a2: Bi9(OH)2l6+ + H2O <=> Bi9(OH)22

5+ + H+

-2.6 ' [1961OLI] T= 298.15 K, 1=0.1 0.1 NaClO. pot

log P6jJ5: 6Bi3+ + 15H2O

-8.63 [1960TOB]

Bi6(OH)153+ + 15H+

T= 298.15 K, 1=1 NaClO, tit1 same values as [1959OLI]

formation of nitrate, chloride or sulfate complexes with electrolyte possible,extrapolated from measurements at 75-300 °C.

4 Bi: 0.01 - 1000 mM. Formation of polynuclear species probable5 Bi: <0.1 mM. Formation of polynuclear species possible6 log K14 values are given in Table 7.17 existence of Bi6(OH)15questionable, see Section 7.1.8

229

JNC TN8400 9 9 - 0 1 1

Table 7.3: Thermodynamic data for the bismuth hydroxide system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

l°g P Reference Comments I (M) Medium

log Pu: Bi3+ + H2O <=> BiOH2+ + H+

-1.09-1.10-1.09-1.4-0.91 '-1.58 2

-1.58 2

-1.10-1.40-1.40

[1976BAE/MES][1976SMI/MAR][1981BAE/MES][1982WAG/EVA][1985BAB/MAT][1985LOV/MEK][1985LOV/MEK][1989SMI/MAR][1993KRA/DEC]n993KRA/DECl

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 296±0.5 K, 1=1T= 296±0.5 K, 1=1

0000000011

NO3"C1O,

log fih2: Bi3*

-0.1-4.00-1.4

2H2O Bi(OH)2+ + 2H+

[1952LAT][1976BAE/MES]ri982WAG/EVA1

T= 298 K, I=dilutedT= 298.15 K, 1=0T= 298.15 K, 1=0

000

log pu: Bi3+ + 3H2O <=> Bi(OH)3° + 3H+

-8.86-8.90-8.86-9.00

log PK4: Bi3+ -

-21.80-21.80-21.20

logP61,-6Bi3-

-0.530.32-0.31 l

0.33 '-0.180.300.30

[1976BAE/MES][1976SMI/MAR][1981BAE/MES][1989SMI/MAR]

(- 4H2O « Bi(OH)4~

[1976BAE/MES][1976SMI/MAR][1989SMI/MAR1

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

+ 4H+

T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0

f + 12H2O <=> Bi6(OH)126+ + 12H+

[1976SMI/MAR][1982WAG/EVA][1985LOV/MEK][1985LOV/MEK][1989SMI/MAR][1993KRA/DEC][1993KRA/DEC]

T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 296+0.5 K, 1=1T= 29610.5 K, 1=1

0000

000

1000111

NO3-

cicv

230

JNC TN8400 99-011

Table

log Pi

-2.61-3.00-3.95-3.01-2.61

7.3:

>.20:

1

1

continued

9Bi3+ + 20H2O <=> Bi9(OH)207+ + 20H+

[1976SMI/MAR][1982WAG/EVA][1985LOV/MEK][1985LOV/MEK][1989SMI/MAR]

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.1

0.1000

0.1

log A>,2/.- 9Bi3+ + 21H2O <=> Bi9(OH)216+ + 21H+

-5.79-6.21-7.16 2

-6.21 2-5.79

[1976SMI/MAR][1982WAG/EVA][1985LOV/MEK][1985LOV/MEK][1989SMI/MAR]

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.1

0.1000

0.1

log P9,2i- 9Bi3+ + 22H2O <=> Bi9(OH)225+ + 22H+

-8.47 [1976SMI/MAR] T= 298.15 K, 1=0.1 0.1-8.84 [1982WAG/EVA] T= 298.15 K, 1=0 0-9.76 2 [1985LOV/MEK] T= 298.15 K, 1=0 0-8.81 2 [1985LOV/MEK] T= 298.15 K, 1=0 0-8.47 ri989SMI/MAR1 T= 298.15 K, 1=0.1 01

log p6J5: 6Bi3+ + 15H2O <=> Bi6(OH)!53+ + 15H+

-8.6 ri982WAG/EVAT T= 298.15 K, 1=0 0_1 calculated with a Afi° of -95.55 kJ/mol for Bi2+ (Section 7.11).2 numbers given by [1985LOV/MEK] in the text and in the table are different

231

JNC TN8400 9 9 - 0 1 1

7.2 Solid bismuth-oxide/hydroxide

7.2.1 a-Bi2O3(cr)

The low temperature modification of solid bismuth oxide is the monoclinic a-Bi2C>3 which isstable up to 715° C where it undergoes a polymorphic transformation to cubic P-Bi2C>3 (or 5-Bi2O3) [1943SCH/RIT, 1981KAL/BOR, 1995WIB]. Unfortunately, in some papers andreviews no description of the exact nature of the Bi2O3 used is given. However,[1943SCH/RIT] and [197IBID] state explicitly that they used a-Bi2O3. [1943SCH/RIT]reported solubility measurements in 0.5 to 2.5 M NaOH which they interpret in terms of oneequilibrium V2 a-Bi2O3(cr) + 3/2 H2O + OH- <=> Bi(OH)4- (Table 7.4). [1971BID] mademeasurements in 1 M NaClO4 and reported a log Kbs4 = -4.39, which are in close agreement tothe data given by [1943SCH/RIT] (see Table 7.4 and Figure 7.6).

Extrapolation of the data of [1943SCH/RIT] to I = 0 with the SIT term gives a log Kb°S4 =-4.28 as shown in Figure 7.6. The extrapolation of the data of [197 IBID] to I = 0 with a Ae of0.03 (Figure 7.6) resulted in a log Kb° 34= -4.39. From the mean of these two values (logKb°S4 = -4.33) and log p°])4 = -18.71, a log Kso*° can be calculated:

2Bi 3 + +3H 2 O <=> a-Bi2O3(cr) + 6H+ log K*°so = -0-76

The value of this constant strongly depends from log P°];4 value. Thus the values given indifferent compilations can differ strongly (Table 7.6). Based on log K*°so = -0.76 and a AfG°of 95.55 kJ/mol for Bi3+ (Section 7.11), a AfG° of -515.99 kJ/mol is obtained for oc-Bi203(cr).

Table 7.4: Experimentally determined equilibrium data compiled for the dissolution ofbismuth oxide. These data were chosen for the evaluation of recommended valuesin the present report. Additional information for the different references seeSection 7.12: 'Comments on selected references'. Method: sol = solubilitymeasurements.

log KSo

log KSo:

-4.27 i-4.28 1-4.31 1-4.30 ]

-4.29 »-4.39

Reference

0.5 a-Bi2O3(cr)+ OH~

[1943SCH/RIT][1943SCH/RIT][1943SCH/RIT][1943SCH/RIT][1943SCH/RIT][197 IBID]

Comments

<=> Bi(OH)4

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

-

K, 1=0K, 1=0K, 1=1K, 1=1K, 1=2K, 1=1

.5

.99

.41

.97

.46

KM,

0.51

1.41.972.46

1

Medium

NaOHNaOHNaOHNaOHNaOHNaCIO,

Method

solsolsolsolsolsol

Table III of [1943SCH/RIT] contains an error in the exponent of K).

232

JNC TN8400 9 9 - 0 1 1

0.5 a-Bi?02 ^ 3

-2

-2.5

-3

O -3.5

+ -4

-5

-5.5

-6

-6.5

-7

COJO

[1943SCH/RIT]

y = -0.03x - 4.28

o o-4.33

{1971 BID]

0 1 2 3

!mi molal

Figure 7.6: Plot of log KbS4 + 0 D vs. Im for the reaction 0.5 a-Bi2O3(cr) + OH- <=>

Bi(OH)4-at 25 °C, calculated from data compiled in Table 7.4 (KbS4 denotes the

solubility product of a-Bi2O3(cr) in equilibrium with the hydroxide OH~concentration. Details of notation are given in Table 2.2 of this report). Thestraight line shows the result of the linear regression of the experimental data of[1943SCH/RIT]: Ae = 0.03; log Kb°S4 = - 4.28. Extrapolation of the data of[1971BID] to 1=0 gives a log Kb°S4 = -4.39, giving a mean log Kb°S4 of-4.33.

7.2.2 Precipitated Bi(OH)3(s)

[1993KRA/DEC] give for freshly precipitated Bi(OH)3(s) a log K*S3 of -4.70 for the reactionBi(OH)3 <=> Bi(OH)3(s) (Table 7.5), indicating a slightly higher solubility of Bi(OH)3(s) incomparison to well crystallized oc-Bi203(cr).

7.2.3 Additional data compiled for solid bismuth oxides/hydroxides

Data calculated from the AfG° values given in different compilations are difficult to interpret, asalready the ArG° values reported in different compilations for the redox equilibria between theBi3+ ion and Bi(cr) differ by 10 kJ/mol. The use of different ArG° values for Bi3+ results in adifference of 3.2 log units in the calculated log K*°so values. Thus, log K*°So values calculatedfrom AfG° values are quite uncertain. Solubility products derived from emf measurements athigher temperature (see Table 7.5) are in the range of log K*so of -4 to -6 (using a ArG° valueof 95.55 kJ/mol or a log P of-16.74 for the redox equilibria between Bi3+ and Bi(cr); see alsoSection 7.11: Bi3+/Bi(cr)).

233

JNC TN8400 9 9 - 0 1 1

Table 7.5: Additional experimentally determined equilibrium data compiled for the precipitation of bismuthhydroxide/oxide. These data were not chosen in the present report for the evaluation ofrecommended stability values. Reasons for not selecting these references are given in Section7.12: 'Comments on selected references'. Method: cal = calculated from AH and AS values, emf= emf measurements at high temperature, sol = solubility.

logK•so Reference Comments KM) Medium Method

log K*so: 2Bii+ + 3H2O <=> a-Bi2O3(cr)+ 6m

-3.78 ''-4.10 3'-5.90 4

[1978CAH/VER][1981GOR/GAV][1982LAP/KOL1

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=diluted

n/an/a

water

emfcalsol

log K*so: 2Bi3+ + 3H2O <z> y-Bi2O3(cr) + 6H+

-5.08 5 [1943SCH/RIT1 T= 298.15 K, 1=1.41 1.4 NaOH sol

log K*S3: Bi(OH)3° <=> 0.5a-Bi2O3(cr) + 1.5H2O

10.80 6

10.68[1971BID]fl971BID]

T= 298.15 K, 1=0.5-2.5T= 298.15 K,I=dil

NaOHwater

calsol

log K*S4: Bi(OH)s o 0.5a-Bi2O3(cr) + 1.5H2O + OH'

-4.32 6 [197IBID] T= 298.15 K, 1=0.5-2.5 NaOH cal

log K*S3: Bi(OH)3° « Bi(OH)3(s)

-4.70 7

-4.80 7[1993KRA/DEC]ri993KRA/DEC1

T= 296±0.5 K, 1=1T= 296±0.5 K, 1=1 cio4

solsol

1 extrapolated from measurements at 1000 K2 calculated with a AjG° of -95.55 kJ/mol for Bi2+ (Section 7.11).3 AG values calculated by [1981GOR/GAV] from measurements of heat capacity and data taken from literature4 extrapolated from measurements at higher temperatures5 calculated with a log P,4 =-20.66 (1=1.41) (See Section 7.1.4: Bi(OH)4")5 calculated by [1971BID] based of the data of [1943SCH/RIT]7 freshly precipitated Bi(OH)3 (t=30 min.)

234

JNC TN8400 99 - O i l

Table 7.6: Thermodynamic data for the precipitation of bismuth hydroxide/oxide system taken fromprevious compilations. As pointed out in Section 2 of this report only experimental data wereused for the present evaluation. The following table serves only for comparison. Medium: Wheredata refer to specific electrolyte solutions, this is indicated.

log Kso Reference Comments I (M) Medium

log Kso: 2Bi3+ + 3H2O <=> a-Bi2O3(cr) + 6H+

-4.44 '-2 H995RIS/HAL1 T= 298.15 K, I=n/a

log Kso: 2Bi3+ + 3H2O <=> Bi2O3(cr) + 6H+

-4.16 2

-4.30 2

-4.16 2

-9.17-4.67 2

-4.68 2

-6.920.18 3

-4.08 3

-5.26 2

-9.2-5.26 2

-6.92-4.71 2

-9.13-4.66 2

-4.67 2

[1952LAT][1954COU][1963WIC/BLO][1968ROBAVAL][1971NAU/RYZ][1973BAR/KNA][1976BAE/MES][1976SMI/MAR][1976SMI/MAR][1977BAR/KNA][1978ROB/HEM2][1979KUB/ALC][1981BAE/MES][1982PAN][1982WAG/EVA][1984VE/TAR][1985LOV/MEK]

T=298.15K, I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n7aT= 298.15 K, I=n/aT=298.15K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/a

001

log Kso: Bi3+ + 2H2O <=> BiOOH(s) + 3H+

-1.562 [1952LAT] T=298.15 K, I=n/a-4.1 [1982WAG/EVA] T= 298.15 K, I=n/a

log Kso: Bi3+ + 3H2O <=> Bi(OH)3(s) + 3H+

-7.49 2 [1952LAT] T= 298 K, I=diluted-5.81 2 [1971NAU/RYZ1 T= 298.15 K, I=n/a1 extrapolated from measurements at higher temperatures2 calculated with a AfG° of -95.55 kJ/mol for Bi2+ (Section 7.11).3 calculated with a log P u =-5.31 (1=0) and -7.44 (1=1) (See Section 7.1.3: Bi(OH)3°)

235

JNC TN8400 9 9 - 0 1 1

7.3 Bismuth chloride system

7.3.1 Bismuth chloride complexes

Bismuth ions form complexes with chloride. Several authors determined stability constants forthe formation of chloride complexes in chloride and perchlorate medium. The log P° valuesgiven in Table 7.7 are extrapolated to 1=0 by using the SIT equation (see Figures 7.7 - 7.12):

1.1 =3.65, Ae = -0.011.2 = 5.85, Ae = -0.161>3 = 7.62, Ae = -0.191,4 = 9.06, Ae = -0.121*5 = 8.33, AE = - 0 . 3 1

log p°i i6 = 7.64, Ae = -0.12

Based on the data available, only tentative values can be given for the formation of BiCls2" andBiCl63". These complexes however, will only become important in concentrated chloridesolution (chloride concentration > 1 M).

Bi3+H

B i 3 + ,

Bi3 +H

Bi3 +H

B i 3 + ^

Bi3+H

hCl-h2Cl-h3Cl-h 4 C hh5Cl-h 6 e l -

<=><=><=><=><=>

BiCl2+

BiCl2+

BiCl3°(aq)BiCl4-BiCl5

2-BiCl63-

loglogloglogloglog

Table 7.7: Experimental equilibrium data compiled for the bismuth(III) chloride system,according to the equilibria Bi3+ + mCl" <=> BiClm

3-m. These data were chosenfor the evaluation of recommended values in the present report. Additionalinformation for the different references see Section 7.12: 'Comments on selectedreferences'. Method: cat = cation exchange, sol = solubility, sp =spectrophotometry, pol = polarography and pot = potentiometry.

log p\,n Reference Comments I (M) Medium Method

log pltl: Bi3+ + CI- <=> BiCl2"

2.361.911.962.092.002.082.182.203.002.162.35

[1957AHR/GRE][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1963MIR/KUL][1969JOH][1970BONAVAU][1970KAN]

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298 K,T= 303.15T= 298.15

K, 1=2K, 1=6K, 1=4K, 1=3K, 1=2.5K, 1=2K, 1=1K, 1=31=4K, 1=2K, 1=5

2643

2.5213425

NaClO4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaClO4, HC1O4

NaClO4, HC1O4

LiClO4

NaClO4

NaClO4

NaCIO,

solpolpolpolpolpolpolpot

solpolSP

236

JNC TN8400 9 9 - 0 1 1


2.822.712.532.36

log fil,2--

3.503.804.584.543.904.044.223.743.504.303.824.404.444.044.663.61

log P1.3:

5.355.906.115.405.305.714.875.806.705.605.455.455.186.324.95

[1974FED/KAL][1974FED/KAL][1974FED/KAL][1990SUG/ONO1]

Bi3+ + 2CI- <=> BiCl2+

[1957AHR/GRE][1957AHR/GRE][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1963MIR/KUL][1969JOH][1970BONAVAU][1970KAN][1974FED/KAL][1974FED/KAL][1974FED/KAL][1990SUG/ONO1]

Bi3+ + 3CI- <=> B1CI30

[1957AHR/GRE][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1963MIR/KUL][1969JOH][1970BONAVAU][1970KAN][1974FED/KAL][1974FED/KAL][1974FED/KAL][1990SUG/ONO1]

T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3T= 298.15 K, 1=1

T= 298.15 K, 1=2T= 298.15 K, 1=2T= 298.15 K, 1=6T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=2.5T= 298.15 K, 1=2T= 298.15 K, 1=1T= 298.15 K, 1=3T= 298 K, 1=4T= 303.15 K, 1=2T= 298.15 K, 1=5T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3T= 298.15 K, 1=1

T= 298.15 K, 1=2T= 298.15 K, 1=6T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=2.5T= 298.15 K, 1=2T= 298.15 K, 1=1T= 298.15 K, 1=3T= 298 K, 1=4T= 303.15 K, 1=2T= 298.15 K, 1=5T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3T= 298.15 K, 1=1

0.5131

22643

2.5213425

0.5131

2643

2.5213425

0.5131

HC1O4

HC1O4

HC1O4

HC1O4

NaClO4

NaC104

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaClO4, HC1O4

LiG104

NaC104

NaC104

NaC104

HC1O4

HC1O4

HC1O4

HC1O4

NaC104

NaC104, HC1O4

NaClO4, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaClO4, HC1O4

NaC104, HC1O4

LiC104

NaClO4

NaClO4

NaClO4

HC1O4

HC1O4

HC1O4

HC1O4

potpotpotcat

potsolpolpolpolpolpolpolpotsolpolsppotpotpotcat

potpolpolpolpolpolpolpotsolpolsppotpotpotcat

237

JNC TN8400 9 9 - 0 1 1


log pli4: Bi3+ + 4CI- & BiCl4-

6.107.696.916.877.477.186.906.756.906.906.656.236.417.93

log P15:

9.298.497.688.046.756.656.727.308.607.29

6.115.958.18

log Pi,6-

7.707.546.567.368.407.066.686.00

[1957AHR/GRE][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1963MIR/KUL][1969JOH][1970BON/WAU][1970KAN][1974FED/KAL][1974FED/KAL][1974FED/KAL]

Bi3+ + 5Ct <=* BiCl52-

[1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1959DES/PAN][1961AHR/GRE][1963MIR/KUL][1969JOH][1970KAN][1974FED/KAL][1974FED/KAL][1974FED/KAL]

Bi3+ + 6CI- <=> BiCl63~

[1959DES/PAN][1959DES/PAN][1961AHR/GRE][1963MIR/KUL][1969JOH][1970KAN][1974FED/KAL][1974FED/KAL]

T= 298.15 K, 1=2T= 298.15 K, 1=6T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=2.5T= 298.15 K, 1=2T= 298.15 K, 1=1T= 298.15 K, 1=3T= 298 K, 1=4T= 303.15 K, 1=2T= 298.15 K, 1=5T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3

T= 298.15 K, 1=6T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=2.5T= 298.15 K, 1=2T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298 K, 1=4T= 298.15 K, 1=5T= 298 K, 1=0.5T= 298 K, 1=1T=298K, 1=3

T= 298.15 K, 1=6T= 298.15 K, 1=4

T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298 K, 1=4T= 298.15 K, 1=5T= 298 K, 1=0.5T= 298 K, 1=3

2643

2.5213425

0.513

643

2.5212345

0.513

642345

0.53

NaC104

NaC104, HC1O4

NaClO4, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

LiC104

NaC104

NaC104

NaC104

HC1O4

HC1O4

HC1O4

NaC104, HC1O4

NaC104) HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104, HC1O4

NaC104

LiC104

NaC104

NaC104

HC1O4

HC1O4

HC1O4

NaClO4, HC1O4

NaClO4, HC1O4

NaC104

LiC104

NaC104

NaClO4

HC1O4

HCIO;

potpolpolpolpolpolpolpotsolpolsppotpotpot

polpolpolpolpolpolpotpotsolsppotpotpot

polpolpotpotsolsppotpot

238

JNC TN8400 9 9 - 0 1 1

Bi3+ + Cl- o BiCI2+

4 6lm, molal

Figure 7.7: Plot of log (3U + 6 D vs. Im for the reaction : Bi3+ + Cl~ <> BiCl2+ at 25 °C. Thestraight line shows the result of the linear regression: Ae = - 0.01; log P0^] =3.65. Calculated from data compiled in Table 7.7.

5.5 •

4 6lm, molal

10

Figure 7.8: Plot of log pi,2 + 10 D vs. Im for the reaction : Bi3+ + 2Ch <=> BiCl2+ at 25 °C.

The straight line shows the result of the linear regression: Ae = - 0.16; log (3° 1,2 =5.85. Calculated from data compiled in Table 7.7.

239

JNC TN8400 99 - O i l

10 -i

9.5

9

CM 8.5 -

+ 8

"7.5

O) 7O

— 6.5 •

6

^ t;

?; -

Bi 3 +

Illillilt

+

m•§i

lift

S3

3Ch

isnp)

lili

<=> BiCI3°

lli-cpiiiSiÊiiill-iS

liliiii'iSiii-

s||||il:|i|||j

'-:;:p::;::i!i:j:r'"^'!::':fi:':W^-:'vf

4 6lm, molal

10

Figure 7.9: Plot of log pii3 + 12 D vs. Im for the reaction : Bi3+ + 3Ch <=> BiCl3° at 25 °C.The straight line shows the result of the linear regression: Ae = - 0.19; log (3° 13 =7.62. Calculated from data compiled in Table 7.7.

QCM

rn

11 -

10.5 i

10 -

9.5 -

9

8.5

8

7.5

7 •

6.5

R

Bi3+ + 4Ch•.;••:;.: •• .,::=:«?;:;*,;.T-.»:i:f:!y:jrk-ty.;:~y,'.:_•?,?

K;.;::^sili!i#;;3i;Sl;blai

"S::i;^li?l;iliiBiii•J^;^:*lji@iBiiiit

<=> BiCLf

; | | | | | | | | | | | | | | : | ; |

IKfe'K.: i I;:' "i::;v::::.:'.;::..;: \,v;-î";.:,;..;-.

2 4 6U, molal

8 10

Figure 7.10: Plot of log p]>4 + 12 D vs. Im for the reaction : Bi3+ + 4C1~ <=> BiCl4- at 25 °C.The straight line shows the result of the linear regression: Ae = - 0.12; log p°] 4 =9.06. Calculated from data compiled in Table 7.7.

240

JNC TN8400 9 9 - 0 1 1

Qo

+"1

CO.

D)

_g

•

11

1

10

oc.

5

1

5 -

10

9

8

7.

5 -

9

5

8

5

7

o o

oo

o

—1

5CI

y

• < = >

8

= 0

BiCI52-

o

o

31x + 8

\ 1-

o

.33

, molar 10

Figure 7.11: Plot of log pi>5 + 10 D vs. Im for the reaction : Bi3+ + 5C1~ o BiCl52- at 25 °C.

The straight line shows the tentative result of the linear regression: Ae = -0.31;P°i>5 = 8.33. Calculated from data compiled in Table 7.7.

BiCL3"

4 6lm, molal

Figure 7.12: Plot of log p l i 6 + 6 D vs. Im for the reaction : Bi3+ + 6C1~ <=> BiCl63- at 25 °C.

The straight line shows the tentative result of the linear regression: Ae = - 0.12;log P°i,6 = 7.64. Calculated from data compiled in Table 7.7.

241

JNC T N 8 4 0 0 99 - O i l

7.3.2 BiOCl(s) and Bi(OH)2Cl(s)

Thermodynamic data for precipitated BiOCl(s) have been determined by different authors (Table7.8). No direct determination of the crystalline structure of the precipitated BiOCl(s) are givenin the different experimental reports.

Table 7.8: Experimental data for the precipitation of BiOCl(s) and Bi(OH)2Cl(s). These datawere chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 7.12: 'Commentson selected references'. Method: sol = solubility, pot = potentiometry.

log Kso Reference Comments I (M) Medium Method

log Kso: Bi3+ + H2O + Or <=> BiOO(cr) + 2H+

8.42 i [1918NOY/CHO] T= 298.15 K, 1=0.001-0.5 0 HC1 pot

7.26 2

7.26 2

7.15 2

7.10 2

7.017.39

[1933FEI][1933FEI][1933FEI][1933FEI][1957AHR/GRE][1969JOH]

T=T=T=

T=T=

298.15298.15298.15298.15298.15298 K,

K, 1=0.25K, 1=0.54K, 1=0.6K, 1=0.9K, 1=21=4

0.250.540.60.924

HC1HC1HC1HC1NaClO4

NaClO4

solsolsolsolpotsol

1 estimated from potentiometric data assuming a log (3 (Bi(cr)/Bi3*) = 16.74(see Section 7.11: BP/Bi(cr))2 corrected in this report for the formation of BiCl2+, BiCl2

+, BiCl3°, BiCI4" using the constants calculated in Section7.3.1. Original values given in Table 7.9.

[1933FEI] determined the solubility of freshly precipitated Bi(OH)2Cl(s). The data compiled inTable 7.9 seem to indicate a higher solubility of Bi(OH)2Cl(s) than of BiOCl(s). [1933FEI],however, did not correct his measurements for the formation of bismuth chloride complexes,thus resulting in a too high solubility. A correction of the measurements of [1933FEI] for theformation of bismuth chloride complexes gives the same solubility product as reported by[1918NOY/CHO, 1957AHR/GRE, 1969JOH] in Table 7.8.

The potentiometric data of [ 1918NOY/CHO] for the formation of BiOCl(s) from Bi(cr) wereextrapolated in this report to I = 0 (See [1918NOY/CHO] in Section 7.12: Comments onselected references) and give a log K*°s of -8.32 for the reaction Bi(cr) + Cl~ + H2O <=>BiOCl(s) + 2H+ + 3e-. Correction with a log (3° = 16.74 (Section 7.11) for the reaction Bi(cr)<> Bi3+ + 3e- gives a K*So of 8.42 for the reaction Bi3+ + H2O+ Cl- <=> BiOCl(s) + 2H+ asgiven in Table 7.8. Extrapolation of the data compiled in Table 7.8 to I = 0 (Figure 7.13) gives:

H2O+C1- *=> BiOCl(s) + 2H+ log K*°so = 8.47

242

JNC TN8400 9 9 - 0 1 1

Bi3++ BiOCI(s)+2H+

10

9.5

9

g s.s < J r c r -

O

8

7.5

7 •

6.5

6

5.5

0

y = 0.18x + 8.47

2 3 4lm, molal

Figure 7.13: Plot of log K*so + 8D vs. Im for the reaction Bi3+ + H2O+ Or <=> BiOCl(s) +2H+ at 25 °C. The straight line shows the result of the linear regression: Ae =-0.18; log K*s° = 8.47. The value at 1=0 is taken from Figure 7.25. Calculatedfrom data compiled in Table 7.8.

7.3.3 BiCl3(s)

Thermodynamic data are compiled for the precipitation of BiCl3 in Table 7.10. The solubilityproduct of ~ 103 indicates that under environmental conditions BiOCl will be the stable solidphase and BiCl3 will not precipitate except in very acidic solutions (pH < 0).

7.3.4 Additional equilibrium data compiled for the bismuth chloride system

Table 7.9: Additional, experimentally determined equilibrium data compiled for the bismuth(III) chloridesystem and the precipitation of BiOCl(s) and Bi(OH)2Cl(s). These data were not chosen in thepresent report for the evaluation of recommended stability values. Reasons for not selecting thesereferences are given in Section 7.12: 'Comments on selected references'. Method: sol =solubility, sp = spectrophotometry, pol = polarography and pot = potentiometry.

log p l l i r Reference Comments I (M) Medium Method

log Pu: Bi3+ + Cl- <=> BiCl2'

2.43 '2.44 2

3.70 3

[1957NEW/HUM][1969CAR][1974FED/KAL1

T= 298.15 K, 1=1T= 298.15 K, 1=0.1-1T= 298 K, 1=0

HC1O4

HC1HC1O,

sppolpot

243

JNC TN8400 99 - Oil

Table

log Pi

4.46 '3.10 2

5.50 3

7.9: continued

.2.' Bi3 + + 2Cl~ <=> BiCl2+

[1957NEW/HUM][1969CAR][1974FED/KAL]

T= 298.15 K, 1=5T= 298.15 K, 1=0.1-1T= 298 K, 1=0

1

0

HC1O4

HC1HC1O4

sppolpot

log pu: Bi3+ + 3Cl~ <=> BiCl30

[1957NEW/HUM][1969CAR]ri974FED/KAL1

5.76 »3.74 2

6.90 3

T= 298.15 K, 1=5T= 298.15 K, 1=0.1-1T= 298 K, 1=0

5

0

HC1O4

HC1HC1O4

sppolpot

log P1A: Bi3+ + 4CI- <=> BiClf

[1953BAB/GOL][1957NEW/HUM][1969CAR][1974FED/KAU|

5.42 l

6.24 !

