SALT EFFECTS ON THE RATE OF EXCHANGE OF THALLOIJS AND TH-FAELIC IONS IN WATER AND

HEAVY WATER

BY S . W. GILKS AND GWYNETH M. WAIND Chemistry Department, Queen Mary College, University of London, Mile

End Road, London, E.l

Received 28th January, 1960

At 25°C in perchlorate solutions the thallous-thallic exchange rate in water is 1.5 times as fast as in heavy water. There is a linear decrease with increase in perchlorate ion con- centration from 1.0 to 5.0 M; in this range replacing protons by sodium or barium ions has little effect on the exchange rate. At an ionic strength of 3.0 the thallic-ferrous re- action is catalyzed by hydroxyl ions ; the thallous-thallic exchange is not. Both reactions are catalyzed by platinum metal.

The thallous-thallic reaction is one of a group for which the hydrogen-ion inhibition is known to change markedly with increase in ionic strength.l.2 Such inhibition is usual for oxidation-reduction reactions of metal salts and has been interpreted as due to the faster reaction of hydrolyzed as compared with unhydro- lyzed metal aquo ions.3-9 This is of particular interest for the ferrous-ferric electron exchange because a mechanism of hydrogen-atom transfer has been based on the much greater catalytic effect of hydroxyl ions (as compared with other small anions) together with the size and existence of the isotope effect when the solvent (water) is replaced by heavy water.loy11 This is an analogous study of the exchange reaction between thallous and thallic ions in perchlorate solutions of ionic strengths greater than one. At these concentrations equilibrium properties show that ionic atmosphere effects are small compared with specific anion-cation interactions.

The thallous-thallic exchange involves a formal two-electron transfer and is catalyzed by platinum metal ;2 we have therefore compared our results with those for the thallic-ferrous oxidation where there is evidence for a two-stage process,l2 and in order to make this comparison more complete, have studied the effect of platinum metal on the rate of this latter reaction.

EXPERIMENTAL MATERIALS

Conductivity water was used throughout. Deuterium oxide (99h0.5 %) was obtained from I.C.I. in sealed glass ampoules.

Thallous perchlorate was prepared by evaporation of a perchloric acid solution of thallous nitrate (obtained from Johnson, Matthey & Co.) and analyzed by titration with potassium iodate solution.13 Thallic perchlorate was prepared in perchloric acid solu- tions from thallous sulphate.12 The thallic concentration was determined gravimetric- ally as oxide 14 and volumetrically with potassium iodide and sodium thiosulphate. The total thallium concentration of these solutions was determined in the same way after reduction with sulphur dioxide, and the formal hydrogen ion concentration by titration with sodium hydroxide after treatment with hydrogen peroxide. This was confirmed by potentiometric titration with standard alkali, correcting for precipitated thallic oxide.

Sodium perchlorate was obtained from B.D.H. and was chlaride-free. Initial diff- culties encountered in attempts to prepare and recrystallize sodium perchlorate are de- scribed in detail elsewhere.15 A.R. 60 % perchloric acid was standardized by titration.

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S. W. GILKS AND G. M. WAXNL) 103

Barium perchlorate solutions were prepared from a recrystallized B.D.H. product and were analyzed for barium gravimetrically as the chromate and as the sulphate.

All the other materials were A.R. grade and were used without further purification. Solvent mixtures of deuterium oxide and water were analyzed using an Abbe refracto-

meter. Stock solutions of radioactive thallous and thallic perchlorate were made by adding

a small quantity of 204Tl perchlorate obtained from the Radiochemistry Centre, Amersham, Bucks.

RATES OF BXCHANGE

These were followed at 25 &O*Ol"C by precipitating thallous ion as chloroplatinate.16 All samples were counted as solids using a G.E.C. G.M. 4-type end window counter. Since samples of the same weight were always compared no absorption or scattering corrections were made. The average deviation between the measured activities of duplicate samples was always less than 1 %. All measurements were made in duplicate.

THALLIC-FERROUS OXIDATION RATES

All solutions were made and standardized exactly as described in ref. (12) but at room temperature. One aliquot was removed from 100ml of reaction mixture in a beaker containing the platinum metal and titrated as in ref. (12).

[ClOil, M FIG. 1.

0 HC104, NaCIO4, X Ba(C104)2 k = 0.5489-0.08794 [ClO,]

RESULTS

The bimolecular rate constants were calcuIated in the usual way 17 and are given in fig. 1 and tables 1 and 2. All measurements were made at 25°C. The line

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1 04 E X C H A N G E O F T H A L L O U S A N D T H A L L I C I O N S

drawn in fig. 1 is the least-squares slope (excluding data below 1.0 and above 5.0 M perchlorate) and corresponds to the equation,

k = 0.5489 - 0.0879[C10~]total. ( l a ) The corresponding equation for 80 % D20 solutions calculated from the data

in tables 1 and 2 is k = 0.392 - O~0628[C10~]tota~. (1 b)

TABLE 1

EI(C104)3] 040770 M, [TIClO4] 0.020 M, ionic strength 1 *20

0.336 0.388 0.377 0279 0.564 0.422 0.690 0.300 0.906 0.441 1 a00 0300 1.134 0.456

