NITRATE SELECTIVE RESINS · 14N nmr spin-lattice relaxation times and a method for interpreting the non-linear concentration dependence of these times, in terms of ion association,

¡$. r-z-.1

I

NITRATE SELECTIVE RESINS

INTERACTION OF MONOMERIC AND POLYMERIC

QUATERNARY AMMONIUM COMPOUNDS WITHNITRATES

by

Gary "Petert' Owens

B.Sc. (Hons.)

Thesis submitted for the degree of

Doctor of Philosophy

1n

The University of Adelaide

Department of ChemistrY

March 1995

Table of Contents

Absract

Statement...

Acknowledgements

Quotation

Abbreviations

1 General Introduction.

1

v

vi

vii

viü

ix

1.1

t.2

Introduction

Nitrates and Health Effects

1

1

1

2

3

3

3

4

6

6

6

7

7

8

1.2.1 Metabolism...........

1.2.2 Methaemoglobinaemia

1.2.3 Nitrosamines and Cancer

I.3 Nitrate Distribution in Nature

1 .3 . 1 The Nitrogen Cycle

1.3.2 Nitrate Distribution in Australia. .. . . .

1.4 Current Methods of Nitrate Removal.

1.4.1 Biological Denitrification.

1.4.2 Blending

1.4.3 Reverse Osmosis.

1.4.4 Ion Exchange.....

1.5 Nitrate Selective Resins

1.5.1 Amidines

L.5.2 Salinogen Resins.

1.5.3 Commercial Resins

1.5.4 Polyallyamine Resins.

1.5.5 Polystyrene-divinylbenzene Resins

... 10

... 11

...13

... 16

...18

...231.6 Conclusion...

ll

2 Synthetic Methods and Characterization.

2.1

2.2

2.3

2.4

2.5

General Experimental .

Monomeric Analogues

..33

28

28

32

35

37

39

46

46

47

47

49

50

52

52

54

54

55

55

58

63

64

64

Linea¡ Polymers

3.1

3.2

2.3.1 Monomers...........

2.3.2 Polymerization.......

Cross-linked Polymers

Interpretation of NMR

3 Solubility, Enthalpy and Entropy-

Introduction .....

The Stn¡cnre of Water.

3.2.1 Flickering Cluster Model

3.2.2 Structure Breaking and Structure Making by Ions

3.2.3 Hydrophobic Interactions..

3 .2.4 Theory of Solution . . . .

3.3 Water-Solute Interactions....

3.3.1 Alkanes

3.3.2 Tetraaþlammonium Salts

Solubility

Enthalpy and Entropy of Solution..........

Halides

Nitrates

3.4 Methods and Experimental

3.4.I Assumptions Regarding Activity Coefficients......

3.5 Trends in the Benzylrialkylammonium Salts .. . .

4.1

4.2

Solubility

Enthalpy and Entropy of Solution...

3.6 Conclusion ....

4 Density and ViscositY.

Introduction .

Theoretical Basis...

Introduction

Theory

5.2.L Nuclei within a Magnetic Field.

5.2.2 Relaxation and Quadrupole Effects ...........

5.2.3 Relaxation and Rotational Correlation Times

5.2.4 Equilibrium Constants.

64

66

74

79

79

79

81

82

83

84

84

85

87

87

89

90

96

4.2.L Density......

Extrapolation to Infinite Dilution

4.2.2 Viscosity

Viscosity in the Presence of Solute

4.3 Experimental Techniques..........

4.3 .l Determination of Density . . . .

4.3 .2 Determination of Viscosity . . .

4.4 Results and Discussion

4.4.L Patial Molar Volumes

4.4.2 Viscosity

4.4.3 Cation Size and Hydration

4.5 Conclusion

5 Nuclear Magnetic Resonance.

5.1

5.2

.. 108

99

100

100

104

105

5.3 ExperimentalMethods 111

lv

5.3.1 Solution Preparation

5.3.2 Density and ViscositY.

5.3.3 Design of the Spin-lattice Relaxation Experiment ..... .

Choice of Pulse Sequence.

Choice of Delay Times

5.3.4 NMR.

5 .4 Results and Discussion . ...... .

5.4.1 Relaxation in Simple Nitrate Salts........

Concentration Dependence.........

5.4.2 Relaxation in Benzyltrialkylammonium Salts.

5 .4.3 Relaxation in Linear Polymer-Nitrate Solutions . . . .

5 . 4.4 Relaxation in Crosslinked Polymer-Nitrate Solutions

5.5 Conclusion

6 Conclusion.

111

6 Conclusion .. . ...

Appendices.

S electivity Palameters

Preliminary Synthesis' of Improved Capacity Resins......

Accuracy of the Spectrophotometer

Enthalpy and Entropy of Tetraalkylammonium Halides

Calibration of Pycnometers............

Experimental Data and Fitting Pammeters ,

AO

A1

A2

A3

A4

A5

..154

111

r12

Lt2

115

115

116

116

t20

130

140

145

t49

A-1

A-6

A-r4

A-16

A-23

A-28

ABSTRACT

The aim of this work was to discover the basic principles that govern the selective

removal of nitrates from drinking water by ion-exchange resins. Nitrate association of

vinylbenzyltrialþlammonium polymers and their monomer analogues were studied using

14N nmr spin-lattice relaxation times and a method for interpreting the non-linear

concentration dependence of these times, in terms of ion association, was presented for

the first time.

When applied to simple nitrate salts (NH¿+, K+, Na+, Ag+, Ca2+,B,a2*,p62+¡ the

method confirmed that the rate of chemical exchange was slow relative to molecular

reorientation of the nitrate ion. Treatment of monomer analogues, such as

benzytrialklyammonium salts, indicated that association for nitrate increased with the

length of the alkyl substituent. The advantage of this new technique was that it could be

used over a wide range of concentrations up to near saturation. 'When applied to linear

polymers it was found that nitrate association increased with the size of the alkyl group

and that the relaxation time was related to hindered rather than site bound nitrate. The

relaxation time was also found to give information on the hindered rotational state of the

nitrate ion in lightly cross-linked gels. Any selectivity observed for monomeric analogues

was enhanced upon polymerization.

Solubility measurements and application of the Van't Hoff isochore enabled the

thermodynamics of the solution process of the monomer analogues to be examined more

closely. The results were interpreted in tenns of significant hydrophobic bonding by these

cations. Density and viscosity measurements of aqueous solutions of these salts

supported the notion that these salts caused an increase in the structure of water in their

vicinity. The degree of hydration increased with the size of the alkyl substituent and this

was attributed to increased clathrate cage formation. In this respect it was concluded that

these salts were not significantly different from tetraalkylammonium salts and that

addition of a benzyl substituent instead of an alkyl group had not significantly altered the

solution properties.

vlt

I would like to thank Dr. Tom Kurucsev for accepting me as his last postgraduate

student as well as for his help, understanding and constructive criticism. I am only sorry

that I did not always live up to his high standards. I would also like to extend my thanks

to include Dr. Barry J. Steel for accepting me as a pseudo student and for his exceptional

insight. I would also like to thank my partner in crime Philippe Guarilloff for comic relief

and other things that occasionally approached chemistry.

I also have to thank all the people that have helped to keep me sane these last few

years and those that loved me even when I was unlovable. My sincere thanks to Miss

Sonya Whitbread for support, encouragement and for proof reading sections of this

thesis. Thanks also to Reina Whitbread, Ramesh Dhillon and Mike Papps for friendship,

distractions and support in recent times. Thanks also to all the technical staff and

students, past and present, of the department of chemistry, who I have met, and said

goodbye to along the way.

It would be remiss of me if I did not thank my three bosses, Graham Bull, Rob

Morris and Peter Roberts. It was a pleasure to have worked with each of them and I am

so very sorry the department has not given them the recognition they so rightly deserve.

An Adelaide University Scholarship for postgraduate research is acknowledged in

the same begrudging manner that it was awarded.

vlll

'What did Dr Thurmer say to you, boy? I understand you had quite a little chat.'

'Yes, we did. We really did. I was in his office for around two hours, I guess.'

'What'd he say to you?'

'Oh ... well, about Life being a game and all. And how you should play it according to

the ru|es. He was pretty nice about it. I mean he didn't hit the ceiling or anything. He just

kept talking about Life being a game and all. You know.'

'Life is a. gamq boy. Life is a game that one plays according to the rules.'

'Yes, sir. I know it is. I know it.'

Game, my ass. Some game. If you get on the side where all the hot-shots are, then it's a

game, all right - I'll admit that. But if you get on the other side, where there aren't any

hot-shots, then what's a game about it? Nothing. No game.

J. D. Salinger

The Catcher in the Rye.

Ncr,c

NCT,S

lx

Abbreviations and Symbols

= Separation factor indicating the relative preference by a resin for nitrate

over chloride.

= Separation factor indicating the relative preference by a resin for nitrate

over sulfate.

= Standa¡d Gibbs free energy of solution.

= Standa¡d Enthalpy of solution.

= Standard Entropy of solution.

= Molar absorptivity.

= Asymmetry factor.

= Mean ionic activity coefhcient.

= Relative viscosity.

= Wavelength maximum.

= Absolute density.

= Rotational correlation time.

= Rotational correlation time for rotation about the Cz axis of the nitrate ion.

= microseconds.

= Ion size parameter in cm.

= Mole fraction of nitrate bound.

= Angstrom.

= 2,2'-azo-bis-isobutyronitrile (free radical initiator).

= Localized binding ratio.

= Aryl substituent, CoHs-

= Analytical grade reagent.

= boiling point.

= butylsubstituent,CH3CH2CH2CH2-

: Benzyl substituent, -CH2C6FI5

= Resin capacity in meq dm-3.

= calorie.

= 95Vo Confidence Interval.

AGosol

ÂHosol

ÂSosol

€

95

T¡

lrIru^

p

Xç

1r

ps

a

A

Å

AIBN

Al

Ar

AR

b.p.

Bu

Bz

C

cal

CI

x

CLVo

CME

CPE

calc.

cond.

COIT.

DADMAC

dec.

DVB

EMF

Et

expt.

FAB

FHIRFT

He

HEXA

hrs

HSP

hvd

Hz

IR

IRFT

i-Pe

KÈ

Kl

K¡

K5

cross-linking percentage.

cross-linked poly(vinylbenzyltrimethylammonium chloride).

cross-linked poly(vinylbenzyltriethylammonium chloride).

calculated.

conductance.

corrected.

Diallyldimethylammonium chloride.

decomposes.

Divinybenzene.

Electomotive force.

ethyl substituent, CH3CH2-

experimental.

Fast atom bombardment.

Freeman-Hill modification of the Inversion Recovery Fourier Transform

pulse sequence.

hexyl s ub stitue n t, CH3 CH2CHzCHzCHzCHz-

1,6 dibromohexane.

hours.

Homospoiling pulse.

hydration.

Hertz.

Infrared.

Inversion Recovery Fourier Transform pulse sequence.

isopentyl substituent, (CHg)zCH C}l2C}f2-

Selectivity coefficient indicating the selectivity of a resin for nitrate relative

to chloride.

= Selectivity coeff,rcient indicating the selectivity of a resin for nitrate relative

to sulfate.

= Thermodynamicassociationconstant.

= Stoichiometricassociationconstant.

XI

Iat

lit.

LPBU

LPET

LPME

LPPE

LPPR

Me

meq

mg

MHG

rnl

m.p.

MS

Mw

n

nm

nlnr =

NSS =

NQCC,1 =

Oc=PDAA =

Pe=pH=PMeDAA =

PMeTAA =

ppm =

PPTDAA =

Pr=ps=

lattice.

literature.

linear poly (vinylbenzyltributylammonium chloride).

linear poly(vinylbenzyltriethylammonium chloride).

linear poly(vinylbenzyltrimethyl ammonium chloride)

linear poly(vinylbenzyltripentylammonium chloride).

linear poly(vinylbenzylripropylammonium chloride)

methyl substituent, CH3-

milliequivalent.

milligram.

Methaemoglobinaemia

millilitre.

melting point.

milliseconds.

Molecular weight.

number of transients.

nanometers.

Nuclear magnetic resonance.

Nitrate to sulphate selectivity index.

Nuclear quadrupole coupling constant.

octyl substituent, CH:(CHz) sCHz-

Polydiallyamine.

pentyl sub stituent, CH¡ CHzCHzCHzCHz-

-log[H+]

Polymethyldiallyamine.

Polymethyltriallyamine.

parts per million.

Polypropyldiallyamine.

propyl sub stituent, CH3CH2CH2-

picoseconds

xll

PS-DVB =

PSFT =

TTAA =

PVA =

QTüXA =R=rad =

ref =

S_

S=SATMA =

S/N =

Sol =

SRFT =

Tt='tirãl

T6=T¡=TLC =

TMS =

T¿=UV=WHO =

Poly styrene-d ivinylbenzene matrix.

Progressive Saturation Fourier Transform pulse sequence.

Polytriallyamine.

Polyvinylalcohol.

1,6-bis(N,N,N-diallymethylammonium)hexane dichloride.

Mole ratio of nitrate to polymer.

radians.

reference.

seconds

Molal solubility.

Spacer arrn trimethylammonium cation.

Signal to noise ratio.

solution.

Saturation Recovery Fourier Transform pulse sequence.

Spin-lattice or lon gitudinal relaxation time.

Acquisition time.

bound relaxation time of the nitrate ion.

free relaxation time of the nitrate ion.

Thin layer chromatography.

Trimethylsilane

Delay or recycle time.

Ultra-violet (spectroscopy).

World Health Organization.

xlll

Spectroscopi c Abbreviati ons

õ chemical shift in ppm

broad

doublet

multiplet, medium

quartet

singlet, strong

riplet

weak

br

d

m

q

s

t

w

lll

1

General Introduction.

1.1 Introduction.

The contamination of drinking water by nitrates is one of the three major water

quality problems facing health organizations today (1). The other problems being bacterial

contamination and pollution due to toxic chemicals from industrial spills.

The main concern in recent years has been possible health effects associated with

the presence of high concentrations of nitrate in potable drinking water. The World Health

Organization (V/HO) has recommended a limit of 11.3 mg of nitrate as nitrogen per litre

(2), which is exceeded in many places, suggesting the need for a method of nitrate

removal from such polluted water supplies.

1.2 Nitrates and Health Effects.

L.2.1 Metabolism.

Nitrates are readily absorbed by the body. The main sources of introduction being

by way of food and drinking water. The exact metabolism of nitrates in humans is not

completely understood, although once absorbed into the bloodstream it plays little role in

normal biological functions and is readily excreted, primarily in the urine. Nitrate is

absorbed in the upper gastrointestinal tract and concentrated ultimately by the salivary

glands in the saliva. However, nitrate itself is not very toxic and the major health

concerns are due to the bacterial reduction of nitrate to nitrite within the body. Bacterial

reduction occurs mainly within the mouth and the stomach.

Introduction

The endogenous production of nitrite in human saliva can be as much as twenty

times the amount ingested. Bacterial reduction of nitrate in the stomach to nitrite does

occur, but conversion is small unless the pH is greater than 4.6. For an adult stomach

where the pH is in the range 1-5, this means there is no significant conversion. However,

in this pH range the produced nitrites can react with secondary and tertiary amines and

amides to form nitrosamines which may be carcinogenic. The rate of reaction is greatest at

pH 3.5 or less. There are therefore essentially two major health hazards,

methaemoglobinaemia and cancer.

1.2.2 Methaemoglobinaemia.

Methaemoglobinaemia (MHG) is a very rare illness found almost exclusively in

infants below four months of age. In extreme cases symptoms include vomiting and

diarrhoea always accompanied by a blueness of the skin, medically referred to as

cyanosis. Unless treated it will result in death.

MHG is usually traced to nitrate polluted water which was used to prepare infant

formulae and food. The nitrate is bacterially reduced to nitrite in the upper intestine or

stomach and the nitrite in turn oxidizes haemoglobin to methaemoglobin.

Methaemoglobin is a pigment which is incapable of acting as an oxygen carrier. Thus

death occurs due ro oxygen deprivation and it is this that results in the distinct blue

colouration. MHG is also known as "blue-baby syndrome" and may be related to cot

death.

There are several reasons why infants are prone to MHG. Physiologically, the

stomach pH of infants is almost neutral and so there is a high degree of nitrate reduction

in the stomach. Up until about 4 months of age the concentration of metheamoglobin

reductase, an enzyme that reduces metheamoglobin, is also low and the infant cannot

adequarely remove produced methoglobin (1). More practically, the fluid intake per unit

body weight of an infant is approximately three times that of an adult and often water

used for preparing baby foods undergo repeated boiling prior to preparation, thereby

increasing the nitrate concentration.

2

Inftoduclion

MHG poses no serious health threat provided it is diagnosed early enough to allow

treatment. The suggested treatment involves intravenous injection of methylene blue, 1-2

mg kg-1 of body weight of a lVo methylene blue solution in normai saline over a 10

minute period (1). The methylene blue converts metheamoglobin back to hemoglobin. In

milder cases the oral administration of ascorbic acid (Vitamin C), which also reduces

metheamoglobin, may be more suitable (1).

1.2.3 Nitrsoamines and Cancer.

Nitrates produced by reducing bacteria in the stomach and mouth can possibly

combine with secondary and tertiary amines in foodstuffs to produce nitrosamines. There

is some evidence that nitrosamines are carcinogenic in animals (3), but disagteement

concerning the toxicity of these nitrosamines in humans (4).

The formation of nitrosamines can easily be inhibited by many antioxidants, such as

Vitamin C (5), which are readily available in fresh fruit and vegetables. However in

developing countries or remote ateas were fresh fruit and vegetables are scarce this mode

of protection may not be possible. The obvious alternative is the removal of nitrates prior

to ingestion.

1.3 Nitrate Distribution in Nature

1.3.1 The Nitrogen Cycle.

Nitrate is widely distributed throughout nature and its conversion from the many

available oxidation states is described by the Nitrogen Cycle (6).

Almost 95Vo of the nitrogen containing compounds found in topsoil are bound in

organic ma6er. The nitrate present in the soil is taken up by crop and plantlife. The

plantlife in turn is used by a variety of animals as a food source. The nitrate being later

returned to the soil when the plants or animals die or through faeces and urine. However,

when the soil is cultivated, the crop or livestock are later removed and cannot be recycled

J

Inlroduction 4

into the soil. The nitrogen must therefore be externally replenished otherwise depletion

will quickly occur. The discovery of the Haber process in 1905 allowed the production of

an unlimited amount artificial nitrate fertilisers to replace the limited stockpiles of natural

fertilizers (7).