3.77 2

7.90 3

T= 298.15 K, 1=2,3 or 3T= 298.15 K, 1=5T= 298.15 K, 1=0.1-1T= 298 K, 1=0

35

0

KNO3

HC1O4

HC1HC1O4

potsppolpot

log Pu: Bi3+ + 5CI- <=> BiCl52-

7.50 ' [1953BAB/GOL]6.72 ' [1957NEW/HUM]7.72 4 [1959AHR/GRE]7.00 3 f ! 9 7 4 F E D / K A L 1

T= 298.15 K, 1=2,3 or 3 3T= 298.15 K, 1=5 5T= 298.15 K, 1=2 2T= 298 K, 1=0 0

KNO3HC1O4

NaClO4

HC1O4

potsppotpot

log Pi,6:

6.42 5

7.56 4

7.30 3

log Kso:

2.25 6

2.56 6

2.49 6

2.53 6

log Kso:

8.63 7

8.13 8

Bi3+ + 6CI- » BiCl63-

[1953BAB/GOL][1959AHR/GRE]T1974FED/KAL1

T= 298.15 K, 1=2,3 or 3T= 298.15 K, 1=2T= 298 K, 1=0

Bi3+ + 2H2O + Ct <=> Bi(0H)2Cl(s) + 2H+

[1933FEI][1933FEI][1933FEI][1933FEI1

T= 298.15 K, 1=0.3T= 298.15 K, 1=0.5T= 298.15 K, 1=0.6T= 298.15 K, 1=0.9

Bi3+ + H2O + Cl- <=> BiOCl(s) + 2H+

[1918NOY/CHO][1969VAS/GRE1

T= 298.15 K, 1=0.001-0.5T= 298.15 K, 1=1-3

320

0.30.50.60.9

00

KNO3

NaC104

HC1O4

HC1HC1HC1HC1

HC1O4

HC1O4

potpotpot

solsolsolsol

potpot

1 as reported by [1970KAN]2 I not constant3 not reported how data were extrapolated to 1=0 by [1974FED/KAL]4 corrected later by [1961AHR/GRE]5 reported by [1957AHR7GRE], Original not available6 formation of Bi chloride species neglected, corrected values in Table 7.87 value corrected and selected by [1918NOY/CHO], Calculated in this report from E° = 0.1599 V and log K (Bi(cr)/Bi3+) =

16.74 (Section 7.11). Additional values by [1918NOY/CHO] are given in section 7.27.8 as reported by [1985LOV/MEK], extrapolated to 1=0 by [1985LOV/MEK], Calculated in this report from E° = 0.1697 V

and log K (Bi(cr)/Bi?+) = 16.74 (Section 7.11)

244

JNC TN8400 9 9 - 0 1 1

Table 7.10: Thermodynamic data for the bismuth(III) chloride system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log P,m

log Pij:

2.202.782.94 '3.12 >3.42 '2.362.202.253.502.402.303.422.33

log Pi.2:

3.503.504.515.504.004.204.993.45

Reference

Bi3+ + Cl- <=> BiCl2+

[1967AHR][1967VAS/LOB][1967VAS/LOB][1967VAS/LOB][1967VAS/LOB][1976SMI/MAR][1976SMI/MAR][1982WAG/EVA][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1985BAB/MAT][1985LOV/MEK1

Bi3+ + 2Cl~ <=> BiCl2+

[1976SMI/MAR][1976SMI/MAR][1982WAG/EVA][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1985BAB/MAT][1985LOV/MEK1

Comments

T= 298.15 K, 1=3T=298.15K, 1=4T=298.15K, 1=5T= 298.15 K, 1=6T= 298.15 K, 1=0T= 293.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0

T= 293.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0

3456023001400

23001400

Medium

LiC104

NaClO4

NaC104

NaC104

NaC104

log p,j: Bi3+ + 3CI- t=> BiCl30

5.405.807.105.206.006.115.28

[1976SMI/MAR][1976SMI/MAR][1982SMI/MAR][1982SMI/MAR][1982SMI/MAR][1985BAB/MAT][1985LOV/MEK1

T= 293.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=4T= 298.15 K, 1=0T= 298.15 K, 1=0

2301400

log J3,i4: Bi3+ + 4CI- t=> BiClf

8.48 2

6.106.806.918.106.40

[1952LAT][1976SMI/MAR][1976SMI/MAR][1982WAG/EVA][1982SMI/MAR][1982SMI/MAR]

T= 298.15 K, 1=0T= 293.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K,I=0T= 298.15 K, 1=0T= 298.15 K, 1=1

023001

245

JNC TN8400 9 9 - 0 1 1


7.308.126.04

[1982SMI/MAR][1985BAB/MAT][1985LOV/MEK1

T= 298.15 K,T= 298.15 K,T= 298.15 K,

1=41=01=0

400

log p]:5: Bi3+ + 5CI- <=> BiCl52~

7.696.707.308.306.65

[1960FRI/SAR][1976SMI/MAR][1976SMI/MAR][1982SMI/MAR][1985LOV/MEK]

T= 298.15 K, I=n/aT= 293.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=0

2340

log pl6: Bi3+ + 6CI- <=> BiCl63~

6.60 [1976SMI/MAR] T= 293.15 K, 1=2 27.40 [1976SMI/MAR] T= 298.15 K, 1=3 37.40 [1982WAG/EVA] T= 298.15 K, 1=0 07.90 [1982SMI/MAR] T= 298.15 K, 1=4 46.50 fl985L0V/MEK1 T= 298.15 K, 1=0 0_

log Kso: Bi3+ + 3Cl~ <=> BiCl3(s)

3.653.472.632.630.732.23

2

2

2

2

[1952LAT][1963WIC/BLO][1977BAR/KNA][1979KUB/ALC][1982WAG/EVA][1985LOV/MEK]

T=T=

T=T=

298.15298.15298.15298.15298.15298.15

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/a

log Kso: Bi3+ + 2H2O + Ct <=> Bi(OH)2Cl(s) + 2H+

ZS [1982WAG/EVA1 T= 298.15 K, I=n/a

log Kso: Bi3+ + H2O + Ct <=> BiOCl(s) + 2H+

8.6 2

8.14 2

8.45 2

8.43 2

-7.80 3

-6.47 3

-6.75 3

6.407.49

[1952LAT][1971NAU/RYZ][1977BAR/KNA][1979KUB/ALC][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1982WAG/EVA]ri985LOV/MEKl

T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/aT= 293.15 K, 1=0T= 293.15 K, 1=1T= 293.15 K, 1=3T= 298.15 K,I=n/aT= 298.15 K,I=n/a

013

1 the values given in [1967VAS/LOB] are extrapolated to I = 4, 5 and 6 from data compiled in [1964SIL/MAR]2 calculated with a A,G° of -95.55 kj/mol for Bi2+ (Section 7.11).3 the data compiled in [1976SM1/MAR] seem to have the wrong sign.

246

JNC TN8400 9 9 - 0 1 1

7.4 Bismuth perchlorate

7.4.1 BiClO42+ or [Bi

[1993KRA/DEC] reported a log pi;i of 3.5 at I = 1 for the formation of BiC1042+. However,

this constant for BiC1042+ seems to be much too large compared to the constants of the bismuthcomplexes with nitrate and chloride. [1993KRA/DEC] present in their report no directexperimental proof for the predominance of the species BiC1042+ (instead of Bi3+) at pH < 2.[1990SUG/ONO] deduced from cation-exchange measurements the dominance of theBi(H2O)6

3+ - C1O4- ion pair in 1 M NaC104 at a pH of 1. Additionally, [1982HAT/SUG] and[1987SUG/SHI] showed with hydrolysis experiments, that the presence of both nitrate andperchlorate had only a small influence while the influence of chloride (and thus also the complexformation with chloride) is more important. Therefore, the value calculated by [1993KRA/DEC]is not recommended in this report.

7.4.2 BiOClO4(precip)

In 1 M perchlorate medium, BiOClO4(s) precipitates below a pH of 6 [1993KRA/DEC]. At I =1, [1993KRA/DEC] calculate a log Ks0* for BiOClO4(precip) of 0.87. Extrapolation to 1=0assuming a Ae of 0 results in a tentative value log Kso*° of -0.78 for the reaction Bi3+ + H2O +CIO4- <=> BiOClO4(precip) + 2H+.

Table 7.11: Experimentally determined equilibrium data compiled for the bismuth(III) perchlorate system andthe precipitation of Bi0C104(s). These data were not chosen in the present report for theevaluation of recommended stability values. Reasons for not selecting this reference is given inthe text. Method: sol = solubility.

log P,m

log pu: Bi3+

3.5

log K*so: Bi3+

0.87

Reference

+ ClOf <=> BiClO4

[1993KRA/DEC]

+ H2O + ClOf <=>

ri993KRA/DEC]

Comments

2 +

T= 296±0.5 K,

BiOClO4 (precip) +

T= 296+0.5 K,

1=1

2H+

1=1

KM)

1

1

Medium

cicv

cicv

Method

sol

sol

247

JNC TN8400 9 9 - 0 1 1

7.5 Bismuth fluoride system

In the presence of fluoride, bismuth fluoride complexes can be formed. Some constants weredetermined by [1967LOM/VAN] and [1969BON] in 1.9 M and 2 M perchloric acid (Table7.12). Unfortunately, there are not enough data available for the extrapolation to I = 0.

From the thermodynamic data given in the compilations of [1977BAR/KNA] and[1979KUB/ALC] the formation of quite insoluble BiF3(s) is expected. Direct measurements ofthe solubility, or details about the formation of this solid species, have not been found in theliterature, however.

No thermodynamic data for the formation of bismuth fluoride complexes or solids arerecommended in this report.

Table 7.12: Experimentally determined equilibrium data compiled for the bismuth(III) fluoride system. Thesedata are not chosen in this report for the evaluation of recommended log (3 values. Reasons fornot selecting these references are given in Section 7.12: 'Comments on selected references'.Method: cat = cation exchange, pol = polarography

log (3

log J5

1.431.41

log ft

1.760.30

logfi

2.70

l.m

1

2

7.2-"

1

2

1

2

Bi3+

Reference

+ HF <=> BiF2+ +

[1967LOM/VAN][1969BON]

+ 2HF <=> BiF2+ H

[1967LOM/VAN][1969BON1

+ 3HF <=> BiF3° 4

[1969BON1

H+

-2H

-3H-

Comments

T= 298,1=1.89T= 303 K, 1=2

T= 298,1=1.89T= 303 K, 1=2

T= 303 K, 1=2

I(M)

1.92

1.92

2

Medium

HC1O4

NaBr, HC1O4

HC1O4

NaBr, HC1O4

NaBr, HC1O4

Method

catpol

catpol

pol>pH<02 pH < 0.3

248

JNC TN8400 99 - O i l

Table 7.13: Thermodynamic data for the bismuth(III) fluoride system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log P,.n Reference Comments I (M) Medium

log pu: Bi3+ + HF <=> BiF2+ + H+

1.701.42

[1980BON/HEF][1976SMI/MAR1

T= 298.15 K, 1=1.89T= 298 K, 1=2

1.9 HC1O4

2

log Pli2:Bi3+ + 2HF <=> BiF2+ + 2H+

2.30 [1980BON/HEF1 T= 298.15 K, 1=1.89 1.9 HC1O,

log Pu: Bi3+ + 3HF <^> BiF3° + 3H+

4.70 [T980BON/HEF1 T= 303 K, 1=2 2 NaBr, HC1O4

log Kso: Bi3+ + 3F~ <=> BiF3(s)

15.4 '14.7 !15.8 '

[1963WIC/BLO][1977BAR/KNA]ri979KUB/ALC1

T= 298.15 K, I=n/aT= 298.15 K, I=n/aT= 298.15 K, I=n/a

1 calculated with a AfG° of -95.55 kJ/mol for Bi2+ (Section 7.11)

249

JNC TN8400 9 9 - 0 1 1

7.6 Bismuth carbonate system

No data concerning dissolved bismuth carbonate complexes have been found in the literature.[1984TAY/SUN] determined with X-ray analysis the conditions under which(BiO)4(OH)2CO3(cr) and (BiO)2CO3(cr) were predominantly formed in aqueous carbonatesolutions (I, pCO2, pH were varied) and extrapolated their results to 1=0 with the Daviesequation. They observed the formation of these solids at room temperature. From the AfG°values given in their report, the following log K*°so values can be calculated (Table 7.14):

CO32-+6H2O

CO32-+2H2O

(BiO)4(OH)2CO3(cr) +(BiO)2CO3(cr) + 4 H+

10 H+ log K*°so = 8.68log K*°so = 14.27

Table 7.14: Experimentally determined equilibrium data compiled for the bismuth(III)carbonate system. These data were chosen for the evaluation of recommendedvalues in the present report. Additional information for the different references seeSection 7.12: 'Comments on selected references'. Method: cat = cation exchange,sol = solubility

log Reference Comments I (M) Medium Method

log KSo*: 4Bi3+ + CO32~ + 6H2O <=> (BiO)4(OH)2CO3(s) + 10 H+

8.68 ' [1984TAY/SUN] T= 298 K, 1=0.3-1 0 KOH sol

log Kso*-- 2Bi3+ + CO32~ + 2H2O <=> (BiO)2CO3(s) + 4 H+

14.27 > [1984TAY/SUN] T= 298 K, 1=0.3-1 0 KOH sol1 [1984TAY/SUN] gives (based on a selected A,G° (Bi2O3) of-493.5 kJ/mol), A ^ 0 of-1678 kJ/mol and -945 kJ/mol for

(BiO)4(OH)2CO3(s) and (BiO)2CO3(s), respectively. This values are corrected for difference in A(G° of Bi2O3 between[1984TAY/SUN] and this report (AfG° = -515.99 kJ/mol, Section 7.2.1).

250

JNC TN8400 9 9 - 0 1 1

7. 7 Bismuth nitrate system

7. 7.1 Bismuth nitrate complexes

Bismuth forms weak complexes with nitrate. Several authors determined the equilibriumconstants for bismuth nitrate complexes. The existence of bismuth nitrate complexes in acidsolutions is well established by anion exchange results and Raman spectra ([1954NEL/KRA],[1968OER/PLA]). Extrapolation of the measurements of [1967KAP/NAP], [1971 FED/KALI],[1990SUG/ONO1] and [1993KRA/DEC] as given in Table 7.15 results in the followingformation constants (Figures 7.14 - 7.17):

Bi3+ + 3NO3-

BiNO32+

Bi(NO3)2+Bi(NO3)3°Bi(NO3)4-

log P°,,, = 1.97log p° l i 2 = 2.95log (3°1,3 = 3.62log ( 3 \ 4 = 3.09

The formation constants for Bi(NO3)s2~ and Bi(NO3)6

3~ have only been determined at highionic strength and show a large difference (see Table 7.17). Thus, no constants for theformation of Bi(NO3)s

2~ and Bi(NO3)63~ are selected in this report. It can be stated that the

complex formation of bismuth with nitrate is much weaker than with chloride.

Table 7.15: Experimental equilibrium data compiled for the bismuth(III) nitrate system,according to the equilibria Bi3++ mN03- <=> Bi (NO3)m

3-m. These datawere chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 7.12: 'Commentson selected references'. Method: cat = cation exchange, pot = potentiometry, sol =solubility.

log P, Reference Comments I (M) Medium Method

log J5U: Bi3+ + NO 3- t=> BiNO32+

0.960.720..810.720.720.920.730.741.20

[1967KAP/NAB][1971FED/KAL1][1971FED/KAL1][1971 FED/KALI][1971FED/KAL1][1971 FED/KALI][1971 FED/KALI][1990SUG/ONO1][1993KR A/DEC]

T= 298.15 K, 1=1T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=1T= 296+0.5 K, 1=1

10.5

1234311

NaClO4

LiClO4

LiClO4

LiClO4

LiClO4

LiClO4

LiClO4

HC1O4

NO,"

catpotpotpotpotpotpotcatsol

25;

JNC TN8400 9 9 - 0 1 1


log Bi3+ + 2NO3- <=> Bi(NO3)2

1.580.940.900.980.961.231.161.22

log Pl,i

1.930.720.200.111.080.881.54

l°g Pi,'

0.58-0.220.400.54

[1967KAP/NAB][1971FED/KAL1][1971FED/KAL1][1971FED/KAL1][1971FED/KAL1][1971 FED/KALI][1971 FED/KALI][1990SUG/ONO1]

T= 298.15 K, 1=1T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=1

j : Bi3+ + 3NO3~ & Bi(NO3)3°

[1967KAP/NAB][1971 FED/KALI][1971FED/KAL1][1971FED/KAL1][1971FED/KAL1][1971 FED/KALI][1990SUG/ONO1]

T= 298.15 K, 1=1T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=3T= 298.15 K, 1=1

t: Bi3+ + 4NO3~ <=> Bi(NO3)4-

[1971FED/KAL1][1971FED/KAL1][1971FED/KAL1][1971 FED/KALI]

T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4T= 298.15 K, 1=3

10.5

123431

1123431

2343

NaC104

LiC104

LiC104

LiC104

LiClO4

LiC104

LiC104

HC1O4

NaClO4

LiC104

LiC104

LiC104

LiC104

LiClO4

HC1O4

LiC104

LiClO4

LiC104

LiC104

catpotpotpotpotpotpotcat

catpotpotpotpotpotcat

potpotpotpot

252

JNC TN8400 9 9 - 0 1 1

2 3lm, molal

Figure 7.14: Plot of log p u + 6D vs. Im for the reaction Bi3+ + NO3- <=>BiNO32+ at 25 °C.

The straight line shows the result of the linear regression: Ae = - 0.07; log (31,1° =1.97. Calculated from data compiled in Table 7.15.

Qo

+CM,-T

CO.

enjo

5

4.54

3.5 •

3

2.5 •

2

1.5 -

'

0.5 •

0

'• :•••". ".yr-y: y: -y/î^r- '-'• '•-'• '•-•'••'h'.:.-;::-.rJ^.''-''-'-^['[

•'. \\ : :;.-Y-. •; "."••••;},': •:.-]•:--l:%]-^^'^.:\^{"\:\v--: •-.••:\-:^:-::;^:-"-.-:-..v^::K™;ii^:..^j::;;:;ria:|;;j;:i:;:::i:;i::ij;iv:E:;-

;?S;C;f -ff^Qw^ i l l l i l lsS

mmmmmmm-yiK^'O-;Sht0î:i"f^fvmbX:K

' - • . . ' . '•: : . i : -i -.-'•'•':: •"-< ':• .'•::.'0 ~ '' ''i\::-"-'-T-Js- :•..:•'• •">£[••

;•• • : ; ; ! - ; ; . : ! : : : r i i ' j ^ v i ' - i i j i " : f: -: i ; i"?^ •-.".*;--:i*j=; I•;•;!!;-:;i i : ; : : ; :

"-: :ii!;:;.: i f iO.::!= J;y?;w j ss ; ! J;::HSl:i;;:y

iifgiijiflnhfgsjjiif;

lillillillli!li3!S3|S||p£,|i^: ;sjj!i:|i.- ij:^.i;:i; :iîii:-

Bi:li'!:^b::.j::i::':;.li:':::';:-h'^.'-:

Kftiij W&i- iBTW:'•:::::;:^p::^|^^j-:'i;v..;;-^>

5 - : ; ' g l ' : : L " : f r :

i::!;:!af!::!g|;;-i:;i;^

".:.h;;,^:;j.s-^;:;i;:i;i:.-::v:.

Hsi i'!i lWh l' j: i:i!:-V;'\:

••'ISljifslJjtiiilSr1:1?1':'

2 3lm, molal

Figure 7.15: Plot of log p u + 10D vs. Im for the reaction Bi3+ + 2NO3- o Bi(NO3)2+ at 25

°C. The straight line shows the result of the linear regression: Ae = - 0.15; logPi,20 = 2.95. Calculated from data compiled in Table 7.15.

253

J N C T N 8 4 0 0 99 - O i l

+ 1

2DCO

CO.

log

4.5

4

3.5

3

2.5

2

1.5

1

0.5 •

n -

Bi3++3NO3-

i!l!iilt|iilJll!ipjjillligll|f;::|;^l||ll|i|i|::

lfll||3j||:llillllllgiii;^:-lli;!fflliIfiliSsl|iSSi§

='••:-•• - . . : . | : ' : ' . - ) . | r . : ; v , - ; - - . . - : v . . . ; : v . ' / f : - ^ - r : ' -

oBi(N03)3°

i|S|l|fi;;;:Jl|f-::::i||:ii;s|:p«^&:S-&»;:-o;ri-:::--:i;-:f?&iiiiffiW^

- ::•-• " " • • • •i: "

;i v '

11 - - "

1 * •

:=. , ! . ' : J ' " " . ; ; : : : r;:

1;;-: .^ . . - • ' ; : : • • . ; • ; • : - .

=:.:"' ;•: ; - : - . [ - ~ : = " : ' - ' : - ~ - . ;••••••;.:"/• • :V • • . : . . . v : : : » : . " - : : ; . •-•::• .. '•:-:•*' •: "

r---: F-'^':1-j::i>:ii;:ii-;=:J!- O ^ ' : •-•::.[;.:: i'=7jii^-iKJ11".:-' :: :

: . ; - , ; ; ; ; i ^ ' ; - : . : . . | ; : . . ; : . ; : . ' ; . | : : ; " v i . " ; ' ; ; ; ': ==-E:-;=§::... -.-\r-.-\ •:•. ] ••,;;•

|||i|l|Il|f|l||i|;;|:

™'f•i''A- AQv^LM^Q: COi;^^':-:-•^ " U * U O A T OiQ&u>

1 2 3lm, molal

Figure 7.16: Plot of log p1>3 + 12D vs. Im for the reaction Bi3+ + 3NO3- <=> Bi(NO3)3° at 25°C. The straight line shows the result of the linear regression: Ae = 0.03; log (3°it3= 3.62. Calculated from data compiled in Table 7.15.

o •

4.5

4

+ 3

^ 2 . 5

D) 2

1.5

0.5 -

n -

Bi3++4NO3"

1-1:

:'-i^r

1T

:ii-ii":i:fi->-r-5:iivi;

;;i;i;:iiuin;ii!iH;:;^:^V^KL

:o.-'

iSfii}i|}|liipiiif|ii|:

^B i (NO 3 ) 4 "

iil|6il;|i!l|S;llllSI:

l!lff-|P|:p3 ;:;!-:::f:||SSB:;f: • " . . • : L - . K : i i r . p i ' . : t : . i i ! 1 . i : - . : ; j i j i . : . ; : • . • : ; : . ' i . ! . ' ; j i ' j 1 " v v ; j : : ' ' - ' ; : : • ; " .

• , > " " i i i j ! . i - ' - i . : : • • : - • • • • ' • • • • ' . " • . ' : ' - l : - : : : - : - ' . ; : - v : . • ; : ' : ; • ; • ' • • ; , -

»•- ""\Ji\J- 1 -jC- *T" »J • U w ''"••••

"' ;'*':"' i': ' ':' ' ''1 ";;': ' "'':v" —

2 3lm, molal

Figure 7.17: Plot of log p l i 4 + 12D vs. Im for the reaction Bi3+ + 4NO3- <=> Bi(NO3)4- at 25°C. The straight line shows the result of the linear regression: Ae = 0.01; log P°i,4= 3.09. Calculated from data compiled in Table 7.15.

254

JNC TN8400 9 9 - 0 1 1

7.7.2 BiONO3(s)

In 1 M nitrate medium BiONO3(s) precipitates below pH 6 [1993KRA/DEC]. The data given in[1951SWI/GAR] were recalculated in this report including a correction for the formation ofBiNO3

2+ complex (see [1951SWI/GAR] and Figure 7.26 in Section 7.12: Comments onselected references). Extrapolation of the data of [1951SWI/GAR] and [1993KRA/DEC] (Table7.16) is shown in Figure 7.18 and result in:

Bi3+ + H2O + NO3- BiONO3(s) + 2H+ log Ks0*° = 2.75; Ae = - 0.09

Bi3++NO3-+H2O<^BiONO3(s)+2H+

4.5 -

4 -•

Q 3.5 |CO

+ 3 -:G

2.5 -o

1 . 5 -••

1

0.5 -

00.5 1

lm, molal1.5

Figure 7.18: Plot of log Ks0* + 8D vs. Im for the reaction Bi3+ + H2O + NO3~ <=> BiONO3(s)+ 2H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae =-0.09; log KSo*° = 2.75. The value at 1=0 results from Figure 7.29. Calculatedfrom data compiled in Table 7.16.

255

JNC TN8400 9 9 - 0 1 1

Table 7.16: Experimental equilibrium data compiled for the precipitation of BiONC>3(s). Thesedata were chosen for the evaluation of recommended values in the present report.Additional information for the different references see Section 7.12: 'Commentson selected references'. Method: sol = solubility.

log KSo Reference Comments I (M) Medium Method

log Kso: Bi3+ + H20 + N03-

2.75 l [1951SWI/GAR]1.20 [1993KRA/DEC]

+ 2H+

T= 298.15 K, 1=0.1-0.025 0 HNO3 solT= 296+0.5 K, 1=1 1 NO," sol

1 recalculated considering the formation of BiNO32+ complexes; See Figure 7.26 ([1955SWI/GAR]) in the Section 7.12:

Comments on selected references)

7. 7.3 Additional equilibrium constants compiled for the bismuth(III) nitrate system

Table 7.17: Additional, experimentally determined equilibrium data compiled for the bismuth(III) nitratesystem, according to the equilibria Bi3+ + mNO3~ <=> Bi (NO3)m

3-m. These data were not chosenin the present report for the evaluation of recommended stability values. Reasons for notselecting these references are given in the text. Method: cat = cation exchange, sol = solubility,pot = potentiometry.

log P,.n Reference Comments I (M) Medium Method

log f}u: Bi3+ + NOf <=> BiNO32+

1.74 ' [1971FED/KAL1][1972BON][1974FED/KAL][1974FED/KAL][1974FED/KAL]ri974FED/KAU]

1.26 2

1.74 3

0.72 3

0.81 3

0.72 3

T=298.15K, 1=0T= 298.15 K, I=n/aT= 298 K, 1=0T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3

0

00.5

13

LiClO4HNO3

HC1O4

HC1O4

HC1O4

HC1O4

potsolpotpotpotpot

log Pu: Bi3+ + 2NO3~

2.55 >2.55 3

0.95 3

0.90 3

0.96 3

[1971FED/KAL1][1974FED/KAL][1974FED/KAL][1974FED/KAL][1974FED/KAL]

T= 298.15 K, 1=0T= 298 K, 1=0T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3

00

0.513

LiClO4

HC1O4

HC1O4

HC1O4

HC1O4

potpotpotpotpot

log PJ

0.72 3

0.11 3

• Bi3+ + 3NO3~ « Bi(NO3)3°

[1974FED/KAL][1974FED/KAL1

T=298K, 1=1T= 298 K, 1=3

HC1O4

HC1O4

potpot

256

JNC TN8400 99 - Oil


log p]i4: Bi3+ + 4NOf <=> Bi(NO3)4

1.99 4

-0.22 3[1967KAP/NAB][1974FED/KAL1

T= 298.15 K, 1=1T= 298 K, 1=3

1 33

NaNO3

HC1O.catpot

log Pu: Bi3+ + 5NOf <=> Bi(NO3)5

1.81 4

-0.10

.2-

[1967KAP/NAB]T1971 FED/KALI]

T= 298.15 K, 1=2T= 298.15 K, 1=3

2 3

3NaNO3

LiClO.catpot

log Bi3+ + 6NOf<=> Bi(NO3)6.3-

1.25 4

-0.40

log Kso: Bi3

2.55 5

[1967KAP/NAB] T= 298.15 K,[1971FED/KAL11 T= 298.15 K,

+ + H2O + NOf <=> BiONO3(s) + 2H+

[1951SWI/GAR1 T= 298.15 K,

1=31=3

1=0

3 3

3

0

NaNO3

LiClO4

HNO,

catpot

sol1 extrapolation by [1971 FED/KALI] to 1=0 with Vasilev equation.2 In [1972BON]. reported from Yatsimirski, 1954 (in Russian, not available). Experimental details not known.3 same data as in [1971 FED/KALI]4 ionic strength not constant, estimated5 original value (not including complex formation of bismuth with nitrate ; recalculated value in Table 7.16).

Table 7.18: Thermodynamic data for the bismuth(III) nitrate system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison.

log PV

log

log

Reference Comments

Bi3+ + NOf <=> BiNO32+

Bi3+ + 2NOf <=> Bi(NO3)2+

1.700.720.810.720.720.92

[1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMI/MAR]

T= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4

00.5

1234

2.500.940.900.980.961.23

[1976SMI/MAR][1976SMI/MAR][1976SMI/MAR][1976SMIMAR][1976SMI/MAR]fl976SMI/MARl

T= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=4

00.5

1234

257

JNC TN8400 9 9 - 0 1 1


log / 3 U : Bi-'+ + 3NO3~ <=> Bi(NO3)3°

0.70 [1976SMI/MAR] T= 298.15 K, 1=1 10.20 [1976SMI/MAR] T= 298.15 K, 1=2 20.10 [1976SMI/MAR] T= 298.15 K, 1=3 31.10 [1976SMI/MAR] T= 298.15 K, 1=4 4_

log PJA: Bi3+ + 4NOf <=> Bi(NO3)4~

0.60 [1976SMI/MAR] T= 298.15 K, 1=2 2-0.20 [1976SMI/MAR] T= 298.15 K, 1=3 30.40 H976SMI/MAR1 T= 298.15 K, 1=4 4_

log Ks0: Bi3+ + H2O + NOf <=> BiONO3(s) + 2H+

2.55 [1976SMI/MAR] T= 298.15 K, 1=0 02.64 ri982WAG/EVA] T= 298.15 K, I=n/a

258

JNC TN8400 9 9 - 0 1 1

7.8 Mixed bismuth nitrate and chloride system

7.8.1 Bismuth chloride nitrate complexes

The formation of mixed bismuth chloride nitrate complexes is also described in the literature.[1974FED/KAL] measured log p values for the formation of mixed BimCln(NO3)0(

3m-n-°)+

complexes (given in Table 7.19). Extrapolation of these values to 1=0 with SIT as shown inFigures 7.19 to 7.23 results in:

BiCl(NO3)+ log P° u , i =5.16, Ae = -0.11BiCl2(NO3)° log p \ 2 ' i = 6.86, Ae = -0.30BiCl3(NO3)- log (3°!^] = 8.09, AE = -0.45BiCl(NO3)20 log p°i'i'2 = 5.28, Ae = -0.27BiCl2(NO3)2- log p° l i 2 i 2 = 5.75, As = -0.67

Thermodynamic formation constants for other bismuth chloride complexes (BiCl(NO3)3~,BiCl2(NO3)3

2-, and BiCl3(NO3)22~) were determined only at I = 3 and could therefore not be

extrapolated to I = 0. These values are compiled in Table 7.20. Again, it can be observed thatthe influence of chloride is stronger than the influence of nitrate.

Table 7.19: Experimental equilibrium data compiled for the bismuth(IH) chloride nitratesystem, according to the equilibria Bi3+ + mCl - + nNO3~ <=> BiClm(NO3)n

3-m-n.These data were chosen for the evaluation of recommended values in the presentreport. Additional information for the different references see Section 7.12:'Comments on selected references'. Method: cat = cation exchange, sol =solubility, sp = spectrophotometry, pol = polarography and pot = potentiometry.

Pl.m,nReference Comments I (M) Medium Method

log pUJ: Bi3+ + Cl- + NO3-

3.403.363.15

[1974FED/KAL][1974FED/KAL][1974FED/KAL]

BiCl(NO3)+

T= 298 K, 1=0.5T= 298 K, 1= 1T= 298 K, 1=3

0.513

HC1O4

HC1O4

HC1CX

potpotpot

log ft

4.605.175.00

• Bi3+ + 2CI- +2NO3- d BiCl2(NO3)°

[1974FED/KAL] T= 298 K, 1=0.5[1974FED/KAL][1974FED/KAL]

T= 298 K, 1=1T= 298 K, 1=3

0.513

HC1O4

HC1O4

HC1O4

potpotpot

259

JNC TN8400 9 9 - 0 1 1


log p1>3J: Bi3+ + 3CI- + NO3- <=> BiCl3(NO3)-

6.30 [1974FED/KAL] T= 298 K, 1=0.56.14 [1974FED/KAL] T= 298 K, 1=16.93 [1974FED/KAL] T= 298 K, 1=3

0.513

HC1O4

HC1O4

HCKX

potpotpot

log $UX- Bi3+ + CI- + 2NO3~

3.19 [1974FED/KAL]3.36 [1974FED/KAL]3.38 [1974FED/KAL]

& BiCl(NO3)20

T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3

0.513

HC1O4

HC1O4

HC1O,

potpotpot

log $

3.854.305.29

2: Bi3+ + 2CI- + 2NO3~ <=> BiCl2(NO3)2-

[1974FED/KAL][1974FED/KAL][1974FED/KAL]

T= 298 K, 1=0.5T= 298 K, 1=1T= 298 K, 1=3

0.513

HC1O4

HC1O4

HC1CX

potpotpot

BiCINO,

1 2lmi molal

Figure 7.19: Plot of log p u > ] + 10D vs. Im for the reaction Bi3+ + Cl" + NO3- <=> BiClNO3+at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.11;log p \ u ° = 5.16. Calculated from data compiled in Table 7.19.

260

JNC TN8400 99 - O i l

BiCI2(NO3)°

1 2lm, molal

Figure 7.20: Plot of log p u , i + 12D vs. Im for the reaction Bi3+ + 2C1~ + NO3- <=> BiCl2NO30 at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.30;log Pi^,!0 = 6.86. Calculated from data compiled in Table 7.19.

101 \J

9.5

9QCM 8.5 •

+ 8

" : 7.5 •

O) 7 •o

6.56 -

5.5 -

5 • — — - — " • • • " ' ' i

BiCI3(NO3)-111 Illlilil'-.vi.-i'vi::- r " i : ' : : ; ^ ; / i : ifVX ••'". . '•-" •

'• i •'•'-':^': ^'.Lii.ii ' i::;-::!.'. >:"-" / " "•'

1 2 3L molal

Figure 7.21: Plot of log p1)3>1 + 12D vs. Im for the reaction Bi3+ + 3C1" + NO 3 - <=>BiCl3(NO3)~ at 25 °C. The straight line shows the result of the linear regression:Ae = - 0.45; log p1 > 3 i l

o = 8.09. Calculated from data compiled in Table 7.19.

261

J N C T N 8 4 0 0 99 - O i l

BiCI(NO3)2

CO.

O) 5 fio

4.5 £4

3.5 |

3 - t - --F-

1 2 3lm, molal

Figure 7.22: Plot of log p 1 > l j 2 + 1 2 D vs. Im for the reaction Bi3+ + C h + 2NO3- <=>BiCl(NO3)2° at 25 °C. The straight line shows the result of the linear regression:Ae = - 0.27; log (3U>2

O = 5.28. Calculated from data compiled in Table 7.19.