[HC10]4, M itHzo, mole-1 1. h-1 [DClO4]*, M k*,,, mole-1 1. h-1

*total D/H = 80 % throughout

TI(C104)3 0.00773 M, [TIC1041 0.020 M, ionic strength 3.0 [HC104], M kHz0, mole-1 1. h-1 DClO4]*, M k*D20, mole-1 1. h-1

0.42 0268 0.40 0.90 0.271 0.70 1-37 0.271 1.03 1.92 0268 1.38 2.39 0.265 1 -75 2.93 0.268 213

[Tl(C104)3] 0.00770 M, [TIC1041 0.020 M, LHC1041, M kHzo, mole-I 1. h-1 DClO41*, M

0.365 0.0827 0.283 0.936 0.0762 0.832 2.01 0.0687 1.75 3.08 0.0558 2.66 4.5 1 0.0485 5.93 0.0439

0.204 0.202 0.198 0.1 99 0.200 0.203

ionic strength 6.0 k*,,,, mole-1 1. h-1

0.0667 0-0563 0.0502 0.0471

TABLE 2

[TlC104] = 0.020 M, solvent 100 % H20 TI(C104)3] M = 0.001345 0,00269 0.00538 0*00807

p = 1.20 HC104, M = 1.11 k = 0.474 0,473 0459 0.437 mole-1 1. h-1

k = 0.271 0.270 0.263 0.257

k = 0.275 0.280 0.273 0.265

k = 0.0907 0.0888 0.0843 0.08 17

k = 0.0627 0.0610 0.0635 0.0561

p = 3.0

p = 3.0

p = 6.00

HClO4, M = 2.499

HClO4, M = 2.910

HClO4, M = 0.473

HClO4, M = 4.28 p = 6.00

Tables 1 and 2 are for solutions at fixed ionic strengths ( p = 1-2, 3.0 and 6.0) and contain the reactants together with a large excess of sodium and hydrogen

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S . W . G I L K S AND G . M. W A l N D 105

perchlorate only. The ionic strength here is therefore effectively the perchlorate concentration. In water at p = 1.2, the rate constant increases from 0-388 to 0-456 ; at p = 3-0 there is no change ; at p = 6-0 the rate constant decreases from 0.083 to 0.044, when in each case the formal hydrogen ion concentration is increased from about 0.4 M to p. Fig. 2 shows that at p = 6.0 the rate is proportional to the molar concentrations of perchloric acid and sodium perchlorate.

I I 1 1 0 I 2 3

[H+l/INa+l FIG. 2.

ionic strength 6.0; 0, present work (H20); e, Harbottle and Dodson kNa+ = 0.470 kHi = 0.517

The curvature of the second-order plot for the thallic-ferrous oxidation observed by Higginson12 after about 60 % reaction and attributed by him to the back reaction FeITr+ TII1=FeI1+ T P is eliminated by platinum metal (fig. 3).

DISCUSSION

The linear decrease in rate with increasing perchlorate concentration (fig. 1) of a reaction between two positive ions which is known to be catalyzed not only by nitrate 2 but also by sulphate ions 14 seems unusual. It has, however, been pre- viously shown that catalysis by sulphate ions can be accurately described 18 and also qualitatively understood 19 if the reaction intermediates are TI+SO2--Tl3+ and SO~-Tl+SO~-T13+SO~--. In these species the reactants niay :or may not, also be separated by water molecules and the reaction therefore [may be either the outer-sphere or the inner-sphere type of Ta~be.20~9 20b It seems likely that the nitrate catalysis is due to Tl3+NO3Tl+, as a structure in which the nitrogen must interact with the thallic ion not only explains the kinetic data but also incor- porates the explanation given by Gutowsky21 of the equal dependence of the chemical shifts of thallous and thallic nuclei on increasing ratio [nitrate]/[thallium]

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106 EXCHANGE OF THALLOUS AND T H A L L I C tONS

t (min)

FIG. 3. - - - - no pt.

1 in. square 0.1 mm thick

[FeII] 0.003884 M bright pt. fnIII] 0.005202 M

[H+] 0567 M p - 0.9 with NaClO4 room temperature

in concentrated nitrate solutions. Although similar n.m.r. (2OsTl resonance) studies show that concentrated thallic perchloric solutions (which cannot be prepared free from thallous ion) give a very broad absorption band ;22 the variation in position and in band width with change in perchlorate concentration and with change in thallous-thallic ratio have not yet been measured. Many properties of concentrated solutions containing tetrahedral oxy-anions have been attributed to either (i) change in solvent structure,23 or (ii) cation-water-anion complexes.24

For the solutions used here there is no information which makes it possible to distinguish between these effects; either or both of which might be expected to cause the observed linear decrease in the exchange rate with increasing perchlorate ion concentration.