In an excellent review by Owen and Jûrgens-Gshwind (4) the distribution of

nitrates has recently been summarised. In general, nitrate movement closely follows the

movement of water, so that rainfall or excess irrigation in winter favours loss of nitrate

from soil due to leaching, but a similar rainfall in summer will result in the water being

taken up by plantlife. Leaching and runoff are enhanced by sandy or loamy soils and a

lack of plantcover. The application of higher concentrations of nitrogen fertilizer also

leads to greater leaching of nitrates from soils to groundwater. These and other effects

help to explain the highly variable concentrations of nitrate groundwater in areas having

similar arable lands.

1.3.2 Nitrate Distribution in South Australia.

The majority of groundwater sources in South Australia, Fig. 1.1. and Table 1.1.,

appear to have nitrate levels well below the WHO limit, with the exception of water from

the Musgrave basin. In this case the high nitrate level was attributed to a seasonal effect,

resulting from the fixation of atmospheric nitrogen by native vegetation (8). The results

should be viewed with caution because the measurements are only the average nitrate

levels from a limited number of observation wells in any particular basin. It is only when

looking at particular bore sites that high levels of nitrate become evident.

Lawrence (9) conducted an exhaustive review of nitrate levels in groundwater

throughout South Australia and identified 138 bores that had nitrate levels greater than 20

ppm. The nitrate levels varied greatly between bore sites, having an upper limit of 490

ppm nea.r Blanche. Although this survey is somewhat old, there is no reason to suspect

that nitrate levels would have diminished and may have actually increased in some areas

due to increased agricultural activity. Information supplied by the Department of Housing

and Construction (10) also indicates that some aboriginal communities which rely on

Unrl Boundary

Sed menlary Bas,d t

MUSGRAVEBLOCK

STUAFfSHELF

YOFKE PENINSULA

Table I I Median Anion Concentrations ofSouth AustrulianCroundu,ater î

Rcgion Chloride Nitrilre ¡ìs N Sulldte

sr. v¡t¡cENtEAâIN

(Jts r\, Ëìcs¡¡ìMumv lJrsinSl VirccDt BnsinMt Lofty/Flindcrs RangesYork Pen¡nsulrEyre Pcninsulr['irie ]'orens ËlasinCrert Atrcsifln BnsinOhrv ,\rcN,lusgrirve ßlock.Strrirn ShclIEuclr Il¡sin

240s803ó0

0680.230.23

448819

130I 2004r0

1200700

ì 700190

| 700.5 I (X)

0,451,05,2

0560.2 3l 8l

t510

068

tó0t9062

170260110160

-5t0r 200

rcltrcncc ll. conccntmtions in nt¡y'L lppnt), No dtttarLr,¡rilrbie on the Ol'l'iccr Brsin

Fig. 1.1. Seclinrentury llrsins unrl lrigh rrrtrute level observittiorl wclls il] SoLrth Austrlliit.

Introduction 6

borewater may experience high levels of nitrate contamination. The isolated nature of

many of these wells highlights the need for a method of nitrate removal that is simple,

portable and relatively inexpensive.

1.4 Current Methods of Nitrate Removal.

Many methods have been suggested for the removal of nitrates from drinking water

(10-15,17-41). All the reported methods have both advantages and disadvantages which

shall be briefly reviewed in this section. It appears that ion-exchange offers the highest

efficiency and cost effectiveness of all the currently available methods.

I.4.1 Biological Denitrification.

In this method bacteria are used to reduce nitrate to nitrogen (Il,I2). The main

disadvantage is that the reduction is sometimes incomplete and produces nitrite which is a

far more hazardous compound. The process also introduces other toxic substances such

as methanol, which acts as electron donor and is therefore essential for the reduction.

There also exists the possibility that the microbiological integrity of the water might be

compromised by such treatment (13). Several pilot plants utilizing microbiological

denitrification are in operation and great promise exists for its dual application with ion-

exchange techniques (14).

I.4.2 Blending.

The simplest method for reducing nitrate concentrations below a nominal value is

by blending polluted water with unpolluted water to give a lower nitrate concentration by

dilution. This is a perfectly acceptable method provided an alternate source of nitrate free

water is available. Unfortunately this is not always the case and from a scientific

standpoint also side steps the issue of selective nitrate removal. It is likely however,

Inlroduction

where applicable, that because of its ease and cost effectiveness, any pilot plant will

include blending in combination with one or more of the other methods discussed here.

L.4.3 Reverse Osmosis.

Reverse osmosis uses a semi-permeable membrane to remove ions from solution.

Like Blending, the main disadvantage is that it is not selective for nitrate and so all ions

are removed. Generally, the nonselective removal of all contaminants is found to be

considerably more expensive. Most membranes are curently prepared from cellulose

acetate or its derivatives. The use of such membranes requires that the wastewater be

extensively pretreated to avoid organic fouling and results in additional treatment costs.

The membranes are also subject to deterioration during extended use. It is believed that

further research into the design of membranes from polymeric materials will solve some

of these problems (15). Finally, the level of sophisticated instrumentation required for

reverse osmosis makes it unsuitable for small-scale water purif,rcation in rural areas.

1.4.4 Ion-Exchange.

Ion-exchange has become one of the most important methods for nitrate removal

and a wide variety of compounds have been reported to selectively remove nitrate. This

selective removal is important because not all ions are harmful and many have beneficial

effects. Unfortunately many compounds which remove nitrate to some degree normally

remove sulfate to a greater degree. It has been an on going problem to find compounds

suitable for use as ion-exchange resins which selectively remove nitrate in preference to

sulfate from potable water.

The two most desirable properties of a commercial ion-exchange resin are a high

capacity for nitrate and significant selectivity for nitrate over other coûìmon anions present

in the water. Given that two resins possess similar such properties the ease with which

the spent resin can be regenerated becomes important in determining commercial viability.

A review of the previous attempts to prepare nitrate selective resins follow. The capacity

7

Introduction 8

and a measure of the suspected nitrate selectivity will be discussed, while the regeneration

efficiency of the resins is, in general, poorly reported and will not be reviewed.

1.5 Nitrate Selective Resins.

There are three commonly used parameters for the discussion of nitrate-sulfate

selectivity the separation factor, the selectivity coefficient and the nitrate-to-sulfate

selectivity factor respectively denoted, o\n, rN and NSS.

Consider first the reaction between a monovalent ion, such as nitrate, and an ion-

exchange resin conditioned with another monovalent ion, such as chloride. The two

monovalent ions will compete for available binding sites on the resin. The equilibrium can

be represented by,

No¡- CI <+ Nor + cl- (1)+

where NO¡- and Cl- denote free anions in solution and NO3 and Cl denote anions bound

to the resin. The nitrate-to-chloride separation factor, o!, it then given by,

oÈ= (2)

which may be rewritten in terms of equivalent ion fractions to yield

oÈ= XN XcXc XN

(3)

where XN and ><C .. the equivalent fractions of nitrate and chloride in the resin phase

and Xg and X¡¡ are the equivalent fractions of ions free in solution. A resin is said to be

nitrate selective over chloride if of > t.

For competition between two monovalent anions separation factors and selectivity

coefficients are identical. However, this is not true for any equilibrium involving a

Introduclion

divalent anion. In the case of competitive binding between a monovalent anion, such as

nitrate, and a divalent anion, such as sulfate, the equilibrium expression may be written

âS'

2NO¡- + SO4 (+ 2NO3 + so42- (4)

9

where SOoz- and NO3- represent free anions in solution and NO3 and SO¿ represent

anions in the resin phase. This equilibrium expression implies that the resin has been

preconditioned in the sulfate form. The corresponding nitrato-to-sulfate separation factor

and selectivity coefficient are then,

tNo:l ISo¿2-loN=tsõ7-l tNo¡-l

KN wo¡lz [Sooz-1

(s)

(6)

tsõ/-l tNo3-12

The nonequivalence of cr$ and K$ is clearly seen from a comparison of equations 5 and

6. Equation 6 may be rewritten in terms of equivalent sulfate ion fractionS, XS, to yield

an equilibrium constant on the molar scale.

Xs (t-Rlz cKN= (7)

(1 - Xs)2 xs C

where X5 and X5 are respectively the equivalent fractions of sulfate in solution and

sulfate in the resin phase, C is the total anion concentration in the resin phase in

meq dm-3, more generally known as the capacity, and C is the total anion concentration

in solution, also in meq dm-3. While the separation factor is often a practically useful

parameter in the study of selectivity, a more valuable parameter, particularly for

monovalent-divalent equilibria is the selectivity coefficient, f$. ttre disadvantages of

separation factors are that they are independent of capacity and the total ionic strength of

the solution. The latter dependence is extremely important because at high anion

Introduction l0

concentrations a monovalent anion will always be preferentially absorbed by the resin due

to "elecüoselectivity" (16). This would occur at a total anion concentration above 52 meq

dm-3 (17). This effect can lead to false reports of nitrate selectivity. Thus, in generd, Kf

is a far more useful parameter since it allows comparison of resin efficiency between

groundwaters of different concentrations. A resin is said to be nitrate selective when K$

is large, but note that in equation 7 the totat anion concentration also effects the fraction of

sulfate bound. Consequently, at high concentrations nitrate will be preferred. This is the

cause of the "elecüoselectivity" described above.

In practice total anion concentrations in groundwaters rarely exceed 10 meq dm-3

and Guter (17) arbitrarily chose to define a nitrate selective resin as "a resin which in

column operation with common groundwaters retains nitrate as the last ion to break

through when exchanging ions at anionic strengths of 10 meq dm-3". He proposed that a

more convenient indicator of selectivity was NSS, the nitrate-to-sulfate selectivity,

defined as

NSS=logK$-logõ+1 (8)

where C was expressed in meq dm-3. A nitrate selective resin was then defined as a resin

for which NSS > 0. It is important to observe that equation 8 is derived directly from

equation 7 by assuming that the equivalent sulfate fraction ratio must be greater than 1 for

a nitrate selective resin and that the total anion concentration, C, is 10 meq dm-3. The

equation merely gives a convenient indicator of selectivity at anion concentrations more

typically found in groundwaters.

1.5.L Amidines.

Initial attempts to make nitrate selective resins involved the incorporation of known

nitrate precipitants within a supporting polymer matrix. The amidines have the general

structure given in Fig. I.2, and being strong bases can exist as ionized salts. The

amidinium salt is more stable as nitrate than either chloride, sulfate or bicarbonate (17b).

Inlroduclion

R-C-NHRtlN-R

HNO3 : .-i-Rr2R o NoP

N-R

l1

+

Fig.1.2. Amidines are basic and form stable amidinium salts with nitrate.

Grinstead and Jones (18) introduced alkyl substituted amidines into porous polymer

beads. The resultant resins had a very high selectivity for nitrate over sulfate, f$ = 16S-

106, but a very low capacity, less than 0.2 eq dm-3. The resins were also too nitrate

selective over chloride, K[ = 15 - 40, and could not be easily regenerated using chloride

salts. All such resins suffered because the amidine precipitant was not covalently bound

within the polymer matrix. Hence the capacity was low and the precipitant was frequently

lost via leaching to the elutant water.

Roubenik (19) was later able to incorporate the amidine group covalently within a

vinyl polymeric matrix cross-linked with divinylbenzene and so overcome the leaching

problem. It was found that as the length of the alkyl group increased nitrate selectivity

increased but regeneration efficiency decreased. The best balance between nitrate

selectivity and ease of regeneration was found when the amidins group contained 5 - 7

carbon atoms; with ethyl or butyl substituents being favoured. Kf values in the tange 7 -

10 were reported, but no K$ values were reported. Enquiries by Guter (17) led him to

believe that the resin was not commercially viable and it may be, like many resins, it

worked too well and could not be regenerated.

L.5.2 Salinogen Resins.

Salinogens are another general group of compounds which form insoluble salts

with nitrate. The best known example is 4,5-Dihydro-1,4-diphenyl-3,5-phenylimine-

1,2,4-tnazole which is more commonly known as Nitron, Fig. 1.3.

Intoduction 12

Fig. 1.3. Idealized structure of Nitron.

Chiou et al. (20) prepared Nitron-polyvinylbenzylchloride polymers by reacting

Nitron derivatives with preformed polymers of p-chloromethypolystyrene. The

introduction of Nitron produced a polymer apparently selective for nitrate as well as other

oxidizing anions. In addition, the polymer was readily and repeatedly regenerated with

ammonium chloride with no loss in capacity. However, the resulting polymers became

tacky upon contact with water requiring that the polymer be supported on either silica gel

or Chromosorb W for testing. This resulted in a decreased capacity. Guter (17c)

esrimared a capacity of only 0.33 meq dm-3 for the silica supported polymer. This is

significantly lower than the capacities of commercial ion-exchange resins; which are

usually greater than 1000 meq dm-3. It is difficult to evaluate the nitrate selectivity of this

material given the absence of detailed information on capacities, selectivity coefficients or

breakthrough curves, presumably due to the handling difficulties.

Similar studies by Walitt and Jones (21) which incorporated salinogen into a cross-

linked polystyrene matrix produced resins which could not be regenerated. Although no

K$ values were reported, Kf selectivity coefficients tended to increase with the size of

the alkyl group substituted on the polystyrene making them commercially unacceptable.

Triazole is a simpler analogue of Nitron. Hauptmann et al. (22) prepated a nitrate

selective resin by attaching 3-amino-I,2,4-îiazole to a preformed polystyrene-

divinybenzene resin. The capacity was 2.3 meqg-l and the separation factors were cr$ =

2 - 3 and oÈ = 5 - 10. This meant rhat while the ¡esin was selective for nitrate over

sulfate, it was also more selective for nitrate over chloride and could not readily be

regenerated using chloride salts.

N

N

N

Intoduction

1.5.3 Commercial Resins.

13

Prior to 1980 the majority of nitrate selective resins were produced only in small

quantities by research laboratories. Although a large number of commercially available

resins already existed, the majority of these resins did not exhibit any preference for

nitrate. However, some selectivity for nitrate over sulfate had been observed for a small

number of these resins when the sulfate loading was high (16,22-25).

Clifford and Weber (23) reviewed the selectivity of thirty commercial anion

exchange resins. Their work is of particular interest because it applied statistical

methodology to a fairly large sample size of the then available resins. Selectivity was

tested for correlation with a variety of resin properties such as matrix, functionality,

porosity, capacity, type and pK¿. They determined that the two most important resin

properties influencing sulfate/nitrate selectivity were functionality and matrix. More

generally they proposed that the primary determinant of divalent/monovalent selectivity

was the distance of fixed-charge separation in the resin'

The uptake of a divalent anion, such as sulfate, requires the presence of two closely

spaced positive charges and indeed all such resins fulf,rlling this criteria were found to be

sulfate selective. While separation of fixed charges enhanced nitrate selectivity. They

found that polystyrene based resins were more nitrate selective than non-polystyrene

based resins and that generally nitrate selectivity was favoured by attaching amine groups

to pendant alms rather than within the polymer chain. A selection of some of the more

nitrate selective resins is given in Table 1.2.

The steric hindrance produced by either functional groups of large size or number

also prevented the close proximity of fixed charges. Thus nitrate selectivity increased in

the order,

1o 1-NH2R) 1Zo (-NHRz) ç 3o (-NR3) < 40 (-NR4+)

Nitrate selectivity was also found to increase with the hydrophobic nature of the resin and

consequently Type I resins were more selective than Type II resins, shown in Fig. 1.4.

Table 1.2. Selected examples of commercial polystyrene based quaternary ammoniumanion exchange resins and their characteristics.

Trade Name Porosity Marix Functionality . Kä a! KN NSS Ref.S

Amberlite IRA-900 Macroporous PS-DVB Type I 1.101.00.9

Amberlite IRA-400 Microporous PS-DVB Type I 1.53r.33

DuoliteA-101D Improved PS-DVB Type I

Amberlite IRA-910 Macroporous PS-DVB Type II

3.4t33a

2.9 a

0.58 23t725

60a90 0.0

0.532.4 87

25 -0.1r

-23-0.2 l'1,26

23t7

t.321.30

1.311.0

0.39

0.31

0.42

2317

AmberliteIRA-410 Microporous

Imac HP555 -

a-

b-

PS-DVB

PS-DVB

Type II

Er3 b

1.40

1.0

36a

23

3000 1.5 25

average approximate values calculated from graphs in reference 17.

Et3 indicates a resin having a structure simila¡ to a Type I resin, but possessing Ethyl instead of Methylsubstituents

Introduction 15

CHr

loR-N-CH2CH2OHt"

CH¡

Fig. 1.4. Structure of the two common Types of commercial ion-exchange resins, where

Type II

R denotes the resin backbone

For the majority of resins studied, porosity was not correlated with selectivity.

However, for the Type I resins it was found that the macroscopic resins, having a higher

degree of crosslinking, were more nitrate selective than the less cross-linked isoporous

resins. Clifford and Weber did not attribute this to a "sieve effect" because the sulfate

anion was not significantly hydrated (23). They attributed the greater sulfate selectivity of

the less cross-linked isoporous gel resins due to the greater flexibility of the polymer

chain, allowing freer movement of the quaternary groups to "pair-up" with divalent

sulfate ions. In all other cases porosity was uncorrelated with selectivity. Similarily,

capacity and pK¿ had little or no effect on nitrate selectivity.

It is imporranr ro note that all of the resins surveyed by Clifford and'Weber had

average sepamtion factors, of, 1".. than one, and were consequently all sulfate selective.

Guter (17) also surveyed a number of commercial resins for their nitrate selectivity. He

supported the major observations of Clifford and V/eber, including the idea that the nitrate

ion became more favoured as the resin became more hydrophobic. He also concluded

that, at the time of writing his final report, the commercially available ion-exchange resins

were not selective enough to be of significant value (17d).

The review of Clifford and Weber (23) was most significant in that it inspired other

resea¡chers to study polystyrene based resins containing quaternary ammonium functional

groups. The variation of selectivity with the size of the attached functional group will be

considered more closely in section 1.5.5.

Recently (25), another cornmercial resin, Imac HP555, has become available. This

resin, which is believed to contain ethyl substituents, is far more selective than any

Introduction 16

previous commercial resins. Its release is a natural consequence of continued research in

this f,reld.