BiCI2(NO3)2

y = 0.67x4-5,75

1 2 3lm, molal

Figure 7.23: Plot of log pi,2,2 + 12D vs. Im for the reaction Bi3+ + 2C1~ + 2NO3~ <=>BiCl2(NO3)2~ at 25 °C. The straight line shows the result of the linear regression:Ae = - 0.67; log Pi,2,2° = 5.75. Calculated from data compiled in Table 7.19.

262

JNC TN8400 9 9 - 0 1 1

7.5.2 Additional equilibrium data compiled for the bismuth(III) chloride nitrate system

Table 7.20: Additional, experimentally determined equilibrium data compiled for the bismuth(III) chloridenitrate system, according to the equilibria Bi3+ + mCl" + nNO3~ <=> BiClm(NO3)n

3~m-n.These data were not chosen in the present report for the evaluation of recommended stabilityvalues. Method: pot = potentiometry.

l°g Pi.m.n Reference Comments I (M) Medium Method

log pUJ: Bi3+ + Cl- + NOf <=> BiCl(NO3)+

5.00 ' r 1974FED/KAL1 T= 298 K, HC1O. pot

log PUJ: Bi3+ + 2Ct- + NOf <=> BiCl2(NO3)°

6.90 ' H974FED/KAL1 T= 298 K, HC1O, pot

log P,,3j: Bi3+ + 3Cl~ + NOf <=> BiCl3(NO3)-

7.80 ' [ 1974FED/KAL1 T= 298 K, 1=0 pot

l°g Pi.i.2: Bi3+ + Cl- + 2N0f <=> BiCl(NO3)2°

5.20 ' ri974FED/KAL1 T= 298 K, 1=0 HC1O4 pot

log Pi.2,2-- Bi3+ + 2CI- + 2NOf <=> BiCl2(NO3)2-

5.60 ' f!974FED/KALl T= 298 K, 1=0 HC1O,. pot

log Pi.J.3-Bi3+ + Cl- + 3N0f <=> BiCl(NO3)f

2.83 [1974FED/KAL1 T= 298 K, 1=3 HC1O. pot

log Pu,3:Bi3+ + 2CI- + 3N0f <=> BiCl2(NO3)32-

4.38 F1974FED/KAL1 T= 298 K, 1=3 HC1O4 pot

log Pi.iX- Bi3+ + 3Cl~ + 2N0f « • BiCl3(NO3)22-

6.10 [1974FED/KAL] T= 298 K, 1=3 HCIO. potNot indicated by [1974FED/KAL] how the data were extrapolated to 1=0.

263

JNC TN8400 9 9 - 0 1 1

7.9 Bismuth phosphate system

From the thermodynamic data given in the compilations of [1971NAU/RYZ] and[1984VTE/TAR] the formation of quite insoluble bismuth phosphate BiPO^s) can be expected.However, direct measurements of the solubility or details about the formation of this solidspecies have not been found in the literature.

BiPC>4(s) is a potentially a solubility determining solid under repository conditions. However,no solubility product can be proposed in this report as experimental data are missing.

Table 7.21: Thermodynamic data for the precipitation of BiPO4(s) recommended previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison.

log K$o Reference Comments

log K*so: Bi3+ + PO/- <=> BiPO4(s)

222222

.4 '

.4 '

.9

[1971NAU/RYZ][1977TAR/VIE][1984VIE/TAR1

T=T=T=

298.15298.15298.15

K,K,K,

I=n/aI=n/aI=n/a

calculated with a AfG° of -95.55 kj/mol for Bi2+ (Section 7.11)

264

JNC TN8400 99 - Oil

7.10 Bismuth sulfate system

7.10.1 Bismuth sulfate complexes

[1971FED/KAL2] determined equilibrium constants for bismuth sulfate complexes in 3 Msulfate/perchlorate solutions at different temperatures (Table 7.22). Unfortunately, as no othervalues are available, these values can not be extrapolated to 1=0.

7.10.2 Bi2(SO4)3(s)

The formation of Bi2(SO4)3(s) is reported in the literature (see values compiled in Table 7.23).Direct measurements of the solubility or details about the formation of this solid species havenot been found in the literature. At higher temperature (800 - 1000 K) mixed precipitates withoxygen (Bi2O2SO4, a-Bi2O(SO4)2 and P-Bi2O(SO4)2 can also be formed [1984JON].

7.10.3 Equilibrium data compiled for the bismuth sulfate system

Table 7.22: Experimentally determined equilibrium data compiled for the bismuth sulfate system. These datawere not chosen in the present report for the evaluation of recommended stability values.Reasons for not selecting these references see text. Method: pot = potentiometry

log Pi.r Reference Comments I (M) Medium Method

log pu: Bi3+ + SO/~ <=> Bi(SO4)+

1.98 [1971FED/KAL2] T= 298.15 K, 1=3.0 HC1O4, LiC104 pot

log j3;i2: Bi3+ + 2SO42~ <=> Bi(SO4)2~

3.41 [1971FED/KAL2] T= 298.15 K, 1=3.0 HC1O4, LiClO4 pot

log pu: Bi3+ + 3SO42' <=> Bi(SO4)3

3-

4.08 [1971FBD/KAL2] T= 298.15 K, 1=3.0 HC1O4, LiClO4 pot

log P1A: Bi3+ + 4SO42- <=> Bi(SO4)4

5~

4.34 ri971FED/KAL21 T= 298.15 K, 1=3.0 HC1O4, LiClQ, pot

log p L 5 : Bi3+ + 5SO42~ <=> Bi(SO4)5

7~

A_J6 fl971FED/KAL21 T= 298.15 K, 1=3.0 HC1O4, LiClO4 pot

265

JNC TN8400 9 9 - 0 1 1

Table 7.23: Thermodynamic data for the bismuth sulfate system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison.

l°g Pi,m Reference Comments I(M)

log pu: Bi3+ + SO42~ <=> Bi(SO4)

+

1.98 [1976SMI/MAR1 T= 298.15 K, 1=3.0

log pl2: Bi3+ + 2SO42~ <=> Bi(SO4)2-

3.41 [1976SMI/MAR1 T= 298.15 K, 1=3.0

log Pu: Bi3+ + 3SO42~ o Bi(SO4)3

3~

4.08 [1976SMI/MAR1 T= 298.15 K, 1=3.0

log pli4: Bi3+ + 4SO42~ <=> Bi(SO4)4

5~

4.34 ri976SMI/MAR] T= 298.15 K, 1=3.0

log Pu: Bi3+ + 5SO42~ <=> Bi(SO4)5

7~

_4.6 [1976SMI/MAR1 T= 298.15 K, 1=3.0

log K*so: 2Bi3+ + 3SO42~ t=> Bi2(SO4)3(s)

29.2 ' [1977BAR/KNA], T= 298.15 K, I=n/a29.2 1 [1979KUB/ALC1, T= 298.15 K, I=n/a1 calculated with a AfG° of-95.55 kJ/mol for Bi2+ (Section 7.11)

266

JNC TN8400 9 9 - 0 1 1

7.11 Bi3+/Bi(cr)

The ArG° values for the redox equilibria between Bi3+ ion and metallic rhombohedral Bi(cr)reported in different compilations differ by 10 kJ/mol. The use of different ArG° values for Bi3+

can result in a difference of 1.6 log units in the calculated log (3° values (compare the valuescompiled in Table 7.26). Potentiometric measurements of [1969VAS/GLA] with a platinumelectrode in perchlorate media are compiled in Table 7.24.

Table 7.24: Experimentally determined equilibrium data compiled for the equilibrium Bi3+ +3e- <=> Bi(cr). These data were chosen for the evaluation of recommended valuesin the present report. Additional information for the different references seeSection 7.12: 'Comments on selected references'. Method: pot = potentiometry

logK Reference Comments Medium Method

log K: Bi3+ + 3e- <=> Bi(cr)

14.8614.7314.7914.7914.6214.6114.6514.7314.6214.6114.72

[1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA][1969VAS/GLA]

T=298.15, 1=1.2T=298.15, 1=2T=298.15, 1=2.3T=298.15, 1=2.3T=298.15, 1=3T=298.15, 1=3.1T=298.15, 1=3.3T=298.15, 1=3.6T=298.15, 1=4T=298.15, 1=4.3T=298.15, 1=4.6

1.22

2.32.63

3.13.33.64

4.34.6

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

HC1O4

potpotpotpotpotpotpotpotpotpotpot

Extrapolation of the potentiometric measurements of [1969VAS/GLA] (compiled in Table 7.24)to 1=0 with the SIT result in (see Figure 7.24):

Bi3 ++ 3e- Bi(cr) logK°= 16.74E° = 0.330 V

corresponding to a AfG° of 95.55 kJ/mol for Bi3+. The data of [1984PIN/GAL] in 1 M HN0 3

and of [1945LIN] in 0.6 M NaCI medium which are compiled in Table 7.25 agree well.

267

JNC TN8400 9 9 - 0 1 1

Bi(cr)

n, molal

Figure 7.24: Plot of log K + 9D vs. Im for the reaction Bi3+ + 3e- <=> Bi(cr) at 25 °C. Thestraight line shows the result of the linear regression: Ae = - 0.05; log K° =16.74. Calculated from data compiled in Table 7.24.

Table 7.25: Additional, experimentally determined equilibrium data compiled for the equilibrium Bi3+ + 3e"<=> Bi(cr). These data were not chosen in the present report for the evaluation of recommendedstability values. Method: pot = potentiometry

logK Reference Comments Medium Method

log K: Bis

16.09 l

15.62 2

14.19 3

14.71 4

+ + 3e- <=> Bi(cr)

[1969VAS/GLA][1975HEI/SCH][1984PIN/GAL][1945LIN1

T=298.15,1=0T=298.15,1=1.5T=298.15,1=1T=298.15,1=0.6

01.51

0.6

HC1O4

H2SO4

HNO3

NaCl, tartaric acid

potpotpotpot

1 extrapolated to 1=0 by [1969VAS/GLA] from the measurements given in Table 7.242 Bi makes strong complexes with SOt

2', exact I not clear3 complexes with NC>3~ are probable4 complexes with Cl are probable

Table 7.26: Thermodynamic data for the Bi3+/Bi(cr) redox equilibrium taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison.

logK Reference Comments

log K: Bi3+ + 3e- <=> Bi(cr)

14.5014.5116.0816.09

[1968 ROB AVAL][1982 WAG/EVA][1985BAB/MAT],[1985LOV/MEK]

T=298.15,1=1T=298.15,1=0T=298.15,1=0T=298.15,I=0

000

268

JNC TN8400 9 9 - 0 1 1


[1918NOY/CHO]: The potentiometric data of [1918NOY/CHO] (compiled in Table 7.27)for the reaction Bi(cr) + Cl~ + H2O o BiOCl(s) + 2H+ + 3e~ can be extrapolated to I= 0 as shown in Figure 7.25, resulting in a log K*°s of -8.32. Correction for thereaction Bi(cr) <=> Bi3+ + 3e- with a log K° of 16.74 (Section 7.11) gives:

H2O+Cl- BiOCl(s)+2H+ log K*°so = 8.42

Further data concerning the solubility of BiOCl(s) are discussed in Section 7.3.2:BiOCl(s) and Bi(OH)2Cl(s) and shown there in Table 7.8 and Figure 7.13.

Table 7.27: Experimental data for the precipitation of BiOCl(s) determined by[1918NOY/CHO]. These data were chosen for the evaluation of recommendedvalues in the present report. Method: pot = potentiometry

log Kso Reference Comments I (M) Medium Method

log K*so: Bi(cr) + Or + H2O <=> BiOCl(s) + 2H+ + 3er

-8.31 1 [1918NOY/CHOJ T= 298.15 K, 1=0.001-8.31 i [1918NOY/CHO] T= 298.15 K, 1=0.002-8.23 1 [1918NOY/CHO] T= 298.15 K, 1=0.003-8.23 i [1918NOY/CHO] T= 298.15 K, 1=0.01-8.24 i [1918NOY/CHO] T= 298.15 K, 1=0.02-8.28 i [1918NOY/CHO] T= 298.15 K, 1=0.03-8.29 ] [1918NOY/CHO] T= 298.15 K, 1=0.1-8.30 i [1918NOY/CHO] T= 298.15 K, 1=0.23-8.30 ' [1918NOY/CHO] T= 298.15 K, 1=0.5

0.0009840.001

0.003180.010.02

0.0460.10.230.5

HC1HC1HC1HC1HC1HC1HC1HC1HC1

potpotpotpotpotpotpotpotpot

values calculated in this report from the E° values and HC1 concentrations given by [1918NOY/CHO].

269

JNC TN8400 9 9 - 0 1 1

Bi(cr)+Cr+H2O o BiOCI(s)+3e+2H+

0.2 0.4lm> molal

0.6

Figure 7.25: Plot of log K*s - ID vs. Im for the reaction Bi(cr) + Cl~ + H2O <=> BiOCl(s) +2H+ + 3e~ at 25 °C. The straight line shows the result of the linear regression: Ae= -0.38; log K*°s = -8.32. Calculated with the data given in Table 7.27. Thisvalue can be corrected for the reaction Bi(cr) <=> Bi3+ + 3e- (log K° = 16.74; seeSection 7.11) to a log K*So° of 8.42 for the reaction Bi3+ + H2O+ Cl~ <=>BiOCl(s) +2H+.

[1951SWI/GAR]: [1951SWI/GAR] determined the solubility of BiONO3(s) in HNO3. Asthe log K*s0 of 2.55 given by [1951SWI/GAR] for 1=0 was criticized by[1993KRA/DEC], this value was recalculated for the present report from the datagiven in [1951SWI/GAR] (Table 7.27 and Figure 7.26). The Bi3+ concentration wascorrected in this report for the formation of BiNO3

2+ (Using the log B j j determinedin Section 7.7.1). Extrapolation of the values compiled in Table 7.28 to I = 0 asshown in Figure 7.26 result in:

Bi3 ++ H2O + NO3- BiONO3(s) + 2H+ log Kso*° = 2.75

This value compares well with the value determined by [1993KRA/DEC] (seeSection 7.7.2: BiONO3(s)).

270

JNC TN8400 99-011

Table 7.28: Experimental equilibrium data compiled for the precipitation of BiONO3(s). Thesedata were chosen for the evaluation of recommended values in the present report.Method: sol = solubility.

log KSo Reference Comments I (M) Medium Method

log Kso: Bi3+ + H2O + NO3- <^BiONO3(s) + 2H+

2.18 1 [1951SWI/GAR]2.20 ] [1951SWI/GAR]2.17 1 [1951SWI/GAR]2.15 1 [1951SWI/GAR]2.13 1 [1951SWI/GAR]2.09 1 [1951SWI/GAR]2.05 1 [1951SWI/GAR]2.00 ! [1951SWI/GAR]1.94 ] [1951SWI/GAR]

T= 298.15 K, 1=0.001 0.025 HNO3

T= 298.15 K, 1=0.03 0.03 HNO3

T= 298.15 K, 1=0.035 0.035 HNO3

T= 298.15 K, 1=0.04 0.04 HNO3

T= 298.15 K, 1=0.046 0.046 HNO3

T= 298.15 K, 1=0.05 0.05 HNO3

T= 298.15 K, 1=0.06 0.06 HNO3

T= 298.15 K, 1=0.07 0.07 HNO3

T= 298.15 K, 1=0.08 0.08 HNO,

solsolsolsolsolsolsolsolsol

1 recalculated in this report including BiNO32+ complex formation

Bi3++H2O+NO3- o BiONO3(s)+2H+

00

4-5 --

4 -:

3.5 -~

g> 2 +1.5 -

1

0.5 +.

00 0.02 0.04 0.06 0.08 0.1

lm, molal

Figure 7.26: Plot of log K*So + 8D vs. Im for the reaction Bi3+ + H2O + NO3- o BiONO3(s)+ 2H+ at 25 °C. The straight line shows the result of the linear regression: Ae =0.34; log K*°s = 2.75. Calculated with the data given in Table 7.28.

[1957OLI]: [Bi] = 0.1 - 50 mM. Olin ([1957OLI], [1959OLI], [1961OLI], [1975OLI]) wasone of the first investigators who measured the hydrolysis of Bi(III) in 0.1 and 3 M

271

JNC TN8400 9 9 - 0 1 1

perchlorate medium. Based on potentiometric measurements he reported for theformation of mononuclear BiOH2+ a log p i ] value of -1.58 and for the polynuclearBi6(OH)i26+ a log p6,i2 value of 0.33 in 3 M NaClO4. His careful experimentalwork encompassed a large range of pH values and bismuth concentrations and hedeveloped a consistent system for the hydrolysis of bismuth. The values determinedgraphically by [1957OLI] were later recalculated with a computer fitting program by[1975OLI].

[1959OLI]: Bi = 0.25 - 4 mM in 0.1 M NaC104. For further comments see [1957OLI].

[1960TOB]: [1960TOB] proposed the existence of a Bi6(OH)]53+ complex (Table 7.2).

Spectrophotometric measurements made by [1972DRA/NIM1] are consistent withthe formation of polynuclear Bi9(OH)2o

7+, Bi9(OH)2i6+, and Bi9(OH)225+

complexes as proposed by [1959OLI].

From the evaluation of his data, [1960TOB] also calculated log p6,l2 value of -0.53for the formation of Bi6(OH)]2

6+. However, as [1960TOB] did not seem to take intoaccount any other complexes (e.g. BiOH2+ or Bi(OH)2

+) his data were not chosen inthis report. [1960TOB] made also an attempt to recalculate the data given in[1957OLI] and found a log (36,i2 value of +0.03 (instead of the 0.33 as given by[1957OLI]). [1960TOB] writes that the cause of the discrepancy between the valuesof the constants obtained from Olin's data is not known. However, the neglect ofmononuclear complexes by [1960TOB] may be reason for this discrepancy.

[1961OLI]: Gives the results already published in [1957OLI] and [1959OLI]. For furthercomments see [1957OLI].

[1971BID]: [197IBID] measured in 0.1 and 1 M perchlorate solution the bismuth hydrolysiswith an organic extractant (dithizone). He was able to show that no (or only little)polynuclear species were present in 0.001-0.01 mM bismuth solutions in the pHrange 9.9 - 11.3. [1971BID] used a log Kw of -13.79 in 0.1 M perchloratesolutions.

[1972DRA/NIM1]: Bi = 0.4-10 mM. Dragulescu and co-workers ([1972DRA/NIM1],[1972DRA/NIM2]) determined bismuth hydrolysis in 0.1 and 1 M perchloratemedium. Their results are comparable to the results of Olin.

[1972DRA/NIM2]: cf. [1972DRA/NIM1].

[1975ANT/NEV]: [1975ANT/NEV] measured bismuth hydrolysis in 0.1, 0.3, 0.5 and 1 MKNO3 spectrophotometrically with an organic ligand. As bismuth tends to formweak complexes with nitrate these values were not included in our calculations.[1975ANT/NEV] excluded in presence of an organic ligand (l-(2-pyridylazo)-2-naphthol) the formation of polynuclear Bi species in their 0.02 mM Bi solutions.

272

JNC T N 8 4 0 0 99 - O i l

[1975HEI/SCH]: [1975HEI/SCH] determined potentiometrically in 1.5 M H2SO4 the redoxpotential of the Bi3+/Bi(cr) and BiO+/Bi(cr) couples. [1975HEI/SCH] tried to correctfor complex formation with sulfate. However, the results are difficult to interpret andare not used for extrapolation.

[1975OLI]: Recalculation of the results already published in [1957OLI] and [1959OLI]. Forfurther comments see [1957OLI].

[1982HAT/SUG]: [1982HAT/SUG] measured in 1 M perchlorate and nitrate solutions thehydrolysis of bismuth with an organic extractant (dithizone) in ~ 0.0001 mMbismuth solutions. [1982HAT/SUG] assumed that the difference between theirresults and the result of [1972BID] may be due to the formation of polynuclearspecies.

[1982LAP/KOL]: The values reported by [1982LAP/KOL] are linearly extrapolated to 25 <Cfrom measurements at 75 - 300 °C. The presence of polynuclear species in presenceof a solid Bi2C>3 phase was not considered by the authors. The measurements are notused for extrapolation to 1=0 in this report.

[1987MIL/ROE]: [1987MHVROE] determined the hydrolysis of Bi(III) in 0.25 perchloratesolutions and at bismuth concentration of 10~12 M.

[1987SUG/ISH]: [1987SUG/ISH] measured in 1 M chloride solutions the hydrolysis ofbismuth with an organic extractant (dithizone) in = 10~n M bismuth solutions. Dueto the formation of chloride complexes the determined hydrolysis constants weresmaller than the ones determined earlier by the same group ([1982HAT/SUG])indicating a quite strong complex formation of bismuth with chloride. Themeasurements are not used in this report for extrapolation with the SIT method.[1987SUG/ISH] give in 1 M chloride a log Kw of 13.62.

[1993KRA/DEC]: [1993KRA/DEC] determined in 1 M perchlorate and nitrate solutions thehydrolysis of bismuth based on solubility measurements (oversaturation). Thehydrolysis constants do not fit well with other data found in the literatureas [ 1993 KRA/DEC] did not consider the possible formation of the polynuclearBi9(OH)2o

7+, Bi9(OH)2i6+, and Bi9(OH)22

5+ complexes in their calculations.

[1993KRA/DEC] reported also a log p i j of 3.5 for the formation of BiC1042+.

However, this equilibrium constant seems to be much too large in comparison withthe constants of the bismuth complexes with nitrate and chloride as given in thisreport. [1993KRA/DEC] present in their report no direct experimental proof for thepredominance of the BiC1042+ (instead of Bi3+) at pH below 2. Additionally,[1982HAT/SUG] and [1987SUG/SHI] showed that nitrate and perchlorate had onlya small influence on bismuth hydrolysis while the influence of chloride (and thus alsothe complex formation with chloride) is more important. In any case, the possible

273

JNC TN8400 9 9 - 0 1 1

existence of BiC1042+ has no influence on the hydrolysis constants of Bi(OH)3 andBi(OH)4~ determined at much higher pH values.

[1993KRA/DEC] also determined constants for the complex formation of Bi(III)with nitrate and the precipitation of BiONO3(s). These constants do not depend onthe presence or absence of the polynuclear Bi9(OH)2o7+, Bi9(OH)2i6+, andBi9(OH)225+ complexes and are therefore included in the present report.

274

JNC TN8400 9 9 - 0 1 1

8 Niobium

A number of oxidation states ranging from -I to +V and +VII have been mentioned in thereview of [1985UDU/VEN]. Thermodynamic data, however, exist only for compounds withthe oxidation states +11, +IV, and +V. Only Nb(V) forms aqueous species. Aqueous alkali-metal niobate solutions contain the hexaniobate anion HmNb6Oi9(8~m)- at pH > 7 inconcentrated solutions [1983POP]. Niobium also forms pentahalogenides with Cl, F, Br and I[1985UDU/VEN, 1995WIB].

8.1 Hydrolysis of niobium(V)

The solubility measurements of [1963BAB/LUK], [1992YAJ^TOB] and [1994YAJ] (seeSection 8.1.1), indicate the presence of a singly, negatively charged species in the pH range 7to 10, while they seem to exclude the presence of higher charged species. As neither[1963BAB/LUK], nor [1992YAJ/TOB and 1994YAJ] determined the predominant speciesunder alkaline conditions, the simplest assumption is that Nb(OH)6~ is the predominant speciesunder these conditions. For the purpose of modeling, only Nb(OH)6~ is used in this report.

However, there is spectrophotometric evidence for the existence of hexaniobate anions:H3Nb6Oi95~, H2Nb6Oi96~, HNb6Oi97~, and NbgO^8-, in saturated niobium solutions in thepH range 9 - 1 4 [1964NEU, 1968SPI, 1994ETX/FER]. An overview of the data available inliterature concerning polynuclear Nb(V) species is given in section 8.1.2.

8.1.1 Monomeric Nb(V) species

[1963BAB/LUK] interpreted their measurements of the solubility of freshly precipitatedNb2O5(s) in terms of the monomeric species Nb(OH)4+, Nb(OH)5° and Nb(OH)6-. As thiswas the only determination of the solubility of Nb2C>5(s) until the very recent measurements of[1992YAJ/TOB] and [1994YAJ], these values were reported in several compilations.[1963BAB/LUK], however, did not indicate any detection limit and [1976BAE/MES]considered the values given by [1963BAB/LUK] as approximate estimates.

[1963BAB/LUK, 1992YAJATOB and 1994YAJ] determined the solubility of freshlyprecipitated MÔsCprecip) in alkaline solutions and fitted their data using Nb(OH)6~ asdominant species (Tables 8.1 and 8.2). Their solubility measurements indicate thepredominance of a singly charged species in the pH range 7 to 10. There is, however, aninconsistency between assuming the predominance of Nb(OH)6~ and the observations of[1960JAN/ERT, 1964NEU, 1968SPI, and 1994ETX/FER] (cf. Section 8.1.2) who observedthe presence of the hexaniobate ion in alkaline solutions. Based on the solubility experiments of[1963BAB/LUK, 1992YAJATOB and 1994YAJ] the predominance of a higher chargedpolymeric species is uncertain and for the purpose of modeling, only Nb(OH)g" is assumed inthis report.

275

J N C T N 8 4 0 0 99 - O i l

[1963BAB/LUK] estimated a log KS6 of 24.5 (Table 8.2) in 1 M KNO3. More recently,[1992YAJ/TOB] and [1994YAJ] extrapolated from measurements in 0.1 M NaCl a log K°s6 of29.2 at 1=0 (Table 8.1) for the reaction:

2Nb(OH)6- + 2H+ <=> Nb2O5(precip) + 7H2O log KS6

29.2

Table 8.1: Experimentally determined stepwise formation constants K compiled for theniobium(V) hydroxide system. These data were chosen for the evaluation ofrecommended values in the present report. Additional information for the differentreferences see Section 8.5: 'Comments on selected references'. Method: sol =solubility measurements, tit = titration (pH). Kb is the stepwise formation constantfor a reaction using OH~~ as a component.

log Kb Reference Comments I (M) Medium Method

log K*S6: 2Nb(OH)6

29.2 1 [1992YAJ/TOB],[1994YAJ]

Nb2O5(precip) + 7H2O

T= 298.15 K, 1=0.1 0 NaCl sol

log Kb6.i9.j: Nb6O19

8~+ H2O & Nb6O19H7' + OH~

-0.36 2

-1.83 3-1.7 3-1.56 3

-1.4 3-0.49 4

[1964NEU][1968SPI][1968SPI][1968SPI][1968SPI][1994ETX/FER]

T= 298.15 K, 1=3T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=3

30.5

1233

KC1KC1KC1KC1KC1KC1

tittittittittittit

log K\19i2: Nb6O]9H7~+ H2O t=> Nb6O19H2

6~ + OH~

-3.28 2

-3.03 3

-3.06 3-3.09 3

-3.143

-4.2 4

[1964NEU][1968SPI][1968SPI][1968SPI][1968SPI][1994ETX/FER]

T= 298.15 K, 1=3T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3T= 298.15 K, 1=3

30.5

123

3

KC1KC1KC1KC1KC1KC1

tittittittittit

tit

276

JNC TN8400 9 9 - 0 1 1


log Kb6j9>3:

-4.19 3

-4.29 3

-4.46 3

-4.61 3-4.77 4

Nb6O]9H26-

[1968SPI][1968SPI][1968SPI][1968SPI]

•+ H20 <^> Nb6O]9H35~ + OH-

T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=2T= 298.15 K, 1=3

[1994ETX/FER] T= 298.15 K, 1=3

0.51233

KC1KC1KC1KC1KC1

tittittittittit

1 corrected to 1=0 with Davies equation by [1992YAJ/TOB, 1994YAJ]2 Nb concentration = 50-140 mM3 Nb concentration = 2.5-20 mM4 Nb concentration = 2.5-25 mM

8.1.2 Polymeric Nb(V) species

Aqueous alkali-metal niobate solutions (Nb concentration = 1-100 mM) appear to contain thehexaniobate anion HmNb6019(8-m)- at pH > 9 [1964NEU, 1968SPI, 1994ETX/FER,1983POP]. Such solutions are prepared by fusion of Nb2Os with excess metal hydroxide orcarbonate, and subsequent dissolution of the melt in water [1960JAN/ERT2]. Also directdissolution of freshly prepared Nb2Os in water is possible [1983POP]. Ion diffusion in water[1960JAN/ERT1] and spectrophotometric measurements [1960JAN/ERT1, 1973GOI//GRA]are consistent with the hexaniobate anion being the predominant species in saturated aqueoussolutions at pH > 9. There is good evidence from potentiometric [1960JAN/ERT3, 1964NEU,1994ETX/FER, 1968SPI] and spectrophotometric studies [1973GOI//GRA, 1974GOI//SPI]that NbgOig8- is protonated below pH = 14 in solution (see Table 8.1).

An overview of the data available in literature is given in the following sections. The data usedfor the calculations of equilibrium constants for the niobium(V) hydroxide system are given inTable 8.1 and Figures 8.1, 8.2 and 8.3. Additional data are compiled in Table 8.2 and 8.3.

8.1.2.1 Profanation at pH > 8: Nb6Oj98~, Nb6O19H

7~, Nb6O]9H26-, and Nb6O19H3

5~

Protonation constants for Nb6O ] 98- have been determined by [1964NEU], [1968SPI], and

[1994ETX/FER] by potentiometric measurements in 0.1 - 3 M KC1 (see Table 8.1). [1968SPI]extrapolated his results graphically to 1=0 (Table 8.2). [1994ETX/FER] extrapolated theirresults and the results of [1964NEU, 1968SPI] to 1=0 with an extended Davies equation.However, for the determination of the first protonation constant (log Kîî : Nb6O]98~ + H+

<=» Nb6O19H7-) [1994ETX/FER] did not include the data of [1968SPI], making theirextrapolation from I = 3 to I = 0 somewhat doubtful. [1964NEU, 1968SPI, and1994ETX/FER] all determined Nb6Oi9

8~ protonation as a function of OH~ concentration andthen converted their constants to hydrolysis constants based on H+ with different log Kw

values. [1968SPI] used for all ionic strengths a log Kw of-14.0. Therefore, in this report for

277

JNC TN8400 9 9 - 0 1 1

extrapolation to 1=0, the original stepwise formation constants Kb (which refer to the reactionwith OH~) were used as given in Table 8.1.

Extrapolation of the stepwise formation constants given in Table 8.1 to I = 0 is shown inFigures 8.1 - 8 . 3 :

Nb6O198-+H+

Nb6Oi9H7- + H+Nb6Oi9H2

6- + H+

Nb6Oi9H7-Nb6Oi9H2

6-Nb6Oi9H3

5-

log K°6,i9,i = 14.29logK°6,19,2= 13.23logK°6,19l3= 11-63

These values, however, are considered as tentative because the solubility measurements indicatethe predominance of a singly charged (but not necessarily mononuclear) species in the pH range7 to 10 (cf. Section 8.1.1).

<=> Nb6O1 9H7 -+OH-

3.5 i

2 -I1.5 §&Bi^^^^^

Figure 8.1: Plot of log Kb6,i9,i + 14 D vs. Im for the reaction Nb6O19

8- + H2O <^>Nb6Oi9H7- + OH- at 25 °C. The straight line shows the result of the linearregression: Ae = -0.73; log Kb°6,i9,i = 0.29; log K°6;19ii = 14.29 for the reactionNb6Oi98- + H+ <> Nb6O19H

7- at 25 °C. Calculated from data in Table 8.1.

278

JNC TN8400 9 9 - 0 1 1

Nb6O19H7+H2O«Nb6O1

12U

6,19

,2

+

cn

c

1 . 5 •

1

0.5

0

-0.5 •

- 1 •

-1.5 -

-2

-2.5

(

;1 :: : " , ; . : < - v : \ . .

• • ; • - . : • • . • . " : ; ; . • : ' > : . : . : i . \.

•• • : : ' : ^ . .! }\y:::^-'\-:-;\:_':_

... • ' " ' : : : ' — - i ! ! ' ; ; j" . ' ! . •;::^f1

;'•••• -- : :v i V " : - ^ ^ . ! 1 ; - ! : : ^ ; " * -

} 1

' '• :;•:• = . : • " • : ; : / 1 ' v:-.:-: '•'..:. •:'„'."•w'yi-'^-'.