HYDROGEN-ION DEPENDENCE

The change in rate with increasing perchloric acid concentration is very different from that of the thallic-ferrous electron transfer where there is a marked acid catalysis above 2 M.25 At constant perchlorate concentration hydrogen ion inhibition is large only in the 6 M solutions. As recent studies of the hydration of protons 26 and of triply charged cations 27 suggest that this may reflect changes in the degree of hydration as well as of the hydrolysis of the thallic ion, and as studies of the ultra-violet absorption of these solutions confirm this,28 we have only calculated specific rate constants for TW and TlOH2+ in the more dilute solutions.

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S. W . G I L K S AND G . M. W A I N D 107

In 3 M perchlorate solution, two independent determinations of the degree of hydrolysis of the thallic ion are in complete agreement.28929 The hydrolysis constant at 25°C is 0.073 and is the same in water and in mixtures of water and heavy water.28 Application of eqn. (2) and (3) to the data in tables 1 and 2 gives :

where (3)

w3 +I ; cIo=- [T10H2 '][H']

[TI3+] [TI"']' K =

ko = 0.267, kl = 0.301 mole-1 1. h-1; and to the data of Higginson 12 for the thallic-ferrous reaction, ko = 0.44 mole-1 1. min-1, Icl = 15.1 mole-1 1. min-1. All these values differ from those quoted in the literature.30~ 31

Although the presence or absence of hydroxyl-ion catalysis can only give definite information about the composition of the transition state, i.e. it either does or does not contain a hydroxyl ion; in view of the much discussed importance of bridging groups in electron transfer reactions and the ability of the hydroxyl ion to penetrate the hydration sheaths of ions, it is interesting that while dimeric ferric hydroxyl complexes are well-known, thallic and mercuric ions are the only cations for which this behaviour has not been postulated.32, 33

D/H ISOTOPE EFFECT

This is independent of ionic strength (table 1) and corresponds to ~ H ~ o / ~ D ~ o = 1.5 for 100 % deuterium oxide. The significance of such an effect has recently been discussed.*os 3 1 ~ 3 4 The importance of water as a constituent in the activated complex for electron transfer reactions has also recently been stressed.35

1 Harbottle and Dodson, J. Amer. Chem. SOC., 1951,73,2442. 2 Prestwood and Wahl, J. Amer. Chem. SOC., 1949, 71, 3137. 3 Anderson and Bonner, J. Amer. Chem. SOC., 1954,76, 3826. 4 Silverman and Dodson, J. Amer. Chem. SOC., 1952, 56, 846. 5 Meier and Garner, J. Physic. Chem., 1952, 56, 852. 6 Sherril, King and Spooner, J. Amer. Chem. Soc., 1943, 65, 170. 7 Keenan, J. Amer. Chem. Soc., 1956,78,2339. 8 Furman and Garner, J. Amer. Chem. SOC., 1952,74,2333. 9 Sutcliffe and Weber, Trans. Faraday Soc., 1956,52, 1225.

10 Davidson and Dodson, J. Physic. Chem., 1952,56, 866. 11 Reynolds and Lumry, J. Chem. Physics, 1955,23,2460. 12 Ashurst and Higginson, J. Chem. Soc., 1953, 3044. 13 Andrews, J. Amer. Chem. SOC., 1903,25,766. 14 Brubaker and Mickel, J. Inurg. Nuclear Chem., 1957,4, 55. 15 S. W. Gilks, Thesis (London, 1960). 16 Challenger and Masters, J. Amev. Chem. Sac., 1956, 78, 3012. 17 McKay, Nature, 1938, 142, 997. 18 Brubaker et al., J. Amer. Chem. Soc., 1957, 79,4641. 19Libby, J. Amer. Chem. SOC., 1940, 62, 1930. 20 Henry Taube, (a) Int. Cunf: Co-ordination Chemistry (Spec. Publ. Chem SOC.,

London, 1959), p. 57 ; (b) Advances in Inorganic Chemistry and Radiochemistry, vol. 1 (Academic Press, 1959).

21 Gutowsky and McGary, Physic. Rev., 1953, 91, 81. 22 B. N. Figgis, private communication. 23 Robinson and Stokes, Electrolyte Sulutiuns, 2nd ed. (Butterworths, 1959), chap. 1. 24 Sykes, J. Chem. SOC., 1959, 2473. 25 Forscheimer and Epple, J. Amer. Chenl. Soc., 1952, 74, 5773.

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108 EXCHANGE O F THALLOUS AND TI-IALLIC I O N S

26 Bell and Bascombe, Faraday SOC. Discussions, 1957, 24, 158. 27 Glueckauf, Trans. Faraday Soc., 1955, 51, 1235. 28 Waind and Rogers, to be published elsewhere. 29 Biederman, Arkiv. Kenzi., 1953, 5, 441. 30 Rossotti, J. Inorg. Ncrcl. Chem., 1955, 1, 159. 31 Basolo and Pearson, Mechanisms of Inorganic Reactions (John Wiley and Sons Inc

32 Gluskova, Russ. J. Inorg. Chem., 1959, 7. 33 Sillen, Acta Chem. Scand., 1954, 8, 299, 1607. 34 Marcus, J. Chem. Physics, 1956, 24, 966. 35 Maddock, Trans. Faraday Sac., 1959,55,1268.

1958), chap. 7.

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