1.5.4 Polyallylamine Resins.

Bolto et al. (27) have prepared cross-linked resins derived from diallylamine by

copolymerization of alkyldiallylamines with the crosslinking agent, 1,6-bis(N,N-

diallylamino)hexane. The resins have the generalized structure given in Fig. 1.5. They

were examined for their nitrate selectivity over common ions present in groundwater. A

complete compilation of results is presented in Appendix 40.

For these diallylamine derivatives, the separation factor "!, yré "alculated as a

ratio of separation factors, of to cr[, which were available in the literature (27). The

nitrate to sulfate separation factors indicated an increased preference for nitrate over

sulfate as the size of the alkyl substituent increased. The calculated separation factors

were however far too small to be of any commercial use. The nitrate to chloride

separation factors exhibited the same trend but the preference for nitrate indicated that

resin regeneration with chloride salts would be inefficient.

It is diff,rcult to evaluate the potential of this type of resin for nitrate removal because

Bolto et al. (27) incorporated the produced microparticles into a cross-linked

polyvinyl alcohol matrix. It is unclear what effect the supporting matrix will have on

selectivity although comparison of this PTAA resin with a later one which was not

incorporated into a PVA matrix suggested that the PVA matrix substantially decreases the

selectivity. It is also not possible to attribute the increase in nitrate selectivity solely due to

an increase in size of the alkyl substituent because the degree of crosslinking, and hence

the hydrophobicity of the resin, also increased with the size.

The effect of quatemization upon the nitrate selectivity of polyallylamine resins was

considered by Jackson et al. (25,26a). In their studies the weak base PTAA resin was

parrially quaternized with dodecyl bromide to various degrees before being fully

quaternized with MeI. Their complete results tabulated in Appendix A0 indicated that the

strong base (quaternized) resins were more nitrate selective than the original weak base

ùrþ

nn n o

Clo oclo NoClo N clo

MeMe

N

l\M"¡l Me

( (

p2

crO clo oO

n p

PTAA (Rl = R2 = H) Poly (QIIEXA) Poly (DADMAC - QHEXA)

Fig. 1.5. Idealized structures of polyallylamine resins. PTAA = Polyriallyamine; QHEXA = Quatemar! (Me) FIEXA;DADMAC = Diallydimethylammonium chloride; HEXA = 1,6, dibromohexane.

N

n

Introduction 18

resin. It was also observed that the preference for nitrate increased with the

hydrophobicity of the resin up to about 607o euaternization with dodecyl bromide.

However, the increase in selectivity was accompanied by a significant loss in capacity by

the resin for nitrate. Additionally, no nitrate to chloride separation factors were reported in

these studies.

Polymers of diallydimethylammonium chloride (DADMAC) have found a variety of

applications (28). Two DADMAC-QHEXA copolymers were prepared by crosslinking

DADMAC with fully methyl quaternized HEXA (26). The resulting copolymers were

sulfate selective and of low capacity. The low capacity was most likely due to the

presence of the bulky QHEXA fraction.

The homopolymers of 1,6-bis(N,N,N-diallymethylammonium)hexane dichloride

(QHEXA) illusrrared in Fig. 1.5. were found to be more nitrate selective than the above

copolymers. The degree of crosslinking for both the homopolymer and copolymer was

estimated as the percentage of solid crosslinking monomer, QHEXA, present prior to

polymerization. For both polymers the preference for nitrate over sulfate was found to

increase with the degree of crosslinking. Jackson et al. (26b) attributed this to the inability

of a highly cross-linked resin to reorientate its binding sites and bind "bidentately" to the

divalent sulfate anion.

Given that the selectivity of the PTAA resin derivatives were shown to increase

with resin hydrophobicity, it was surprising that neither the QHEXA or DADMAC type

resins were quaternized with any substituent larger than a methyl group. Indeed only

methyl derivatives have been considered in detail and the potential of poly(allylamine)

type resins for water purification remains unfulfilled for want of a systematic study. To

this end a preliminary synthesis of DADMAC monomer analogues containing larger alþl

substituents is presented in Appendix 41.

1.5.5 Polystyrene-divinylbenzene Resins.

The review of commercial resin selectivities given in section 1.5.3 indicated that the

resin mafrix was one of the primary determinants of nitrate selectivity. It was observed

Introduction 19

that resins containing a polystyrene-divinylbenzene (PS-DVB) matrix were the most

nitrate selective. This realization led many researchers to concentrate on changing the

functionality while maintaining the PS-DVB matrix.

Jackson et aI. (25,26,29) prepared weak base resins of the type illustrated in

Fig. 1.6. by amination of chloromethylated polystyrene beads (gel type - Permutit).

Their results indicated that selectivity for nitrate over sulfate increased with the

hydrophobic narure of the functional group. Although no K[ were reported the resins

could be readily regenerated with NaOH. However, the resins were not completely

regenerable and Jackson and Bolto (25) found that they contained a large percentage of

strong base groups which undoubtedly contributed to the observed selectivity.

The selectivity of polystyrene quaternary ammonium resins (Fig. 1.7.) have been

extensively studied (16,29,30). Selected selectivity coefficients have been tabulated in

Appendix 40. It was found that, in general, as the size of the alkyl group attached to the

nitrogen increases so does the preference for nitrate. The nitrate selectivity was greatest

for the butyl derivative.

Replacement of one alkyl substituent with the more hydrophilic group,

-CH2CH2OH, resulted in a lower preference for nitrate. This confirmed that the

selectivity for nitrate was related to the hydrophobicity of the resin.

A charged hydrophilic resin will swell by absorption of water to a limit determined

by crosslinking. Swelling will continue until the internal strain equals and opposes further

hydration. Since sulfate is highly hydrated, it contributes to swelling strain. Guter (17e)

suggested that the addition of the polar hydroxyethyl goups were able to contribute to the

stabilization of the fixed charge. Consequently, resins having hydroxyethyl groups

contained lower amounts of water normally required for this purpose and were therefore

able to accommodate more of the highly hydrated sulfate anions. The moisture content of

the resins studied by Guter did indeed decrease upon hydrophilic group substitution in

agreement with the above argument.

The increased preference for nitrate by higher alkyl derivatives has alternatively

been attributed to "water structure enforced ion pairing" (31).This th_eo-ry__was first¿.¿-

suggested by Diamond (32).It suggests that large polarizable anions enforce the stmcture

Inlroducrion 20

cH- c}lz

p1

cH- C}lz

n

p2

of the structure of the weakbase resins possessing a

one. No attempt to represent crosslinking due to theas been made.

n

aternary ¿tmmonrum resmsattempt to represent crosslinking

Intoduction 2l

of water in the vicinity of the anion and cause the anion to associate more readily with a

given large nonpolar cation. This theory would favour ion pair formation of nitrate over

sulfate and would predict an increased preference for nitrate as the alkyl group of the

quaternary ammonium cation became larger and more hydrophobic. This is in perfect

agreement with the observed trends.

Subramonion and Clifford (33) have also suggested that ion-solvent interactions

play an imporrant role in determining selectivity. They found that the unhydrated ionic

radii of monovalent anions were highly correlated with their affinity for Type I resins.

This seems reasonable since larger monovalent anions have a smaller charge density and

tend to be less hydrated. Larger anions therefore disrupt the structure of water and are

rejected by the water phase resulting in a greater preference for the anion by the resin

phase. These phenomena with specific regard to the the benzyltrialklyammonium cation

will be discussed in greater detail in Chapter 3.

Although the capacities of the strong base resins are consistently lower than their

weak base analogues they are much higher than any previously measured resin type.

They also exhibit a substantial increase in selectivity over the corresponding weak base

resin. This trend is followed only up to the butyl derivative and the nitrate selectivity of

strong base pentyl resins are much the same as that of their weak base analogues. None

of the strong base resins studied by Jackson and Bolto (25) could be regenerated

effectively. Although the weak base resins a.re the less selective resins, they will find

some application simply because they can be regenerated, and a.re therefore more cost

effective. It is a common problem that resins that are strongly nitrate selective are

precisely the resins which are difficult to regenerate by conventional means.

Barron and Fritz (31,34) studied the selectivity of low capacity PS-DVB resins for

both monovalent and divalent anions using single anion chromatography. Their results

are tabulated in Appendix 40. In principle it should be possible to obtain estimates of

selectivity coefficients from such retention times. Similar calculations have been

performed by Haldna (35). In practice such calculations required specific information on

the experimental conditions and could not be performed. However, the general trends

observed are in agreement with previous measurements on high capacity resins. That is,

Introduction 22

retention rimes, and hence selectivity for nitrate, increased with alkyl length. Once again,

substitution of hydroxyethyl groups decreased the preference for nitrate. An increase in

capacity was also shown to amplify the selectivity effect.

V/arth and Fritz (36) have also used column chromatography to examine the effect

of spacer arm groups on anion selectivity. Low capacity anion exchange resins of the type

given in Fig. 1.8. were examined.

(PS-DVB)

CHc

I

CH-(CHÐ"-N(CH3)3 Cl

Fig. 1.8. Idealized structure of the polystyrene-divinylbenzene spacer aImtrimethylammonium resin, SATMA (n=n).

Relative retenrion times (Appendix A0) indicated that as the length of the spacer ann

increased the resins became less selective for nitrate, while the preference for sulfate

remained essentially unchanged. The theory of "water-structure enforced ion pairing" was

again used to interpret the results. It was suggested that the close proximity of the benzyl

and ammonium functional groups in the conventional resin promoted the sffucture of

water and was responsible for its enhanced selectivity. The introduction of spacer arm

groups increased the distance between these two functional groups and hindered this

effect.

Recently, Lockeridge (37) has reported selectivity coefficients for both

benzytrialkylammonium and benzyltrialkyphosphonium salts. His results, presented in

Table 1.3. are considered separately because they utilize a different concentration scale

and a¡e therefore not strictly comparable with previous tesults, although exactly the same

trends in selectivity are maintained. As the size of the alkyl chain increases the selectivity

for nitrate over sulfate increases for all the phosphonium and ammonium salts considered.

The phosphonium salts are all considerably more selective for nitrate than the comparable

ammonium salts. Lockeridge attributed this to increased steric hindrance at the anion-

exchange site. Thus while nitrate could readily approach the exchange site, the approach

of highly solvated anions, such as sulfate, was hindered. Lockeridge also demonsffated

Introduction 23

the need of one sulfate anion to occupy two exchange sites on the resin regardless of the

size of the alkyl substituent.

Table 1.3. Selectivity coefhcients of polystyrene-divinylbenzene resins

of quaternary ammonium and phosphonium functionality. a

Resin KNC(meq/g)

-NMe3-NEt3-NPr3-NBug

-PMe3-PEt3-PPr3-PBu3-PPe3

3.6r2.39r.52r.23

2.222.7 |r.991.68r.27

6009000

26700102000

200019500

120000660000840000

a- reference 37, determined from mlvVg and mlWml.

Warth et al. (38) were the first to study anion selectivity of low capacity

tributylphosphonium resins. It was this work that inspired Lockeridge to consider higher

capacity phosphonium resins for ion-exchange. As observed previously, the larger

polarizable anions had consistently higher retention times. The results are compiled in

Appendix 40. The increased preference of the phosphonium resin was interpreted in

terrns of "water enforced ion pairing". This was to be expected since the phosphonium

moiety being larger and more polarizable would interact more strongly than would the

quaternary ammonium group.

1.6 Conclusion.

There is no conclusive evidence to suggest that there is any real need for a nitrate

limit to reduce the risk of health effects. However, these limits have been imposed world

wide and must be legally upheld, requiring the need for a method of nitrate removal.

Introduction 2/+

Of all the methods currently available the dual ion-exchange and bacterial reduction

A4'e''of nitratefperhaps the best. Current ion-exchange methods suffer because of high costs

associated with regeneration and low capacities of ion-exchangers for nitrate. It is a

general aim that by studying the basic principles that govern selective ion-exchange some

insight may be gained into the process which will allow the design of a commercial ion-

exchange resin that overcomes these problems.

Many studies of ion-exchange phenomena have been conducted on resins prepared

by the introduction of the desired functional group into a preformed polymer network.

The disadvantage of this method is that the homogeneity of the resulting polymers may

vary. It is not uncommon that the characteristics of a particular commercial resin will vary

not only from one manufacturer to another, but also from one resin batch to another even

though it be produced by the same manufacturer. In this work, the selectivity of some

quaternary ammonium resins will be studied for polymers prepared directly from the

derivative monomers.

Few previous studies have been thoroughly systematic. This study will be

systematic in that the selectivity of monomer analogues, their linear polymers and finally

lightly cross-linked resins wilt be examined in turn. The role of selectivity will be

explored with an effort to produce highly nitrate selective polymers while maintaining a

high capacity for nitrate.

Þv'c

Introduction

Literature Cited.

1. C. J. Johnson and B. C. Kross, Amer. J. Ind. Med., L8, 449, (1990).

2. Nitrates, Nitrites and N-nitoso Compounds.,Envtronmental Helath Criteria 5

(World Health Organization: Geneva 1978).

3. D. C. Bouchard, M. K. Williams and R. Y. Surampalli, /. Am. Water Works

Assoc., 84(9), 85, (1992).

5. T.R. Owen and S. Jürgens-Gschwind, Fert. Res.,10, 3, (1986).

6. R. V/. Raiswell, P. Brimblecombe, D. L. Dent and P. S. Liss, Environmental

Chemistry,17 (Edward Arnold Limited: London 1984).

7 . R. Chang, Chemisty 4th edn.,577 (McGraw-Hill: New York 1991).

8. Water Resources Inventory,Engineering and Water Supply Department Library

Reference 86/35, South Austrtalian Engineering and Water Supply Department,

(1e87).

9 . C. R. Lawrence, Nitrate-rich groundwaters of Australi¿ Technical Paper No. 79,

Australian Water Research Council, EWS Serials (Australian Government

Publishing: Canberra 1983).

10. South Australian Department of Housing and Construction (Aboriginal Works

Unit), private communications, ( 1989).

1 1. R. B. Mellor, J. Ronnenberg, W. H. Campbell and S. Diekmann, Nature,355,

7t7, (1992).

12. J. Liessens, R. Germonpré, S, Beernaert and W. Verstraete, J. Am' Water Works

Assoc., 85(4), I44, (1993).

13. J. Liessens, R. Germonpré, I. Kersters, S, Beernaert and W. Verstraete, J. Am.

Water Works Assoc., 85(4), 155, (1993).

74. J. P. van der Hoek and A. Klapwjik, Wat. Supply,6, 57, (1988).

15. D. Clifford, S. Subramonian and T. J. Sorg, Environ. Scí.Technol., 20(lL),

t072, (1986).

16. F. G. Helffench,Ion Exchange,156 (McGraw-Hill: New York 1962).

25

Introduction 26

17 . G. A. Gurer, Removal of Nitate from ContaminatedWater Supplies for Public

tJse,EPA-600/2-82-042 (March 1982). b 90, c 95, d 10, e 133

18. R. R. Grinstead and K. C. Jones,Nitrate RemovalfromWastewaters by lon

Exchange,EPA report 17010FSJ01/7 I (January 197 l).

19. L. Roubinek, U.S. Patent 4,134861(16 January 1979).

20. S. J. Chiou, T. Gran, C. E. Meloan and W. C. Dann, Anal. Lett., 14(411), 865,

(1e81).

21. A. L. V/alitt and H. L. Jones, Basic Salinogen Ion-exchange Resins for Selective

Nitrate Removal From Potable and Eff'luentWaters, U.S. EPA, Cincinnati,

Advanced Waste Treatment Laboratory, L970, U.S. GPO, Washington, DC.

22. R. Hauptmann, P. Froelich, H. Weber, Ger. (East) DD 217,528 (16 January

1985); Chem. Abstr., 103,124434, (1985).

23. D. Clifford and V/. J. Weber, React. Polym.,1,77, (1983).

24. M. Cox, R. C. Haries, D. V. Nowell and R. cla¡k, chem. Ind.,16l, (1981).

25. M. B. Jackson and B. A. Bolto, React. Polym', L2,2'77, (1990).

26. M. B. Jackson and L. J. Vickers, Effect of structure on nitrate selectiviry. Part 2.

Resins with either sryrenelDVB, PECH or allylamine backbones., CSIRO Report

No. DB-127, (April 1987)., a 10, b 13,

27 . B. A. Bolto, M. B. Jackson and R. V. Siudak, Isr' J. Chem.,26, 77, (1985).

28. B. Bolto, Chem. Aust.,61(8), 431, (1994).

29 . M. B. Jackson, Effect of structure on nitrate selectiviry. Part I. Resins with

undecanoic groups., CSIRO Report No. DAIB-I11, (August 1986).

30. G. A. Guter, U. S. Paten¡.4,479,877 (30 October 1984).

3I. R. E. Barron and J. S. Fritz, J. Chromatogr.,284, 73, (1984)'

32. R. M. Diamond, J. Phys. Chem., 67,2513, (1963).

33. S. Subramonian and D. Clifford, J' Sol. Chem.,18(6), 529, (1989).

34. R. E. Barron and J. S. Fritz, J. Chromatogr.,316,20l, (1984).

35. Ü. Haldna, J. Chromatogr.,604,282, (1992)-

36. L. M. Warth and J. S. Fritz, J. Chromatogr. 5ci.,26, 630, (1988).

37. J. E. Lockeridge, Ph.D. Dissertion, Iowa State University, (1990).

Introduction

38. L. M. Vy'arth, R. S. Cooper and J. S. Fritz, J. Chromatogr., 479, 401, (1989).

39. G. Solt, Chem. Ûng.,436,33, (1982).

40. J. P. van der Hoek, W. F. van der Hoek and A. Klapwjk, Water, Air and Soil

P ollut., 37,41, ( 1988).

41. J. L. Cox, R. T. Halten and M. A. Lilga, Envíron. Sci. Technol., 28,423,

(19e4).

27

28

2

Synthetic Methods and Charac terization.

2.1 General Experimental.

All solvents used were purified using standard methods (1,2).'Water was

obtained from a Milli-Q purification system and had a resistance greater than 15 MQ.

Melting points were recorded on a Gallenkamp melting point apparatus and are

uncorrected. The melting points of the vinyl monomers could not be determined

accurately due to their hygroscopic nature. Even after prolonged drying, the proton

spectra of some vinyl monomers still showed a singlet around 4'8 ppm due to H2O.

1g NVIR were either recorded on a Jeol JNM-PMX60 or Gemini 200

spectrometer. 13C NMR were either recorded on a Bruker CXP-300 spectrometer

operaring ú 75.47 MHz or on a Gemini 200 spectrometer at 50.289 MHz. Unless

otherwise stated 13C spectral shifts were taken relative to the deuterated solvent used or to

an external trimethylsilane (TMS) standard. lH spectral shifts were taken relative to TMS.