2lmi molal

C'"J :; :- i •••• ' . • • ] : . - l . : i . : i i :^.; i ; ;1h.::: :V.

iliiii3 4

Figure 8.2: Plot of log Kb6,i9i2 + 12D vs. Im for the reaction Nb6Oi9H7- + H2O o

Nb6Oi9H26- + OH- at 25 °C. The straight line shows the result of the linear

regression: Ae = -0.08; log Kb°6,i9,2 = -0.77; log K°6,i9i2 = 13.23 for thereaction Nb6Oi9H7- + H+ «• Nb6Oi9H2

6- at 25 °C. Calculated from datacompiled in Table 8.1.

Nb6O19H26+H2O<^Nb6O19H3

&-+OH-

Qo

+

O

-0.5

-1

-1.5

- 2 •

-2.5 -

-3

-3.5 •

- 4 •

•4.5 -

-5 -

1 2lm, molal

Figure 8.3: Plot of log Kb6,i9,3 + 10D vs. Im for the reaction Nb6Oi9H2

6- + H2O <=»Nb6Oi9H3

5- + OH" at 25 °C. The straight line shows the result of the linearregression: Ae = -0.06; log Kb°6,i9,3 = -2.37; log K°6,i9,3 = 11.63 for thereaction Nb6O19H2

6- + H+ <=> Nb6Oi9H35- at 25 °C. Calculated from data

compiled in Table 8.1.

279

JNC TN8400 9 9 - 0 1 1

8.1.2.2 Very alkaline solutions: Nb4O]2(OH)48~, Nb4O16

12~ and NbO2(OH)43~

On the basis of Raman and UV spectroscopy, [1973GOI//GRA] and [1974GOI//SPI] suggestedthat in very alkaline solutions (1-12 M KOH) NbgOi98- is degraded into tetrameric andmonomeric species: Nb4Oi2(OH)4

8-, Nb4Oi612~ and NbO2(OH)4

3- (Table 8.2). Theproportions of the latter two species depend upon niobium concentration. As I is not constant inthese experiments (1-10 M KOH), extrapolation to I = 0 is not possible. These species are notimportant under the pH conditions of natural waters.

8.1.2.3 Neutral and acidic solutions: H6Nbj20366', H5Nb12036

7~, H4Nb]20368~, and

Different authors proposed further polymerization of the N b ^ O ^ ^ 5 " species in neutral andacidic solution. [1960JAN/ERT1] found evidence from diffusion measurements in water for theformation of a polymer with roughly the threefold molecular mass of NbgO^Hs5- at pH < 8.[1968SPI] proposed the presence of dodecaniobates (HeNhnC^6", H5Nbi2O367~,H4Nb]2O368~, H3Nbi2C>369~) and determined by pH titration constants for the formation ofH5Nb ]2O36

7-, H4Nbi2O368-, H3Nbi2O369- in 1 M KC1 solutions (Table 8.2). However, it

cannot be excluded that the solutions were oversaturated with respect to Nb2Os(s). Thus, nodirect proof exists of the presence of dodecaniobates in equilibrium with

8.1.3 Additional equilibrium data compiled for the niobium(V) hydroxide system

Table 8.2: Additional, experimentally determined equilibrium data compiled for the niobium(V) hydroxidesystem. These data were not chosen in the present report for the evaluation of recommendedstability values. Method: sol = solubility measurements, sp = spectrophotometry, and tit =titration (pH).

log Kbm,n0 Reference '

log K*S6

24.5 '

log Kl4

-0.6 '

log KIi6

-7.40 '

Comments

.- 2Nb(OH)6~ + 2H+<=>Nb205(precip) + 7H2O

[1963BAB/LUK1 T= 298.15 K, 1=1

: Nb(OH)5° + H+ <=> Nb(OH)4+

[1963BAB/LUK]

.• Nb(OH)5° + H2O » Nb(OH)

[1963BAB/LUK]

+ H2O

T= 292 K, 1=1

T=292K, 1=1

I (M) Medium

1 KNO,

1 KNO,

1 KNO,

Method

sol

sol

sol

280

JNC TN8400 9 9 - 0 1 1


log Kb6J9J: Nb6O19

8-<=>Nb6O,9H7-+ OH~

-2.02 2

-2.1 3

2.11 4

log Kh6J92:

-2.99 2

-3 3

-2.14 4

log K\19J:

-4.06 2

- 4 3

-2.06 4

[1968SPI][1968SPI][1994ETX/FER1

Nb6O!9H7~<=> Nb6O19H2'

[1968SPI][1968SPI]R994ETX/FER1

Nb6O19H26-<=>Nb6O19H

[1968SPI][1968SPI]ri994ETX/FERl

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0

*-+ OH-

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0

/ - + OH-

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0

0.100

0.100

0.100

KC1KC1KC1

KC1KC1KC1

KC1KC1KC1

tittittit

tittittit

tittittit

log Kb4il2A: 2Nb6O,9

8-+ 8OH~ <=> 3Nb4O12(OH)48~

-0.61 s [ 1973GOI/GRA] T= 298.15 K, 1= 1 -10 1MKC1, 1-10 MKOH sp

log Kbh6A: Nb4O12(0H)/-+ 40H~ + 4H2O « NbO2(OH)/-

-11.96 5 [1973GOI7GRA1 T= 298.15 K, 1=1-10 1 M KC1, 1-10 M KOH sp

log Kbuzo: Nb4OJ2(OH)/-+ 40H~<=> Nb4O16'

2- + 4H2O

-3.23 5 ri973GOI/GRA1 T= 298.15 K, 1=1-10 1 MKC1, 1-10 MKOH sp

log KiJ2M: H3Nb60195-

7.0 6

195 + H+

H968SPI]

0.5H4Nb12O368~+ H2O

T= 298.15 K, 1=1 1 M KC1 tit

log K3J2i36: H4Nb120368~ <=> H3Nb12036

9-+ H+

-7.8 6 [1968SPI] T= 298.15 K, 1=1 1 M KC1, tit

log K4J2J6: H5Nb120367- « - H4Nb12O36

8- + H+

-6.34 6 [1968SPI1 T= 298.15 K, 1=1 1 1 M KC1, tit' Detection limit not determined, determined in presence of freshly precipitated Nb2O5 (s).2 1 probably not constant (which is rather difficult at 1=0.1 and pH 13).3 extrapolated graphically to 1=0 by [1968SPI]4 extrapolated I = 0 by [ 1994ETX/FER] (see also Section 8.1.1)5 I not constant6 oversaturation possible

281

JNC TN8400 9 9 - 0 1 1

Table 8.3: Thermodynamic data for the niobium(V) hydroxide system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log Kb619 m Reference Comments I (M) Medium

log Kb619J: Nb6O19

8-<=>Nb6O19H7- + OH~

-0.36-0.36

[1976BAE/MES]fl976SMI/MARl

log K\J9X. Nb6O19H7-&Nb6O,9H2

-3.28-3.28

log KlA.

-0.6-0.45-0.59-0.44-0.44

log KL6:

-7.40-7.34-7.40-7.33-7.34

[1976BAE/MES][1976SMI/MAR1

Nb(OH)5° + H+ <=> Nb(OH)4

[1976BAE/MES][1982WAG/EVA][1985CHA/DAV][1985UDU/VEN][1992PEA/BER1

T= 298.15 K, 1=3T= 298.15 K, 1=3

fi- + OH-

T= 298.15 K, 1=3T= 298.15 K, 1=3

,+ + H2O

T= 292 K, 1=1T= 298.15 K, 1=1T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K,I=n/a

Nb(OH)5° + H2O <=> Nb(OH)6~ + H+

[1976BAE/MES][1982WAG/EVA][1985CHA/DAV][1985UDU/VEN][1992PEA/BER1

T= 292 K, 1=1T= 298.15 K, 1=1T= 298.15 K,I=n/aT= 298.15 K,I=n/aT= 298.15 K, I=n/a

33

33

11

11

KC1KC1

KC1KC1

KNO3

KNO3

282

JNC TN8400 9 9 - 0 1 1

5.2 Solubility of solid niobium pentoxide

Freshly precipitated Nb2Os(precip) has an amorphous structure. Crystallization to Nb2C>5(cr)occurs upon heating around 700 °C when all water molecules are removed [1989VAN/POU].The data used for the calculations of the solubility of NbiOsCprecip) are given in Table 8.4.Additional data that were not chosen for the evaluation in this report are compiled in Tables 8.5and 8.6.

[1963BAJB/LUK] determined the solubility of freshly precipitated Nb2Os(precip) in 1 MKNO3. In the neutral pH range, [1963BAB/LUK] determined a minimum solubility of 1.4 x10~5 M Nb, corresponding to a logK*so of 4.85 for the theoretical equilibrium Nb(OH)5° <=>0.5Nb2O5(precip) + 2.5H2O (Table 8.5). However, as no detection limit is indicated in thework of [1963BAB/LUK], one may suspect that the measured concentrations below pH 7correspond to the detection hmit and that the real concentration may be smaller. Recently,[1992YAJ/TOB] and [1994YAJ] determined in 0.1 M NaCl a minimum solubility < 1 x 10~8 MNb, corresponding to a logK*so > 8 for the theoretical equilibrium Nb(OH)5° o0.5Nb2C>5(precip) + 2.5H2O in both over- and undersaturation experiments (Table 8.4). Fromthese data and the discussion concerning the hydrolysis of niobates (Section 8.1), it can beconcluded that in the neutral pH range, niobium solubility is smaller than 1 x 10"8 M.

Table 8.4: Experimentally determined equilibrium data for the dissolution ofThese data were chosen for the evaluation of recommended values in the presentreport. Method: sol = solubility measurements.


log K*S5: 2Nb(OH)5° <^Nb2O5(precip) + 5H2O

> 16.0 1 [1992YAJ/TOB], T= 298.15 K, 1=0.1 0 NaCl sol[1994YAJ]

1 corrected to 1=0 with Davies equation by [1992YAJ/TOB, 1994YAJ]

It is not possible, due to the lack of experimental data, to give an exact thermodynamic constantfor the solubility of niobium in neutral or acidic solutions. However, an upper solubility limit isknown from the experiments of [1992YAJ/TOB] and [1994YAJ] and for modeling purposes itis practical to express this upper solubility limit in the neutral and acidic pH range through anequation involving the uncharged, hypothetical Nb(OH)s° as 'dummy species'. Using the datagiven in [1992YAJ/TOB] and [1994YAJ] one calculates:

283

JNC TN8400 9 9 - 0 1 1

2Nb(OH)5° <=> Nb2O5(precip) + 5 H2O log K*S5= 16.0

From this also a constant for the equilibrium between the hypothetical species Nb(OH)s° andNb(OH)6- (cf. Section 8.1.1) can be calculated

Nb(OH)5° + H2O <=> Nb(OH)6- + H+ log K*5,6 = -6.6

It should be stressed that the 'species' Nb(OH)$0 (and possibly Nb(OH)<5~) is an hypotheticalspecies which allows one to calculate a solubility limit for niobium in the neutral and acidicrange but does not necessarily correspond to a real species.

[ 1993KUL/HAK] determined the solubility of niobium in groundwater in the pH range 6 to 13.The results of their overs aturation experiments show an increase of Nb solubility in the pHrange 6 - 10 and are consistent with the results of [1963BAB/LUK, 1992YAJ/TOB and1994YAJ]. At pH 10 - 13, [1993KUL/HAK] observed no further pH dependency of niobiumsolubility (« 10-3 to 10"2 M). Similarly, [1992YAJ/TOB] and [1994YAJ] also observed at pH> 10 a rather constant Nb solubility in the range of 10-4 to 10~3 M.

In undersaturation experiments using Nb2O5(cr?) [1993KUL/HAK] observed after 4 months aniobium solubility in the range of <1X10~9 M (ultrafiltration experiments) to 5x10~8 M(centrifugation experiments) at pH 8 and 13. [1990PIL/STO] determined in undersaturationexperiments in concrete water (pH 11.8 - 12.5) a niobium solubility of <2xlO~7 M (after 1week) to 2x10-3 M (after 18 month). The results of [1990PIL/STO] show that the dissolutionof commercially available (probably calcined) Nb2Os(cr?) is a much slower process than thedissolution of freshly prepared Nb2O5(precip). [1992YAJ/TOB] and [1994YAJ], on the otherhand, found after four weeks equilibration time no difference between the solubility ofNb2C>5(cr) in undersaturation experiments and the solubility of precipitated Nb2O5(am) inoversaturation experiments. The difference between the observations of the threeundersaturation studies is probably due to the different Nb2Os(cr) used. Unfortunately, nodetailed information can be found regarding the nature of the solids used by [1993KUL/HAK]and [1990PIL/STO].

284

JNC TN8400 9 9 - 0 1 1

8.2.1 Additional equilibrium data compile

Table 8.5: Experimental determined equilibrium data compiled for the precipitation of Nb2O5. These datawere not chosen in the present report for the evaluation of recommended stability values.Method: sol = solubility measurements.


log K*S5: 2Nb(OH)5° <=> Nb2O5(precip) + 5H2O

9.7 ' [1963BAB/LUK] T= 298.15 K, 1=1 1 KNO, sol1 Detection limit not determined, freshly precipitated Nb2O5 (s).

Table 8.6: Thermodynamic data for the precipitation of Nb2O5 taken from previous compilations.Precipitation of Nb2O5(precip) according to the hypothetical equilibrium (see text) 2Nb(OH)5°<=> Nb2O5(precip) + 5H2O. As pointed out in Section 2 of this report only experimental datawere chosen for the present evaluation. The following table serves only for comparison.Medium: Where data refer to specific electrolyte solutions, this is indicated.

log r s o

log K*S0:

9.69.6610.069.66

Reference Comments

2Nb(OH)5° <=> Nb2O5(precip) + 5H2O

[1976BAE/MES][1982WAG/EVA][1985UDU/VEN][1992PEA/BER]

T=292K, 1=1T= 298.15 K, 1=1T= 298.15 K,I=n/aT= 298.15 K,I=n/a

KM)

11

Medium

KN03

285

JNC TN8400 9 9 - 0 1 1

8.3 Solid niobium phases: redox equilibria

Niobium pentoxide, Nb2O5(cr), is a dense white powder and chemically relatively inert. Acomprehensive review of the thermochemistry and oxidation potential of niobium has beenmade by [1971HIL/WOR] and [1985UDU/VEN]. Nb(cr) is insoluble in water, but reacts inpresence of oxygen above 300 °C to the pentoxide Nb2O5(s). Also the solids NbO(s) andNbO2(s) are known. Both solids are thermodynarnically unstable in the presence of waterwhere the pentoxide Nb2Os(s) is formed under reduction of water ([1985UDU/VEN],[1995WTB]). From the thermodynamic data given in [1985UDU/VEN] and other compilationsequilibria between the different solid niobium phases can be calculated (Table 8.7).

Table 8.7: Thermodynamic data for the redox equilibrium of Nb(cr), NbO and NbO2 with Nb2O5 taken fromprevious compilations. As pointed out in Section 2 of this report only experimental data wereused for the present evaluation. The following table serves only for comparison.

log Pi,0

l°g Pl,0:

-51.12-51.18-51.02-50.65-50.82-50.82-50.82-50.83-50.82-51.01

l°i PJ.O-

-6.66-4.30-4.38-4.30-4.41-4.41-4.19-4.19-4.41-4.30

log A.o-

-25.82-23.70-23.70-26.01-26.05-23.70-26.23

Reference Comments

0.5Nb2O5 + 5e~ + 5H+ <=> Nb(cr)

[1954COU][1963WIC/BLO][1971HIL/WOR][1971NAU/RYZ][1978ROB/HEM2][1979KUB/ALC][1982PAN][1982WAG/EVA][1985CHA/DAV]ri985UDU/VENl

0.5Nb2Os + 3e~ + 3H+ <=> NbO

[1952WOR][1954COU][1963WIC/BLO][1971HIL/WOR][1978ROB/HEM2][1979KUB/ALC][1982PAN][1982WAG/EVA][1985CHA/DAV][1985UDU/VEN1

0.5Nb2O5 + e~ + H+ <^> NbO2+

[1971HILAVOR][1978ROB/HEM2][1979KUB/ALC][1982PAN][1982WAG/EVA][1985CHA/DAV][1985UDU/VEN]

+ 2.5H2O

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

+ 1.5H2O

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

0.5H2O

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15




286

JNC TN8400 9 9 - 0 1 1

5.4 Other niobium( V) complexes and compounds

A few measurements concerning the formation of Nb(V) fluoride complexes are reported in theliterature (Table 8.8). Only the presence of mononuclear Nb(V) is assumed in these papers. Nothermodynamic values for the formation of Nb(V) fluorides are recommended in this report.

The standard Gibbs energy of formation, AfG°, of the solids NbCl5(s), NbOCl3(s),NaNbO3(s), Na3Nb04(s), KNbO3(s), K3Nb04(s), and Ca(NbO3)2(s) are compiled in Table8.9. No solubility products are calculated from these values and no thermodynamic values forthe formation of these Nb(V) solids are recommended in this report.

Table 8.8: Experimentally determined equilibrium data compiled for the antimony(V) system. These datawere not chosen in the present report for the evaluation of recommended stability values.Reasons for not selecting these references are given in the text in Section 8.4: Method: pot =potentiometry.

log (3 Reference Comments I (M) Medium Method

log KUJ: NbOF2+ + F~ <=> NbOF3°

3.78 ri972LAN/OSB] T= 298 K, 1=0.5 0.50 Na, HC1O4 pot

log KUA: NbOF3° + F~ <=> NbOFf

4.30 [1972LAN/OSB] T= 298 K, 1=0.5 0.50 Na, HCI04 pot

log KUi5: NbOF4- + F~ <=> NbOF52'

2.514.51

log K1.1.6:

4.67

log Ki.o,7:

11.41

log Kli0_8:

3.08

log K109:

4.0

[1969NEU][1972LAN/OSB]

NbOF52~ + F- <=> NbOF6

3~

[1972LAN/OSB1

NbOF63~ + F~ + 2H+ <=> NbF7

2~ A

[1972LAN/OSB1

NbF72~ + F~ <=> NbF8

3-

[1972LAN/OSB]

NbF/- + F~ <=> NbF94-

[1972LAN/OSB1

T= 298 K, 1=3T= 298 K, 1=0.5

T= 298 K, 1=0.5

vH2O

T= 298 K, 1=0.5

T= 298 K, 1=0.5

T= 298 K, 1=0.5

3.00.50

0.50

0.50

0.50

0.50

KC1Na, HC1O4

Na, HCIO,,

Na, HCIO,

Na, HC1O4

Na, HCIO,

potpot

pot

pot

pot

pot

287

JNC TN8400 9 9 - 0 1 1

Table 8.9: Standard Gibbs energy of formation, A{G°, of the solids NbCl5(s), NbOCl3(s), NaNbO3(s),Na3Nb04(s), KNbO3(s), K3Nb04(s), and Ca(NbO3)2(s) taken from previous compilations.

Solid

NbCl5(c)

NbOCl3(c)

NbF5(c)

NbF6(c)

NaNbO3(c)

Na3Nb04(c)

KNbO3(c)

K3NbO4(c)

Ca(NbO,)?(c)

AfG°[kJ/mol]

-700.7-682.7-691.8-687.3-683.2-684.1-682.7

-790.9-784.8-782

-1699-1700-1700-1699-1699

-1699

-1242-1234-1233

-1811

-1261-1249-1251

-1825

-606.2

Reference

[1963WIC/BLO][1971HEL/WOR][1971NAU/RYZ][1979KUB/ALC][1982WAG/EVA][1985CHA/DAV][1985UDU/VEN]

[1971HIL/WOR][1979KUB/ALC][1982WAG/EVA]

[1971HILAVOR][1977BAR/KNA][1979KUB/ALC][1982WAG/EVA][1982WAG/EVA]

[1971NAU/RYZ]

[1971NAU/RYZ][1981LIN/BES][1982WAG/EVA]

[1981LIN/BES]

[1971NAU/RYZ][1981LIN/BES][1982WAG/EVA]

[1981LIN/BES]

[1971NAU/RYZ]

Comments

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

T= 298.15T= 298.15T= 298.15

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15

T= 298.15

T= 298.15T= 298.15T= 298.15

T= 298.15

T= 298.15T= 298.15T= 298.15

T= 298.15

T= 298.15


K, I=n/aK, I=n/aK, I=n/a

K, I=n/aK, I=n/aK, I=n/aK, I=n/aK, I=n/a

K, I=n/a


K, I=n/a


K, I=n/a

K, I=n/a

288

JNC TN8400 9 9 - 0 1 1


[1968SPI]: [1968SPI] determined the protonation constants of Nb6Oi98- by potentiometricmeasurements in 0.1 - 3 M KC1. [1968SPI] extrapolated his results graphically to1=0.

[1994ETX/FER]: [1994ETX/FER] determined the protonation constants of Nb6Oi98- by

potentiometric measurements in 3 M KC1. [1994ETX/FER] extrapolated their resultsand the results of [1964NEU and 1968SPI] to 1=0 with an extended Davies equation.However, for the determination of the first protonation constant of NbgOig8-(Nb6Oi9

8- + H+ o Nb6O19H7-) the data of [1968SPI] were not included, makingthe extrapolation made by [1994ETX/FER] from 1=3 to 1=0 somewhat doubtful.

[1992YAJ/TOB]: see [1994YAJ].

[1994YAJ]: Recently, [1992YAJ/TOB] and [1994YAJ] determined in 0.1 M NaCl aminimum Nb(V) solubility < 1 x 10~8 M or a log K*so > 8 for the theoreticalequilibrium Nb(OH)s° <=> 0.5Nb2O5(precip) + 2.5H2O. From these data it can beconcluded that in the neutral pH range, niobium solubility is smaller than 1 x 10~8 M.It is not possible, due to the lack of experimental data, to give a reliablethermodynamic constant for the solubility in neutral or acidic solutions.

289

JNC TN8400 9 9 - 0 1 1

9 Palladium1

Palladium is a typical B-metal (,,soft" character). In contrast to the actinides, for example, Pd2+

forms weak complexes with fluoride but strong complexes with chloride, bromide and iodide.Pd2+ has also a strong affinity to nitrogen donors such as ammonia, ethylenediamine, pyridineand its derivatives, etc.

9.1 The redox pair Pd2+/Pd(s)

Palladium is a noble metal, which tends to precipitate as Pd(s) from uncomplexed or weaklycomplexes Pd(II) solutions. The redox potential of the half cell

Pd2+ + 2e" = Pd(s) (1)

has been measured by Templeton, Watt and Garner [1943TEM/WAT] in 4 M HC1O4 solutionand by Izatt, Eatough and Christensen [1967IZA/EAT]. No other experimental examinations ofthe redox reaction (1) have been found. The ionic strength of the latter investigation wasprobably 0.1 M, so the pH cannot be lower than 1, and Pd2+ must be suspected to have beenpresent in partly hydrolyzed form, see comments in Section 9.7. Since the hydrolysis constantsof these authors are doubtful, their reported redox potential of £°(1, 298.15 K) = (0.915 ±0.005) V cannot be recommended. The measurement of Templeton, Watt and Garner[1943TEM/WAT] in 4 M HC1O4 solution is thus the only credible redox experiment availablefor Reaction (1). Although the larger part of their investigation was performed in hydrochloricacid medium, they did two measurement in HC1O4. The resulting mean value is

£(1, 4 MHC1O4, 298.15 K) = 0.987 V.

It is difficult to extrapolate this value to / = 0, because we have no reliably experimentalinformation at lower ionic strengths. If we wanted to make an attempt to extrapolate this value to/ = 0, we can convert it to log K for convenience, log K{\, 4 M HC1O4, 298.15 K) = 33.37,and by using the SIT equation, log K{2) + 4D- log K°(2) - Ae(2)7 for the reaction

Pd2+ + 2 e = Pd(s) (2)

with D = 0.261 (/ = 4.89 m) and As = -0.3 (from e(M2\ C1O4") « 0.3, cf. [1992GRE/FUG,Appendix B]), we obtain

log K°(2, 298.15 K) = 32.86and £°(2,298.15 K) = 0.972 V.

This chapter was partly written by Hans Wanner, HSK, Villigen, Switzerland

290

JNC TN8400 9 9 - 0 1 1

The uncertainty in the resulting potential is quite large due to the uncertain ion interaction

parameter. If we assume a maximum uncertainty of ±0.1 in Ae, cf. [1992GRE/FUG, B . I .3 ] ,

and an initial uncertainty of ±0.007 V in the reported redox potential {i.e., ±0.24 in log K), an

uncertainty of ±0.54 is calculated in log K°{2), which corresponds to ±0.016 V in E°{2).

Table 9.1: Experimental equilibrium data compiled for the redox pair Pd2+/Pd(s), according

to the equilibria Pd2+ + 2e~ <=> Pd(s). These data were chosen for the

evaluation of recommended values in the present report. Additional information

for the reference see Section 9.7. Method: pot = potentiometry.


log K: Pd2+ + 2er <=> Pd(s)

33.37 [1943TEM/WAT] T= 298.15 K, 1=4 HC1O4 pot

Table 9.2: Additional, experimentally determined equilibrium data compiled for the redox pair Pd27Pd(s),according to the equilibria Pd2+ + 2e~ <=> Pd(s). These data were not chosen in the presentreport for the evaluation of recommended stability values. Reasons for not selecting thesereferences are given in the text. Method: pot = potentiometry.


log K: Pd2+ + 2e <=> Pd(s)

30.93 n967IZA/EAT1 T= 298.15 K, 1=0. Pd(ClO4)? pot

Table 9.3: Thermodynamic data for the redox pair Pd27Pd(s) taken from previous compilations. As pointedout in Section 2 of this report only experimental data were used for the present evaluation. Thefollowing table serves only for comparison.

logK

log K: Pd2* -

33.3531.0730.9330.9230.9330.9230.92

Reference

h 2e <=> Pd(s)

[1952LAT][1968GOL/HEP][1980BE1OTEA][1982WAG/EVA][1985BAB/MAT][1985COL]fl988PHI/HALl

Comments

T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T= 298.15T=298.15

K, 1=0K, 1=0K, 1=0K, 1=0K, 1=0K, 1=0K, 1=0

291

JNC TN8400 9 9 - 0 1 1

9.2 Hydrolysis ofpalladium(II)

9.2.1 Hydrolysis ofpalladium(II)

The Pd2+ aqua ion is a fairly strong acid: Non-complexing solutions of Pd2+ are stable at pH = 0but start hydrolyzing at pH = 0.7 [1984WAN]. However, the reproducibility of the visibleabsorption spectra around pH = 0.7 and higher is poor, indicating that the reaction may not be asimple, mononuclear hydrolysis reaction leading to PdOH+, Pd(OH)2°, etc. In contrast, thehydrolysis of Pd2+ most probably involves polynuclearization and the subsequent formation ofcolloidal species. Papers reporting hydrolysis constants of Pd2+ are thus to be regarded withgreat care (see discussions of the papers mentioned hereafter in the Section 9.7 ,,Comments onselected references"). The published hydrolysis data of Izatt, Eatough and Christensen[1967IZA/EAT] and Nabivanets and Kalabina [1970NAB/KAL] are unreliable. Wood[1991 WOO] observed an increase of the solubility of Pd metal at pH values above 11, but thesolution composition varied strongly. On this basis it is not possible to select any hydrolysisconstants for palladium(II). The suggestion of Byrne and Kump [1993BYR/KUM] to relate thestepwise formation constants of the Pd(II) hydroxide complexes to those of the Pd(II) chloridecomplexes, i.e., Kj+l/K- is similar for OFT and Cl" as a ligand, is unlikely to be correct due tothe strong tendency of the Pd(II) hydroxide complexes toward polymerization, contrary to thePd(II) chloride complexes. Predicted hydrolysis constants are unlikely to be any better, exceptperhaps at extremely low Pd(II) concentration where polynuclearization can be expected to benegligible. However, it is at present unknown at what Pd(II) concentration this will be the case.

Several hydrolysis studies of Pd(II) have been carried out in seawater or otherwise usingchloride as a background electrolyte [1967KAZ/PTI, 1984MIL/BUG, 1989KUM/BYR,1991TAI/JAN]. Of course, due to the high stability of the chloro complexes of Pd(II),hydrolysis of Pd2+ (more correctly: PdCl4

2") starts at considerably higher pH values under theseconditions than in the absence of any complexants. In seawater (0.558 M Cl") hydrolysis ofPd(II) starts at pH values between 7 and 8 [1989KUM/BYR, Figure 2]. From the spectroscopicwork of Kump and Byrne [1989KUM/BYR] we can derive a constant for the formation ofPdCl3OH2~, and a limiting constant for the formation of PdCl2(OH)2

2", both from PdCl42~, see

comments on Ref. [1989KUM/BYR] below. These constants are valid at / = 0.7 M, but thecorrection factor to infinite dilution (/ = 0) should be small because of the well-balanced(isocoulombic) charge pattern of the reaction.

= PdCl3OH2-+ Cl" \ogK°= 4.8

= PdCl2(OH)22-+ 2 Cr \ogK°< 9.3

Using the hydrolysis notation, we obtain (with log A w° = -14.0):

PdCl42" + H2O(1) = PdCl3OH2- + H+ + Cl" log *K° = -9.2

PdCl42" + 2 H2O(1) = PdCl2(OH)2

2-+ 2 H++ 2 Cr log *K° < -18.7

292

JNC T N 8 4 0 0 99 - O i l

The hydrolysis data base of Pd(II) selected in this review is quite incomplete. It should be notedthat these data allow to predict Pd(II) hydrolysis with confidence only at chloride concentrationsof 0.5 M and higher, and at pH values lower than about 9 to 9.5. At lower chlorideconcentrations other, less chlorinated hydrolysis species might become dominant.