The 13C specrra of the linear polymers required acquisitions of at least 13 hrs on the

Gemini to obtain reasonable resolution. Fast atom bombardment mass spectra (FAB)

were recorded on a Vacuum Generators ZAB 2Íß spectrometer. Ultraviolet spectra were

recorded on a Cary 2200 spectrophotometer.

2.2 Monomeric Analogues.

The monomeric analogues were either available commercially or were prepared by

the reaction of benzyl chloride with the appropriate alkyl amine (Scheme I). The reaction

conditions were similar to those described previously (3).

Synthetic Methods

cH2cl-NR3Cl+R¡N:

Scheme I.

Benzyltrimethylammon¡um chloride - was available commercially from Fluka as

'purum' grade material and was recrystallized from acetone/ether and vacuum dried.

m.p. = 239 oC dec. Gravimetric analysis as silver chloride indicated a purity of 100.15 +

O.05 7o. The molar absorptivity at the maximum of 262.2 nm was 427 + 2 crfl mmol-l.

Benzyltriethylammonium chloride - was available commercially from Fluka as


m.p.= 185 0C dec. lHNMR: õ(CDCI¡) 7.57-7.40 (m, 5H, ArH); 4.7\ (s,2H,

ATCHzN); 3.38 (q, 6H, NCHz);1.44 (t,9H, CH:). 13C NMR: õ(CDCI¡) 131.53 (s, Ar

C-1); 129.53 (s, Ar ortho); 128.30 (s, Ar meta); 126.45 (s, Ar para); 59.97 (s'

ATCHzN); 51.84 (s, NCH2);7.39 (s, CH3). Gravimetric analysis as silver chloride

indicated a purity of 99.80 + 0.16 7o. The molar absorptivity at the maximum of 262.4

nm was 312 + 2 cmL mmol-I.

Benzyttripropylammonium chloride - Tripropylamine (10.53 9,74 mmol) and

benzyi chloride (12.03 g, 95 mmol) were refluxed for 46 hrs at 70 oC with constant

stirring. The resultant orange solid was dissolved in an acetone/ethanol mixture and was

twice recrystallized using anhydrous ether to yield 1 1.60 g (59Vo) of fluffy white product.

m.p. = I94-L95 oC, lit. 189-192 oC (4).Microanalysis for CroHzsNCl, Found: C,70.7;

H, 10.3; N,5.1. Calc: C,7L.2;H,10.4; N,5.2Vo.lHNMR: ô(CDCI:) 7.49-7.46(m,

5H, ArH); 4.82 (s,2H, ATCHzN); 3.25-3.21 (bt t, 6H, NCHz); 1.89 (br m, 6H, C}lù;

1.00 (t, 9H, CH3). 13C NMR: õ(cDCl¡) 131.85 (s, Ar C-1); 130.04 (s, Ar ortho);

128.71(s, Ar meta); 126.87 (s, Ar para); 62.03 (s, ATCH2N); 59.27 (s, NCHz); L5.45

(s, CHZ); 10.08 (s, CH3). Mass spectrum m/2234 (C1OHZSN+). Gravimetric analysis as

29

Synthetic Methods 30

silver chloride indicated a purity of 100.38 + 0.07 7o. The molar absorptivity at the

maximum of 262.6nm was 366+2 cmL mmol-l.

Benzyltributylammonium chloride - was available commercially from Fluka as


m.p.= 164-165 oC, lit. 185 oC (5). Gravimetric analysis as silver chloride indicated a

purity of 99.82+ 0.18 Vo.The molar absorptivity at the maximum o1262.6 nm was 362

-f2 cm2 mmol-l.

Benzyltripentylammonium chloride - Tripentylamine (29.25 g, I29 mmol) and

benzylchlonde (20.75 g , L64 mmol) were refluxed at 69 oC for 19.5 hrs with constant

stirring. The orange solution was dissolved in acetone and reprecipitated with anhydrous

ether. The resultant white solid was recrystallized twice more from acetone/ether to yield

11.55 g (25Vo) of fine white powder. m.p. = 140-142 oc, lit. 141 oc (6). Microanalysis

for CzzH¿,gNCl, Found: C,74.6; H, 11.7; N, 4.0. Calc: C,74'6: H, 11'4; N,4'0Vo'

lU NUR ô(CDClg) 7.58-7.42 (m, 5H, ArH); 4.92 (s,2H, ATCHzN); 3.39-3.31 (br t,

6H, NCHz); 1.S0 (br s, 6H, NCH2CIFIù;1.34 (br s, 12H, CIJù;0'90 (t, 9H' CH3)'

13C NMR: õ(CDCI¡) 131.88 (s, Ar C-L); 129.93 (s, Ar ortho); 128.57 (s, Ar meta);

127.O7 (s, Ar para); 62.49 (s, ATCH2N); 58.09 (s, NCH2);27'75 (s, CHz); 2L'66 (s'

CIclù; 21.5I (s, CH2); 13.19 (s, CH3). Mass spectrum mlz 318 (CzzH+oN+)'

Gravimetric analysis as silver chloride indicated a purity of 100.165 + 0.007 7o.The

molar absorptivity at the maximum of 262.6nm was 363 + 2 cm2 mmol-l.

Benzyltrihexylammonium chtoride - Trihexylamine (22.26 g, 83 mmol) and

benzylchloride (14.59 g, 115 mmol) were dissolved in acetone (10 ml). The mixture was

refluxed for 6 hrs to yield a ginger brown solution. Anhydrous ether was added in excess

to precipitate a solid which was recrystallized twice from acetone/ether to yield 5.10 g

(16To) of pure white solid. m.p.= 111-112oC. Microanalysis for CZSH¿ONCI, Found:

C,74.0; H, 11.9; N,3.4. Calc: C,75.8; H, 11.7; N, 3.57o.lHNtr'tR õ(CDCI¡) 7'43-

7.56 (m,5H, ArH);4.95 (s,2H, ATCHzN); 3.3t-3.39 (br t,6H, NCHz); 1'80 (brs,

Synthetic Methods 3l

6H, NCHzCEù;1.34 (br s, 18H, ClHù 0.91 (t, 9H, CH3). I3CNMR: õ(CDCI¡)

131.98 (s, Ar C-1); 130.03 (s, Ar ortho): 128.66 (s, Ar meta);127.16 (s, Ar para);

62.62 (s, ATCHzN); 58.22 (s, NCH2); 30.60 (s, CH2); 25.47 (s, CHz); 22.03 (s' CH2);

2I.82 (s, CHz); 13.27 (s, CH3). Mass spectrum mlz 360 (CzSH¿ON+). Gravimetric

analysis as silver chloride indicated a purity of 100.45 + 0.05 Vo.

Benzyltrialkylammonium nitrates - The chloride salts were converted to the

corresponding nitrate salts using B.D.H. Amberlite IRA-400 (Cl) ion-exchange resin

which had been converted to the nitrate form by overnight exchange with 2M KNO3. The

purity of the nitrate salts was confirmed using UV absorbance measurements. EMF

measurements using a nitrate selective electrode confirmed the UV results. The two

higher alkyl derivatives, benzyltripentylammonium and benzyltrihexylammonium nitrate,

were sufficiently insoluble that they could be prepared by precipitation from a chloride

solution using potassium nitrate. The solid was washed with wate and recrystallized

twice from acetone/ether.

Benzyltrialkylammonium sulfates - All of these salts were prepared using the ion-

exchange techniques as described for the preparation of the nitrates. The resin was

converted to the sulfate form using lM NazSO¿ in a batch process'

Benzyltrialkylammonium bromides - The majority of the bromide salts were

available commercially and were reprecipitated from acetone/ether in a manner similar to

that of the chlorides. Bromides that were not available commercially were prepared in a

manner similar to that of the least soluble nitrates, that is; by precipitation from a chloride

solution by addition of saturated NaBr.

Benzyltrialkylammonium iodides - The iodides were either available commercially

and were recrystallized in a manner similar to that of the chlorides, or were precipitated

from a chloride solution by addition of saturated NaI.


Recently, several benzyltriethylammonium salts were prepared by heating

benzyltriethylammonium chloride with the appropriate silver salt under reflux (7). Should

this method be applicable to other benzyltrialkylammonium salts it would offer

advantages in time and general applicability over the resin methods described above.

2.3 Linear Polymers.

The linear polymers were prepared in two steps. Initially monomers were

prepared in a manner similar to that described for the monomeric analogues and purified

by repeated reprecipitation. The monomers were then polymenzed in aqueous solution

using 2,2'-azo-bis-isobutyronitrile (AIBN) as an initiator as previously described (8).

The complete reaction is illustrated in Scheme II. Vinylbenzylchloride was obtained from

Kodak as an isomeric mixture of para and meta isomers and was inhibited with

t-butylcatechol.

CH CHz CH CHz

+ R¡N

cH2cl

-CH"t'NClI

R3

CH"

t"NClI

R3

Scheme II.

Synlhetic Methods

2.3.I Monomers.

Vinylbenzyltrimethylammonium chloride - Vinylbenzylchloride (28.04 g, 184

mmol) was added dropwise to a rapidly stirred aqueous solution of trimethylamine 25Vo

(w/w) (54.53 g,231mmol). The initially yellow solution was refluxed at3l-42 oC for

23 hrs, during which time it steadily da¡kened to yield a burgundy solution. The solution

was washed with ether (4 x 50 ml) and the lower layer evaporated over silica gel to yield

a dark burgundy gel-like material. This gel was recrystallized with heating from

ethanol/ether to yield 16.75 g (43EùI of crude product. The product was further purified

by repeated reprecipitation from ethanol/ether and vacuum dried for 7 hrs at 50 oC.

lH NMR: ô(CDEOD) 7.59-7.45 (m, 4H, ArH); 6.82-5.32 (m, 3H, ArCH=CHz); 4.60,

4.56 (s, 2H, ATCHzN); 3.77, 3.16 (s, 9H, NCH3). Mass spectrum m/z 176

(CrzHtsN+).

Vinylbenzyltriethylammonium chloride - Triethylamine (45.10 9,446 mmol) was

dissolved in methanol (30 ml) and vinylbenzylchloride (44.82 g,294 mmol) was added

dropwise to the rapidly stirred solution. The mixture was heated under reflux for 38 hrs at

38 oC. The solution darkened on heating. The cooled solution was poured into excess

ether. The lower oily yellow layer was vacuum dried at 45 oC to yield 58.02 g QTVo) of

crude beige product. The sample was purihed by recrystallization from 2-propanol/ether.

Microanalysis for CtsHz¿NCl-H2O, Found: C,66.3; H,9'6; N, 5.2. Calc: C, 61.2;H,

9.3; N, 5.17o. lHNMR: ô(DzO) 7.57-7.27 (m, 4H, ArH); 6.79-5.29 (m, 3H,

ArCH=CH);4.23 (s,2H, ATCHzN); 3.09 (q,2H, NCHù; l-29 (t,9H, CHs).

l3CNMR: ð(DzO) t43.89, 142.80 (s, ArCH=CHz); 140.11, 140.18 (s, Ar C-1);

137.86-137.16 (m, Ar); 120.81, 120.20 (s, ArCH=-Hù; 64.95, 64.22 (s, ATCHzN);

56.77 (s, NCHz); 11.55 (s, CHg). Mass spectrummlz2lS (CrsH24N+).

T ,l yi"l¿ of 84Vo has been reporred by doubling the reflux time and diluting the aqueous 257o

trimethylamine in methanol (9).

JJ


Vinylbenzyltripropylammonium chloride - Vinylbenzylchloride (25.38 E,

166 mmol) was dissolved in methanol (30 ml) and added dropwise to tripropylamine

(28.14 g,I97 mmol) with rapid stirring. The solution was heated under reflux for 24 hrs

at 30 oC with constant stirring to produce a yellow solution. The cooled solution was

poured into a large excess of ether and the lower layer was further washed with ether (2 x

50 ml) and dried over silica gel until a jetty-like consistency was obtained. This sample

was reprecipitated from ethanol/ether to give an off white solid which was washed briefly

with ether and dried over silica gel under nitrogen to yield 7.3L g (15Vo) of crude

material. The sample was further purif,red by recrystallization from ethanol/ether before

use. lH NMR: ô(cD:oD): 7.64-1.31 (m, 4H, ArH); 6.85-5.36 (m, 3H, ArcH=cHz);

4.56,4.54 (s, 2H, ATCHzN); 3.19-3.10 (m, 6H, NCHz); 1.51 (m,2H, CHz); 1.06 (t,

9H, CH¡). Mass spectrumm/z 260 (CrsH3oN+).

Vinylbenzyltributylammonium chloride - Vinylbenzylchloride (10.00 g, 66 mmol)

in AR MeOH (10m1) was added dropwise to a vigorously stirred solution of tributylamine

(1636 g, 90 mmol) in MeOH (10m1) over thr. The mixture was stirred at ambient

temperatures for 24 hrs. The solution was then heated under reflux at 32 oC for 169 hrs.

The cooled solution was poured into excess ether (300m1). The lower yellow layer was

recovered by rotary evaporation. The resultant yellow gel was dissolved in 2-propanol

and reprecipitated using AR ether. The solid was repeatedly washed with AR ether,

vacuum dried and finally dried over phosphorous pentoxide to constant weight to yield

1 1.16 e 607o) of solid producr. lH NMR: ð(CD¡OD) 7 .61-7 .42 (m, 4H, ArH); 6.87-

5.33 (m, 3H, ArCH=CHz); 4.58, 4.55 (s, 2H, ATCHzN); 3.23-3.15 (br t, 6H, NCHz);

1.83 (br m,6H, CHù;1.41 (m,6H, CH2); 1.04 (t,9H, CH3). I3CNMR: õ(CD¡OD)

I4l.4g, 140.20 (s, ArÇH=CIF¡ù; 137.07, 136.98 (s, Ar C-1); 133.90-128.06 (m, Ar);

1,16.64, 115.96 (s, ArCH=ÇHz); 63.18, 62.9I (s, ATCH2N); 59.52 (s, NCH2); 25.09

(s, CHz); 20.71(s, CHz); 13.98 (s, CH3). Mass spectrumm/2302 (CztH36N+).

Vinylbenzyltripentylammonium chloride - Tripentylamine (55.13 g, 242 mmol)

was dissolved in methanol (35 ml) and vinylbenzylchloride (23.92 g, I57 mmol) was

Synrheüc Merhods 35

added dropwise to the rapidly stirred solution. The yellow solution was heated under

reflux for 60 hrs at 30 oC, then heated at 50 oC for a further 2 hrs. The cooled solution

was poured inro excess ether. The lower yellow layer was vacuum dried at 45 oC to yield

40.20 g (68Vo) of crude beige product. The product was purified by repeated

recrystallizarion from acetone/ether. lH NMR: ô(CD¡OD) 7.65-7.38 (m, 4H, ArH);

6.88-5.28 (m, 3H, ArCH=CH2); 4.59,4.56 (s, 2H, ATCHzN); 3.25-3.12 (br t, 6H,

NCHz); 1.85 (br m, 6H, CHz); 1.43 (br m, 12H, CHz); 0.99 (t, 9H, CH3). 13C NMR:

õ(cD¡oD) 141.32, 140.11 (s, AreH=CHz)i 136.94, 136.84 (s, Ar C-1); 133.70-

127 .96 (m, Ar); 116.53, 115.84 (s, ArCH=ÇHù; 63.O7 , 62.79 (s, ATCHzN); 59.44 (s'

NCHz); 29.39 (s, CHz); 23.16 (s, CH2); 22.76; (s, CHz); 14.20 (s, CH3). Mass

spectrum mlz 344 (CZ¿H¿zN+).

2.3.2 Polymerization.

Free radical polymerization of the vinyl monomers was typically conducted in

aqueous-ethanol mixtures at 70 oC using the initiator AIBN. The reaction conditions are

briefly summarized in Table 2.1.

Table 2.1. Conditions for the polymerization of linearpoly(vinylbenzyltrialkylammonium chlorides).a

Polymer Cn'ono'n.. Solvent Temp

(mol dm-3) (oc)Time(hrs.)

LPMELPETLPPRLPBULPPE

0.330.310.790.880.36

WO /EIOHIl2orbo IEIOH}lzO IEIOHIlzO /MeOH

7010706070

5.54.59.0

13.511.0

a - Initiator: AIBN

Polymer purifications were conducted at room temperature and under a nitrogen

atmosphere as required. In general, a change in physical characteristics of the product

when compared to that of the corresponding monomer was considered sufficient to


confirm polymerization. This was in part due to the high purity of the monomers used.

For some polymers, nmr was also used to confirm the purity of the products. For

example, the absence of the vinyl carbons in the 13C spectra clearly indicated complete

polymerization. Likewise lH nmr spectra indicated that the polymers contained negligible

numbers of double bonds. The nmr of such compounds are discussed in more detail in

section 2.5.

Poly(vinylbenzyltrimethylammonium chloride) - Trimethylvinylbenzylchloride

(3.10 g, 15 mmol) and AIBN (0.04 g, 0.24 mmol) were dissolved in a water/ethanol

mixture (40m1/5ml) and were heated under reflux for 5.5 hrs at 70 oC. The solution was

cooled to room temperature and precipitated into excess acetone. The white precipitate

was washed with acetone and dried under vacuum for 4 hrs at 50 oC to yield 0.85 g

(27Vo) of crude product. The product was further purified by repeated recrystallization

from erhanol/ether. lH NMR: õ(DzO) 7.94 (br s, ArH); 6.64 (br s, ArH); 4.25 (br s,

ATCHzN); 2.82 (br s, NCH¡); 0.98 (br s, backbone -CH-CH2-).

Poly(vinylbenzyltriethylammonium chloride) - Triethylvinylbenzylammonium

chloride (2.1O g,5.52 mmol) and AIBN (0.02 g, 0.14 mmol) were dissolved in water

(18 ml) and heated under reflux for 4.5 hrs at 70 oC. The resultant solution was

evaporated under vacuum and the solid reprecipitated with heating from acetone/ether to

yield 0.46 g (22Vo) of crude product. The product was further recrystallized from

acetone/ether to obtain a fluffy white solid which was vacuum dried before use. lH

NMR: ô(DzO) 7.05 ( br s, ArH);6.64 (br s, ArH); 4.10 (br s, ATCH2N);2.88 (br s,

NCHz); 1.50 (br s, backbone -CH-CHz); l.l2 (br s, CH3). 13C NMR: õ(DZO) 150.62

(br m, Ar C-1); 133.70 (br m, Ar C-2,C-3,C-5,C-6); I3I.7l (s, Ar C-Ð; 64.45 (br m,

ATCHZN);56.92 (s, NCH2); 48.55 (br m, CH backbone); 45.58; (br m, CH2

backbone); 11.83 (s, CHg).

Poty(vinylbenzyltripropylammonium chloride) -

Tripropylvinylbenzylammonium chloride (10.21 g, 38 mmol) and AIBN (0.10 g' 0.62


mmol) were dissolved in water/ethanol (40 ml/8 ml) and heated under reflux for t hrs at

70 oC. The resultant solution was reprecipitated into dry ether and washed with acetone.

The product was recrystallized from acetone/ether and vacuum dried at 50 oC for 6 hours

to yield 7 .96 g (787o) of snowy white product.

Poly(vinylbenzyltributylammonium chloride) - Tributylvinylbenzylammonium

chloride (14.17 g, 42 mmol) and AIBN (0.13 g, 0.78 mmol) were dissolved in

water/methanol (40 ml/8 ml) and heated under reflux for 13.5 hrs at 60 oC. The cooled

solution was precipitated into, and washed with, excess acetone. The final product was

recrystallized from ethanol/ether and vacuum dried for 6 hours at 50 oC to yield 10.02 g

(7l7o) of white product. The product was later observed to have a greenish tinge.

Microanalysis for CzrHgoNCl-2HzO, Found: C,67.9; H, 10.1; N, 3.9. CaLc: C,67-4:

H, 10.8; N, 3.7 Vo. rH NMR: õ(DzO) 7.16 (br s, ArH); 6.61 (br s, ArH); 4.10 (br s,

ATCHzN); 2.93 (br s, NCHz); 1.60 (br s, NCH2CHz); 1.19 (br s, CHz); 0.83 (br s'

CH¡). 13C NMR: õ(DZO) 149.81 (br m, Ar C-1); 134.2I (br m, Ar C-2,C-3,C-5,C-6);

131..92 (s, Ar C-Ð; 65.74 (v br m, ATCH2N); 62.29 (s, NCH2); 48.75 (br m, CH

backbone); 45.39; (br m, CH2 backb one); 27 .56 (br s, CH2); 23.92 (s, CH2); 17.70 (s,

cH¡).

Poly(vinylbenzyltripentylammonium chloride) - Tripentylvinylbenzylammonium

chloride (13.66 g,36 mmol) and AIBN (0.11 g, 0.67 mmol) were dissolved in

warer/methanol (100 ml) and heated under reflux for 11 hrs at 70 oC. The cooled solution

was precipitated into acetone (200 ml) and washed twice with acetone. The final product

was recrystallized from ethanol/ether and vacuum dried at 50 oC for 5 hours to yield

8.33 g (6lEo) of a fluffy white solid.

2.4 Cross-Linked Polymers.

Cross-linked polymers were prepared using procedures simila¡ to those described

for the linear polymers, except that methanol was used as a solvent and divinylbenzene


was used as a cross-linker. Divinylbenzene was supplied by Aldrich as a 55Vo solution

inhibited with p-tertbutylcatechol. The inhibitor was removed by washing into the basic

layer with a I}Vo sodium hydroxide solution, and rinsing with water to remove traces of

the base. A typical preparation is given below, while a summary of the reaction

conditions for all preparations appear in Table 2.2. The final product was purified by

repeated swelling in methanol. As with the linear polymers, the change in the physical

properties of the products; most notably their insolubility in water, was taken as evidence

that a cross-linked polymer had been produced. The extent of cross-linking was estimated

from the mole ratio of DVB to polymer.

Cross-linked poly(vinylbenzyltriethylammonium chloride) - AIBN (0.01 g'

0.09 mmol) was added to a solution of triethylvinylbenzylammonium chloride (0.34 g,

1.33 mmol) in methanol (3.20 g). A solution of DVB (0.02 g, 0.13 mmol) in methanol

was then added and the sealed vial left to stand overnight. A white solid was noted and

the sample was evaporated in air. Solvent was completely removed with mild heating.

The resultant solid was purified by repeated swelling in methanol to give a waxlike solid

in almost quantitative yields.

Table 2.2.Polymerization conditions for the preparation of cross-

linked poly(vinylbenzyltrialkylammonium chlorides).a

Sample Crono,n". CnnN(mol dm-3) (mol dm-3)

Crvs(mol dm-3)

Vo CL

CME-1CET-1CET-2CET-3CET-4CET-5CET-6

0.26690.1 8050.29290.253r0.21970.12040.2786

0.01850.01210.01960.01670.01540.00770.0206

0.01130.01700.01060.01450.01730.02600.0116

4.29.43.85.77.9

2r.64.2

a - solvent was methanol, initiator was AIBN and cross-linkerwas DVB. The nominal degree of cross-linlcrng, To CL, wasdefined as the mole ratio of DVB to monomer.

CME = cross-linked poly(vinylbenzyltrimethylammonium chloride),CPE = cross-linked þoly(vinylbenzyltriethylammonium chloride).

Synthetic Methods

2.5 Interpretation of NMR.

The characterization of the monomer analogues by nmr was fairly straight

forwa¡d. In contrast, the spectra of the vinyl monomers were slightly more difficult to

interpret due to the presence of two isomeric products. This resulted directly from the use

of an isomeric mixture of vinylbenzylchloride in the initial preparation. Representative lH

and 13C nmr specrra of the ethyl derivative are given in Fig. 2.I. andFig.2.2. The ethyl

derivative was chosen to illustrate the general results because this allowed clearer

assignment in proton spectra without the clutter which arose from the longer alkyl

resonances.

39

The proton spectra of the vinyl monomers clearly indicates that the three vinyl

protons are nonequivalent. What was previously simply assigned as a multiplet of relative

area three in section 2.3.L., is actually a series of twelve visible lines. This was attributed

to an ABC type spectra. The coupling constants were assigned as Jab = 11.0 Hz , Jbc =

5¡¡z,Iac= 17.7 Hz,andwereusedtogeneratethesplittingpatterninFig.2.l.,which

closely resembled that of the experimental spectra. The possibility of additional long

range coupling between geminal vinyl protons and those of ATCH2N has not been ruled

out.

In the 13C nm¡ spectra, Fig.2.2., doublets for the methine and methylene carbons

at approximately 142 and 120 ppm respectively, were attributed to the presence of para

and, meta isomers in solution. This was also true of the carbon corresponding to the

ATCH2N group around 64 ppm. This assumption was confirmed by the presence of eight

distinct aromatic carbon resonances in addition to that of the C-1 resonance at 140.18

ppm. The remaining low field resonances were assigned to alkyl carbons.

The lH nrm spectra of the linear polymers were sufficiently broad that integration

of the peak area was prone to a large uncertainty and no area assignment was

consequently made. Resolution into distinct isomeric components was also not observed

because of this broadening. Representative polymer spectra with peak assignment are

given in Fig. 2.3. andFig.2.4.

oa b

H

C

H

d

ocbl e c

þH,-cH,¡$-

fH. ÌL Hb

e

t7.7 Hz

6 lt0tlz 5Hz

c b

a

PPN

Fig. 2.1. Representative 200 MHz lH nmr spectra of benzylvinyltriethylammonium chloride. Recorded at arnbient temperature in D2O. The assignment ofcoupling constants is discussed in the text.

'7

cL p p

CHz

2b

CHz

2

c[

4b

1

6b

5b5

2

J

{cHr-cHr)67

3b

-R1.",-cHl:47b 67

I

N

5

@

6

)

2,3,4,2b,3b,4b,5b, 6b

p

5,7b

G

Fig.2.z.Representative l3c broad band decoupled nmr spectra (50 MHz) of benzylvinyltriethylammonium chloride' Recorded at ambient temperature ln-piO. fn. tp..ru is complicated due to the presence of meta andparc isomers'

PPHr60

ab

-{ cHr)-n

CH_oE

c¡ndd candd

e

{ctr-cHJ ¡fg" " " "'t " " " "'J iñ o

f

c,d ârb

Fig.2.3.Representative 200 MHz lH nmr spectra of poty(benzylvinyltriethylammonium chloride). Recorded at ambient tempemtule in DzO' The inset is

th; lH specra of poly(benzylvinyltrimethylammonium chloride) recorded under identical conditions.

e

'1

-{ß

ör,)-n

c[

cH-I

4

2

J

,)

5

{cHr-cHJ,67

6

4ct p

Fig. 2.4. Represenrative l3C broad band decoupled nmr spectm (50 MHz) of poly(benzylvinyltriethylammonium chloride). Recorded at ambienttemperature in D2O.

5

PPII

Synthetic Merhods 44

The broad proron resonances in Fig. 2.3. near 7.0 and 6.6 ppm are in the ratio of

approximarcIy 3:2. Bovey (10) had suggested that in the case of polystyrene these two

peaks corresponded to meta-para and ortho protons respectively. Although the situation is

likely to be more complicated for the disubstituted ring considered here, the observed

resolution indicates that the polymer contains two distinct types of aromatic protons.

Although, it is difficult to say with any certainty how exactly this arises. If, as suggested

by the spectra of the vinyl monomers, only meta and para isomers are present, a total of

eight hydrogens, the 3:2 ratio suggests that these isomers cannot be present in equal

amounts. Although not conclusive, the unequal relative magnitudes of the isomeric

resonances in the 13C spectra of the vinyl monomers also suggested that one isomer was

more prevalent. However, due to the uncertainty in the integrals it was not possibly to

unambiguously identify which of these isomers was predominate. In the spectra of the

ethyl derivative the backbone resonances near 1.5 ppm were partially obscured by the

broad methyl resonance at 1.1 ppm. For the methyl derivative, the inset in Fig. 2.3., the

methyl protons which are now deshielded due to their closer proximity to the nitrogen are

shifted further downfield and the backbone resonances at 1.4 ppm can clearly be

resolved. This in conjunction with the absence of any peaks attributable to vinyl

resonances indicated complete polymerization. For the higher alkyl derivatives the

backbone resonances became totally obscured by those of the alkane resonances in the

functional group.

A represenrative 13C spect.a of the polyethyl derivative is given in Fig. 2.4. Once

again, there are no resonances attributable to unreacted vinyl groups. The backbone

resonances at 48.6 and 45.6 ppm were attributed to cr and B protons respectively. The

assignment of the spectra is in agreement with that proposed by Ford (11) for the spectra

of polystyrene.

Synthetic Methods

Literature cited.

11.

D. D. Perrin, V/. L. F. Armarego and D. R. Perrin, Purification of Laboratory

Chemicals 2nd edn. (Pergamon Press: Oxford 1980).

A. I. Vogel,Vogel's Textbook of Practical Organic Chemistry 4th edn.

(Longman: London 1978).

B. J. Steel, A. S. Kayaalp, T. Kurucsev, A. D. Ward, and M. B. Jackson, Aust'

f . Chem., 43, 1983, (1990).

M. Makosza, and B. Serefin, Rocz. Chem.,39, L233, (1965); Chem. Abstr., 64,

t2595, (1966).

Dictionary of Organic Compounds. 5th edn., (New York: Chapman andHall

1990).

A. S. Kayaalp, unpublished results.

D. A. Scola and N. A. Edelson, J. Chem. Eng. Data,13(3)' 453' (1968).

T. Ikeda, S. Tazuke, and Y. Suzuki, Makromol Chem',185, 869, (1984).

P. Guarilloff, unpublished results.

F. A. Bovey,HíghResolutionNMR of Macromolecules ChapterVI, 118

(Academic Press: New York L972).

W. T. Ford and T. Balalrishnan, Macromolecules,14,284, (1981).

45

1

2

J

4

5

6

7

8

9

10

46

3

Solubility, Enthalpy and EntroPY.

3.1 Introduction.

The main reason for studying the solubility of the benzyltrialkylammonium salts

was to estimate, using the van't-Hoff isochore (1), the enthalpy and entropy of solution

and thereby gain some insight into the structure of the corresponding aqueous solutions.

This method has been chosen in preference to calorimetric determination because of its

simplicity in that it does not require specialized equipment even though, admittedly, it is

not as accurate, in general, as calorimetric measurements. However, in this particular case

we are concerned chiefly with the trends in enthalpy and entropy for a series of alkyl

groups rather than with the absolute values and a wider tolerance was considered,

therefore, acceptable.

When solutes are dissolved in water the measured thermodynamic properties are

usually interpreted in how the solute alters the structure of water in its vicinity. For this

reason, I shall first review the current thoughts about the structure of water in the pure

state and the effect that simple solutes have on this bulk structure.

3.2 The Structure of Water.

Surprisingly, some conÍoversy still exists over the bulk sructure of water. While it

is generally agreed that the structure of water possesses some "icelikeness", especially at

low temperature undþr.rrure, and that heating causes some disruption to this structure,

neither the exact nature of the disruption nor the exact model that best describes liquid

water is known with any certaintY.

b

Solubility 47

The best conjectures, reviewed by Franks (2), suggest that liquid water is indeed

substantially icelike; leading to the proposition that each water molecule is tetrahedrally

hydrogen bonded to four others and favouring the interstitial models to describe liquid

water. The latter assumes that in some portions of water there exist free water molecules

encapsulated by a highly sffuctured water lattice, exemplified by, the cluster or clathrate

models. It is further suggested that heating causes the breakdown of the lattice structure

by either bending or breaking hydrogen bonds between the water molecules.

3.2.1 Flickering Cluster Model.

According to this model there are short lived liquid clusters of extensively hydrogen

bonded molecules interspersed between non hydrogen bonded water molecules. The half-

life of these clusters is thought to be about 10-11 s (3) so that only a single peak is

observed in nmr studies. It is thought that an equilibrium between clustered and non

clustered water is set up,

(HzO) free ê (HzO) clusrered

such that upon addition of solute the equilibrium is shifted either to the left or to the right

depending on the nature of the solute.

3.2.2 Structure Breaking and Structure Making by lons.

In the case of simple electrolyte systems hydration can be considered to be divided

into three distinct regions which are roughly spherical and concentric about an ion,

(Fig. 3.1.) for example, an alkali metal ion such as lithium.

In the inner sphere the water molecules are oriented by the electric field of the ion

through ion-dipole interactions. This effect, known as hydration of the first kind,

becomes less important as the size of the ion increases. The region furthest from the ion

remains uninfluenced by the presence of the ion and maintains the normal bulk structure

\-

Fig. 3.1. Schematic representation of the three hydration regions around a simple ion.

The ovaloids in region 4., which are radially orientated about the central ion, represent

water molecules posessing dipole moments.

Solubility 50

This effect is equivalent to hydrophobic hydration, structure promotion, water

structure enforced ion pairing, iceberg formation, hydration of the second kind or the

clathrate effect, where the molecule is completely encased in a cage of hydrogen bonded

water.

Note that although there is no direct evidence for iceberg or clathrate cages around

solutes in solution, it seems plausible, since solid clathrates are already known to exist

(s).

Regardless of the details of the various theories currently being advanced ions in

solution can be generally classified into three distinct categories.

a) Structure Breakers which disrupt the structure of water around the ion.

b) Electrostrictive Stucture Makers which order the water molecules around the

ion due to the ions electric field via ion-dipole interactions.

c) Hydrophobic Stucture Makers which induce more hydrogen bonding in water

near a non polar surface of the ion. This effect increases with the size of the hydrophobic

portion and usually results in large negative enthalpies and entropies of soluúon as well as

a positive heat capacity.

It is to this last group that the majority of the tetraalkylammonium salts belong and

where we shall direct our interest in the discussion of the benzyltrialkylammonium salts.

3.2.4 Theory of Solution.

Solution may be considered to occur in a two step dissolution, solvation process.

Initiaily the bonds between the ions are broken and the compound enters the aqueous

phase. This process is endothermic since energy is required to break bonds,

MXls¡ (â MXiaql AH1> 0

In solution bonds will be formed between water molecuies and the ions, that is hydration

occurs. This process is exothermic as bonds are formed,

Solubility

ND(1aq¡ (+ M+(aq) + X-1aq) AH2< 0

The enthalpy of solution will be the sum of these two processes

AHsol= AHI +L¡12 (1)

and could be either positive or negative depending on the relative magnitudes of

enthalpies from the two contributing processes.

The overall equilibrium expression for solution is

ND{¡s¡ <+M+(aq) +XlaÐ

where the equilibrium constant K is,

l(= a- (2)

and ap¡ç is equal to 1 since it is in its standard state.

K = Ír+ v_= y+2 52

AGo = -RT ln K = -RT ln (S2yt2 )

= -2RT ln (Syt)

= aHo - TASo

AHO ASO=) ln Slt =¿RT *-ZR

51

(3)

Thus it follows that from a plot of lnSyt versus 1Æ the slope can be used to

estimate AHo and the intercept ASo.

Solubility 52

The main assumption made in this treatment is that the enthalpy of solution is

temperature independent. This is generally not the case. The enthalpy of solution is

related to its heat capacity (6) using,

T2

AH(Tz)=AH(T')+ JCPdTT1

(4)

where Cp is the heat capacity at constant pressure and AH(T2) and AH(Tr) are the

enthalpies of solution at the corresponding temperatures T2 and T1. However, provided

the temperature change is fairly small, AH may be taken to be essentially constant.

3.3 Water-Solute Interactions.

3.3.1 Alkanes.

Due to their structural simplicity; the alkane series proves a useful starting point for

the discussion of water-solute interactions. Indeed, much work has been done on their

solubility and thermodynamics in aqueous media. Recently the thermodynamics have

been reviewed by Abraham (7) and the data reported there have been used to produce

F|g.3.2.

The Gibbs free energy of solution increases with the increasing length of the alkyl

group, as might be expected from the increasingly unfavorable solubility. The enthalpy,

however, is large and negative and becomes more so as the alkyl group increases. The

positive Gibbs free energy, therefore results from an unfavorable entropy contribution.

The entropy of solution becomes more negative as the alkyl group increases. This has

been explained by proposing that the lower alkanes form clathrate-like hydrates in

solution. This decreases the due to the increased order around the

hydrocarbon, relative to pure water. This effect increases as the size of the alkane group

increases presumably due to the larger area of contact between solute and water.

The decrease in entropy has also been explained by suggesting that with increased

alkane length the formation of "flickering clusters" becomes more dominant. This may be

AG

osol

(kJ

mol

-1).