9.2.2 Additional equilibrium data compiled for the hydrolysis ofpalladium(II)

Table 9.4: Experimentally determined equilibrium data compiled for the palladium(II) hydroxide system,according to the equilibria Pd2+ + mH20 <=> Pd(OH)m

2~m + mH+. These data were not chosen inthe present report for the evaluation of recommended stability values. Reasons for not selectingthese references are given in the text and in Section 9.7. Method: sol = solubility measurements,sp = spectrophotometry, pot = potentiometry, and tit = titration (pH)

log P,.m

log fii.f

-1.38 '-1.00 2

-1.60 2

-2.28 3

-9.23 4

-9.3 4

-9.35 4

-9.39 4

-9.45 4

-9.61 4

log j3!2:

-2.36 '-2.20 2

-1.50 2

-4.42 3

-7.14 5

-16.57 3

-19.14 5

Reference Comments

P ^ + + H2O <=> PdOH++ H+

[1967IZA/EAT][1967IZA/EAT][1967IZA/EAT][1970NAB/KAL][1984MII7BUG][1984MEL/BUG][1984MIL/BUG][1984MIL/BUG][1984MIL/BUG][1984MIL/BUG1

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 290 K, 1=0.1T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=1.5T= 298.15 K, 1=2T= 298.15 K, 1=2.5T= 298.15 K, 1=3

P&+ + 2H2O <=> Pd(OH)2°+ 2H+

[1967IZA/EAT][1967IZA/EAT][1967IZA/EAT][1970NAB/KAL][1991 WOO]

T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 290 K, 1=0.1T= 298.15 K, 1=0-0.35

PcP+ + 3H2O <=> Pd(OH)f+ 3H+

[1970NAB/KAL][1991WOO1

T= 290 K, 1=0.1T= 298.15 K, 1=0-0.35

KM)

0.100

0.10.5

11.52

2.53

0.100

0.1

0.1

Medium

Pd(C104)2

Pd(ClO4)2

Pd(ClO4)2

HC1O,NaClNaClNaClNaClNaClNaCl

Pd(ClO4)2

Pd(ClO4)2

Pd(C104)2

HC1O4

n/a

HC1O4

n/a

Method

tittitspsolpotpotpotpotpotpot

tittitspsolpot

solpot

log Pi.4- pd2+ + 4H2® <=> Pd(OH)42~+ 4H+

-29.57 3 n970NAB/KALl T= 290 K, 1=0.1 0.1 HC1O4 sol

293

JNC TN8400 9 9 - 0 1 1


log p4t4:

-28.81 4

-29.1 4

-29.63 4

-29.88 4

-30.35 4

-30.51 4

4PJ 2 + + 4H2O <=> Pd4(OH)4"-+ 4H+

[1984MIL/BUG][1984MIL/BUG][1984MIL/BUG][1984MIL/BUG][1984MIL/BUG][1984MIL/BUG1

T= 298.15 K, 1=0.5T= 298.15 K, 1=1T= 298.15 K, 1=1.5T= 298.15 K, 1=2T= 298.15 K, 1=2.5T= 298.15 K, 1=3

0.51

1.52

2.53

NaClNaClNaClNaClNaClNaCl

potpotpotpotpotpot

1 log Kw used -13.782 extrapolated to I = 0 by [1969IZA/EAT] with extended Debye-Huckel equation3 log Kw used by [1970NAB/KAL]: -14.04 Pd concentration = 2.5 - 40 raM. Formation of chloro complexes.5 Pd concentration = 0.0001-0.1 raM.I not constant. Calculated in this report with a log K Pd(cr)/Pd2

Section 9.1).of 32.86 (cf

Table 9.5: Thermodynamic data for the palladium(II) hydroxide system taken from previous compilations.As pointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log P,,m

log p,j: Pd2

-1.30-2.00 '-1.38 2

-1.00-2.10-1.3-1.2-1.0 3

-2.1-2.3-2.3-5.7 4

-1.87 5

log pu: Pd2

-1.90-4.00 '-2.36 2

-2.20-1.9-3.42-2.2 3

-4.6-4.4-4.4-12.2 4

-3.80 5

Reference Comments

+ + H2O <=> PdOH++ H*

[1968GOL/HEP][1976BAE/MES][1976SMI/MAR][1976SMI/MAR][1981BAE/MES][1985BAB/MAT][1987BRO/WAN][1988MOU/WOO][1988PHI/HAL][1989WOO/MOU][1992WOO/MOU][1993BYR/KUM]|"1992PEA/BER1

T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.5-1T= 298.15 K, 1=0

+ + 2H2O <=> Pd(0H)2°+ 2H+

[1968GOL/HEP][1976BAE/MES][1976SMI/MAR][1976SM1/MAR][1985BAB/MAT][1987BROAVAN][1988MOUAVOO][1988PHI/HAL][1989WOO/MOU][1992WOO/MOU][1993BYR/KUM][1992PEA/BER1

T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=0.1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.5-1T= 298.15 K, 1=0

I (M) Medium

01 HC1O4

0.100000000

0.7 seawater0

01 HC1O4

0.10000000

0.7 seawater0

294

JNC TN8400 9 9 - 0 1 1


log Pu: Pd2+ + 3H2O <=> Pd(OH)f+ 3H+

- 1 2 '-3.6 3

-16.6-16.6-19.9 4

-15.94

log P,,

-13 '-5.2 3

-29.6-29.6-28.5 4

-29.36

[1976BAE/MES][1988MOUAVOO][1989WOO/MOU][1992WOO/MOU][1993BYR/KUM]

5 [1992PEA/BER1

T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.5-1T= 298.15 K, 1=0

,: Pd2+ + 4H2O <=> Pd(OH)/-+ 4H+

[1976BAE/MES][1988MOU/WOO][1989WOO/MOU][1992WOO/MOU][1993BYR/KUM]

5 fl992PEA/BER]

T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.5-1T=298.15K, 1=0

1000

0.70

1000

0.70

HC1O,

seawater

HC1O4

seawater

1 estimated by [1976BAE/MES] based on [1970NAB/KAL]2 log Kw used -13.783 P,",and P12 values selected from [1967IZA/EAT], p13and p u estimated by [1988MOUAVOO]4 [1993BYR/KUM] recalculated log B values based on [1970NAB/KAL] and [1991 WOO]; (log Kw used -13.76)5 values from [1970NAB/KAL] corrected to 1=0 with Davies equation by [1992PEA/BER]

295

JNC TN8400 9 9 - 0 1 1

9.3 Solid palladium(H)-oxide/hydroxide

In presence of oxygen Pd(cr) is oxidized to PdO(cr) at 600 CC. In solutions which containPd(II) an amorphous, yellow-brown palladium-oxide-hydrate precipitates, PdO-H2O(precip)(or Pd(OH)2(precip)) [1955GLE/PEU, 1995WIB]. In absence of water, Pd(OH)2(precip)dehydrates to PdO(cr) above 90 °C, while in presence of water above 100 °C, a PdO(s) solidwith disturbed lattice is produced, indicating the incorporation of H2O in the lattice[1955GLE/PEU].

9.3.1 Pd(OH)2(precip)

Only two experimental determinations of the solubility of Pd(OH)2(precip) in aqueous solutionshave been found. [1970NAB/KAL, 1971NAB/KAL] measured a constant Pd(II) concentrationof 4xlO"6 M between pH 3 and 11 in 0.1 M perchlorate media, corresponding to a log K*s2 of5.4 for the reaction Pd(OH)2° <=> Pd(OH)2(precip). They did not indicate any detection limitand we suspect that the measured minimum Pd(II) concentration reflects the detection limit ofthe analytical method used.

[1991WOO] determined a constant Pd(II) concentration of approx. 9xl0~8 M between pH 8 and11 in diluted solutions. Unfortunately, [1991 WOO] was not able to show whether Pd(s), whichhe used as a starting material, or Pd(OH)2(precip), which he expected based on the Ehmeasurements, was the solubility limiting phase present. However, it is interesting to note thatabove a pH of 12 the curves determined by both [1970NAB/KAL, 1971NAB/KAL] and[1991 WOO] agree well.

No solubility product for Pd(OH)2(precip) is proposed in this report.

9.3.2 PdO(cr)

The standard Gibbs molar energy of formation of PdO(s) has been determined at 700 - 1000 Kby [1982LEV/NAR] and [1983MAL/SRE]. Both authors obtained after extrapolating theirmeasurements to 298 K a AfG° of approx. -83 kJ/mol from which a log K°*so of ~ 6 can becalculated for the reaction Pd2+ + H2O & PdO(cr) + 2H+ (Table 9.6). [1983MAL/SRE] definedtheir product with X-ray analysis as PdO(cr). This agrees well with the observations of[1955GLE/PEU] who observed that, in absence of water, PdO(cr) crystallizes above 90 °C.The precipiation product from aqueous solutions, however, is PdO-H2O(precip) orPd(OH)2(precip) and not PdO(cr) [1955GLE/PEU].

296

JNC TN8400 9 9 - Oil

Table 9.6: Experimentally determined equilibrium data compiled for the formation of the solid Pd(II)-oxide/hydroxide. These data were not chosen in the present report for the evaluation ofrecommended stability values. Reasons for not selecting these references are given in the text.Method: emf = emf measurements, sol = solubility.

logK

log K*S2:

5.40 2

7.00 3

log K'SQ:

5.98 4

5.86 4

Reference Comments

Pd(OH)2° <=>Pd(OH)2(precip)

[ 1970NAB/KAL] T= 290 K,[1991WOO1 T= 298.15

Pd2+ + H2O <=>PdO(crt

[1982LEV/NAR][1983MAL/SRE]

) + 2H+

T= 298.15T= 298.15

1=0.1K, 1=0-0.002

K, I=n/aK, I=n/a

KM)

0.10

Medium

HCIO4

NaOH

Method

solsol

emfemf

1 Extrapolated to 1=0 with Debye-Hiickel2 P d = l . l m M3 Pd = 0.0001-0.1 mM; solubility limiting phase may Pd(s) and not Pd(OH)2(precip)4 extrapolated by the respective authors from measurements at 700-1000 K; Calculated in this report with a log

K of 32.86 for the reaction Pd2+ + 2er = Pd(s) (Section 9.1)

Table 9.7: Thermodynamic data for the for the formation of the solid Pd(II)-oxide/hydroxide taken fromprevious compilations. As pointed out in Section 2 of this report only experimental data wereused for the present evaluation. The following table serves only for comparison. Medium: Wheredata refer to specific electrolyte solutions, this is indicated.

log K Reference Comments I (M) Medium

log K*so: Pd2+ + 2H2O <=> Pd(OH)2(s) + 2H+

2.98 [1952LAT] T=298.15 K, I=n/a2.65 [1967IZA/EAT] T= 298.15 K, 1=0 0 Pd(C104)2

0.80 [1968GOL/HEP] T= 298.15 K, 1=0 00.60 [1976BAE/MES] T= 298.15 K, 1=0.1 0.1 HC1O4

0.50 [1976SMI/MAR] T= 298.15 K, 1=0 00.80 [1981BAE/MES] T= 298.15 K, 1=0 00.79 [1988PHI/HAL] T= 298.15 K, 1=0 01.6 [1992PEA/BER] T= 298.15 K, I=n/a

log K*so: Pd2+ + H2O <=>PdO(s) + 2H+

2.34 [1952LAT] T=298.15 K, I=n/a3.63 ' [1954COU] T= 298.15 K, I=n/a14.14 ' [1967WAR] T= 298.15 K, I=n/a4.41 [1968GOL/HEP] T= 298.15 K, 1=06.68 ' [1971NAU/RYZ] T= 298.15 K, I=n/a2.82 ' [1973BAR/KNA] T= 298.15 K, I=n/a4.40 [1976BAE/MES] T= 298.15 K, 1=05.74 ' [1977BAR/KNA] T= 298.15 K, I=n/a5.74 ' [1979KUB/ALC] T= 298.15 K, I=n/a4.70 2 [1987BRO/WAN] T= 298.15 K, 1=0 04.24 [1988PHI/HAL] T= 298.15 K, 1=04.40 [1992PEA/BER1 T= 298.15 K, I=n/a1 calculated in this report with a log K of 32.86 for the reaction Pd2+ + 2e" = Pd(s) (Section 9.1)2 reported in [1987BROAVAN] for 'Pd(OH)2(cr)'

297

JNC TN8400 9 9 - 0 1 1

9.4 Chloride complexes of palladium(H)

9.4.1 Chloride complexes

Among the other ligands considered in the PNC-TDB project, reliable formation constants areonly available for the chloride complexes of palladium(H): PdCl+, PdCl2(aq), PdCl3" andPdCl4

2~. We concur with the review of Wood, Mountain and Pan [1992WOO/MOU] that themost reliable study of Pd(II) chloride complexes is the spectrophotometric investigation byElding [1972ELD] a t /= 1 M (HC1OJ. The results of other experimental studies do not differgreatly from Elding's values. The extrapolation from / = 1 to / = 0 is not straightforward. Weuse the specific ion interaction equation (cf. NEA-TDB [1995SIL/BID]) and estimate Ae asfollows: We assume e(Pd2+, C1O4") - e(Co2+, C1O4") = 0.34 and e(PdCl+, C1O4") - e(CdCl\C1O4") = 0.25. For the anionic complexes, we only have an indication of the interactioncoefficients with Na+, not with H+, cf. Table B.4 [1995SIL/BID]. We use average values of therespective charge patterns, e(PdCl3~, Na+) » e(M", Na+) = 0.00 and e(PdCl4

2", Na+) « e(M2",Na+) = -0.10. In order to compensate the use of the Na+ interaction coefficient for the anioniccomplexes, we use e(Cl~, Na+) instead of £(C1~, H+) for the last two reactions. In this way weobtain the following constants at / = 0 (with errors between ±0.1 and ±0.2):

Pd2+-Pd2+-Pd2+-Pd2+-

i-crh2cr1-3 crH4cr

<=><=><^<=>

PdcrPdCl2°PdCl3-PdCl2"

loglogloglog

p,° =p2° =(33° =

P4° =

5.1,8.3,

10.9,11.7,

Ae = -0.21Ae = -0.58Ae = -0.43Ae = -0.56

The bromide complexes measured by Elding [1972ELD] are equally reliable, but bromide is nota priority ligand in the present project. The formation constants of Pd(II) with many otherligands, such as carboxylic acids and numerous nitrogen donors, have been determined, butcomplexes with carbonate, phosphate or sulfate are unknown to our knowledge. It should bementioned that carbonate complexes of Pd(II) may simply not form at all because of theenormous competition by hydroxide.

298

J N C T N 8 4 0 0 99 - O i l

Table 9.8: Experimental equilibrium data compiled for the palladium(II) chloride system,according to the equilibria Pd2+ + m Cl~ <=> PdClm

2"m. These data were chosenfor the evaluation of recommended values in the present report. Additionalinformation for the reference see text. Method: sp = spectrophotometry.

log [3 l.mReference Comments I (M) Medium Method

log Pu: Pd2+ + Cl- <=> PdCl+

4.47 [1972ELD] T= 298.15 K, 1=1 Na,HC104

log p1>2: Pd2+ + 2Ct «• PdCl2°

1.16 [1972ELD] T= 298.15 K, 1=1 Na,HClOd

log pli3: Pd2+ + 3CI- & PdCl3-

10.17 [1972ELD] T= 298.15 K, 1=1 Na,HC104 sp

log PiA: Pd2+ + 4Ct <=> PdCl42-

11.54 [1972ELD] T= 298.15 K, 1=1 Na,HClO4

9.4.2 Additional equilibrium data compiled for palladium chloride complexes

Table 9.9: Additional, experimentally determined equilibrium data compiled for the palladium(II) chloridesystem, according to the equilibria Pd2+ + m Cl" <=> PdClm

2~m. These data were not chosen in thepresent report for the evaluation of recommended stability values. Reasons for not selecting thesereferences are given in the text. Method: sp = spectrophotometry, pot = potentiometry.

log !*,.„ Reference Comments I (M) Medium Method

log pu.- Pd2* + cr <=> Pdcr

4.70 ' [1969GEL/KIS]4.47 2 R991TAI/JAN1

T= 298 K, I=n/aT= 292 K, 1=0.01-0.5

n/aNaCl

potsp

log Pi.2: Pd2+ + 2CI- <=> PdCl2°

7.70 '8.01 2

[1969GEL/KIS][1991TAyjAN]

T= 298 K, I=n/aT= 292 K, 1=0.01-0.5

n/aNaCl

potsp

299

JNC TN8400 9 9 - 0 1 1


log Pu: Pd2+ + 3Cl~ <=> PdCl

10.30 •10.61 2

[1969GEL/KIS]T1991TAI/JAN1

T= 298 K, I=n/aT= 292 K, 1=0.01-0.5

n/aNaCl

pot

log PL4: Pd2+ + 4CI- <=>

11.92 '11.84 2

[1969GEL/KIS][1991TAI/JAN]

T= 298 K, I=n/aT= 292 K, 1=0.01-0.5

n/aNaCl

potsp

log K1A: PdClf + Cl- <=> PdCl42-

1.4331.5931.77 3

2.01 3

1.27 4

[1968LEV][1968LEV][1968LEV][1968LEV][1973GUL/SCH]

T=298K, 1=1T= 298 K, 1=2T= 298 K, 1=3T= 298 K, 1=4T= 298 K, 1=0.04-0.3

12340

LiC104

LiC104

LiC104

LiC104

NaCl

spspspspsp

' I not given, probably not constant2 I not constant3 Pd = 0.1 mM; only log K4 measured at different I;4 Pd = 1.2 mM I not constant, only log K4 measured at different I

Table 9.10: Thermodynamic data for the palladium(II) chloride system taken from previous compilations. Aspointed out in Section 2 of this report only experimental data were used for the presentevaluation. The following table serves only for comparison. Medium: Where data refer to specificelectrolyte solutions, this is indicated.

log p l i i r Reference Comments Medium

log pu: Po

6.084.476.105.08 '3.946.007.37 2

5.004.983.97

F+ + Cl- <=> PdCl+

[1967AHR][1976SMI/MAR][1976SMI/MAR][1980KRA][1982WAG/EVA][1985BAB/MAT][1985COL][1987BRO/WAN][1988PHI/HAL][1992PEA/BER]

T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/a

010000000

CIO4-

300

JNC TN8400 9 9 - 0 1 1


log (Sj,2

7.7410.708.16 '7.5110.609.44 2

8.777.717.51

: Pd2+ + 2CI- <=> PdCl2°

[1976SMI/MAR][1976SMI/MAR][1980KRA][1982WAG/EVA][1985BAB/MAT][1985COL][1987BRO/WAN][1988PHI/HAL][1992PEA/BER]

T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/a

10000000

log pu: Pd2+ + 3CI- <=> PdClf

10.2013.1010.58 '10.3312.94 2

11.4610.9910.32

[1976SMI/MAR][1976SMI/MAR][1980KRA][1982WAG/EVA][1985COL][1987BRO/WAN][1988PHI/HAL][1992PEA/BER]

T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0 .T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/a

1000000

log j3L4: Pd2+ + 4CI- <=> PdCl42-

12.2614.15 2

10.9411.5015.4011.46 '9.5012.0413.96 2

13.0911.8312.04

[1952LAT][1963GRI/KIS][1968GOL/HEP][1976SMI/MAR][1976SMI/MAR][1980KRA][1982SMI/MAR][1982WAG/EVA][1985COL][1987BROAVAN][1988PHI/HAL][1992PEA/BER]

T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=1T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0.5T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K, 1=0T= 298.15 K,I=n/a

010100

0.50000

C1O4-

cio4-

C1O4-

1 extrapolated with SIT from data from different sources [1980KRA]2 Calculated in this report assuming a log K of 34.32 for the reaction Pd2+ + 2e+ = Pd(s) (Section 9.1), corresponding

to a AfG° of 187.57 kJ/mol.

301

JNC TN8400 9 9 - 0 1 1

9.5 Amino complexes ofpalladium(II)

Although nitrogen containing ligands are not among the ligands selected by PNC, it may beuseful to have recommended constants for the amino complexes, because nitrogen containingligands form very stable complexes with Pd(II). Rasmussen and J0rgensen [1968RAS/J0R]determined the consecutive formation constants of all four amino complexes of Pd(II) usingvisible absorption spectroscopy. Like the spectroscopic study of Elding [1972ELD], this studyis reliable and can be recommended. Due to the isocoulombic equilibria, the ionic strengthdependence will be very small, and we believe that we can neglect it. The constants reported byRasmussen and J0rgensen [1968RAS/J0R] are thus recommended at any ionic strength (witherrors also between ±0.1 and ±0.2):

Pd2+ + NH3(aq) = PdNH32+ log (3,° = 9.6, As = 0

Pd2+ + 2 NH3(aq) = Pd(NH3)22+ log (32° = 18.5, Ae = 0

Pd2+ + 3 NH3(aq) = Pd(NH3)32+ log (33° = 26.0, Ae = 0

Pd2++ 4 NH3(aq) = Pd(NH3)42+ log P4° = 32.8, Ae = 0

9.6 Conclusions

It is unfortunate that the hydrolysis reactions cannot be quantified in a reliable way. The mainproblem that precludes any reliable quantification of the Pd(II) hydroxide system at this time isthe strong tendency toward polymerization. It is likely that the composition of such polymers,changes with pH and Pd(II) concentration.

However, the solubility of Pd(II) is very small in the absence of strong complexants such asorganic ligands containing amino donors. On the other hand it is known that the affinity ofmetal ions to oxide-type surfaces is related to their hydrolysis behavior, and it can thus beexpected that Pd(II) may form very stable surface complexes on these types of solids.

302

JNC TN8400 9 9 - 0 1 1

9. 7 Comments on selected references:

[1943TEM/WAT]: In this careful work, Templeton, Watt and Garner authors determinedthe redox potential of the couple Pd27Pd(s) in 0 to 4 M hydrochloric acid and in 4 Mperchloric acid at 25°C. Two measurements were performed in HC1O4 (both at / = 4M). The resulting potentials, corrected by the authors to refer to the molal hydrogenelectrode, are E = 0.9849 V and 0.9895 V for Pd2+ concentrations of 9.73 mM and2.62 mM, respectively, which gives a mean value of 0.9872 with a simple standarddeviation of ±0.0033 V. The reported ionic strength was 4.02 M (= 4.89 m) and thehydrogen concentration in the experiments was 4.007 M.

[1967IZA/EAT]: Izatt et al. investigated the hydrolysis of Pd2+ at 25°C by pH titration andspectroscopy at pH values between 0.5 and 2.3. The ionic strength was notmaintained constant and hence varied between 0.005 and 0.3 M. The concentrationof Pd(II) was in the millimolar range. Although 5 hours were required to establishconstant pH values, which might indicate polymerization reactions, only monomerichydrolysis species were considered in the reported data set. From the spectroscopicexperiments the authors reported errors as large as 30% and 85% in the extinctioncoefficients of ,,PdOH+" and ,,Pd(OH)2", respectively. In addition, only the choice ofsome selected data points resulted in positive values for the hydrolysis constants. Theauthors used an extended Debye-Hiickel equation to extrapolate the constants to / = 0and reported the following values from the two methods (pH titrations andspectroscopic measurements):

Pd2+ + OH" = PdOH+ log p°, = 13.0 (pH); 12.4 (sp)Pd2+ + 2 OH" = Pd(OH)2° log p°2 = 25.8 (pH); 26.5 (sp)

These constants cannot be accepted in this review due to the reasons mentionedabove. In addition, these constants do not follow the usual pattern of formationconstants, in particular those derived from the spectroscopic measurements: log p2

o is1.7 orders of magnitude larger than 2 x log p,°, which casts large doubts on thecredibility of the reported values.

In addition, the authors determined the redox potential for the reaction Pd2+ + 2 e~ =Pd(s). However, no raw data are given in any form. The only reported parameter is astandard half cell potential of (0.915 ± 0.005) V that had been extrapolated to I = 0by use of an extended Debye-Htickel equation. It is, however, not clear at what ionicstrength the redox measurements had been performed. The authors explained thatthey used ,,some of the data obtained in the determination of the formation constant ofthe tetracyanopalladate(II) ion" [1967IZA/WAT]. However, the purpose of that paperwas not the determination of a redox potential, and it is not indicated explicitly atwhat ionic strength the measurements were done. However, the authors corrected forhydrolysis of Pd2+ using the hydrolysis constants from [1967IZA/EAT], whichindicates that they did not work at pH = 0. We suppose that they worked at pH = 1or higher, as the ionic strengths reported in the paper [1967IZA/WAT] was / = 0.1 M

303

JNC TN8400 9 9 - 0 1 1

or lower. As a result of the insufficient credibility of the hydrolysis constants from[1967IZA/EAT], see assessment above, the reported redox potential cannot berecommended here.

[1970NAB/KAL]: Nabivanets and Kalabina used batch solubility measurements ofPd(OH)2(s) in the pH range from 0.7 to 13.5, at 17°C, and in 0.1 M perchloratesolutions to examine the complex formation of Pd2+ with the hydroxide ion. Theymeasured a constant Pd(II) concentration of 4x10~6 M from pH 3 to 11. Thereproducibility of the experiments was not verified. The determination of the Pd(II)concentration was done spectrophotometrically with xylenol orange. The minimumconcentration of Pd(II) required for this method, as reported by Otomo [1963OTO],was between 0.2 and 0.8 ppm, i.e., between 2X10"6 and 8X10"6 M Pd(II). Hence,the Pd(II) concentration in the solution to be analyzed needs to be even higher due tosubsequent dilution caused by the addition of acid and xylenol orange. The measuredminimum Pd(II) concentration of 4x10"6 M is thus very likely to reflect the detectionlimit of the analytical method used.

The authors reported the following constants, valid at 7 = 0.1 M (H+/Na+, C1O4~)-For the ion product of water, the authors used K^ = 10"14, which may be somewhaterroneous at the ionic strength of the system. The consecutive constants (log Kn)below refer to the same medium:

Pd2+ + OH"Pd2+ + 2 OH" =Pd2+ + 3 OH' =Pd2+ + 4 OH" =Pd(OH) 2°

PdOH+

Pd(OH)2°Pd(OH)3-Pd(OH)4

2"Pd(OH)2(s)

logloglogloglog

ppppK

, ( ' =

2 a=ii=

0.1 M)= 11.70.1 M)= 23.60.1 M)= 25.40.1 M)= 26.40.1 M)= 5.4

loglogloglog

K\ =K2 =K3 =K4 =

11.711.9

1.81.0

Looking at these values, it is not plausible why, on one hand, log K2 is larger thanlog Kx, and on the other hand, log K3 and log K4 so much smaller than log K2,although a steric hindrance is inconceivable. Quite certainly, something is suspectabout these data. It is likely that the minimum solubility measured by the authors wasequal to the detection limit of the method, and that the actual solubility of thePd(OH)2(s) used in the experiments was in fact much lower. In addition, extensiveformation of colloids would also induce significant changes in the above constants.However, there are other shortcomings in this paper.

Comparative dialysis with Co2+ as a reference ion in the pH range between 1.1 and2.0, where the solubility of Pd(OH)2(s) increases strongly, showed a relativedecrease of the Pd(II) concentration after dialysis with decreasing pH. The authorspostulated a general polymerization according to

p Pd(II) = [Pd(II)],

304

JNC TN8400 9 9 - 0 1 1

and calculated the ,,polymerization constants" from the dialysis experiments at pH = 1and Pd(II) concentrations ranging from UxlO"4 M to 3.4x10"* M. Assuming that nopolymerization occurs at a Pd(II) concentration of 6.3x10"5 M, the authors calculatedthe polymerization constants for different polymerization degrees (p = 2 to 6) at eachpalladium concentration and found more or less constant values only for p = 6 in theconcentration range of 1.3x10^ M to 2.4X10"4 M. The fact that only a slight increaseof the Pd(II) concentration to 3.4x10"4 M reduces the polymerization constant by afactor of as much as 100, gives rise to serious doubts about the reliability of thereported results. The calculation of log K2 (= log P2 - log P,) from the point at whichthe solubility of Pd(II) is twice the minimum solubility of 4x10^ M, presumes thatthe only monomeric aqueous species are Pd(OH)2(aq) and PdOH+. However,according to the resulting constants Pd2+ would also be present at significantconcentrations at these pH values. For the derivation of log P2 (at pH = 1.55 to 2),the data were corrected for the formation of the polymer (p = 6), although thepolymer data were obtained at a constant pH of 1. However, the pH dependency ofthe polymer formation, which is expected to be significant in this case, was notexamined by the authors. In addition, the suspicion that the minimum solubility of4x10~6 M Pd(II) merely reflects the detection limit of the analytical method used bythe authors casts additional doubt on the credibility of the reported results.

The authors also derived log P3 and log P4 from very few points (see[ 1970NAB/KAL, Figure 1]: 4 points in total at pH values of about 12.2, 12.6, 13.0and 13.6). The basis is too weak for these two constants to be selected.

[1984MIL/BUG]: Milic and Bugarcic investigated the hydrolysis of Pd2+ in sodium chloridesolutions. They used 0.0025, 0.005, 0.01, 0.02 and 0.04 mM Pd2+ solutions in 3 MNaCl, and 0.01 M Pd2+ solutions in 0.5, 1, 1.5, 2, 2.5 and 3 M NaCl solutions.They found that hydrolysis started near a pH of 8, and they postulated the formationof PdOH+ and Pd4(OH)4

4+. Hydrolysis increased with increasing Pd(II)concentration, and decreased with increasing NaCl concentration, but the maximumhydrolysis degree in the experiments was only 0.13. The authors qualified theobserved strong NaCl concentration dependence of the Pd(II) hydrolysis as ,,mediumeffect", although the formation of chloride complexes of Pd2+ had been shownspectroscopically by several authors, for example Elding [1972ELD]. The study ofMilic and Bugarcic [1984MIL/BUG] is not suitable to derive any hydrolysisconstants of Pd2+ and has to be discarded.

[1989KUM/BYR]: Kump and Byrne measured the UV absorbance of Pd(II) in seawater at25°C between pH = 7.1 and pH = 8.7. They used a surface seawater taken from theGulf of Mexico with a salinity of 36%o. The total chloride concentration was 0.558M, and the Pd(II) concentrations applied were 10~5 and 5xl0~6 M, respectively.From the change in the UV absorbances at four different wavelengths (290, 300, 310and 320 nm), the authors derived the following constants:

305

JNC TN8400 9 9 - 0 1 1

f[PdOH+] + [PdClOH0] + [PdCl2OH"] + [PdCLOH2"]) xn _ _V / s-i\

1 [Pd2+ ] + [PdCl+ ] + [PdCl5 ] + [PdClg ] [PdClJ" ]

([Pd(OH)5 ] + [PdCl(OH)2 ] + [PdCl2 (OH)!" l ) x t H + fD — O)

[Pd2+] + [PdCl+] + [PdCl§] + [PdCl3 ] + [PdCl2" ]

From their spectroscopic measurements, the authors obtained the values B] = (2.1 ±0.6)xl0~9 and B2 < 2.4x10~18. At the high chloride concentration of theseexperiments, we can expect the Pd(II) complexes to be saturated with chloride orhydroxide ligands, i.e., to contain no unhydrolyzed H2O ligands. Hence, theconcentrations of PdOH+, PdClOH0 and PdCl2OH' in Eq. (1), of Pd(OH)2° andPdCl(OH)2" in Eq. (2), as well as of Pd2+, PdCl+, PdCl2

0 and PdCl3" in bothequations, are assumed to be negligible. The derived constants Bx and B2 thencorrespond to the following equations:

PdCl42" + H2O(1) = PdCl3OH2-+ H++ Cr (3)

PdCl42" + 2 H2O(1) = PdCl2(OH)2

2" + 2 H++ 2 Cr (4)

The equilibrium constants for the above reactions then correspond to K{Z) = 5,x[Cl ]and K(4) = 52x[Cl~]2, respectively. The chloride concentration in the seawater usedwas 0.558 M, and the ionic strength was reported later [1993BYR/KUM] as / = 0.7M. It is preferable to convert these constants to refer to isocoulombic reactions, i.e.,reactions in which reactants and products have the same charge pattern, because theirdependence on the ionic strength is minimal:

= PdCl3OH2- + Cl" (5)4 = PdCl2(OH)2

2-+ 2 Cr (6)

The constants for these reactions are K(5) = K{3)/Kw and K{6) = K(A)/KW2. The

constants K(5) and ^(6) are assumed to be independent of the ionic strength. We arefurther forced to assume that reactions occurring due to the possible presence of othercomplexants in the natural surface seawater used, have a negligible influence on thevalues of Bx and B2. The ion product of water at 25°C and / = 0.7 M (NaCl) is aboutlog ^w = -13.72 [1976BAE/MES]. The selected constants for the mixed chloro-hydroxo complexes of Pd(II) are thus:

log K°(5, 298.15 K) = log Bl + log [CT] - log Kw = 4.8log K°(6, 298.15 K) = log B2 + 2 log [Cl"] - 2 log Kw < 9.3

[1991TAI/JAN] Tait et al. used Raman and UV/VIS spectroscopy to determine the speciationof Pd(II) chloro complexes at varying pH and chloride concentration. From theRaman spectra they assigned typical absorption peaks to PdCl4

2~, PdQ3(H2O)~ and

306

J N C T N 8 4 0 0 9 9 - O i l

PdCl2(H2O)2(aq). They concluded that in seawater with a chloride concentration of0.558 M, they assumed that they had mixed chloro-hydroxo complexes and solidphases of Pd(II). However, they claimed that this would only be possible ,,whenOH" serves as a limiting reagent (i.e., only when not enough OH" ligand is present tocompletely replace the chlorides)". This claim is surprising, because hydroxide ion asa ligand is always available in aqueous solution. In their opinion, in natural systems,mixed chloro-hydroxo complexes would not be stable ,,because OH" readily replacesCl" over a negligible pH change". Experimental confirmation of this statement is notgiven. It should be mentioned that Byrne and Kump [1993BUR/KUM] contradictedthis statement and claimed that mixed-ligand complex formation would be animportant aspect of Pd(II) hydrolysis in natural solutions, cf. comments under[1993BYR7KUM].