Iu

t9

N)

b,J

f..)

(,(-

¡t

O\

{ O

O

\O

O

5 o\N)

I 1..)

I u)I Þ

oo

I b..) l'.)

I þ Þ

oo

AS

osol

(kJ

mol

-l;.

tttll

9VV

VV

þL;L

;O

coO

\Þb.

)

aHos

ol (

kJ m

ot-1

¡

N) Þ o\

N) è o\

æ

tr

TD

o

Þtr

o

()tr

tr

Sotubility 54

a more appropriate explanation since larger hydrocarbons may be too large to fit within

the available clathrate cavities. The "flickering cluster" model involves an equilibrium

which is shifted by the presence of solute, either promoting or disrupting hydrogen

bonding. As such, it does not require a clath¡ate cavity of a given size to accommodate the

particular salt under consideration. However, the possibility of partial enclosure of the

alkane by the cavity still exists.

3.3.2 Tetraalkylammonium Salts.

It is now of interest to discuss trends in the tetraalkylammonium salts; because they

differ only from the benzyltrialkylammonium salts studied here, by the replacement of

one alkyl group with a benzyl $oup. The benzyl group is typically hydrocarbon-like and

is very nearly the same size as a methyl group such that the replacement is not expected to

alter the physical properties signihcantly. This assumption may be erroneous, but there

are no similar compounds reported as extensively as the tetraalkylammonium salts, from

which comparisons with the benzyltrialkylammonium compounds can be drawn.

Solubility.

The solubility of the tetraalkylammonium salts are unusual because, unlike

hydrocarbons, they do not exhibit a monotonic decrease in solubility as the size of the

alkyl substituent increases. Their solubilities have been summarized in a recent review

(8), where the nitrate solubilities were reported for the first time. For a given anion the

solubility changes in a zig-zag fashion as the size of the alkyl substituent increases, and

decreases sharply for pentyl and higher groups. For a given cation the solubility

decreases as the size of the anion increases. At high alkyl length, normally hexyl and

above, a two phase liquid equilibrium is formed, especially at elevated temperatures.

Solubility

Enthalpy and Entropy of Soluti<¡n.

No complete reviews have been written conceming the thermodynamic properties

of the tetraalkylammonium salts, but because of their relevance to this work a review of

such properties seemed warranted. All of the relevant data and calculations used to

produce figures have been summarized in Appendix 43, and only the major trends shall

be discussed here.

Halides.

The standard enthalpies and entropies of solution of the halides of the

tetraalkylammonium cations are shown in Figures 3.3. and 3.4., respectively, as

functions of the number of ca¡bon atoms in the alkyl chain. The initial trend for all three

halides shown is a decrease in both enthalpy and entropy of solution with increasing size

of the alkyl group. Such a trend may be interpreted as reflecting an increase in water

structure with increasing hydrocarbon content of the cations due to the hydrophobic

effect. There is a resultant decrease in enthalpy since formation of hydrogen bonds in

developing the clathrate-like water structures around the alkyl residues is an exothermic

process and a decrease in entropy due to the resultant increase in order.

However, this trend persists only with the chlorides up to the pentyl group. For the

bromides ÂHsel reaches a minimum at the butyl derivative and at the propyl derivative for

the iodides. These observations imply a decrease in hydrophobic effect with increasing

chain length and cannot be explained with any certainty.

It may be that the larger alkyl groups are unable to be accommodated fully within

the clathrate-like cages of water and so perturb cage formation leading to the higher

observed enthalpies of solution. Another possibility is that the long alkyl chains are

"curled" somewhat so that they appear to have the thermodynamic characteristics of a

shorter alkyl chain. Such "curling" is known to occur in the alkane series, but usually

only for twenty or more ca¡bon atoms.

55

60

40

20oE

hxoØ

oÉ{ 0

-20

-401 2 J 4 5 6 7

n

Fig. 3.3. Standard enthalpy of solution, ÀHosol GJ mol-l¡ versus n, the number of

carbon atoms in the aþl chain of the corresponding tetraalkylammonium salt.

tr Chlorides Â Bromides o lodides

0

oÊ

J¿

oøoU)

100

0

-1000 1 2 3 4 5 6 7

n

Fig. 3.4. Standard entropy of solution, ASosol ftJ mol-l) versus n, the number of ca¡bon

atoms in the alþl chain of the corresponding tetraalkylammonium salt.

tr Chlorides Â Bromides o lodides

Sotubility 58

The enrropy of solution data for the halide salts may be regarded as supporting the

increase in the hydrophobic effect with increasing chain length. In general, for the

majority of tetraalkylammonium salts considered here the entropy of solution becomes

smaller as rhe size of the alkyl group increases (Fig. 3.a.). This is reasonable since the

more hydrophobic the compound becomes the more unfavorable the solution process, a

positive contribution to AGo when ASo is negative.

Two of the larger salts, tetrapentylammonium bromide and tetrabutylammonium

iodide show an unexpected increase in the entropy of solution. This may be attributed to

the "curling" of long alkane chains previously mentioned. It cannot be attributed to the

formation of clathrate-like structures a¡ound the compound because this would cause a

decrease in entropy. It is of interest that the chloride salt which does show such a

decrease in entropy is the only salt not expected to form clathrate hydrates in solution

because the chloride ion is too hydrated (9).

The anion contribution to the enthalpy and enffopy of solution has not been

extensively considered in the literature. For the tetraalkylammonium halides, at least, it

appears that the enthalpy of solution increases with the size of the anion (Fig. 3.3.). This

is essentially also tn¡e of the entropy of solution (Fig. 3.a.) but to a lesser extent. Simple

halide salts which possess fairly large cations also exhibit the same trend in enthalpies.

This could readily be attributed to an increase in "structure breaker" chalacter of the

anion. However, if the major contribution to the enthalpy is due to the hydrophobic

cation, the presence of the anion should only have a minor effect. It is surprising then that

such changes are not swamped by the contribution from the hydrophobic cation, unless,

cooperative cation-anion interactions are important in these solutions.

Nitrates.

Enthalpies of solution of tetraalkylammonium nitrates have been reported by

Nakayama et al. (8). However, we have been unable to reproduce values quoted by these

authors from their solubility data either by the use of a linear plot of ln m5s1 versus 1Æ or

by fitting of the data to a quadratic and using the tangent to this curvo to estimate AHsot.

Sotubitity 59

Because of the relevance to this work, the reported solubilities of the tetraalkylammonium

nitrates were refitted to the van't Hoff relation to obtain the estimates of AHo and ASo in

Table 3.1.

For several reasons, the results obtained by this treatment should be viewed with a

high degree of caution. Firstly, the solubility data are available only in molalities and no

density data are available to allow conversion to the molarity scale as required for the

van't Hoff relation. Secondly, no activity coeff,rcients a¡e available from the literature and

hence no allowance for them has been made. Thirdly, and most importantly, the van't

Hoff relation holds soluble salts and the lower alkyl derivatives a¡e far

too soluble to provide reliable estimates of AHo and ÄSo. However, in the absence of

more reliable data this treatment provides the best approximation of the enthalpy and

entropy of solution of tetraalkylammonium nitrates yet available.

Table 3.1. Enthalpy and entropy of solution of some teraalkylammonium salts. a

Compound AHosol (kJ mol-1) ÂSosol (J mol-1¡

Ref,rtted Nakayama et al. b Refitted

MeaN-NO3EgN-NO3PraN-NO3BuaN-NO3PeaN-NO3HeaN-NO3OcaN-NO3

PeaN-IHe4N-I

(0.4)(1.7)(1.7)(rr.2)(1s.1)(1.7)(r1.2)

(8.0)(28.3)

85.396.3

r29.9181.130.0

-48.5-r29.6

7.2-32.2

t5.714.426.237.3

6.59.69.7

25.627.5

13.010.515.5

(2.6)(6.e)(6.0)

(3e.1)(s1.6)(6.5)

(38.7)

(8.4)t7.438.3 (30. I )

a- Values in brackets denote 95Vo Conftdence Intervals determined usingCI = t x s, where s is the standard error of the fit and t is a tabulated functionof the degrees of freedom. The refitted values are discussed further in the text

reference 8.b-

Solubility 60

Initially the enthalpy of solution decreases; as might be expected from the formation

of clathrate-like cages around the molecules. The higher alkyl derivatives are also of

lower enthalpy than the methyl derivative. The propyl and butyl derivatives however,

exhibit an unusually high enthalpy of solution (Fig. 3.5.). This would seem to indicate

cage disruption with respect to the other salts. The nitrate ion is a strong "structure

breaker" and nitrate salts might be expected to exhibit similar trends to those of the

iodides since the iodide anion is also a strong "structure breaker". Comparison of

Fig. 3.3. with Fig.3.5., for example indicates that this is not the case. Nakayama

suggested that the tetrabutylammonium chloride, bromide and nitrate salts all form stable

clathrate-like hydrates (10,11). This is not supported by this treatment, since cage

formation should be exothermic, in agreement with the sign of the enthalpy of solution

obtained by Kudryavtseva et al. (12).

It is interesting to note that the enthalpy of solution of simple alkali nitrates are

always much more endothermic than the corresponding halides. It may be that there is

some special interaction between the cations and nitrate ions that results in their unusual

thermodynamic behaviour.

The recalculated entropies of solution appearing in Table 3.1., like the recalculated

enthalpies are unusual. As can be seen in Fig. 3.6. the entropy steadily increases from the

methyl to the butyl derivative where it is maximum before steadily decreasing for the

remaining salts. The lower entropy of the octyl and hexyl nitrates could be explained by

micelle formation as both these salts are oils. The lower entropy of the pentyl nitrate

relative to butyl could be attributed to clathrate cage formation. The initial increase in

entropy for the lower alkyl groups is as yet unexplained, but is in line with structure

breaking effects implied from enthalpy and contrary to accepted theory.

I

oE

g20oØo

Hr¿{

40

30

r0

08620 4

n

Fig 3.5. Standard enthalpy of solution, AHosol (kJ mol-l) versus n, the number ofca¡bon atoms in the alþl chain of the corresponding teraalþlammonium nitrate,

recalculated from van't Hoff plots using the solubility data of Nakayama et. al. (8).

I

oEhJ4

oo(t)

200

100

0

-100

-2000 862 4

n

Fig. 3.6. Standard entropy of solution,Asosol (kJ mol-l¡ versus n, the number ofcarbon atoms in the alkyl chain of the corresponding tetraalkylammonium nitrate,

recalculated from van't Hoff plots using the solubility data of Nakayama et al. (8).

Solubility

3.4 Methods and ExPerimental.

Saturated solutions of the desired compounds were obtained by the classical method

of shaking a solution in the presence of excess solute. A specially designed water bath

was constructed for this purpose with the temperature being controlled to within

0.001 0c.

A sample of the aqueous phase was withdrawn, diluted and analysed by UV

absorbance using the previously determined maxima and absorptivities (Table 3.2.) The

concentration of the original sample was then expressed in units of mol dm-3. No

allowance for density was made in these dilutions. For concentrated solutions this could

lead to a significant error. All UV measurements were performed on a Cary 2200

spectrophotometer thermostated to within 0.1 oC. There was no significant spectral

modification when the anion was changed. All dilutions were made by mass.

63

Table 3.2. V/avelength maxima and molar absorptivity ofbenzyltrialkylammonium chlorides.

Compound l,¡¡¿¡ç (nm) € (M-1cm-1¡

BzMe3N-ClBzEr3N-ClBzPr3N-ClBzBu3N-ClBzPqN-ClBzHe3N-Cl a

262.20262.35262.60262.60262.55262.60

427372366362363363

a-

The reported solubilities in Table 3.3. were the average of at least three

measurements at each of the six temperatures and in several cases the solubilities were

determined using duplicate or triplicate samples. The enthalpy and entropy of solution

were estimated from a linear fit to the van't Hoff relation after allowance for activity

coefficients.

The wavelength maxima and molar absorptivity werenot determinèd but were assumed to be similar to thehigher alkyl derivatives.

Solubility

3.4.L Assumptions Regarding Activity Coefficients.

The activity coefficients of the benzyltrialkylammonium salts have not been reported

in the literature. The activity coeff,rcients were estimated using,

- A./s (s)log yt = 1+Ba S

where the constants A and B have been tabulated as a function of temperature, and a is the

ion size parameter in cm (13). The ion size parameter of the nitrates, sulfates and

chlorides have previously been reported (14) up to and including the pentyl derivatives.

The ion size parameters of the bromides and iodides were estimated by adapting the

chloride ion parameter to allow for the increased size of the appropriate ion using ionic

radii of that ion. The hexyl parameters were estimated from a linear regression of the first

five analogues in the series. The a parameters used in this work have been summarized in

Table 3.4.

The error in the ion size pammeter will signifrcantly effect the reported lnSyt values

when this value is small, that is when the solubility is large. For sparingly soluble salts a

2O7o change in the a pammeter will cause less than aITo change in lnSy+.

However, even for the more soluble salts a 207o change in the a parameter will

cause only a I-3Vo change in the estimated values of AHo and ASo.

The ion size parameter is also assumed to be independent of temperature in the

range 25 - 50 oC.

3.5 Trends in the Benzyltrialkylammonium Salts.

Solubility

The trends in the solubility of the benzylriaikylammonium salts were similar to that

of the tetraalkylammonium compounds. It was of interest to note however, that even

though the benzyltrialkylammonium compounds had the benzyl group in place of the

&

Solubility

Table 3.3. Temperature dependence of molar solubilities of benzyltrialkylammonium salts.

compound 25oC 30 oc 35 oC 40 oc 45 oc 50 oc

65

BzPr3N-NO3BzPr3N-I

BzBu3N-BrBzBu3N-NO3BzBu3N-IBzBu3N-IBzBu3N-IBzBugN-SO¿

BzPe3N-BrBzPe3N-NO3BzPe3N-I

BzHe3N-BrBzHe3N-NO3BzHe3N-IBzHe3N-I

r.23590.0549

0.75342.48300.t6970.r7970.1749

0.16910.03320.0051

0.01710.00820.00230.0023

1.48230.0605

0.93s72.63780.r7940.r8730.18221.8581

0.15760.03580.0063

0.01800.00880.00270.0024

r.71060.0668

0.97632.67410.19050.19800.t9441.8918

0.14120.03950.0068

0.01940.00650.00270.0025

1.90030.0760

r.53262.67780.20590.2t280.20831.9972

0.13590.04430.0079

0.02160.00750.00280.0026

1.90490.0863

1.60412.74850.22070.22830.224r2.1525

0.13010.05050.0089

0.02330.01020.00320.0028

r.92970.0993

r.79042.70340.23470.24210.23732.3109

0.12760.05020.0102

0.02380.01060.00340.0030

Table 3.4. Ion size parameters of the benzyltrialkylammonium salts.

Atlq/l Chloride a Nirate a Sulphate a Bromide b Iodide b

MethylEthylPropylButylPentylHexyl

59160563t818828901

637827722903930t002

573s03568577569578

606620646833843916

630644670857867940

a - previously determined in reference 14. The units of a are inpicometers, (1 pm = 1 ¡ 16-12 -¡.

b - determined here, see text.

Solubitity 66

atkyl group, the benzyltrialkylammonium compounds solubility was not always less than

that of the corresponding tetraalkylammonium salts as might have been expected

(Table 3.3.).

Nakayama et al. (8) suggested that mru¡ would change monotonely with increasing

alkyl group length, since lrs and ¡r5o were expected in turn to change monotonely with the

alkyl group. lr5 and lrso are respectively, the chemical potentials of the solid salt and the

salt dissolved in water to the standard state. They have attributed any anomalous increase

in solubility to the formation of a clathrate-like cage structure around the sample. Under

this criterion and by examination of Fig. 3.7 . BzBu3NI and BzBu3NNO3 appear to form

clathrate-like compounds but apparently BzBu3NBr does not. A possible explanation may

be that in this case the bromide ion is too hydrated, but this seems unlikely since

tetraalkylammonium bromide is thought to form clathrate-like structures in water and in

section 3.3.2 no thermodynamic anomalies were observed for this salt. However, the

benzyltrialkylammonium bromide considered here has an inexplicably high entropy of

solution. The higher solubility therefore results from a large entropy contribution in this

case. This result highlights the anion dependence of some thermodynamic properties in

these salts and will be discussed in more detail in the next section.

Enthalpy and Entropy of Solution.

The enthalpies and entropies of solution were estimated from the slope and intercept

of a plot of lnSy+ versus 1Æ (Fig 3.8, 3.9, 3.10), using the van't-Hoff equation. The

plots were quite linear and the results are presented in Table 3.5. The linearity of the plots

confirms the assumption made previously that the enthalpy of solution is essentially

temperature independent.

In broad terms, for the samples considered here, the enthalpy of solution increased

with the size of the alkyt group. This is essentially the same trend found for the

tetraalkylammonium salts when the alkyl groups were equal to or larger than a propyl

group.

1

cîÉ'doÊ

U)äo(JJ

0

1

-3

-2

4 5

Fig. 3.7. l,og S (mol dr¡-3¡ of various benzyltrialkylammonium salts versus n, the

number of carbon atoms in the alkyl chain. Where S denotes the molar solubility at

25 oC; circles (o): iodides; filled circles (o): nitrares; triangles (A): bromides.

7632

+tU)

J

-2

-3

-4

-5

-6

-73.0 3.1 3.2 3.3 3.4

1000/r (K)

Fig. 3.8. Ln Sy+ versus 1Æ (K) for four benzyltrialkylammonium iodides. The solidlines represent a line of best frt obtained by linear regression.

c propyl

o buryl 1

tr pentyl

a hexyl 1

1

0

1

+tV)

J2

-3

-4

5

3.0 3.I 3.2 3.3 3.4

1000/T (K)

Fig. 3.9. Ln Sy+ versus lff (K) for three benzyltrialkylammonium bromides. The

solid lines represent a line of best ht obtained by linear regression.

o butyl

tr p€ntyl

a hexyl

o

1

+tc/)

J

0

1

a

-3

-4

-5

-63.0 3.r 3.2 J.J 3.4

1000/r (K)

Fig. 3.10. Ln Sy+ versus lÆ (K) for four benzyltrialkylammonium nirrates. The solidlines represent a line of best fit obtained by linear regression.

o propyl

o butyl

tr pentyl

a hexyl

A ô

Solubility 7l

Table 3.5. Enthalpy and Entropy of Solution of variousBenzyltrialkylammonium Salts. a

Compound AHsot (kJ mol-l) ÂSsot (J mol-l)

BzBugN-SO+

BzBu3N-BrBzPe3N-BrBzHe3N-Br

BzPr3N-NO3BzBu3N-NO3BzPegN-NOgBzHe3N-NO3

BzPr3N-IBzBu3N-I b

BzPe3N-IBzHe3N-I c

14 (6) 40 (2r)

56-1822

264

28t6

(20)(s)(4)

(15)(4)(7)

(36)

(4)(4)(s)(s)

176-95

4

872434

-30

666849-37

(48)(14)(22)

(1 17)

(14)(14)(r7)(1e)

(65)(18)(14)

35384l19

a-

b - average of three different samples.

c - average of two different samples.