[1991WOO] Wood investigated the hydrolysis behavior of Pd(II) and Pt(II) by solubilitymeasurements of the respective metals in 0.0004 M to 10.0 M NaOH solutions. Forthe solubility measurements Pd metal (,,Pd shot") and Pt wire were both put incontact with the respective NaOH solution and heated to 85°C for 18 days. Thetemperature was then held at 70°C and 60°C for 14 days each, and then lowered to25°C. Samples were analyzed after a total of 319 days and 465 days, respectively.The whole procedure was carried out in a glove box under nitrogen atmosphere.Prior to analysis by graphite furnace atomic absorption, the authors separated Pd andPt by coprecipitation with tellurium. At the end of the experiments, they examined thesolid phases microscopically and by X-ray diffraction, and they found no evidencefor the presence of Pd(II) hydroxide. However, they detected a white precipitatecontaining Na, C and O, and they concluded that CO2 had diffused into the reactionsolutions, or that some organic matter might have been leached from the polyethylenebottles at higher temperatures. The authors derived the following constants from theirmeasurements:

Pd(s) + 2OH" = Pd(OH)2(aq) + 2 e" log p2 = -12.0Pd(s) + 3 OH" = Pd(OH)3- + 2 e" log p3 = -10.0 (log K, = 2.0)

Unfortunately, the ionic strength was not held constant during the experiments. Infact, the ionic strength varied by as much as four orders of magnitude, which per seprecludes the recommendation of these data in the selected data set. It should,however, be noted that the authors suspected their own Eh measurements to be inerror, because their measured points fall into the stability range of Pd(OH)2(s)(instead of Pd metal) in the Eh-pH diagram. In our view it might be equally justifiedto question the correctness of the Eh-pH diagram that [1989WOO/MOU] used,which is based on the more than doubtful data of Izatt et al. [1967IZA/EAT] andNabivanets and Kalabina [1970NAB/KAL].

[1993BYR/KUM] This is a comment paper on the publication of Tait et al. [1991TAI/JAN]including a review of the Pd(II) chloro complexes and hydroxo complexes, as wellas extended comments on the formation of mixed ligand complexes. In particular,

307

JNC TN8400 9 9 - 0 1 1

they contradicted the statement in [1991TAI/JAN] that hydroxide would act as alimiting reagent. They correctly stated that in hydrolysis equilibria ,,hydroxide cannever be a limiting reagent". Byrne and Kump [1993BYR/KUM] reported, fromtheir earlier spectroscopic data in seawater (0.558 M Cl") [1989KUM/BYR], thefollowing constant valid for seawater conditions:

PdCl42- + OIT = PdCl3OH2- + Cr log K(I = 0.7 M) = 4.8

This means that in seawater, hydrolysis of Pd2+ (more correctly: PdCl42 ) starts at a

pH between 7 and 8. This finding is consistent with the experimental findings ofKazakova and Ptitsyn [1967KAZ/PTI] and Milic and Bugarcic [1984MIL/BUG].

Byrne and Kump [1993BYR/KUM] presented a model to statistically predict theformation constants of simple mixed ligand complexes. In the case of chloride andhydroxide as ligands, the following equation is used to predict the formation ofmixed complexes:

(7)

Here, a{j accounts for the ,,promotion" of mixed ligand complex formation throughstatistical effects, and in the present case where (/ + j) = 4, the term a,y = -0.402[1982BEL/KOL]. The term log 8,y accounts for ,,ligand effects" which are usuallyconsidered negligible for the substitution of equally-charged ligands. In this way, byusing log (34(PdCl4

2-) = 11.32 and log p4(Pd(OH)42~) = 26.4, they predicted the

following constants for the mixed chloro-hydroxo complexes according to Eq. (7):log p31 = 15.69, log P22 = 19.66, and log p13 = 23.23.

Finally, Byrne and Kump [1993BYR/KUM] recognized the fact that insufficientexperimental information is available to obtain reliable hydrolysis constants of Pd2+,and they estimated them by making the following two assumptions: 1) The value oflog P4(Pd(OH)4

2") from [1970NAB/KAL] is correct, and 2) The ratios of thestepwise formation constants of the Pd(II) hydroxide complexes are the same asthose of the Pd(II) chloride complexes. In this way they obtained the followingconstants: log p,(Pd(OH)+) = 8.1, log P2(Pd(OH)2°) = 15.34, log P3(Pd(OH)3") =21.4 and log P4(Pd(OH)4

2") = 26.5. However, according to this set of constants,Pd2+ would be present as unhydrolyzed aqua ion below a pH of about 5, which is incontradiction to all experimental evidence obtained hitherto in non-complexingaqueous solutions.

308

JNC TN8400 9 9 - 0 1 1

10 References

[1883SCH] Schulze, H. (1883), J. prakt. Chemie, 27, 320 (as cited in [1952LAT].

[1906GOL/ECK] Goldschmidt, H. and Eckhardt, M. (1906), Uber die Reduktion vonNitrokorpern durch alkalische Zinnoxydullosungen. Zeitschrift phys. Chemie, 56,385-452.

[1918NOY/CHO] Noyes, A.A. and Chow, M. (1918), The potentials of the bismuth-bismuthoxychloride and the copper-cuprous-chloride electrodes. J. Am. Chem.Soc, 40, 739-763.

[1922APP/REI] Applebey, M.P. and Reid, R.D. (1922), The isomerism of metallic oxides.Part I. Lead monoxide. J. Chem. Soc, 121, 2129-2136.

[1922GER] Gerke, R.H. (1922), J. Am. Chem. Soc, 44, 1684.

[1923GRU/SCH] Grube, G. and Schweigardt, F. (1923), Uber das elektrochemischeVerhalten von Wismut und Antimon in alkalischer Losung. Z. elektrochem. angew.phys. Chemie, 29, 257-265.

[1923SMI] Smith, D.F. (1923), J. Am. Chem. Soc, 45, 360 (as cited in [1952LAT]).

[1923SMI/WOO] Smith, D.F. and Woods, H.K. (1923), The free energy and heat offormation of lead monoxide. J. Am. Chem. Soc.45, 2632-2637.

[1923SWI] Swift, E.H. (1923), J. Am. Chem. Soc. 45, 371 (as cited in [1952LAT]).

[1924SCH] Schuhmann, R. (1924), The free energy of antimony trioxide and the reductionpotential of antimony. J. Am. Chem. Soc, 46, 52-59.

[1926GRU/LIN] Gruner, J.W. and Lin, S.C. (1926), Solution of tin minerals studied.Engineering Mining Journal-Press, 121, 924.

[1928PRY] Prytz, M. (1928), Komplexbildung in Stannochlorid- und Stannobromid-losungen. Z. anorg. allg. Chemie, 172, 147-167.

[1928PRY2] Prytz, M. (1928), Hydrolysemessungen in Stannosalzlosungen. Z. anorg. allg.Chemie, 174, 355-375.

[1928RAN/SPE] Randall, M. and Spencer, H.M. (1928), Solubility of lead monoxide andbasic lead carbonate in alkaline solution. J. Am. Chem. Soc, 50, 1572-1583.

[1929CAR] Carmody, W.R. (1929), Studies in the measurements of electromotive force indilute aqueous solutions. I. A study of the lead electrode. J. Am. Chem. Soc, 51,2905-2909.

309

JNC TN8400 9 9 - 0 1 1

[1929MIL/JOW] Millet, H. and Jowett, M. (1929), The solubilities of lead phosphates. J.Am. Chem. Soc, 51, 997-1005.

[1929MIL] Millar, R. (1929), The heat capacities at low temperatures of the oxides of tin andlead. J. Am. Chem. Soc., 51, 207-223.

[1930RAN/MUR] Randall, M. and Murakami, S. (1930), The free energy of stannoushydroxyl chloride and the activity coefficient of stannous chloride and stannous ion.J. Amer. Chem. Soc, 52, 3967-3971.

[1930RIG/DAV] Righellato, E. and Davies, C.W. (1930), The extent of dissociation of saltsin water. Part II. Uni-bivalent salts. Trans. Faraday Soc., 26, 592-600.

[1932JOW/PRI] Jowett, M. and Price, H.I. (1932), Solubilities of the phosphates of lead.Trans. Faraday Soc., 28, 668-681.

[1933FEI] Feitknecht, W. (1933), Gleichgewichtsbeziehungen bei den schwerloslichenbasischen Salzen. Helvetica Chimica Acta, 16, 1302-1315.

[1934HUE/TAR] Huey, C.S. and Tartar, H.V. (1934), The stannous-stannic oxidation-reduction potential. J. Am. Chem. Soc, 56, 2585-2588.

[1935AKE/TUR] Akerlof, G. and Turck, H.E. (1935), The solubility of some strong highlysoluble electrolytes in methyl alcohol and hydrogen peroxide - water mixtures at25°. J. Am. Chem. Soc, 57, 1746-1750.

[1937HOA] Hoar, T.P. (1937), The corrosion of tin in nearly neutral solutions. Trans.Faraday Soc, 33, 1152-1167.

[1939BLO] Bloom, M.C. (1939), The mechanism of the genesis of polymorphous forms.Am. Mineralogist, 24, 281-292.

[1939GAR/VEL] Garrett, A.B., Vellenga, S. and Fontana, G.M. (1939), The solubility onred, yellow, and black lead oxides (2) and hydrated lead oxide in alkaline solutions.The character of the lead-bearing ion. J. Am. Chem. Soc., 61, 367-373.

[1939GOR] Gorman, M. (1939), Hydrolysis of stannous ion in stannous perchloratesolutions. J. Am. Chem. Soc, 61, 3342-3344.

[1941GAR/HEI] Garrett, A.B. and Heiks, R.E. (1941), Equilibria in the stannous oxide-sodium hydroxide and in the stannous oxide-hydrochloric acid systems at 25°.Analysis of dilute solutions of stannous tin. J. Am. Chem. Soc, 63, 562-567.

[1942GOR/LEI] Gorman, M. and Leighton, P.A. (1942), Solubility of stannous oxide inperchloric acid. J. Am. Chem. Soc, 64, 719-721.

[1942KOL/PER] Kolthoff, I.M., Perlich, R.W. and Weiblen, D. (1942), The solubility oflead sulfate and of lead oxalate in various media. J. Phys. Chem., 46, 561-570.

310

J N C T N 8 4 0 0 9 9 - O I L

[1943SCH/RIT] Schumb, W.C. and Rittner, E.S. (1943), Polymorphism of bismuthtrioxide. J. Am. Chem. Soc, 65, 1055-1060.

[1943TEM/WAT] Templeton, D.H., Watt, G.W. and Garner, C.S. (1943), The formalelectrode potentials of palladium in aqueous hydrochloric and perchloric acidsolutions. Stability of chloropalladite ion. J. Am. Chem. Soc, 65, 1608-1613.

[1945LIN] Lingane, J. (1945), Coulometric Analysis. J. Am. Chem. Soc., 67, 919-925.

[1945PED] Pedersen, K.J. (1945), The acid dissociation of the hydrated lead ion and theformation of polynuclear ions. Matematisk-fysiske meddelelser, 22, 3-29.

[1947BRO/DEV] Broene, H.H. and DeVries, T. (1947), J. Am. Chem. Soc, 69, 1644.

[1947GRA/SIL] Graner, F. and Sillen, L.G. (1947), Acta Chem. Scand., 1, 631.

[1948TOU/MOU] Tourky, A.R. and Mousa, A.A. (1948), 150. Studies on some metalelectrodes. Part V. The amphoteric properties of antimony tri- and pent-oxide. J.Chem. Soc, 2, 759-763.

[1949RIC/POP] Riccoboni, L., Popoff, P. and Arich, G. (1949), Comportamentopolarografico delle soluzione di stagno stannoso. Gazzetta Chimica Italiana,LXXIX, 547-587.

[1950DUK/COU] Duke, F.R. and Courtney, W.G. (1950), The stannous chlorideequilibrium. Iowa State J. Research," 24, 397-403.

[1951DUK/PIN] Duke, F.R. and Pinkerton, R.C. (1951), The role of halide ions in theferric-stannous reaction. J. Am. Chem. Soc, 73, 3045-3049.

[1951SWI/GAR] Swinehart, D.F. and Garrett, A.B. (1951), The equilibria of two basicbismuth nitrates in dilute nitric acid at 25°. J. Am. Chem. Soc., 73, 507-510.

[1952GAY/GAR] Gayer, K.H. and Garrett, A.B. (1952), The equilibria of antimonousoxide (rhombic) in dilute solutions of hydrochloric acid and sodium hydroxide at25°. J. Am. Chem. Soc, 9, 2352-2354.

[1952LAT] Latimer, W.M. (1952), The oxidation states of the elements and their potential sinaqueous solutions, Prentice-Hall Inc., New York, 392 pp.

[1952VAN/RHO] Vanderzee, C.E. and Rhodes, D.E. (1952), Thermodynamic data on thestannous chloride complexes from electromotive force measurement. J. Am. Chem.Soc, .74, 3552-3555.

[1952WOR] Worrell, W. (1952), The free energy of formation of niobium dioxide between1100 and 1700°K. J. Phys. Chem., 68, 952-953.

3J1

JNC TN8400 9 9 - 0 1 1

[1953BAB/GOL] Babko, A.K. and Golub, A.M.(1953), Sbornik Statej Obsh. khim. SSSR,1, 64. As cited in [1957AHR/GRE].

[1953HAI] Haight, G.P. Jr. (1953), Polarography of tripositive antimony and arsenic.Cathodic reduction of antimonous in strong hydrochloric acid and anodic oxidationof arsenite and stibnite in strong sodium hydroxide. J. Am. Chem. Soc, 75, 3848-3851.

[1953HEM] Hemley, J.J. (1953), A study of lead sulfide solubility and its relation to oredeposition. Economic Geology Bull. Soc. Economic Geologists, 48, 113-138.

[1953HER/SMI] Hershenson, H.M., Smith, M.E. and D.N., Hume (1953), A polarografic,potentiometric and spectrophotometric study of lead nitrate complexes. J. Am.Chem. Soc., 75, 507-511.

[1954COU] Coughlin, J.P. (1954), Contributions to the data on theoretical metallurgy: XII.Heats and free energies of formation of inorganic oxides. US Bureau MinesBulletin, 542, 1-80.

[1954FAU2] Faucherre, J. (1954), No. 57 - Sur la constitution des ions basiquesmetalliques. Bulletin Soc. Chim. France, 252-267.

[1954FAU] Faucherre, J. (1954), 29. - Sur la constitution des ions basiques metalliques.Bulletin Soc. Chim. France, 128-142.

[1955BIG/PAR] Biggs, A.I., Parton, H.N. and Robinson, R.A. (1955), The constitution ofthe lead halides in aqueous solution. J. Am. Chem. Soc., 9, 5844-5849.

[1955BRU] Brubaker, C.H. (1955), The hydrolysis of tin(IV) in sulfuric acid. J. Am.Chem. Soc, 77, 5265-5269.

[1955DEL/ZOU] Deltombe, E., Zoubov, N. and Pourbaix, M. (1955), Comporementelectrochimique de l'etain. Diagramme d'equilibre tension-pH du systeme Sn-H2O,a 25°C. Corrosion de l'etain; etamages electrolytique et chimique. Rapport Tech. Rt25 du CEBE LCOR, 1-24.

[1955GLE/PEU] Glemser, O. and G., Peuschel (1955), Beitrag zur Kenntnis des SystemsPdO/H2O. Z. anorg. allg. Chemie, 281, 44-53.

[1955NAN] Nancollas, G.H. (1955), Thermodynamics of ion association. Part I. Leadchloride, bromide, and nitrate. J. Chem. Soc, 4, 1458-1463.

[1956BAL/DAV] Bale, W.D., Davies, E.W. and Monk, C.B. (1956), Spectrophotometricstudies of electrolytic dissociation. Trans. Faraday Soc, 52, 816-823.

[1956ROB/THE] Robin, M. J. and Theolier, M.A. (1956), Preparation et proprietls del'hydroxide de plomb (Note de laboratoire). Bull. Soc. Chim. France, 116, 680-681. As cited in [1995MAR/MAC].

312

JNC TN8400 9 9 - 0 1 1

[1956KIV/RIN] Kivalo, P. and Ringbom, A. (1956), Polarographic determination of thesolubility products ot the sulfides of lead and cadmium. Suomen Kemistilehti, 29,109-13.

[1957AHR/GRE] Ahrland, S. and I., Grenthe (1957), The stability of metal halidecomplexes in aqueous solution. Acta Chem. Scand., 11, 1111-1130.

[1957EGG] Egger, K. (1957), University of Berne, Berne, Switzerland, Lie. Arbeit, as citedin [1963FEI/SCH].

[1957NEW/HUM] Newman, L. and Hume, D.N. (1957), J. Am. Chem. Soc, 79, 4576.Ascitedin[1970KAN].

[1957OLI] Olin, A. (1957), Studies on the hydrolysis of metal ions; 19. The hydrolysis ofBismuth (III) in perchlorate medium. Acta Chem. Scand., 11, 1445-1456.

[1957PIT/POU] Pitman, A.L., Pourbaix, M. and Zoubov, N. (1957), Potential-pH diagramof the antimony-water system. J. Electrochem. Soc, 104, 594-600.

[1958JOH/KRA] Johnson, J.S. and Kraus, K.A. (1958), Hydrolytic behavior of metalions. IX. Ultracentrifugation of Sn (IV) in acidic chloride and perchlorate solutions.J. Phys. Chem., 63, 440-441.

[1958TOB] Tobias, T.S. (1958), Studies on the hydrolysis of metal ions. 21. The hydrolysisof the tin(II) ion, Sn2+. Acta Chem. Scand., 12, 198-223.

[1959DES/PAN] Desideri, P. and Pantani, F. (1959), La riduzione del bismuto (III)all'elettrodo a goccia nelle soluzioni cloridriche. Ricerca Scientifca, 29, 1436-1445.

[1959FAU/BON] Faucherre, J. and Bonnaire, Y. (1959), Electrochemie - Sur la constitutiondes carbonates complexes de cuivre et de plomb. Comptes rendus hebdomadairesde seances de l'academie des Sciences, 3705-3707.

[1959JAG] Jager, L. (1959), Bestimmungen der Loslichkeit und der Aktivitatskoeffizientenvon Bleisulfat in Salpetersaure bei 25°C mittels markierter Atome 3 5S. CollectionCzechoslov. Chem. Commun., 24, 1703-1705.

[1959JOH/KRA] Johnson, J.S. and Kraus, K.A. (1959), Hydrolytic behavior of metalions. X. Ultracentrifugation of lead (II) and tin (IV) in basic solution. J. Phys.Chem., 81, 1569-1572.

[1959OLI] Olin, A. (1959), Studies on the hydrolysis of metal ions; 23. The hydrolysis ofthe Ion Bi6(OH)6+

12 in perchlorate medium. Acta Chem. Scand., 13, 1791-1808.

[1959PAN/DES] Pantani, F. and Desideri, P.G. (1959), II comportamento polarograficodell'antimonio (III) in soluzioni chloridriche. Acta Chim. Italiana, 89, 1360-1372.

313

JNC TN8400 9 9 - 0 1 1

[1959UGG] Uggla, R. (1959), A study of the corrosion and passive states of zinc, lead andtin in the system water oxygen-nitrogen-carbon dioxide, Ann. Acad. Sci. Fennicae,97, 1-74.

[1960CAR/OLI] Carell, B. and Olin, A. (I960), Studies on the hydrolysis of metal ions ;31. The complex formation between Pb2+ and OH" in Na+ (OH", C1O"4) medium.Acta Chem. Scand., 14, 1999-2008.

[1960FRI/SAR] Fridman, Y.D., Sarbaev, D.S., Sorochan, R.I. (1960), Equilibrium insolutions of heterogeneous complex compounds of metals: Mixed lead andcadmium halides, Russ. J. Inorg. Chem., 5, 381-387.

[1960JAN/ERT1] Jander, G. and Ertel, D. (1960), Uber Niobsauren und wasserloslicheAlkaliniobate - I: Lichtabsorptions- und Diffusionsmessungen an Akkaliniobat-losungen. J. Inorg. Nucl. Chem., 14, 71-76.

[1960JAN/ERT2] Jander, G. and Ertel, D. (1960), Uber Niobsauren und wasserloslicheAlkaliniobate- II: Praparativ-analytische Untersuchungen. J. Inorg. Nucl. Chem.,14, 77-84.

[1960JAN/ERT3] Jander, G. and Ertel, D. (1960), Uber Niobsauren und wasserloslicheAlkaliniobate - III: Konduktiometrische Titrationen und vergleichende rontgeno-graphische Untersuchungen. Das Hydrolyseschema der Isopolyniobate. J. Inorg.Nucl. Chem., 14, 85-90.

[1960KAZ] Kazantsev, A.A. (1960), The solubility of Pb(NO3)2. Russ. J. Inorg. Chem., 5,773-775.

[1960NAS/MER] Nasanen, R. and Merilainen, P. (1960), Thermodynamic properties oflead hydroxide iodidePb(OH)I. Suomen Kern. B, 33, 149-151.

[1960OLI1] Olin, A. (1960), Studies on the hydrolysis of metal ions. Acta Chem. Scand.,14, 126-150.

[1960OLI2] Olin, A. (1960), Studies on the hydrolysis of metal ions. Acta Chem. Scand.,14, 814-822.

[1960TOB] Tobias, R.S. (1960), Studies on hydrolyzed bismuth(III) solutions. Part I.e.m.f. titrations. J. Am. Chem. Soc, 82, 1070-1072.

[1961AHR/GRE] Ahrland, S. and I., Grenthe (1961), Correction to "The stability of metalhalide complexes in aqueous solution III. The chloride, bromide and iodidecomplexes of bismuth". Acta Chem. Scand., 4, 932.

[1961CON/PAU] Connick, R.E. and Paul, A.D. (1961), The fluoride complexes of silverand stannous ions in aqueous solution. J. Phys. Chem., 65, 1216-1220.

314

JNC TN8400 9 9 - 0 1 1

[1961GAT/RIC] Gate, S.H. and Richardson, E. (1961), Some studies on antimonic acid. -I: Some properties, effect of H2O2, and reaction with polyhydroxy compounds. J.Inorg. Nucl. Chem., 22, 257-263.

[1961NAS/MER] Nasanen, R., Merilainen, P. and Leppanen, K. (1961), Potentiometricdetermination of the solubility product of lead carbonate. Acta Chem. Scand., 15,913-918.

[1961OLI] Olin, A. (1961), Studies on the hydrolysis of metal ions. Svensk KemiskTidskrift, 73, 482-500.

[1961RAB/MOO] Rabideau, S.W. and Moore, R.H. (1961), The application of high-speedcomputers to the least squares determination of the formation constants of thechloro-chomplexes of tin (II). J. Phys. Chem., 65, 371-373.

[1961RAM/STE] Ramette, R.W. and Stewart, R.F. (1961), Solubility of lead sulfate as afunction of acidity. The dissociation of bisulfate ion. J. Phys. Chem., 65, 243-246.

[1962CAR/OLI] Carell, B. and Olin, A. (1962), Studies on the hydrolysis of metal ions ;42. A thermodynamical study of hydrolysed Pb(C104)2 solutions. Acta Chem.Scand., 16, 2350-2356.

[1962HAI/ZOL] Haight, G.P., Zoltewicz, J. and W., Evans (1962), Solubility studies onsubstituted ammonium salts of halide complexes. Acta Chem. Scand., 16, 311-322.

[1962PAJ/OLI] Pajdowski, L. and Olin, A. (1962), Studies on the hydrolysis of metal ions.Acta Chem. Scand., 16, 983-991.

[1963BAB/LUK] Babko, A.K., Lukachina, V.V. and Nabivanets, B.I. (1963), Solubilityand acid-base properties of tantalum and niobium hydroxides. Russ. J. Inorg.Chem., 8, 957-961.

[1963FEI/SCH] Feitknecht, W. and Schindler, P. (1963), Solubility constants of metaloxides, metal hydroxides and metal hydroxide salts in aqueous solution. Pure Appl.Chem., 1, 130-199.

[1963GRI/KIS] Grinberg, A.A., Kiseleva, N.V.and Gel'fman, M.I. (1963), Instabilityconstants of palladium complexes: compounds of the type K2PdX4. Dokl. Akad.Nauk SSSR, 153, 1327-1329.

[1963LEF/MAR] Lefebvre, M.J. and Maria, H. (1963), Etude des equilibres dans lessolutions recente des polyantimoniates. Comptes rendus hebdomadaires desSeances de l'Academie des Sciences, 256, 3121-3125.

[1963MEI/SCH] Meites, L. and Schlossel, R. (1963), Kinetics of the reaction betweenantimony (III) and ferricyanide in alkaline media. J. Am. Chem. Soc, 67,

315

JNC TN8400 9 9 - 0 1 1

[1963MES/HUM] Mesaric, S. and D.N., Hume (1963), A polarograhic study of leadfluoride complexes and solubility. Inorg. Chem., 2, 788-790.

[1963MIR/KUL] Mironov, V.E., Kul'ba, F. Y., Fedorov, V.A. and Nikitenko, T.F.(1963), A potentiometric study of chloro-complexes of bismuth. Russ. J. Inorg.Chem., 8, 964-967.

[1963OTO] Otomo, M. (1963), The spectrophotometric determination of palladium(H) withxylenol orange, Bull. Chem. Soc. Japan, 36, 889-892.

[1963SHA/DAV] Shah, S.N. and Davies, D.E. (1963), The anodic behaviour of tin inalkaline solutions - I. 0.1 M sodium borate solutions. Electrochim. Acta, 8, 663-678.

[1963WIC/BLO] Wicks, C.E. and Block, F.E. (1963), Thermodynamic properties of 65elements,their oxides, halides, carbides and nitrides. US Bureau Mines Bulletin,605, 146 p.

[1963WIL/BRO] Wilcox, D.E. and Bromley, L.A. (1963), Computer estimation of heat andfree energy of formation for simple inorganic compounds. Ind. Eng. Chem., 55,32-39.

[1964BAK] Baker (1964), Mineral equilibrium studies of the pseudomorphism ofpyromorphite by hinsdalite. Am. Mineralogist, 49, 607-613.

[1964HIR] Hirayama, C. (1964), Thermodynamic properties of solid monoxides,monosulfides, monoselenides and monotellurides of Ge, Sn and Pb. J. Chem.Eng. Data, 9, 65-68.

[1964HUG] Hugel, R. (1964), No 253. - Etude de l'hydrolyse de l'ion Pb2+ dans lessolutions de perchlorate de sodium. Bulletin Soc. Chim. France, 7, 1462-1469.

[1964LIN/TU] Ling, C.-C. and Tu, Y.-M. (1964), Calculation of the stability constants ofmononuclear and polynuclear complexes. Russ. J. Inorg. Chem., 9, 727-732.

[1964MIR/KUL] Mironov, V.E., Kul'ba, F.Y., Fedorov, V.A. and Fedorova, A.V.(1964), Chloro-complexes of lead (II). Russ. J. Inorg. Chem., 9, 1155-1157.

[1964NEU] Neumann, G. (1964), On the hydroysis of niobates in 3 M K(C1) medium. ActaChem. Scand., 18, 278-281.

[1964SIL/MAR] Sillen, L.I. and Martell, A.E. (1964), Stability Constants of Metal-IonComplexes, Chemical Society, London.

[1965BAR/BAR] Baranova, N.N. and V.L., Barsukov (1965), Transport of lead byhydrothermal solutions in the form of carbonate complexes. Geokhimiya, 9, 1093-1100.

316

J N C T N 8 4 0 0 99 - O i l

[1965BOT/CIA] Bottari, E. and Ciavatta, L. (1965), On the complex formation between lead(II) and fluoride ions. J. Inorg. Nucl. Chem., 27, 133-141.

[1965FRI/VER] Fridman, Y. D., Veresova, R.A. and Lukyanents, A.N. (1965), Neorg.Fiz. Khim., 13, as cited in [1968MIS/GUP].

[1965GAR/CHR] Garrels, R.M., Christ, C.L. (1965), Solutions, Minerals and Equilibria,Freeman, Cooper & Co, San Francisco, 450p.

[1965HUG1] Hugel, R. (1965), No 174. - Etude de 1'hydrolyse de l'ion Pb2+ dans lessolutions de nitrate de sodium. Bulletin Soc. Chim. France, 5, 968-971.

[1965HUG2] Hugel, R. (1965), No 175. - Complexes nitrato de l'ion Pb2+ . I. - Etudepotentiometrique. Bulletin Soc. Chim. France, 1, 971-975.

[1965HUG3] Hugel, R. (1965), No 312. - Complexes nitrato de l'ion Pb2+ . II. - Etudepolarographique. Bulletin Soc. Chim. France, 2017-2019.

[1965JAN/HAR1] Jander, G. and Hartmann, H.-J. (1965), Uber Reaktionen vonAntimon(m) in wassriger Losung. I. Lichtadsorption- und Diffusionsversuchesowie die Bestimmung des Ionenladungsvorzeichens im sauren Bereich. Z. anorg.allg. Chemie, 339, 239-247.

[1965JAN/HAR2] Jander, G. and Hartmann, H.-J. (1965), Uber Reaktionen vonAntimon(III) in wassriger Losung. II. Praparativ-analytische Untersuchungen. Z.anorg. allg. Chemie, 339, 248-255.

[1965JAN/HAR3] Jander, G. and Hartmann, H.-J. (1965), Uber Reaktionen vonAntimon(III) in wassriger Losung. III. Polarographische Messungen. Z. anorg.allg. Chemie, 339, 256-261.

[1965MES/IRA] Mesmer, R.E. and Irani, R.R. (1965), Metal complexing by phosphoruscompounds - VIII: Complexing of tin (II) by poly phosphates. J. Inorg. Nucl.Chem., 28, 493-502.

[1967AHR] Ahrland, S. (1967), 36. Enthalpy and entropy changes by formation of differenttypes of complexes. Helvetica Chimica Acta, 50, 306-318.

[1967BAR] Baranova, N.N. (1967), Zhur. Neorg. Khim., 12, 1438. As cited in[1969BAR].

[1967IZA/EAT] Izatt, R.M., Eatough, D.and Christensen, J.J. (1967), A study of Pd2+(aq)hydrolysis. Hydrolysis constants and the standard potential for the Pd, Pd2+

couple. J. Chem. Soc. (A), 1301-1304.

[1967IZA/WAT] Izatt, R.M., Watt, G.D., Eatough, D., Christensen, J.J.(1967),Thermodynamics of metal cyanide co-ordination: Part VII. Log K, AH°, and AS°

317

JNC TN8400 9 9 - 0 1 1

values for the interaction of CN" with Pd2+. AH° values for the interaction of Cl~and Br" with Pd2+, J. Chem. Soc. (A), 1304-1308.

[1967KAP/NAB] Kapantsyan, E.E. and Nabivanets, B.I. (1967), Determination of thecomposition and stability of bismuth nitrate complexes by ion-exchangechromatography. Soviet Progress Chem., 33, 961-964.

[1967KAZ/PTI] Kazakova, V.I. and Ptittsyn, B.V. (1967), Hydrolysis of halogeno-complexes of palladium. Russ. J. Inog. Chem., 12, 323-326.

[1967LOM/VAN] Loman, H. and van Dalen, E. (1967), On the use of cation exchangers forthe study of complex systems-II; The system Bismuth (Ill)-bromide, iodide andfluoride. J. Inorg. Nucl. Chem., 29, 699-706.

[1967SCH/ING] Schorsch, G. and Ingri, N. (1967), Determination of hydroxide ionconcentration by measurements with a lead amalgam electrode. Plumbate and borateequilibria in alkaline 3.0 M NaCl-medium: Absence of monoborate (-2) and (-3)ions. Acta Chem. Scand., 21, 2727-2735.

[1967VAS/LOB] Vasil'ev, V.P. and Lobanov, G.A. (1967), Influence of temperature andionic strength on the heat of formation of monohalogeno-complexes of bismuth inaqueous solution. Russ. J. Inorg. Chem., 12, 463-466.

[1967VAS/YUS] Vasil'ev, L.N. and Yustus, Z.L. (1967), Behavior of antimony (III) by themethod of vector polarography with a stationary mercury electrode. SovietElectrochem., 3, 842-845.

[1967WAR] Warner, J.S. (1967), The free energy of formation of palladium oxide,Electrochem. Soc. J., 114, 68-71.

[1968AND] Andersson, L.H. (1968), On the separation and determination of fluorine. III.Composition and some properties of lead chloride fluoride precipitates. Arkiv forKemi, 30, 57-69.

[1968CHA/FLE] Charette, G.G. and S.N., Flengas (1968), Thermodynamic properties ofthe oxides of Fe, Ni, Pb, Cu, and Mn, by emf mesurements. J. Chem.Thermodyn., 115,796-804.

[1968GOL/HEP] Goldberg, R.N. and Hepler, L.G. (1968), Thermochemistry and oxidationpotentials of the platinum group metals and their compounds. Chem. Rev., 68,229-251.

[1968HAL/SLA] Hall, F.M. and Slater, S.J. (1968), Determination of the stability constantsof the fluoride complexes of tin(II) using the fluoride electrode. Aust. J. Chem.,21, 2663-2667.

[1968HSE/REC] Hseu, T.-M. and Rechnitz, G.A. (1968), Analytical study of a sulfide ion-selective membrane electrode in alkaline solutions. Anal. Chem., 40, 1054-1060.

318

JNC TN8400 9 9 - 0 1 1

[1968JOH/OLI] Johansson, G. and Olin, A. (1968), Acta Chem. Scand., 22, 3197 (as citedin [1976BAE/MES]).

[1968LEV] Levanda, O.G. (1968), Influence of ionic strength on the stability constant of thetetrachloropaladate (II) ion in water. Russ. J. Inorg. Chem., 13, 1707-1709.

[1968MIS/GUP] Mishra, S.K. and Gupta, Y.K. (1968), Spectrophotometric study of thehydrolytic equilibrium of Sb (III) in aqueous perchloric acid solution. Indian J.Chem., 6, 757-759.

[1968NOR/KAZ] Norakidze, I.G., Kazakov, V.A. and Vagramyan, A.T. (1968),Equilibrium potential of antimony electrode in strongly acid chloride solution.

[1968POP/DAL] Pope, M.T. and Dale, B.W. (1968), Isopoly-vanadates, -niobates, and -tantalates. Quart. Rev. Chem. Soc, 22, 527-549.

[1968RAS/JOR] Rasmussen, L. and Jorgensen, C.K. (1968), Palladium(II) complexes. I.Spectra and formation constants of ammonia and ethylenediamine complexes. ActaChem. Scand.22, 2313-2323.

[1968ROB/WAL] Robie, R.A. and Waldbaum, D.R. (1968), Thermodynamic properties ofminerals and related substances at 298 K (25.0 C) and one atmosphere (1.013 bars)pressure and at higher temperatures. U.S. Geological Survey Bull. No. 1259, 256

P-

[1968SPI] Spinner, B. (1968), Etude quantitative de l'hydrolyse des niobates de potassium.Revue Chim. Min., 5, 839-868.

[1968SUS/KHO] Sushchevskaya, T.M. and Khodakovskiy, I.L. (1968), Mode of origin oftin minerals in hydrothermal deposits. Doklady Akademii Nauk SSSR, 181, 1476-1479.

[1969AWA/ELH] Awad, S.A. and Elhady, Z.A. (1969), Behaviour of lead as a metal-metalphosphate electrode and mechanism of its corrosion inhibition by phosphate ions.J. Electroanal. Chem., 20, 79-87.

[1969AWA/KAS] Awad, S.A. and Kassab, A. (1969), Behaviour of tin as metal-metalphosphate electrode and mechanism of promotion and inhibition of its corrosion byphosphate ions. J. Electroanal. Chem., 20, 203-212.

[1969BAR] Baranova, N.N. (1969), Investigation of the Carbonatocomplexes of lead at 25°and 200°C. Russ. J. Inorg. Chem., 14, 1716-1720.

[1969BON] Bond, A.M. (1969), An application of rapid polarographic techniques and thederivation of an equation for the polarographic study of the fluoride chmplexes ofbismuth (III) in acid media. Electroanal. Chem. Interfacial Electrochem., 23, 269-276.

319

JNC TN8400 9 9 - 0 1 1

[1969BRY/IOF] Bryukhanov, V.A., Iofa, B.Z. and Semenov, S.I. (1969), Investigation ofthe hydrolysis of antimony(V) in solutions of hydrochloric acid with the aid of theMossbauer effect. Soviet Radiochem., 11, 356-357.

[1969CAR] Carpentier, J.-M. (1969), Etude par electrophorese de la complexation del'uranium(VI), de l'etain(II) et du bismuth(III). Bulletin Soc. Chim. France, 11,3850-3855.

[1969CHA] Chartier, P. (1969), Oxydes de plomb: Solubilites en milieu alcalin (2e partie).Bulletin Soc. Chim. France, 7, 2253-2255.

[1969DYR/IVA] Dyrssen, E., Ivanova, K. and Oren, K. (1969), Solubility curves ofcalcium, strontium, and lead sulfates. Moscow Univ. Chem. Bulletin, 24, 32-35.

[1969GEL/KIS] Gel'fman, M.I. and Kiseleva, N.V. (1969), Stability Constants of chloro-complexes of Palladium (II). Russ. J. Inorg. Chem., 14, 258-261.

[1969HEL] Helgeson, H.C. (1969), Thermodynamics of hydrothermal systems at elevatedtemperatures and pressures. Am. J. Sci., 267, 729-804.

[1969JOH] Johansson, L. (1969), The complex formation of bismuth (III) with chloride inaqueous solution. A solubility study. Acta Chem. Scand., 23, 548-556.

[1969NEU] Neumann, G. (1969), A potentiometric study of the system Nb(V)-OH"-F- in3M K(C1) medium. Arkiv for Kemi, 32, 229-247.

[1969VAS/GLA] Vasil'ev, V.P. and Glavina, S.R. (1969), Thermodynarnic characteristicsof the Bi3+ ion in aqueous solution. Soviet Electrochem., 5, 374-378.

[1969VAS/GRE] Vasil'ev, V.P. and Grechina, N.K. (1969). Elektrokhimiya, 5, 426. Ascited in [1985LOV/MEK],

[1970BAR/KLI] Barsukov, V.L. and Klintsova, A.P. (1970), Solubility of cassiterite inwater and aqueous NaOH at 25°C. Geokhimiya, 10, 1268-1272.

[1970BON/HEF] Bond, A.M. and Hefter, G. (1970), Use of Ion-selective Electrodes in theEvaluation of stability constants of sparingly soluble salts. Application to the lead(Il)-fluoride system in aqueous solution. Inorg. Chem., 9, 1021-1023.

[1970BON/TAY] Bond, A.M. and Taylor, R.J. (1970), Polarographic studies of thefluoride complexes of tin (II) in neutral and acidic media. J. Electroanal. Chem.,28, 207-215.

[1970BON/WAU] Bond, A.M. and Waugh, A.B. (1970), ac polarography and itsapplication to overcome the problem of DC polarographic maxima in the study ofcomplexes ions. Electrochim. Acta, 15, 1471-1482.

320

JNC TN8400 9 9 - 0 1 1

[1970BON] Bond, A.M. (1970), Study of the fluoride complexes of antimony (III) in acidicmedia by rapid a-c polarography. Journal of Electrochemical Society 117, 1145-1151.

[1970DAW/WIL] Dawson, J.L., Wilkinson, J. and Gillibrand, M.I. (1970), Antimonyspecies in aqueous sulphuric acid solution. J. Inorg. Nucl. Chem., 32, 501-517.

[1970FED/SAM] Fedorov, V.A., Samsonova, N.P., Mironov, V.E. (1970), Binuclearhalogeno-and thiocyanato-complexes of lead (II), Russ. J. Inorg. Chem., 15,1325-1326.

[1970GAR/NAN] Gardner, G. and Nancollas, G.H. (1970), Complex Formation in leadsulfate solution. Anal. Chem., 42, 794-795.

[1970GOL/POL] Goleva, G.A., Polyakov, F.A. and Nechayeva, T.P. (1970), Distributionand migration of lead in ground waters. Geochemistry international 7, 256-268.

[1970KAN] Kankare, J.J. (1970), Computation of equilibrium constants for multicomponentsystems from spectrophotometric data. Anal. Chem., 42, 1322-1326.

[1970KUR/BAR] Kuril'chikova, G.E. and Barsukov, V.L. (1970), Stability ofhydroxystannate complexes and experimental crystallization of cassiterite underhydrothermal conditions. Geokhimiya 1, 35-42.

[1970LIT/NAN] Little, D.M.S. and Nancollas, G.H. (1970), Kinetics of crystallization anddissolution of lead sulphate in aqueous solution. Trans. Faraday Soc, 66, 3103-3111.

[1970NAB/KAL] Nabivanets, B.I. and Kalabina, L.V. (1970), State of palladium (II) inperchlorate solutions. Russ. J. Inorg. Chem., 15, 818-821.

[1971ADA/DOW] Adams, C.J. and Downs, A.J. (1971), Features of the co-ordinationchemistry of B-metals. Part I. Antimony (III) fluoride complexes. J. Chem. Soc.A, 1534-1542.

[1971BID] Bidleman, T.F. (1971), Bismuth-dithizone equilibria and hydrolysis of bismuthion in aqueous solution. Anal. Chim. Acta 56, 221-231.

[1971BON1] Bond, A.M. (1971), Some suggested calculation procedures and the variationin results obtained from different calculation methods for evaluation ofconcentration stability constants of metal ion complexes in aqueous solution.Coordin. Chem. Rev., 6, 377-405.

[1971BON2] Bond, A.M. (1971), Use of Rapid a.c. polarography for the evaluation ofcomplexes of sparingly soluble salts. Anal. Chim. Acta, 53, 159-167.

[1971FED/KAL1] Fedorov, V.A., Kalosh, T.N. and Mironov, V.E. (1971), Nitrato-complexes of tervalent bismuth. Russ. J. Inorg. Chem., 16, 539-542.

321

JNC TN8400 9 9 - 0 1 1

[1971FED/KAL2] Fedorov, V.A., Kalosh, T.N., Chernikova, G.E. and Mironov, V.E.(1971), Sulphato-complexes of bismuth(III). Russ. J. Phys. Chem., 45, 106.

[1971HEG/BAK] Hegedus, A.J., Bakcsy, G. and Chudik-Major, L. (1971), Thermo- undrontgenanalytische Untersuchung des Sb-0 systems im Bereich SbOj 5.2- ActaChimica Acad. Scient. Hungaricae, 77, 227-247.

[1971HIL/WOR] Hill, J.O., Worsley, I.G. and Hepler, L.G. (1971), Thermodynamic andoxidation potentials of vanadium, niobium, and tantalum. Chem. Rev., 71, 127-137.

[1971KUR/BAR] Kuril'chikova, G.Y. and Barsukov, V.L. (1971), Effects of CO2 and ofsodium and potassium bicarbonates and carbonates on the formation of Sn (IV)complexes in solution. Geokhimiya, 6, 642-653.

[1971MAB/KAL] Nabivanets, B.I., Kalabina, L.V. and Kudriskaya, L.N. (1991),Solubility of the hydroxides of palladium (II) and platinum (IV) and the ionic stateof the elements in perchlorate, chloride, and sulphate solutions. Russ. J. Inorg.Chem., 16, 1736-1738.

[1971NAB/KAL] Nabivanets, B.I., Kalabina, L.V. and Kudritskaya, L.N. (1971)Solubility of the hydroxides of palladium(II) and platinum(iV) and the ionic state ofthe elements in perchlorate, chloride, and sulphate solutions. Russ. J. Inorg.Chem., 16, 1736-1738.

[1971NAU/RYZ] Naumov, G.B., Ryzhenko, B.N. and Khodakovskiy, I.L. (1971),Handbook of thermodynamic data, Moscow: Atomizdat, in Russian; Engl. transl.:Report USGS-WRD-74-001 (Soleimani, G.J., Barnes.L, Speltz, V.,eds.), U.S.Geological Survey, Menlo Park, California, USA, 1974, 328p.

[1971NAZ/ANT] Nazarenko, V.A., Antonovich, V.P. and Nevskaya, E.M. (1971),Spectrophotometric determination of the hydrolysis constants of tin (IV) ions.Russ. J. Inorg. Chem., 15, 980-982.

[1971VAS/GLA] Vasil'ev, V.P. and Glavin, S.R. (1971), Thermodynamic properties of thePb^+ ion in aqueous solutions. Soviet Electrochem., 7, 1352.

[1972BON/HEF] Bond, A.M. and Hefter, G. (1972), Stability constant determination inprecipitating systems by rapid alternating current polarography. Electroanal. Chem.Interfacial Electrochem., 34, 227-237.

[1972BON] Bond, A.M (1972), The ac and dc polarographic reduction of bismuth (III) inacidic halide and other media. Electrochim. Acta, 17, 769-785.

[1972DRA/NIM1] Dragulescu, C , Nimara, A. and Julean, I. (1972), Contributions to thebismuth hydrolysis study. I. Spectrophotometric and polarographic investigationson bismuth perchlorate hydrolysis. Chemica Analyt., 17, 631-641.

322

JNC TN8400 9 9 - 0 1 1

[1972DRA/NIM2] Dragulescu, C , Nimara, A. and Julean, I. (1972), Contributions to thebismuth hydrolysis study. II. Spectrophotometric and polarographic investigationson bismuthyl perchlorate hydrolysis. Revue Roumaine Chimie 7, 1181-1190.

[1972ELD] Elding, L.I. (1972), Palladium (II) halide complexes. I. Stabilities and spectra ofpalladium (II) chloro and bromo aqua complexes. Inorg. Chim. Acta, 6, 647-651.

[1972FED/ROB] Fedorov, V.A., Robov, A.M., Grigor, T.I. and Mironov, V.E. (1972),Lead (II) nitrate complexes. Russ. J. Inorg. Chem., 17, 990-993.

[1972FED/SHI] Fedorov, V.A., Shishin, L.P., Likhacheva, S.G., Federova, A.V. andMironov, V.E. (1972), Influence of the ionic strength of the solution on theformation of the bromochloro-complex of lead (II). Russ. J. Inorg. Chem., 17, 41-43.

[1972HEF] Hefter, G. (1972), The use of ion-selective electrodes for the determination ofmixed stability constants. Electroanal. Chem. Inerfacial Electrochem., 39, 345-355.

[1972LAN/OBS] Land, J.E. and Osborne, C.V. (1972), The formation constants of theniobium fluoride system. Journal Less-Common Metals, 29, 147-153.

[1972NEK/LAD] Nekrasov, I.Y. and Ladze, T.P. (1972), Solubility of cassiterite in silicicchloride solutions at 300° and 400°. Doklady Akademii Nauk SSSR, 213, 933-936.

[1972NRI] Nriagu, J.O. (1972), Lead orthophosphates. I. Solubility and hydrolysis ofsecondary lead orthophosphate. Inorg. Chem., 11, 2499-2503.

[1972RYH] Ryhl, T. (1972), Thermodynamic properties of palladium (II) chloride andbromide complexes in aqueous solution. Acta Chem. Scand., 26, 2961-2962.

[1972VAS/SHO] Vasil'ev, V.P. and Shorokhova, V.I. (1972), Determination of thestandard thermodynamic characteristics of the antimonyl ion SbO+ and antimonyoxide by a potentiometric method. Soviet Electrochem., 8, 178-183.

[1972VIE] Vierling, F. (1972), No 644 - Interpretation plus elaboree des equilibres entre lesions Pb2+ et Cl-solubilite de PbCl2 a 25°C dans les solutions Na+(CiO4-,Cl-) 4M.Bull. Soc. Chim. France, 11, 4096-4099.

[1972ZIR/YAM] Zirino, A. and Yamamoto, S (1972), A pH-dependent model for thechemical speciation of copper, zinc, cadmium, and lead in seawater. LimnologyOceanography, 17, 661-671.

[1973BAR/KNA] Barin, I. and Knacke, O. (1973), Thermochemical properties of inorganicsubstances, Springer-Verlag, Berlin, 921 p.

[1973BEH/ROS] Behrens, R.G. and Rosenblatt, G.M. (1973), Vapor pressure andthermodynamics of orthorhombic antimony trioxide (valentinite). J. Chem.Thermodyn., 5, 173-188.

323

JNC TN8400 9 9 - 0 1 1

[1973BIL/STU] Bilinski, H. and Stumm, W. (1973), Pb (Il)-species in natural waters.EAWAG News, 1.

[1973BON/HEF] Bond, A.M. and Hefter, G. (1973), Influence of anion-induced adsorptionon half-wave potentials and other polarographic characteristics. Electroanal. Chem.Interfacial Electrochem., 42, 1-23.

[1973GAB/SRI] Gabe, D.R. and Sripatr, P. (1973), Anode bahaviour of tin during alkalinestannate plating. Trans. Inst. Metal Finishing, 51, 141-144.

[1973GOI/GRA] Goiffon, A., Granger, R., Bockel, C. and B., Spinner (1973), Etude desequilibres dans les solutions alcalines du niobium (V). Revue Chim. Min., 10, 487-502.

[1973GUL/SCH] Gulko, A. and Schmuckler, G. (1973), Accurat determination of thefourth stability constant of palladium (Il)-halide complexes. J. Inorg. Nucl. Chem.,35, 603-607.

[1973HUT/HIG] Hutchinson, M.H. and Higginson, W.C.E. (1973), Stability constants forassociation between bivalent cations and some univalent anions. J. C. S. Dalton,1247-1253.

[1973JOH/OHT] Johansson, G. and Ohtaki, H. (1973), An X-ray investigation of thehydrolysis products of tin(II) in solution. Acta Chem. Scand., 27, 643-660.

[1973KLI/BAR] Klintsova, A.P. and Varsukov, V.L. (1973), Solubility of cassiterite inwater and in aqueous NaOH solution at elevated temperatures. Geokhimiya, 5,701-709.

[1973NRI1] Nriagu, J.O. (1973), Lead orthophosphates-HI. Stabilities offluoropyromorphite and bromopyromorphite at 25°C. Geochim. Cosmochim. Acta,37, 1735-1734.

[1973NRI2] Nriagu, J. (1973), Lead orthophosphates-II. Stability of chloropyromorphite at25°C. Geochim. Cosmochim. Acta, 37, 367-377.

[1973VAS/GLA] Vasil'ev, V.P. and Glavina, S.R. (1973), Izv. Vys. Ucheb. Zaved.SSSR, Khim. i Khim. Tekhnol, 16, 39 (as cited in 1989COX/WAG).

[1973VAS/SHO1] Vasil'ev, V.P. and Shorokhova, V.I. (1973), Potentimetric investigationof alkaline solution of antimony (III). Soviet Electrochem., 9, 953-957.

[1973VAS/SHO2] Vasil'ev, V.P. and Shorokhova, V.I. (1973), Determination of thethermodynamic characteristics of antimony (III) in alkaline solutions by a solubilitymethod. Russ. J. Inorg. Chem., 18, 161-164.

324

JNC TN8400 9 9 - 0 1 1

[1974AHR/BOV] Ahrland, S. and Bovin, J.O. (1974), The complex formation ofantimony (III) in perchloric acid and nitric acid solutions. A solubility study. ActaChem. Scand., A 28, 1089-1100.

[1974BLA/BUR] Blandamer, M.J., Burgess, J. and Peacock, R.D. (1974), Solubility ofsodium hexahydroxoantimonate in water and in mixed aqueous solvents. J.C.S.Dalton, 1084-1086.

[1974FED/KAL] Fedorov, V.A., Kalosh, T.N. and Shmyd'ko, L.I (1974), Mixedchloronitratobismuth (III) complexes. Russ. J. Inorg. Chem., 19, 991-993.

[1974GOB] Gobom, S. (1974), The complex formation between tin (II) and acetate ions.Acta Chem. Scand., A 28, 1180-1182.

[1974GOI/SPI] Goiffon, A. and Spinner, B. (1974), Spectres Raman des solutionsaqueuses de niobates de potassium. Rev. Chim. Min., 11, 262-268.

[1974JAC/CHA] Jacob, K.T. and Chan, J.C. (1974), Electrochemical determination of thestability of mono- and dicalcium stannates. J. Electrochem. Soc, 121, 534-537.

[1974KOL/SHI] Kolonin, G.R., Shironosova, G.P. and Laptev, Y.V. (1974),Experimental checking of thermodynamic diagrams of the stability of W, Mo, andBi minerals under hydrothermal conditions. Inst. Geology Geophysics,Novosibirsk, USSR, 52, 161-167.

[1974MIL] Mills, K.C. (1974), Thermodynamic data for inorganic sulphides, selenides andtellurides, Butterworths, London, 845 pp.

[1974SHO/MAB] Shoji, H., Mabucchi, H. and Saito, N. (1974), Solvent extraction studiesof the hydrolysis of antimony (III) in tracer concentration. Bull. Chem. Soc. Japan,47, 2502-2507.

[1974VAD] Vadasdi, K. (1974), On determining the composition of species present in asystem from potentiometric data. J. Phys. Chem., 78, 816-820.

[1975ALY/ABD] Aly, H.G., Abdel-Rassoul, A.A. and Zakareia, N. (1975), Use ofzirconium phosphate for stability constant determination of uranium and antimonychlorocomplexes. Z. physik. Chemie, 94, 11-18.

[1975ANT/NEV] Antonovich, V.P., Nevskaya, E.M., Shelikhina, E.I. and V.A.,Nazarenko (1975), Spectrophotometric determination of the hydrolysis constants ofmonomeric bismuth ions. Russ. J. Inorg. Chem., 20, 1642-1645.

[1975BIE/ZIE] Biernat, J., Ziegler, B. and Zralko, M. (1975), Variability of somepolarographically determined stability constants ; A new kind of coordination: Thedual complexation. J. Electroanal. Chem., 63, 444-449.

325

JNC TN8400 9 9 - 0 1 1

[1975ERN/ALL] Ernst, R., Allen, H.E. and Mancy, K.H. (1975), Characterization of tracemetal species and measurement of trace metal stability constantes by electrochemicaltechniques. Water Research, 9, 969-979.

[1975FED/BOL] Fedorov, V.A., Bol'shakova, I.M. and Moskalenko, T.G. (1975),Formation of mixed bromo/chloro-complexes of tin(II) in aqueous solutions. Russ.J. Inorg. Chem., 20, 859-861.

[1975HEI/SCH] Hein, K. and Schulz, U. (1975), Untersuchungen zum elektrochemischenVerhalten von Antimon und Wismut in schwefelsauren flusssaurehaltigenElektrolyten. Neue Hutte, 20, 25-29.

[1975HEN/LON] Hentz, F.C. and Long, G.G. (1975), Synthesis, roperties, and hydrolysisof antimony trichloride. J. Chemical Education, 52, 189-190.

[1975KLI/BAR] Klintsova, A.P., Barsukov, V.L., Shemarykina, T.P. and Khodakovskiy,I.L. (1975), Measurement of the stability constants for Sn (IV) hydroxofluoridecomplexes. Geochem. Intern., 12, 207-215.

[1975KRA] Kragten, J. (1975), The complexometry of tin (IV). Talanta, 22, 505-510.

[1975NEL/AMI] Nelson, K.G. and Amin, K.N. (1975), Determination of stability constantsof stannous fluoride complexes by potentiostatic titration. J. Pharmaceutical ScL,64, 350-353.

[1975OLI] Olin, A. (1975), A thermochemical study of hydrolysed Bi(CIO4)3 solutions. ActaChem. Scand., A 29, 907-910.

[1976BAE/MES] Baes, C.F. and Mesmer, R.E. (1976), The Hydrolysis of Cations, KriegerPublishing, Malabar, USA, 499 pp.

[1976BIL/HUS] Bilinski, H., Huston, R. and Stumm, W. (1976), Determination of thestability constants of some hydroxo and carbonato complexes of Pb (II), Cu (II),Cd (II) and Zn (II) in dilute solutions by anodic stripping voltammetry anddifferential pulse polarography. Anal. Chim. Acta, 84, 157-164.

[1976GOB] Gobom, S. (1976), The hydrolysis of tin(II) ion. Acta Chem. Scand., A 30,745-750.

[1976HEM] Hem, J.D. (1976), Geochemical controls on lead concentrations in stream waterand sediments. Geochim. Cosmochim. Acta, 40, 599-609.

[1976LEE1] Lee, Y. (1976), The complex equilibria of Pb2+ with 3-Bromo-5-sulfosalicylateions. Acta Chem. Scand., A 30, 593-598.

[1976LEE2] Lee, Y. (1976), The pH equilibria of the 3-Bromo-5-sulfosalicylate ion inalkaline solution. Acta Chem. Scand., A 30, 586-592.

326

JNC TN8400 9 9 - 0 1 1

[1976NAS/LIN] Nasanen, R. and Lindell, E. (1976), Studies on lead (II) hydroxide salts.Part I. The solubility product of Pb(OH)Cl. Finnish Chemistry Lett., 95-98.

[1976NRI] Nriagu, J.O. (1976), Phosphate-clay mineral relations in soils and sediments.Can. J. Earth Sci., 13, 717-736.

[1976RAO] Rao, S.V.C. (1976), Physico-chemical studies of calcium-lead hydroxylapatites.Part III. X-ray, electron-microscopic and solubility equilibria data. Indian Chem.Soc, LIII, 352-354.

[1976SAM/LYA] Samoilenko, V.M., Lyashenko, V.I. and Poltoratskoya, T.V. (1976),Halogeno-and thiocyanato-complexes of tin (II) in protonic and aprotic donorsolvents. Russ. J. Inorg. Chem., 21, 1804-1807.

[1976SMI/MAR] Smith, R.M. and Martel, A.E. (1976), Critical Stability Constants. Vol. 4:Inorganic Complexes, Plenum Press, New York, 257 pp.

[1976TAR/GAR] Tardy, Y., Garrels, R.M. (1976), Prediction of Gibbs energies offormation: I. Relationships among Gibbs energies of formation of hydroxides,oxides and aqueous ions, Geochim. Cosmochim. Acta, 40, 1051-1056.

[1976VAS/KOK] Vasil'ev, V.P., Kokurin, N.I. and Vasil'eva, V.N. (1976), Enthalpy offormation of the Sn2+and SnCl+ions in aqueous solution. Russ. J. Inorg. Chem.,21, 218-221.

[1977ANT/NEV] Antonovich, V.P., Nevskaya, E.M. and Shelikhina, E.I. (1977),Calculation of the hydrolysis constants of monomeric tin(FV) and lead(IV) ions.Russ. J. Inorg. Chem., 22, 197-199.

[1977ANT/NEV] Antonovich, V.P., Nevskaya, E.M. and Suvorova, E.N. (1977),Spectrophotometric determination of the hydrolysis constants of monomericantimony(III) ions. Russ. J. Inorg. Chem., 22, 696-699.

[1977BAR/KNA] Barin, I., Knacke, O. and Kubaschewski, O. (1977), Thermochemicalproperties of inorganic substances (supplement), Springer-Verlag, Berlin, 861 pp.

[1977KEP/TAL] Kepak, F. and Talla, V. (1977), Self-diffusion of 125Sb(III) in aqueoussolutions. Coll. Czecholov. Chem. Commun., 42, 1472-1477.

[1977MAR] Mark, W. (1977), Hydrolysis of the tin (II) ion, Sn2+, in alkaline solution. ActaChem. Scand., A 31, 157-162.

[1977NAS/LIN] Nasasen, R. and E., Lindell (1977), Studies on lead (II) hydroxide salts.Part II. The solubility product of Pb(OH)Br. Finnish Chemistry Lett., 183-186.

[1977PAU] Paul, A. (1977), Chemical durability of glasses; a thermodynamic approach. J.IMaterials Sci., 12, 2246-2268.

327

JNC TN8400 9 9 - 0 1 1

[1977SHA] Shamsuddin, M. (1977), Thermodynamic studies on lead sulfide. MetallurgicalTransactions B, 8 B, 349-352.

[1977SIP/VAL] Sipos, L., Balenta, P., Nurnberg, H.W. and M., Branica (1977),Voltammetric determination of the stability constants of the predominant labile leadcomplexes in sea water. Lead Maritime Environ., 61-76.

[1977TAR/GAR] Tardy, Y. and Garrels, R.M. (1977), Prediction of Gibbs energies .offormation of compounds from the elements: II. Monovalent and divalent metalsilicates. Geochim.Cosmochim. Acta, 41, 87-92.

[1978CAH/VER] Cahen, H.T., Verkerk, M.J. and G.H.J., Broers (1978), Gibbs freeenergy of formation of Bi2O3 from emf cells with 8-B12O3 solid electrolyte.Electrochim. Acta, 23, 885-889.

[1978COD] CODATA (1978), CODATA Recommended Key Values for Thermodynamics.Committee on Data for Science and Technology (CODATA), Bulletin Nr. 28, Paris:International Council of Scientific Unions, 6p.

[1978DIC/LOT1] Dickinson, R. and Lotfi, S. (1978), The nature and standard potential ofthe stannite ion in sodium hydroxide solution. Electrochim. Acta, 23, 995-999.

[1978DIC/LOT2] Dickinson, T. and Lotfi, S. (1978), The anodic dissolution of tin insodium hydroxide solution. Electrochim. Acta, 23, 513-519.

[1978FAT/ROU] Fatouros, N., Rouelle, F. and Chemla, M. (1978), Influence de laformation de complexes chlorures sur la reduction electrochimique de SnIV enmilieu perchloroique acide. J. Chim. Phys. Physico-Chimie Biol., 75, 477-483.

[1978LIN] Lind, C.J. (1978), Polarographc determination of lead hydroxide formationconstants at low ionic strength. Environ. Sci. & Tech., 12, 1406-1410.

[1978ROB/HEM2] Robie, R.A., Hemingway, B.S. and Fisher, J.R. (1978),Thermodynamic properties of minerals and related substances at 298.15 K and 1bar pressure and at higher temperatures. US Geological Survey Bulletin, 1452, 456P-

[1978TOP/EVD] Toptygina, G.M., Evdokimov, V.I., Eliseeva, N.A. and Badanin, V.S.(1978), Precipitation of tin(IV) from hydrochloric acid solutions by calciumhydroxide. Russ. J. Inorg. Chem., 23, 810-813.

[1979GIO/BAR] Giordano, T.H. and Barnes, H.I. (1979), Ore solution Chemistry. VI.PbS solubility in bisulfide solutions to 300°C. Economic Geology, 74, 1637-1646.

[1979KUB/ALC] Kubaschewski, O. and Alcock, C.B. (1979), Metallurgicalthermochemistry, Pergamon Press, Oxford, 449 pp.

328

JNC TN8400 9 9 - 0 1 1

[1979PAT/OBR] Patterson, J.W. and O'Brien, J.E. (1979), Control of lead corrosion. J.AWWA, 264-271.

[1979VAS/GLA] Vasil'ev, V.P., Glavina, S.R. and Tschorochova, V.I (1979),Potentiometric determination of energy of formation of tin (IV) ion in an aqueoussolution (in Russian). Izv. Vyssh. Ucheb. Zaved. SSSR, Khim. i Khim. Tekhnol.,22, 1083-1085.

[1980AND/SAM] Andreev, A.I., Samsonova, N.P., Robov, A.M. and Fedorov, V.A.(1980), Formation of nitrate complexes of tin (II) in aqueous solutions.Koordinationsonnaya Khimiya, 5, 1325-1327.

[1980BEN/TEA] Benson, L.Vand Teague, L.S. (1980), A tabulation of thermodynamic datafor chemical reactions involving 58 elements common to radioactive waste packagesystems, Lawrence Berkeley Laboratory, Berkeley, USA, Report, LBL-11448,97p.

[1980BON/HEF] Bond, A.M. and Hefter, G.T.. (1980), Critical survey of stabilityconstants and related thermodynamic data of fluoride complexes in aqueoussolution. IUPAC Chemical Data Series, 27, Pergamon Press, Oxford, 67 p.

[1980CLE/JOH] Clever, H.L. and Johnston, F.J. (1980), The solubility of some sparinglysoluble lead salts: An evaluation of the solubility in water and aqueous electrolytesolution. J. Phys. Chem. Ref. Data, 9, 751-784.

[1980KAW/ISH] Kawai, T., Ishiguro, S. and Ohtaki, H. (1980), A termodynamic study onhydrolytic reactions of lead (II) ion in an aqueous solution and dioxane-watermixtures. I. A potentiometric study. Chem. Soc. Japan, 53, 2221-2227.

[1980KRA] Kragten, J. (1980), An evaluation of the stability constants of the chloro-complexes of palladium(II). Talanta, 27, 375-377.