The numbers in brackets are the 95Vo confidenceintervals, CI = t xs, where s is the standard error and tis a statistical parameter equal to 2.78 for 4 degrees offreedom.

Solubitity 72

The notable exceptions to the trend of increasing enthalpy with alþl group were the

hexyl compounds which all showed a decrease. The other two compounds which showed

a decrease were BzPe3NBr, which was an oil above 30 oC, and BzBu3NNO3 which was

a solid throughout the temperature range studied. The lower enthalpy of solution of the

BzBu3NNO3 could be attributed to clath¡ate-like cage formation by this sample.

In light of the unusually low enthalpy of solution of the BzBu3NNO3 sample the

enrhalpy of the divalent salt BzBu3NSO¿ was investigated. Even when compared to this

salt the enthalpy of the BzBu3NNO3 salt was unusually low. This further highlights the

unusual solution properties of this salt.

In general the entropy of solution decreased as the size of the alkyl substituent

increased. This was consistent with an increase in order of the solvent due to increased

hydrogen bonding. In the case of oils this could be attributed to micelle formation.

Again, as was the case for the enthalpy trends, the most apparent exceptions were

found for BzPe3NBr, the oil, and the hexyl compounds. The BzBu3NNO3 salt also had a

lower entropy than the propyl derivative as expected due to clathrate cage formation.

Both the enthalpy and entropy of solution exhibit an unexpected anion dependence.

For example, the unusually low entropy of solution for BzBu3NNO3 when compared to

the corresponding halides, and the larger entropy of solution of BzBu3NBr previously

mentioned. It was expected that, like the tetraalkylammonium halides, a series of salts

differing only by their anion would exhibit similar trends offset slightly due to the nature

of the anion. It was surprising that any anion contribution at all was observed since the

major contribution was expected to come from the hydrophobic cation (15). Any anomaly

in a monotone trend could possibly be attributed to a cooperative contribution from anion-

cation interactions.

unfortunately, conductance measurements suggested that there was no significant

ion pair formation in solution (14). However, in that study the concentrations were

necessarily restricted to dilute solutions, while in this study the solubility measurements

were conducted upon saturated solutions.

As an aqueous solution is concentrated there may come a point where there is a

significant contribution to the solution properties due to the cosphere overlap of two or

Solubitity 73

more of the structurally effected regions around the ions (4). Vaslow (16) suggested

cosphere overlap between two "structure making" ions (of low entropy) would result in a

net increase in entropy because a smaller volume of water would be effected, while

overlap of two "structure breaking" ions (of high entropy) would result in a net dec¡ease

in enropy for the same reason. The cosphere overlap of a "structure breaking" anion and

a "structure making" cation, as is most likely the case for BzBu3NBr, was not

considered. Following the argument above, cosphere overlap would most likely result in

a net increase in entropy due to the disruption of clathrate-like structure around the cation.

Since the overlap involves bond breakage the process would also be endothermic. Both a

large enthalpy and entropy of solution are observed for BzBu3NBr in agreement with the

above hypothesis. The reason that a similar effect is not observed for BzBu3NI is that in

this case the salt is sparingly soluble and cosphere overlap is unlikely.

Three other salts were sufficiently concentrated to possibly experience cosphere

overlap. Lacking further information on the salts to either side in the appropriate series , it

is not possible to conclusively decide whether BzBu3NSO4 or BzPTNNO3 support this

idea, but certainly the low enthalpy of solution for BzBu3NNO3 cannot be explained in

terrns of cation-anion cosphere overlap. However, a possible explanation may lie in

consideration of cosphere overlap between two "Structure making" cations.

Krishnan and Freidman (4) reported an enthalpy of solution for

Bu¡N(CH2)3NBu3BrZ (diBuZ+Br2), ÂH = -33.1 kJ mol-l, approximately four times

smaller than that of Bu4NBr, AH = -8.8 kJ mol-l, (Appendix A3). The ion diBu2+ is a

good model compound for two tetrabutylammonium ions experiencing cosphere overlap

and this result indicated that cooperative enhancement of enthalpy was possible. Wen and

Saito (17) have also suggested that at high concentrations, ions that promote cagelike

formations, link up and cause larger "flickering clusters" which stabilise the solution.

This is a possible explanation for the low observed enthalpy of solution of BzBu3NNO3.

The reason that BzBu3NBr does not also exhibit significant cation-cation interactions is

because in this case the Gibbs free energy of solution indicates that such interactions may

be hindered by greater ion pair formation than is observed in BzBu3NNO3.

Solubility 74

In Fig. 3.11. the Gibbs free energy of solution at 25 oç for the

benzyltrialkylammonium salts together with those of the tetraalkylammonium salts is

plotted versus the number of carbon atoms in the alkyl chain. Gibbs free energy can be

considered a measure of ion association and because of the way the equilibrium was

def,rned,

MXqs¡ (â M+1aq¡ + X-1aq), ^Go

- -RT ln K

an increase in ÂG implies an increase in ion association.

Although similar it is clear that BzBu3NBr has a slightly larger free energy than

BzBu3NNO3. Thus carion-cation interactions would be reduced in BzBu3NBr due to

competitive ion pair formation with bromide, whereas less ion pair formation in

BzBu3NNO3 would allow for greater cation-cation interactions.

The tetraalkylammonium chlorides depicted in Fig. 3.11. are consistently some 10 -

20 kJ lower in energy than the corresponding bromides. If this also holds tn¡e for the

benzyltrialkylammonium chlorides, it would imply that all these salts would have a lower

AGo than rhe nitrates and consequently all benzyltrialkylammonium nitrates would be

expected to pafiicipate in a greater number of ion pairs than the corresponding chloride

salts.

3.6 Conclusion.

V/hile the van't Hoff treatment of solubility data is prone to a large elror, the error

is largely systematic and comparison of results from a comrnon alkyl series should enable

the discovery of anomalous solutes even though the accuracy of the absolute enthalpies

and entropies is relatively low.

It is clear for the samples considered here that BzBu3NNO3 and BzBu3NI are

exceptional, indicating that their structure in water is likely to be very different from the

other salts. In particular, the anomalous solution behavior of the benzyltrialkylammonium

nitrate salts points to the formation of a clathrate-like hydrate or at the very least an

60

40

20

0

I

oE

J¿

oØoo

-20

-100 4

Fig. 3.11. Standard Gibbs free energy of solution, AGosol (kJ mol-l¡ versus n, the

number of ca¡bon atoms in the alþl chain. Graph A: tetraalkylammonium salts, Graph

B: benzyltrialkylammonium salts. Filled circles (o): nitrates; triangles (Â): bromides;

circles (o): iodides; squares (o): chlorides.

30

20

l0

0

862

n

Solubility 76

increase in the ordering of water around this salt compared to the other salts. This would

also seem to be true for BzBu3NI but to a lesser extent. In addition, it can be clearly seen

from Fig. 3.11. that both the tetrabutylammonium and benzyltributylammonium nitrates

have lower than expected free energies of solution which suggested that at high

concentrations BzBu3NNO3 may participate specifîcally in cation-cation interactions.

The high salt concentrations used for some of the compounds studied here could be

comparable to that experienced at the polymer-anion interface of an ion-exchange resins.

In such resins the quaternary ammonium groups are physically restrained close to each

other along the resin backbone and so cooperative structural enhancement of the

surrounding water seems likely. This effect may pafiially explain why polystyrene resins

possessing quaternary ammonium groups have different ion-exchange properties to the

majority of commercial resins.

The enthalpy of solution can be considered as arising from contributions from both

the lattice enthalpy, AHtat and the enthalpy of hydration ÄH¡y¿ (18)

AHsot-+AH¡¿¡+AH¡y¿

Since the enthalpy of hydration of the benzyltributylammonium cation was not expected

to be significantly different from the other cations considered here, another possible

explanation for benzyltributylammonium nitrates lower enthalpy of solution may result

from an unusual lattice energy when compared to the other salts. If this were the case the

solid state structure of the benzyltributylammonium nitrate salt would be significantly

different. It is hoped that investigations of the solid state structure of these salts will shed

further light on these anomalies.

(6)

Solubility

9.

10

11

12

13.

14.

15.

16.

77

1

2

J

4

5

7

I

6

Literature Cited.

For example, I. N. Levine, Physícal Chemístry,178 (McGraw-Hill: New York

1988).

H. S. Frank, "structural Models", inWater a Comprehensive Treatise.(Ed.F.

Franks),Vol. 1, Chapter 14, (Plenum Press: New York 1972)-

T. S. Sarma and J. C. Ahluwalia, Chem. Soc' Rev.,2,203, (1973).

C. V. Krishnan and H. L. Freidman, J. Phys. Chem.,74(lI),2356, (1970).

W-Y. Wen, "Aqueous Solutions of Symmetrical Tetraalkylammonium Salts", in

Water and Aque ous S olutions, S ¡uc ture, T hermo dy nami cS, and Transport

Processes. (Ed. R. A. Horne), Chapter 15 (Wiley-Interscience: New York1972).

P. W. Atkins, Physical Chemistry 3rd edn. (Oxford University Press: Oxford

1986).

M. H. Abraham, J. Am. Chem. Soc', 104(8), 2085, (1982).

H. Nakayama, H. Kuwata, N. Yamamoto,Y. Akagi and H. Matsui, Bull. Chem.

Soc. Jpn., 62,985, (1989).

S. Lindenbaum and G. E. Boyd,I. Phys. Chem.,68(4), 9lt, (1964).

H. Nakayama, BuIl. Chem. Soc. Jpn', 56, 877, (1983).

H. Nakayama, Bull. Chem. Soc. Jpn., 54,3717, (1981).

I. V. Kudryavtseva, K. P. Mishchenko, and G. M. Poltoratskli, Zh. Struct.

Khim., 13(6), 995, (1972).

R. A. Robinson and R. H. Stokes, Electrolyte Solutions, Revised Edn.

(Butterworths: London 1965).

B. J. Steel, A. S. Kayaalp, T. Kurucsev, A. D. Ward and M. B. Jackson, Aust. J.

Chem., 43, 1983, (1990).

S. Lindenbaum,,I. Phys. Chem.,70(3), 814, (1966).

F. Vaslow, "Thermodynamics of Solutions of Electrolytes", inWater and Aqueous

S olutions, S tucture, T he rmo dynamic S, and Tr ansp ort P ro c e sses. (Ed. R. A.

Horne), Chapter 12 (Wiley-Interscience: New Yorkl9l2).

W-Y. Wen and S. Saito, J. Phys. Chem.,68(9), 2639, (1964)-t7

Solubility

18. M. F. C. Ladd, Z.Phys. Chem. (Frankfurt).,J2,91, (1970).

78

19

4

Density and Viscosity.

4.1 Introduction.

There are two reasons why it was of interest to study the viscosity and density of

the benzyltrialkylammonium salts. Firstly, measurements of the nitrate salts were

essential for the correction of l4N nuclear magnetic resonance relaxation times reported in

Chapter 5. Secondly, the viscosity of only two benzyltrialkylammonium chlorides have

been reported in the literature (1,2,3) and extension of measurements to the higher alkyl

derivatives seemed warranted to ensure they exhibited no unusual behavior that could

influence the relaxation rate of bound nitrate ions. Comparison of the nitrate and chloride

salts also gave some insights into the solution properties of these salts by providing

information on ion-solvent and ion-ion interactions.

4.2 Theoretical Basis.

4.2.L Density.

Use of the density of a solution to make inferences about the solution process is

usually carried out in terms of partial molal volumes. The partial molal volume in a two-

component system, V Z is defined as the partial derivative of the total volume with

respect to the number of moles of solute n2, at fixed temperature T, pressure P, and

number of moles of solvent n1.

i r=(#) r,P,n1 (1)

Density and Viscosity 80

In accordance wirh this formal definition, the partial molal volume is the change in

volume when one mole of solute is added to sufficient volume of solvent so that the

concentration remains substantially unchanged.

The partial molal volume is conveniently determined from the apparent molal

volume, Qy, which is dehned as,

(2)

where V is the volume of solution containing n1 moles of solvent and n2 moles of solute,

and V to is the volume of one mole of pure solvent.

In the case of an aqueous solution where the concentration is expressed on the

molality scale it follows that in 10009 of solvent the following relations are valid,

PoP m2

When the concentration is expressed on the molarity scale n2 = C, V = 1000 cm3 and a

similar equation can be derived.

V1o=!!un¿ u=(!eÌ@) (3)

where m2 is the solute molality, M1 is the molecular weight of the solvent, M2 is the

molecular weight of the solute, po is the density of pure solvent and p is the measured

density of the solution.

Substituting into equation 2. enables the apparent molfVolume to be expressed in

the measurable quantities of concentration, m2 and solution density, p.

Qu=tvbPo

1000î2=m2; nl = Ml ;

(4)

(s)

Density and Viscosity

Extrapolation to Infinite Dilution.

81

(6)

(7)

At infinite dilution the apparent molar volume becomes equivalent to the partial

molar volume, ouo = V zo. T*o major extrapolation equations have been suggested. The

mosr commonly cited relation is that of Masson (4,5) who found that in dilute solutions

Qu varies linearly with the square root of the molarity.

0u = quo + Su* {õ

where ouo is the apparent molar volume at infinite dilution and Su* is the experimentally

determined slope.

The main criticism of this equation is that it does not conform with the Debye-

Hückel limiting law which requires that in dilute solutions the slope should approach a

limiting value. For a 1:1 electrolyte in water this value should be 1.868 (6). Redlich and

Meyer's (7) extrapolation equation does take this into account,

Qv=Quo+Svr/C+bvC

where Sy is now the theoretical limiting value of the slope and by is an experimentally

determined fitting parameter.

When equations 6 or 7 are combined with equation 5 one obtains the concentration

dependence of the density as,

p = po. (f4ïuä9") c (#) c vz

p = po. (r4iuuä9") c (ffi) ç3D - (*orô) .,

The first equarion is identical to that derived by Root (8). The quantities Quo and Sy are

always additive while Sy* and by appear only to be additive for simple systems (6,9).

(8)

(e)


4.2.2. Viscosity.

82

(10)

(1 1)

Viscosity is a measure of a liquid's fluidity. It is defined as the proportionality

constant relating the force per unit area F, also known as the shear stress, to the velocity

gradient in a fluid such that,

dvf = ndy

This is a general relationship for a Newtonian fluid, that is a fluid which has a viscosity

which is independent of the velocity gradient. An example of a non-Newtonian fluid

might be one containing rod-like molecules that become oriented by the flow so that they

slide past each other more freely as the velocity gradient increases. However, the majority

of simple solvents can be considered to be Newtonian.

This equation was used by Poiseuille (10) to derive the relationship describing the

flow of Newtonian liquids within a capillary.

n _ n r4(po - pr) t8Vl

qpt1o Poto

where r is the radius of the capillary, (po - pr) is the pressure drop across the capillary, t

is the flow time, I is the length of the capillary and V is the efflux volume. This equation

provides the basis for the measurement of the absolute viscosity of any liquid. In most

applications the relative viscosity, l."l rather than the absolute viscosity, q is measured.

1ìrel = (t2)

Where t and to are the flow times of the sample and calibration solvent, q and no are the

sample and solvent viscosities and p and ps are the corresponding densities.


Viscosity in the Presence of Solute.

Io

83

(14)

(16)

In the presence of solute the viscosity of a solution will be modified with respect to

that of the pure solvent. It is usual to express the variation of relative viscosity t'¡¡s¡ with

molarity C in terms of the Jones and Dole equation.

lì1.lrel=-!-=1+A{C+BC (13)

The A coefficient depends on long range coulombic forces. At concentrations above

0.002M it is usually swamped by the higher contribution from the B coefhcient. The A

coefficient is always positive and can be computed using the Falkenhagen equation (11).

A = o.7536 (ìr1l\4' 4\ 1o¡'o

o_\4.41\ro\zo(\1o +\zo)

Where \10 and \2o are the ionic equivalent limiting conductivities of the cation and the

anlon

The B coeff,rcient is an empirical constant and can be either positive or negative. It is

solvent dependant and therefore incorporates ion-solvent interactions. Application of the

Einstein relation for spheres (12),

1'ìrel= l+25þ (1s)

where 0 is the volume fraction of solute, to the Jones and Dole equation, shows that the

B coefhcient can be related to ion size in solution by

B=aVh

where Vn is the total volume occupied by the solvated ion in dm3 moll and a is related to

the shape of the ion, having a value of 2.5 for sphericai ions and a higher value for other

Densiry and Viscosity U

shapes (13). This equation assumes that the contribution due to A.Iõ can be ignored

without significant error.

The Jones and Dole equation is limited to fairly low concentrations. For

concentrations greater than 0.1 M the extended Jones and Dole equation may be used.

lìflrel='1 - I +Ar/e+B C +DC2 (r7)

Tlo

It is unclear at present what importance if any can be placed on the D coefficient, which is

generally found to be positive. Desnoyers and Perron (14) have suggested it may result

from solute-solute interactions, but it is also likely to include contributions unaccounted

for by the A and B coefhcients.

4.3. Experimental Techniques.

4.3.I Determination of DensitY.

Density measurements were made using the tare method described by Mulcahy (15)

in conjunction with 30 ml Long Stem pycnometers. The pycnometers were initially

calibrated using Milli-Q water (Appendix A4). Solutions were prepared by dissolving a

known weight of solid salt into a known weight of water to produce a stock solution on

which all subsequent dilutions were made. All dilutions were made by weight and the

temperature was controlled using a water bath accurate to 25 + 0.005 oC.

The pycnometers were cleaned between samples using chromic acid, washing four

times with Milli-Q water, then rinsing with AR Methanol and oven dried before use.