[1980LOV/BRA] Lovric, M. and M., Branica (1980), Application of ASV for trace metalspeciation IV: Determination of lead-chloride stability constants by rotating mercurycoated glassy carbon electrode. Croatica Chem. Acta, 53, 503-508.

[1980MAN/DEU] Mann, A.W. and Deutscher, R.L. (1980), Solution geochemistry of leadand zinc in water containing carbonate, sulphate and chloride ions. ChemicalGeology, 29, 293-311.

[1980PAT] Patterson, J.W. (1980), Effect of carbonate ion on precipitation treatment ofcadmium, copper, lead and zinc. Proc. Industrial Waste Conference, 36, 579-602.

[1980PRA/PRA] Prassad, A.K. and B., Prassad (1980), Dissociation of PbCl+ in aqueoussolutions and related thermodynamic quantities. Indian Chem. Soc., LVII, 155-159.

329

INC TN8400 9 9 - 0 1 1

[1980SCH] Schock, M.R. (1980), Response of lead solubility to dissolved carbonate indrinking water. J. AWWA, 72, 695-704.

[1980SYL/BRO] Sylva, R.N. and Brown, P.L. (1980), The hydrolysis of metal ions. Part3. Lead (II). J. Chem. Soc, 7, 1577-1581.

[1980WAL/SIN] Wallace, J.R. and Singer, P.C. (1980), Precipitation of lead from a storagebattery manufacturing wastewater. Proc. Industrial Waste Conf., 35, 702-717.

[1981BAE/MES] Baes, C.F. and Mesmer, R.E. (1981), The thermodynamics of cationhydrolysis. Am. J. Sci., 281, 935-962.

[1981BYR/YOU] Byrne, R.H., Young, R.W. and Miller, W.L. (1981), Lead chloridecomplexation using ultraviolet molar absortivity characteristics. J. Solution Chem.,10, 243-251.

[1981BYR] Byrne, R.H. (1981), Inorganic lead complexation in natrual seawater determinedby UV spectroscopy. Nature, 290, 487-489.

[1981CIA] Ciavatta, L. (1981), The specific interaction theory in evaluating ionic equilibria.Annali di Chimica, 70, 551-567.

[1981DAD/SOR] Dadze, T.P., Sorokhin, V.I. and Nekrasov, I.Y. (1981), Solubility ofSnO2 in water and in aqueous solutions of HC1, HC1+KC1, and HNO3 at 200-400°C and 101.3 MPa. Geokhimiya, 10, 1482-1492.

[1981GOR/GAV] Dorbunov, V.E., Gavrchev, K.S., Sarakhov, O.A. and Lazarev, V.B.(1981), Thermodynamic functions of Bi2O3 in the temperature range 11-298 K.Russ. J. Inorg. Chem., 26, 546-547.

[1981HAA/WIL] Haacke, D.F. and Williams, P.A. (1981), Stability of plumbonacrite. J.Inorg. Nucl. Chem., 43, 406.

[1981HEL/KIR] Helgeson, H.C., Kirkham, D.H. and Flowers, G.C. (1981), Theoreticalprediction of the thermodynamic behavior of aqueous electrolytes at high pressuresand temperatures: IV. Calculation of activity coefficients, osmotic coefficients, andapparent molal and standard and relative partial molal properties to 600 C and 5 kb.Am. J. Sci. 281, 1249-1516.

[1981ISH/OHT] Ishiguro, S. and Ohtaki, H. (1981), A thermodynamic study on hydrolyticreactions of lead(II) ions in an aqueous solution and dioxane-water mixtures. II. Acalorimetric study. Bull. Chem. Soc. Japan 54, 335-342.

[1981JOH/PAP] Johnson, G.K., Papatheodorou, G.N., Johnson, C.E. (1981)Theenthalpies of formation of SbFsQ) and Sb2S3(c) and the high-temperaturethermodynamic functions of Sb2S3(c) and Sb2S3(l), J. Chem. Thermodyn., 13,745-754.

330

iwr TNR400 9 9 - 0 1 1

[1981KAL/BOR] Kalinchenko, F.V., Borzenkova, M.P. and Novoselova, A.V. (1981),The Bi2O3-BiF3 System. Russ. J. Inorg. Chem., 26, 118-120.

[1981KOG/OKA] Kogure, K., Okatmato, M. and Kakihana, H. (1981), Solvent deuteriumisotope effect on hydrolysis of Pb(II). J. Inorg. Nucl. Chem., 43, 1561-1564.

[1981LIN/BES] Lindemer, T.B., Besmann, T.M. and Johnson, C.E. (1981),Thermodynamic review and calculations - alkali-metal oxide systems with nuclearfuels, fission products, and structural materials. J. Nuclear Mat., 100, 178-226.

[1981PET/MIL] Pettine, M., Millero, FJ. and Macchi, G. (1981), Hydrolysis of tin (II) inaqueous solutions. Anal. Chem., 53, 1039-1043.

[1981SHA/MIS] Shamsuddin, M. and Misra, S. (1981), Thermodynamic investigations oflead chalcogenides. Chem. Metallurgy, 241-256.

[1981STU/MOR] Stumm, W. and Morgan, J.J. (1981), Aquatic Chemistry. An introductionemphasizing chemical equilibria in natural waters, John Wiley and Sons, NewYork, 780 pp.

[1981TUR/WHI] Turner, D.R., Whitfield, M. and Dickson, A.G. (1981), The equilibriumspeciation of dissolved components in freshwater and seawater at 25 C and 1 atmpressure. Geochim. Cosmochim. Acta, 45, 855-881.

[1982BEL/KOL] Belevantsev, V.I., Kolonin, G.R., Peshchevitskiy, B.I. (1982), Use ofstepwise ligand replacement in estimating the stability constants of mixedcomplexes in geochemically important systems, Geochem. International, 19(1),169-179.

[1982BEN/MEU] Bendiab, H., Meullemeestre, J., Schwing, M-J. and Vierling, F. (1982),Thermodynamic constants and electronic spectra of lead (II) chloro-complexes inaqueous solutions. J. Chem. Research, 280-281.

[1982BIL/SCH] Bilinski, H. and Schindler, P. (1982), Solubility and equilibrium constantsof lead in carbonate solutions (25°C, / = 0.3 mol dm"3). Geochim. Cosmochim.Acta, 46, 921-928.

[1982HAT/SUG] Hataye, I. , Suganuma, H., Ikegami, H. and Kuchiki, T. (1982), Solventextraction study on th hydrolysis of tracer concentrations of Bismuth (III) inperchlorate and nitrate solutions. Chem. Soc. Japan, 55, 1475-1479.

[1982LAP/KOL] Laptev, Y.V. and Kolonin, G.R. (1982), Hydrolysis of bismuth (III) inhigh-temperature solutions. Russ. J. Inorg. Chem., 27, 2515-2520.

[1982LEV/NAR] Levitski, V.A., Narchuk, P.B., Kovba, M.L. and Y.Y., Skolis (1982),Solid electrolytes in thermodynamic studies. The thermodynamic properties ofPdO. Russ. J. Phys. Chem., 56, 1474-1479.

331

INC TN8400 9 9 - 0 1 1

[1982PAN] Pankratz, L.B. (1982), Thermodynamic properties of elements and oxides. USBureau Mines Bulletin, 672, 509 p.

[1982PAU] Paul, A (1982), Chemistry of Glasses, Chapman and Hall, London, 292 pp.

[1982ROH] Rohl, R. (1982), Trace metal speciation in sea water - a paper electrophoreticapproach. Anal. Chim. Acta, 135, 99-110.

[1982SMI/MAR] Smith, R.M. and Martel, A.E. (1982), Critical Stability Constants. Vol. 5:First Supplement, Plenum Press, New York, 604 pp.

[1982WAG/EVA] Wagman, D.D., Evans, W.H., Parker, V.B., Schumm, R.H., Halow,I., Bailey, S.M., Churney, K.L. and Nuttall, R.L. (1982), The NBS tables ofchemical thermodynamic properties: Selected values for inorganic and Cl and C2organic substances in SI units. J. Phys. Chem. Ref. Data 11, suppl. No. 2, 1-392.

[1983LAN] Langmuir, D. (1983), Private communication, Colorado School of Mines,Golden, CO (as cited in NEA database).

[1983MAL/SRE] Mallika, C , Sreedharan, O.M. and Gnanamoorthy, J.B. (1983),Determination of the standard free energy of formation of PdO(s) from the solidoxide electrolyte emf. J. Less-Common Metals, 95, 213-220.

[1983POP] Pope, M.T. (1983), Heteropoly and isopoly oxometalates, Springer-Verlag,Berlin, 180 pp.

[1983SAN/BAR] Sangameshwar, S.R. and Barnes, H.L. (1983), Supergene processes inzinc-lead-silver sulfide ores in carbonates. Economic Geology, 78, 1379-1397.

[1983SCH/GAR] Schock, M.R. and Gardels, M.C. (1983), Plumbosolvency reduction byhigh pH and low carbonate-solubility relationships. J. AWWA, 75, 87-91.

[1984BAN] Bannister, M.J. (1984), The standard molar Gibbs free energy of formation ofPbO. Oxygen concentration-cell measurements at low temperatures. J. Chem.Thermodyn., 16, 787-792.

[1984BER/BRE] Berry, F.J. and Brett, M.E. (1984), Studies of antimony oxides formed bydehydration of antimony supensions in nitric acid. Inorg. Chim. Acta, 83, 167-169.

[1984BYR/MIL] Byrne, R.H. and Miller, W.L. (1984), Medium composition dependenceof lead (II) complexation by chloride ion. Am. J. Sci., 284, 79-94.

[1984HOU/KEL] House, C.I. and Kelsall, G.H. (1984), Potential-pH diagrams for thesn/H2O-Cl system. Electrochim. Acta, 29, 1459-1464.

332

JNC TN8400 9 9 - 0 1 1

[1984JON] Jones, W.M. (1984), Equilibrium pressures over the systems bismuth trisulfate-dibismuthmonoxydisulfate and dibismuthmonoxydisulfate-dibismuthdioxymono-sulfate. Slow transformation between two crystalline forms of dibismuthmonoxydi-sulfate. J. of Chem. Phys., 80, 3408-3419.

[1984KOC/TAP] Kochetkova, N.V., Toptygina, G.M., Karpov, I.K. and Evdokimov, V.I.(1984), Thermodynamic analysis of equilibria in the SnS2-CaCl2-H2O system.Russ. J. Inorg. Chem., 29, 460-463.

[1984LOZ/SCH] Lozar, J., Schuffenecker, L., Cudey, G. and Bourdet, J.B. (1984),Determination a 25° des proprietes thermodynamiques des complexes chlores duplomb divalent a partir de mesures de solubilite dans des solutions aqueuses deforce ionique inferieure a 1 mol kg"1. Thermochim. Acta, 79, 171-186.

[1984MIL/BUG] Milic, N.B. and Bugarcic, Z. (1984), Hydrolysis of the palladium(II) ionin a sodium chloride medium. Transition Metal Chem., 9, 163-173.

[1984MIL/BYR] Millero, F.J. and Pyrne, R.H. (1984), Use of Pitzer's equations todetermine the media effect on the formation of lead chloro complexes. Geochim.Cosmochim. Acta, 48, 1145-1150.

[1984NRI2] Nriagu, J.O. (1984), Chapter 10: Formation and stability of base metalphosphates in soils and sediments, In: Phosphate Minerals (J.O. Nriagu and P.B.Moore Ed.), Springer-Verlag, 318-329.

[1984PIN/GAL] Pingarron Carrazon, J.M., Gallego Andreu, R. and Sanchez Batanero, P.(1984), Determination potentiometrique de constantes de stabilite de complexes dubismuth (III). Analysis, 12, 358-363.

[1984SEW] Seward, T.M. (1984), The formation of lead (II) chloride complexes to 300°C: Aspectrophotometric study. Geochim. Cosmochim. Acta, 48, 121-134.

[1984SVE] Sverjensky, D.A. (1984), Prediction of Gibb's free energies of calcite-typecarbonates and the equilibrium distribution of trace elements between carbonatesand aqueous solutions. Geochim. Cosmochim. Acta, 48, 1127-1134.

[1984TAY/LOP] Taylor, P. and Lopata, V.J. (1984), Stability and solubitity relationsphipbetween some solids in the system PbO-CO2-H2O. Can. J. Chem., 62, 395-402.

[1984TAY/SUN] Taylor, P., Sunder, S. and V.J., Lopta (1984), Structure, spectra, andstabitity of solid bismuth carbonates. J. Am. Chem. Soc, 82, 2863-2873.

[1984UHL/HEI] Uhler, A.D. and Helz, G.R. (1984), Solubility product of galena at 298 K:A possible explanation for apparent supersaturation in nature. Geochim.Cosmochim. Acta, 48, 1155-1160.

333

J N C T N 8 4 0 0 9 9 - O i l

[1984VAS/SHO] Vasil'ev, V.P., Shorokhova, V.I., Raskova, O.G. and Klepikova, L.I.(1984), Thermodynamic properties of antimony (III) in hydrochloric acid solutions.Russ. J. Phys. Chem., 58, 1640-1642.

[1984VIE/TAR] Viellard, P. and Tardy, Y. (1984), Chapter 4: Thermochemical Propertiesof Phosphates, In: Phosphate Minerals (J.O. Nriagu and P.B. Moore Ed.),Springer-Verlag, 171-198.

[1984WAN] Wanner, H. (1984) Bildung und Hydrolyse von Palladium-Komplexen mitPyridin, 2,2'-Bipyridil and 1,10-Phenantrolin. Ph.D. Dissertation, ETH Zurich,Switzerland.

[1985BAB/MAT] Babushkin, V.I., Matveyev, G.M. and Mchedlov-Petrossian, O.P.(1985), Thermodynamics of Silicates, Springer-Verlag, Berlin, 459 pp.

[1985CHA/DAV] Chase, M.W., Davies, C.A., Downey, J.R., Frurip, D.J., McDonald,R.A. and Syverud, A.N. (1985), JANAF thermochemical tables, 3rd ed. J. Phys.Chem. Ref. Data 14, supl. 1,

[1985COL] Collom, F. (1985), Palladium and platinum, In: Standard potentials in aqueoussolution (A.J.Bard, R. Parsons and J. Jordan Ed.), Marcel Dekker, New York,339 - 367.

[1985GAL] Galus, Z. (1985), Carbon, silicon, germanium, tin and lead, In: Standardpotentials in aqueous solution (A.J.Bard, R. Parsons and J. Jordan Ed.), MarcelDekker, New York, 189-236.

[1985JAC/HEL2] Jackson, K.J. and Helgeson, H.C. (1985), Chemical and thermodynamicconstraints on the hydrothermal transport and deposition of tin: II. Interprtetation ofphase relations in the Southeast Asian tin belt. Economic Geology, 80, 1365-1378.

[1985LOV/MEK] Lovrecek, B., Mekjavic, I. and Metikos-Hukovic, M. (1985), V.Bismuth, In: Standard potentials in aqueous solution (A.J.Bard, R. Parsons and J.Jordan Ed.), Marcel Dekker, New York, 180-187.

[1985MUL] Muller, A.B., Fritz, B., Wyman, A. and Snellman, M. (1985), Chemicalthermodynamic data set for minerals associated with granite. Mat. Res. Soc. Symp.Proc 50, 165-175.

[1985PAS] Past, V. (1985), IV. Antimony, In: Standard potentials in aqueous solution(AJ.Bard, R. Parsons and J. Jordan Ed.), Marcel Dekker, New York, 172-179.

[1985RAI/RYA] Rai, D., Ryan, J.L. (1985), Neptunium(IV) hydrous oxide solubility underreducing and carbonate conditions, Inorg. Chem., 24, 247-251.

[1985SED/SIM] Sedova, A.A., Simonov, L.N. and Mel'chakova, N.V. (1985), Thedetermination of the hydrolysis constants of monomeric bismuth (III) ions by the

334

J N C T N 8 4 0 0 99 - O i l

competing reaction involving the formation of its complex with bismuthiol-II.Russ. J. Inorg. Chem., 30, 805-807.

[1985UDU/VEN] Udupa, H.V.K., Venkatesan, V.K. and Krishan, M. (1985), Niobiumand tantalum, In: Standard potentials in aqueous solution (A.J.Bard, R. Parsonsand J. Jordan Ed.), Marcel Dekker, New York, 526-537.

[1986AZA/PAN] Azad, A.M., Pankajavalli, R. and Sreedharan, O.M. (1986),Thermodiynamic stability of Sb2O3 by a solid-oxide electrolyte e.m.f. technique. J.Chem. Thermodyn., 18, 255-261.

[1986DAD/SOR] Dadze, T.P. and V.I., Sorokin (1986), Solubility of SnO2 in water at 200°to 400°C and 1.6 to 150 MPa. Doklady Akademii Nauk SSSR, 286, 426-428.

[1986FLE/JOH] Fleer, V.N. and Johnston, R.M. (1986), A compilation of solubility anddissolution kinetics data on minerals in granitic and gabbroic systems, AtomicEnergy Canada Ltd, Pinawa, Canada, TR-328-2, 170p.

[1986ITA/NIS] Itagaki, K. and Nishimura, T. (1986), Thermodynamic properties ofcompounds and aqueous species of VA elements. Metallurgical Rev. MMIJ, 3, 29-48.

[1986KOV/RYZ] Kovalenko, N.I., Ryzhenko, B.N., Barsukov, V.L., Klintsova, A.P.,Velyukhanova, T.K., Volynets, M.P. and Kitayeva, L.P. (1986), The solubility ofcassiterite in HC1 and HC1 + NaCl (KC1) solutions at 500°C and 1000 atm underfixed redox conditions. Geokhimiya, 2, 190-205.

[1986MYE] Myers, R.J. (1986), The new low value for the second dissociation constant forH2S. J. Chem. Ed., 63, 687-690.

[1987BRO/WAN] Brown, P.L. and Wanner, H. (1987), Predicted formation constantsusing the unified theory of metal ion complexation, OECD Nuclear Energy Agency,Paris, 102 pp.

[1987BRU] Brubaker, K.L. (1987), Solid and aqueous species of significance for the sitingof a high level nuclear waste repository in rock salt: an evaluation of data needs forlead, tin, strontium, and selenium, Argonne National Laboratoy, Argonne, Illinois,Report, ANL/EES-TM-336.

[1987DUB/RAM] Dubessy, J. and al., et (1987), Physical and chemical controls (fO2, T,pH) of the opposite behaviour of U and Sn-W as examplified by hydrothermaldesposits in France and Great-Britain, and solubility data. Bulletin Mineral., 110,261-281.

[1987FER/GRE] Ferri, D., Grenthe, I., Hietanen, S. and Salvatore, F. (1987), Studies onmetal carbonate equilibria. 18. Lead (II) carbonate complexes in alkaline solutions.Acta Chem. Scand., A 41, 349-354.

335

JNC TN8400 9 9 - 0 1 1

[1987MIL/ROE] Milanov, M , Roesch, F., Khalkin, V.A., Henniger, U. and Hung, T.K.(1987), Electromigration of ions of radionuclides without carriers in electrolytes.Hydrolysis of Bi(IIII) in aqueous solutions. Soviet Radiochem., 29, 18-25.

[1987PAN/SRE] Pankajavalli, R. and Sreedharan, O.M. (1987), Thermodynamic stabilityof Sb2O4 by a solid oxide electrolyte e.m.f. method. J. Mat. Sci., 22, 177-180.

[1987SUG/SHI] Suganuma, H., Shimizu, I. and Hataye, I. (1987), A solvent-extractionstudy of the hydrolysis of tracer concentrations of bismuth (III) in chloridesolutions. Chem. Soc. Japan, 60, 877-883.

[1988BAR/SHA] Barybin, V.I., Sharygin, P.M., Gonchar, V.F. and Moiseev, V.E.(1988), Solubility of hydrated tin dioxide in water. Izvestiya Akademii NaukSSSR, Neoganicheskie Materialy, 23, 1162-1165.

[1988BYR/KUM] Byrne, R.H., Kump, L.R. and Cantrell, K.J. (1988), The influence oftemperature and pH on trace metal speciation in seawater. Marine Chem., 25, 163-181.

[1988KRU] Krupp, R.E. (1988), Solubility of stibnite in hydrogen sulfide solutions,speciation, and equilibrium constants, from 25 to 350°C. Geochim. Cosmochim.Acta, 52, 3005-3015.

[1988MOU/WOO] Mountain, B.W. and Wood, S.A. (1988), Chemical controls on thesolubility, transport, and deposition of platinum and palladium in hydrothermalsolutions: A thermodynamic approach. Economic Geology, 83, 492-510.

[1988PHI/HAL] Phillips, S.L., Hale, F.V., Silvester, L.F. and Siegel, M.D. (1988),Thermodynamic tables for nuclear waste isolation, an aqueous solutions database,Vol. 1, Lawrence Berkeley Laboratory, Berkeley, California, USA, ReportNUREG/CR-4864, LBL-22860, SAND87-0323.

[1988VOV/PER] Vovk, S.M., Perekhozheva, T.N. and Sharygin, L.M. (1988), Sorption ofcalcium on hydrated tin dioxide. Russ. J. Phys. Chem., 62, 1601-1602.

[1989BAY/EWA] Bayliss, S., Ewart, F.T., Howse, R.M., Lane, S.A., Pilkington, N.J.,Smith-Briggs, J.L. and Williams, S.J. (1989), The solubility and sorption ofradium and tin in a cementious near-field environment, Mat. Res. Soc. Symp.Proc. Mat. Res. Soc. Symp. Proc, 127, 879-885.

[1989COX/WAG] Cox, J.D., Wagman, D.D. and Medvedev, V.A.. (1989), CODATA KeyValues for Thermodynamics, Hemisphere Publishing Corp., New York, 271 pp.

[1989DOR/MAR] Dorange, G., Marchand, A. and Franco, A. (1989), Solubilite de lacerrusite et constantes de stabilite de PbOH+, PbCO3° et PbHCO3+. Tribune del'eau, 42, 53-59.

336

JNC TN8400 9 9 - 0 1 1

[1989KEL] Kellogg, H.H. (1989), Critical evaluation of the thermochemical properties oflead sulfates. Metallurgical Trans. B, 20B, 77-85.

[1989KUM/BYR] Kump, L.R. and Byrne, R.H. (1989), Palladium chemistry in seawater.Environ. Sci. Tech., 23, 663-665.

[1989NYH/WIK] Nyholm, L. and Wikmark, G. (1989), Precise, polarographicdetermination of the stability constants of cadmium and lead with oxalate andsulphate. Anal. Chim. Acta, 223, 429-440.

[1989SMI/MAR] Smith, R.M. and Martel, A.E. (1989), Critical Stability Constants. Vol. 6:Second Supplement, Plenum Press, New York, 643 pp.

[1989VAN/POU] Vandenborre, M.T., Poumellec, B. and Livage, J. (1989), Etude EXAFSde la formation d'oxyde Nb2O5 par hydrolyse-condensation de chloroethoxyde deniobium. J. Solid State Chem., 83, 105-114.

[1989WOO/MOU] Wood, S.A. and Mountain, B. (1989), Thermodynamic constraints onthe solubilty of platinium and palladium in hydrothermal solutions: reassessment ofhydroxide, bisullfide, and ammonia complexing. Economic Geology, 84, 2020-2028.

[1990PIL/STO] Pilkington, N.J. and Stone, N.S. (1990), The solubility and sorption ofnickel and niobium under high pH conditions, Harwell Laboartory, UKAEA,Report, NSS/R186.

[1990SHI/ZOT] Shikina, N.D. and Zotov, A.V. (1990), Thermodynamic parameters ofSb(OH)3°(sol) up to 723.15 K and 100 bar. Geochem. Intern., 28, 97-103.

[1990SUG/ONO1] Suganuma, H., Ono, K. and Hataye, I. (1990), A cation-exchangestudy of stability constants of complexes formed between bismuth (III) and nitrateor chloride ions. J. Radioanal. Nuclear Chem., 145, 167-173.

[1990SUG/ONO2] Suganuma, H., Ono, K. and Hataye, I. (1990), A cation-exchangestudy of tracer concentrations of bismuth (III) in perchlorate solutions. Radiochim.Acta, 51, 5-10.

[1991FED/KAL] Fedorov, V.A., Kalosh, T.N. and Chernikova, G.E. (1971), Sulphato-complexes of Bismuth (III). Russ. J. Phys. Chem., 45, 106.

[1991MGAS/FUE] Mgaidi, A., Fiirst, W. and Renon, H. (1991), Representation of thesolubility of lead chloride invrious chloride solutions with Pitzer's model.Metallurgical Transactions B, 22B, 491-498.

[1991TAI/JAN] Tait, CD., Janecky, D.R. and S.Z., Rogers (1991), Speciation of aqueouspalladium (II) chloride solutions using optical spectroscopies. Geochim.Cosmochim. Acta, 55, 1253-1264.

337

JNC TN8400 9 9 - 0 1 1

[1991WOO] Wood, S.A. (1991), Experimental determination of the hydrolysis constants ofPt2+ and Pd2+ at 25°C from the solubility of Pt and Pd in aqueous hydroxidesolutions. Geochim. Cosmochim. Acta, 55, 1759-1767.

[1992GRE/FUG] Grenthe, I., Fuger, J., Konings, R.J.M., Lemire, R.J., Muller, A.B.,Nguyen-Trung, C. and Wanner, H. (1992), Chemical Thermodynamics. Vol. 1:Chemical Thermodynamics of Uranium, Elsevier, Amsterdam, 714 pp.

[1992LOZ/SCH] Lozar, J., Schuffenecker, L. and Moliner, J. (1992), Determination desproprietes electrochimiques et thermodynamiques de Pb2+, PbCl+ et PbCl2 a partirde mesures de conductivite de solutions aqueuses de chlorure de plomb a 25°C.Electrochim. Acta, 37, 2519-2522.

[1992PEA/BER] Pearson, F.J. Jr., Berner, U. and Hummel, W. (1992), NAGRAthermochemical database. EL Supplemental data, NAGRA, Wettingen, Switzerland,Technical Report 91-18.

[1992RAG] Raghavan, S. (1992), Thermodynamics of formation of high calcium niobatesfrom e.m.f. measurements. J. Alloys Compounds, 179, L25-L27.

[1992SEA] Seal, R.R. (1992), Superambient heat capacities of synthetic stibnite, berthiertite,and chalcostibite: Revised thermodynamic properties and implications for phaseequilibria. Economic Geology, 87, 1911-1918.

[1992WOO/MOU] Wood, S.A., Mountain, B. and Pan, P. (1992), The aqueous chemistryof platinum, palladium and gold: recent experimental constraints and a re-evaluationof theoretical predictions. J. Mineral. Assoc. Canada, 30, 955-982.

[1992YAJ/TOB] Yajima, T., Tobita, S. and Ueta, S. (1992), Solubility measurements ofniobium in the system Nb-O-H under CO2-free condition. Abstract 1992 FallMeeting Atomic Energy Society Japan 341 (in Japanese).

[1993CRU/VAN] Cruywagen, J.J. And van de Water, R.F. (1993), The hydrolysis oflead(II). A potentiometric and enthalpometric study. Talanta, 40, 1091-1095.

[1993DEL/MIL] De Lisi, R., Milioto, S., Alonzo, G. and Saiano, F. (1993),Thermodynamic and 19F NMR studies of antimony trifluoride in water. J. SolutionChem., 22, 489-505.

[1993BYR/KUM] Byrne, R.H., Kump, L.R. (1993), Comment on "Speciation of aqueouspalladium(II) chloride solutions using optical spectroscopies" by C. D. Tait, D. R.Janecky, and P. S. Z. Rogers, Geochim. Cosmochim. Acta, 57, 1151-1156.

[1993KRA/DEC] Kragten, J. and Decnop-Weever, L.G. (1993), Mixed hydroxide complexformation and solubility of bismuth in nitrate and perchlorate medium. Talanta, 40,485-490.

338

JNC TN8400 9 9 - 0 1 1

[1993KUL/HAK] Kulmala, S. and Hakanen, M. (1993), The solubility of Zr, Nb and Ni ingroundwater and concrete water, and sorption on crushed rock and cement, YJTReport, Helsinki, YJT-93-21.

[1993MAC/PAG] Macchi, G., Pagano, M., Santori, M. and Tiravanti (1993), Batteryindustry wastewater: Pb removal and produced sludge. Water Research, 27, 1511-1518.

[1994AKI/ZOT] Akinifiyev, N.N., Zotov, A.V. and Shikina, N.D. (1994), Experimentalstudies and self-consistent thermodynamic data in the Sb(III)-S(II)-O-H system.Geochem. Intern., 31, 27-40.

[1994ETX/FER] Etxebarria, N., Fernandes, L.A. and Madariaga, J.M. (1994), On thehydrolysis of niobium (v) and tantalum (v) in 3 mol dm-3 KC1 at 25°C. Part 1.Construction of a thermodynamic model for Nbv . J. Chem. Soc. Dalton Trans.3055-3059.

[1994YAJ] Yajima, T. (1994), Solubility measurements of uranium and niobium, NuclearEngineering Research Laboratory, Faculty of Engineering, University of Tokyo,Yayoi Kenkyukai Report, UTNL-R 0331, pp. 127-144 (in Japanese).

[1995DJU/JEL] Djurdjevic, P., Jelic, R., Djokic, D. and Veselinovic, D. (1995),Hydrolysis of tin(II) ion in sodium chloride medium. J. Serb. Chem. Soc. 60, 785-795.

[1995MAR/MAC] Marani, D., Macchi, G. and Pagano, M. (1995), Lead precipitation in thepresence of sulphate and carbonate: testing of thermodynamic predictions. WaterResearch, 29, 1085-1092.

[1995RIS/HAL] Risold, D., Hallsted, B., Gauckler, L.J., Lukas, H.L. and Fries, S.G.(1995), The bismuth-oxygen system. J. Phase Equilibria, 16, 223-234.

[1995SIL/BID] Silva, R.J., Bidoglio, G., Rand, M.H., Robouch, P.B., Wanner, H. andPuigdomenech, I. (1995), Chemical Thermodynamics. Vol. 2: ChemicalThermodynamics of Americium, Elsevier, Amsterdam, 2033 p.

[1995WIB] Wiberg, N. (1995), Lehrbuch der Anorganischen Chemie, Walter de Gruyter,Berlin, pp.

[1996KAS/KAS] Kasenova, Sh.B., Kasenov, B.K. and Mustafin, E.S. (1996), Heatcapacity and thermodynamic functions of MSbO3 (M-Na, K, Cs) in the temperaturerange 298.15 - 673 K. High Temperature, 34, 481-483.

[1996KAS/MUK] Kasenov, B.K., Mukhanova, M.A., Kasenova, Sh.B. and Mustafin,E.S. (1996), The thermodynamic properties of alkaline-earth metal antimonates.Russ. J. Phys. Chem., 70, 24-26.

339

JNC TN8400 9 9 - 0 1 1

[1996STU/MOR] Stumm, W. and Morgan, J J . (1996), Aquatic Chemistry: Chemical

Equilibria and Rates in Natural Waters, John Wiley & Sons, New York.

[1997AMA/CHI] Amaya, T., Chiba, T., Suzuki, K., Oda, C , Yoshikawa, H. and Yui, M.(1997), Solubility of Sn(IV) oxide in dilute NaClO4 solution at ambient temperature.

Mat. Res. Soc. Symp. Proc, 465, 751-758.

[1997KOL/HEN] Koladima, A., Henglein, A. and Matijevic, E. (1997), CoUoidalhydrolysis products of SbCl3 in acidic solutions. Colloid Polym. Sci., 275, 972-

978.

[1997SAL/FER] Salvatore, F., Ferri, D., Trifuoggi, M., Manfredi, C. andVasca, E. (1997),

On the hydrolysis of the tin(II) ion. Annali di Chimica, 87, 477-481.

[1998ODA/AMA] Oda, C. and Amaya, T. (1998), Effects of ligands on the solubility of tin,

JNC Technical Report, Japan, JNC TN8400 98-001.

340

Documents

Thermodynamic Data for the Speciation and Solubility of Pd, Pb, … · 2004. 12. 21. · JNC TN8400 99-011 JP9955327 42 Thermodynamic Data for the Speciation and Solubility of Pd,