Experimentally it was possible to determine the molality m2, and density of any

given solution whereas all the extrapolation equations (eqn.'s 6,7,8,9) required the

concentration in molarity, C. Having determined the density of the solution the

corresponding molarity was calculated using,

omtc(M)=1+oiloft2r\42 (18)


Then the molarity and corresponding density were fit to the Root equation, enabling the

apparent molar volume Quo and Su* to be calculated from the resultant fitting parameters

(Appendix A5).

There was no significant gain by fitting the data to the Redlich-Meyer equation and

so the Root equation was used as this enabled more direct comparison with literature

coefficients, which are usually reported as resulting from a Root fit.

The reported densities, tabulated in Appendix 45, are the average of two or more

determinations at each concentration. The method was found to be extremely sensitive to

the measured density below about 0.01 M, and a float method would be required to

extend measurements to much lower concentrations.

4.3.2 Determination of Viscosity.

The viscosity was determined using Ubbelohde viscometers (Fig 4.1.) having flow

times of approximately 380 seconds for pure water at25 oC. The water bath temperature

was accurare to within 0.005 oC and all dilutions were made by weight. Sample flow

times were repeated until three values within 0.1 seconds were obtained. Generally, in

practice much better precision was obtained. Timing was accomplished using an optical

fibre sysrem in conjunction with an electronic timing device designed here (16). Because

the viscosity is highly sensitive to impurities, particular attention was taken in cleaning the

viscometer and purification of the samples. The viscometer was cleaned between runs in a

manner similar to that described for the cleaning of pycnometers and samples were thrice

recrystallised from ethanoVether (Chapter 2).

Stock solutions of the benzyltrialkyammonium salts suitable for viscosity

measurements were prepared by weighing out solid salt and dissolving in a known

weight of water. Solutions were filtered through sintered glass funnels before

introduction into the viscometer. All additions and dilutions were made with Milli-Q water

by weight.

Fig. 4.1. Ubbelohde type viscometer.


As was the case with the density measurements, a solution of known molality was

converted to the molarity scale using equation 18. The density required to carry out this

calculation was not determined experimentally, but was calculated using coefficients

obtained from a molality fit of the previously determined densities to the Root equation.

These calculated densities were also used to correct the observed flow times for density,

to obtain the relative viscosity, using equation 12.

The resultant viscosities and corresponding concentrations are tabulated in

Appendix 45, together with the fitting parameters. These data were fit to the Jones and

Dole equation using a non-linear least squares programme (17). The A coefficient was

calculated using the Falkenhagen equation (eqn. 14) from the available limiting ionic

conductivities (18) and was considered for fitting purposes to be a constant.

4,4 Results and Discussion.

4.4.L Partial Molar Volumes.

The density data were fitted to the Root equation (eqn. 8) using the fitting

programme referred to in the previous section (17). From the obtained coefficients the

experimental slope; Su*and the partial molal volume at infinite dilution Quo, were

calculated and are tabulated in Table 4.1.

Fitting of the apparent molar volumes, as calculated from a given density and

concentration using equation 4 or 5, gave extrapolated apparent molar volumes at infinite

dilution Ouo, in good agreement with that obtained by fitting the density to the Root

equation.

Both the chloride and nitrate salts (i9) showed a steady increase in apparent molal

volume with increasing alkyl group length, as expected from an increase in intrinsic

volume. Where comparisons with literature values were possible (2,3) there was good

agreement.- The apparent molar volumes of the cations were calculated by subtraction of the

apparent molar volumes of the anions, taken to be V Z(NOg) = 34.4 cm3 mol-l and


Table 4.1. Calculated coefficients from a fit of density to the Root equation (eqn. 8) foraqueous solutions of benzyltrialkylammonium electrolytes at 25 oC.

Compound u QuO (cm3 mel-l) QuO (lit.) Su* Su* (lit.) ouocation

88

BzMe3N-ClBzEr3N-ClBzPr3N-ClBzBu3N-ClBzPe3N-Cl

BzMe3N-NO3BzEr3N-NO3BzPr3N-NO3BzBu3N-NO3BzPe3N-NO3

170.50 r 0.062tt.4 + 0.3263.8 + 0.3308.4 r 0.1354.6 + 0.2

t82.4 + 0.r226J + 0.2279.3 r 0.3

318.63 t 0.08364.5 + 0.2

tTtb2t6b

-3.86-2.2-7.7-5.9-3.1

+ 0.09+ 0.6+ 0.4+ 0.1+ 0.4

-3.0 b r47.3188.2240.6285.233r.4

148.0t91.7244.9284.3330.1

1.8 + 0.5

2.6 + 0.L4.8 + 0.5

1.6 + 0.11.3 + 1.1

a- The error refers to one standard deviation of the mean. All nitrates coeff,rcientswere determined from data supplied by Guarilloff (19).

b - reference2.

Table 4.2. Viscosity coefhcients of the Jones and Dole equation for aqueous

solutions of benzyltrialkylammonium electrolytes at 25 oC.

Compound a 6b B B (lit.) c Bcadon

BzMe3N-ClBzEt3N-ClBzPryN-ClBzBu3N-ClBzPe¡N-Cl

BzMe3N-NO3BzMe3N-NO3BzEt3N-NO3BzPr3N-NO3BzBu3N-NO3BzPe3N-NO3

0.00740.00740.00780.00830.00860.0089

0.370 + 0.0010.571 + 0.0060.954 + 0.003

1.30 + 0.01t.7t + 0.02

0.331 + 0.0010.338 + 0.002

00.020.0020.010.05

00000

007 1

0075007900820084

0.4130.623

0.3710.5840.961I.307I.717

0.3770.3840.5660.9521.2961.726

0.52 +.906 +t.25 +1.68 +

a - The majority of the nitrate coefhcients were determined from data suppliedby Guarilloff (19).

b - determined from conductance data (8) using the Falkenhagen equation.

c - reference 5.


V z(Cf) = 23.2 cm3 mol-l (6). The additivity between chlorides and nitrates was within

the order expected from a Masson fit, * 3 cm3 mol-l (6), although it was of interest to

note that the apparent molar volumes of the nirate associated cations became slightly less

than those of the corresponding chloride cations as the size of the alkyl group was

increased.

This could be attributed to hydrophobic hydration. As the size of the alkyl group

increases, the nitrate salts become more hydrophobic in nature, leading to an increased

clathrate-like cage formation relative to the chloride salts which are not expected to form

clathrate-like cages. The resulting increase in water structure around the ion is expected to

result in a smaller pafiial molar volume than expected due to the loss of free space around

the ion.

The negative value of Sy* for the chlorides is at present apuzzle since generally Sy*

is found to be positive for simple electrolytes. In the case of some salts, such as sodium

carbonare, the negative coefficient was attributed to substantial hydrolysis (9).

However, many tetraalkyammonium halides also exhibit negative coefficients

(1,2,3,15,20). Wen (20) has attributed this to a deficiency of the Debye-Hückel limiting

law, which does not take into account solute-water interactions. It appears that such

interactions are important for tetraalkylammonium salts in dilute solutions and by analogy

may also be important for the benzyltrialkylammonium salts considered here.

It must however, be realised that in the Masson relation the coefficient Su* is little

more th¿m a fitting parameter and no great inferences can as yet be made regarding either

it's sign or magnitude. The values determined here do however, agree with the sign and

trends previously reported from such fits.

4.4.2 Viscosity.

The viscosity data appearing in Appendix 45, corresponding to concentrations

strictly below 0.1 M, were fit to the Jones and Dole equation. The results for the

chlorides appear in Table 4.2. together with the unpublished nitrate coefficients

determined from the data of Guarilloff (19) and the calculated Falkenhagen coefhcients.


The B coefficients are of the expected magnitude, being slightly larger than the

corresponding teraalkylammonium chlorides (14). They also exhibit the correct trend, of

increasing magnitude with the size of the compound.

The B coefficients of the cations were calculated assuming additivity. The scale

B(Cl-) = B(K+) = -0.007 was assumed and the B coefficient of the nitrate ion was

calculated from the B coefficient of potassium nitrate (21) assuming additivity to give

B(NO¡-) = -0.0467. The agreement in Bcadon for the two salts is excellent and where

comparisons were possible the values are in good agreement with those previously

reported (1).

The nitrate ion is known to be a structure breaker in water and the slightly lower

Bcarion for the nitrate salts could be a result of this, although the anionic contribution to

the B coefficient should be extremely small when compared to the contribution due to the

large hydrophobic structure making cation.

4.4.3 Cation Size and Hydration.

The Einstein model assumes ions to be spheres in a solvent continuum. If these

assumptions are indeed correct for the benzyltrialkylammonium-water system, B should

be directly proportional to the volume of the ion in solution. From a plot of B versus V as

in Fig. 4.2., the B values appear to be somewhat higher than expected from Einstein's

relation (eqn. 16).

Of course real liquids are not continuous and do contain a certain amount of space

between molecules. Thus movement through the solvent without friction with the solvent

molecules can occur (22). This would cause the contribution to the B coefficient to be less

than that to the apparent molar volumes, and thus we might expect points to be below the

theoretical line. The two methyl derivatives are the only samples which are, and then only

marginally.

One possible explanation for the larger B coefhcients of the higher alkyl derivatives

could be due to departure from sphericity. Space filling models suggest that the

compounds can be considered to be largely spherical in nature and that slight departure

2

I

o

c.)É()

Éo'tt()É

1

0100 200 300 400

V cation (cm3 mol-l)

Fig. 4,2. B coefficients of the cations versus the corresponding apparcnt molar

volumes at infinite dilution. The full line is in accordance with eqn. 16 having an a

value of 2.5 for spherical ions. Circles (o): benzyltrialþlammonium chlorides;

Filled circles (o) : benzyltrialþlammonium nitrates.

(¡

a

t

$


from sphericity occurs with the increasing size of the alkyl chain. However, Tuan (22)

has shown that in nonaqueous solutions, similar compounds such as the

tetraalkylammonium salts obey Einstein's relation well. It is tempting, therefore to

suggest that the departure could be due to some specific effect in water, such as

hydration.

In water it is reasonable to assume that the apparent molar volumes are a true

measure of the ion's volume in solution because the value V o it fairly insensitive to

hydration (14). This is because the positive contribution due to hydrophobic hydration

will be closely balanced by a negative contribution to V o du" to the loss of free space

near the non-polar site. Thus the overall change due to hydration on V o will be small and

V o "un

be considered to be a measure of the intrinsic or "true" volume of the bare ion

(14).

Such a compensation is not possible for the B coefficients. Hydrophobic hydration

results in a more structured medium around the ion which hinders movement through the

solvent and results in a positive contribution to B.

Rearrangement of the Einstein relation (eqn. 16) gives,

vny¿ = 4ooB (1e)

where V¡y¿ is the hydrated volume of the ion in cm3 mol-l. It is now possible to

approximate the volume of water around each ion by subtraction of the partial molar

volume of the solution at infinite dilution.

Vw=Vhyd- vzo (20)

The volume of one water molecule was taken to be 30 Ã3 QÐ equivalent to a volume of

18 cm3 mol-l. The approximate hydration number h, for the cation, was then calculated in

a manner similar to that of Kurucsev QÐ. The results of such calculations appear in

Table 4.3.


Table 4.3. Shape and hydration numbers of aqueous solutions ofbenzyltrialkylammonium salts at 25 oC.

Compound Bcarion Quocotion Vw

(cm3 mol-l) (cm3 mol-l)vny¿

(cm3 mol-l¡

h

BzMe3N-ClBzEt3N-ClBzPr3N-ClBzBu3N-ClBzPe3N-Cl

BzMe3N-NO3BzEr3N-NO3BzPr3N-NO3BzBu¡N-NOsBzPe3N-NO3

0.3770.5840.9611.302I.717

0.3770.5660.9521.29ø1.726

150.8233.6384.4522.8686.8

150.8226.4380.8518.4690.4

147.3188.2240.6285.233t.4

148.079t.7244.9284.3330.1

3.545.4

t43.8237.6355.4

2.834.7

135.9234.r360.3

02-3

8l320

0281320

Table 4.4. Estimates of Cation Radii

compound R. (Å) u R, b Re (Å) c Ru (Å) d

BzMe3N-ClBzEr3N-ClBzPr3N-ClBzBu3N-ClBzPe3N-Cl

BzMe3N-NO3BzE13N-NO3BzPr3N-NO3BzBu3N-NO3BzPe3N-NO3

2.843.4r4.124.695.1 1

3.404.t24.695.1 1

4.O34.304.6r4.915.19

4.O34.304.6r4.915.19

3.9r4.525.345.926.48

3.904.485.325.906.49

3.884.214.574.845.08

3.894.244.604.835.08

2.84

a - R, is the radius calculated from the available conductance data fromreference 18 using Stokes' Law. (eqn.22).

b - Modified in the manner of Robinson and Stokes (23).

c - Rg is the radius calculated from the B coeffrcients.

d - Ru is the radius estimated from the apparent mola¡ volumes.


The calculated hydration numbers are quite reasonable when compared to those of

simple electrolyte solutions (23,24) and the agreement between the chlorides and nitrates

is excellent. The methyl derivatives are apparently not hydrated and therefore agree well

with expected theory. As the size of the alkyl chain increases, however, the extent of

hydration increases causing the resultant departure from Einstein's relation.'

From the two calculated volumes in Table 4.3. it is possible to calculate the

corresponding radii of the cation if sphericity is assumed. It is also possible to calculate

approximate radii from the Stokes-Einstein relation using the limiting conductance of the

ions reported by Steel et al. (18). The Stokes-Einstein equation relates

coefficients to the effective hydrodynamic radius of a spherical particle,

dir(,f"" {¡

lrl

and leads to,

D= kr6nqRg

(2r)

(22)

where e is the elementary charge, F is Faraday's constant, \ao is the limiting ionic

conductivity of the cation in S m2 equiv-l and q is the viscosity of water taken to be

8.904 xl}-2 poise. This is useful because it provides three independent methods for the

calculation of radii and a good method of checking the accuracy of the results, since the

radius of the cation should remain largely independent of the method of determination.

The Stokes radii are usually much less than the radii estimated by other methods.

This is because the Stokes-Einstein relation (eqn. 21) is not valid for the motion of

smaller ions (23,25). Stokes' law seems to be correct only for ions having a radius

greater than about 5 ,&, and Robinson and Stokes have used the tetraalkylammonium salts

in water to produce a correction curve for ions with a Stokes radius smaller than this (23).

Their data were fit to a quadratic and the coefficients used to obtain the "modified" Stokes

radii in Table 4.4.

It can be seen that all the radii a¡e of the same magnitude and agree closely with the

radii expected from space filling models. The radii determined from viscosity being

slightly larger due to hydration. This is clearly seen from a plot of Rg versus Ry as in

o<

Êq

ú

7

6

5

4

3

76322

4 5

nv (Å)

Fig. 4.3. Cation radii determined from viscosity B coefficients versus the

corresponding radii calculated from apparent molar volumes. The full line

represents the theoretical 1:1 relationship between the ¡wo radii.

r Tetraalkylammonium salts in acetonitrile (reC.22).

tr Tetraalkylammonium salts in water (refs.22 and23).

o Benzyltrialkylammonium chlorides in water (this work).

d

o

I

o

t


Fig. 4.3. The tetraalkylammonium salts in acetonitnle (22) show excellent a'greement

since hydration is not possible, while both the tetraalkylammonium salts and the

benzyltrialkylammonium salts show clear departure from linearity.----

4.5 Conclusion //

The trends observed in the apparent molar volumes and B coefficients for the

benzyltrialkyammonium salts are not markedly different from the tetraalkylammonium

salts. It appears that the salts exhibit increasing hydration with the increasing size of the

alkyl group. This is consistent with the increasing clathrate-like cage structure proposed

in Chapter 3 to explain the unusual solubility results. Such a monotone increase was also

observed in the association constants obtained from conductance measurements (18). The

combination of these two results suggest that hydrophobic hydration of the cation plays a

role in determining association.

â+I


Literature Cited.

10.

11.

12.

t3.

C. Yatome and Y. Takase, Sen-l Gakkaishi,30(l), 54, (1974).

C. Yatome, Y. Yamaguichi and Y. Takase, Sen-l Gakkaishi,3l(4),55, (1975).

C. Yatome and Y. Takase, Sen-l Gakkaíshi,32(4),69, (1976).

D. O. Masson, Phí\. Mag., 7(8), 2I8, (1929).

A.F. Scott, J. Phys. Chem.,35,2315, (1931).

F. J. Millero, "The Partial Molal Volumes of Electrolytes in Aqueous Solutions",

in W ater a nd A que o us S o luti o ns : S tuc tur e, T he rmo þ natni c s, a nd T r ansp or t

Processes. (Ed. R.A. Horne), Chapter 13 (\Miley-Interscience: New YotkL972).

O. Redlich and D. M. Meyer, Chem. Rev.,64(3),221, (1964).

W. C. Root, J. Amer. Chem..Soc., 55, 850, (1933).

H. S. Harned and B. B. Owe n,The Physical Chemistry of Electrolyte Solutions,

Amer. Chem. Soc. Monogr. No. 137, (Reinhold Publishing: New York 1958).

J. Poiseuille, Mém. Savants Étangers,7,l05, (1841) ; 109, 385, (1860).

H. Falkenhagen and E.L. Vernon, Z. Physik.,33, 140, (1932).

A. Einstein , Ann. Phys., L9,289, (1906) i 34, 59L, (1911).

P. V/. Atkins, Physical Chemistry 3rd ed. (Oxford University Press: Oxford

1986).

J. E. Desnoyers and G. Perron, J. Sol. Chem.,l(3), 199, (1972).

D. E. Mulcahy, PhD Thesis, University of Adelaide, (1967).

B. J. Steel, J. Sci. Instr., 42,751, (1965).

T. Kurucsev, J. Chem. Educ., 55, I28, (1978).

B. J. Steel, A. S. Kayaalp, T. Kurucsev, D. Wa¡d and M. B. Jackson, Aust. J'

Chem., 43, 1983, (1990).

P. Guarillofl PhD Thesis, University of Adelaide, (1994).

W-Y. Wen and S. Saito, J. Phys. Chem.,70(5), 1473, (1966).

t4.

15.

16.

17.

18.

t9

97

1

2

3

4

5

6

7

8

9

20


21. R. H. Stokes and J. Mills, "Viscosity of Electrolytes and Related Properties",

inThe Internationnl Encyclopedia of Physical Clumistry and Chemical Physics

(Eds. E. A. Guggenheim, J.E. Mayer and F. C. Tomkins) 3/16, (Pergamon Press:

Oxford 1965).

22. D. F. Tuan and R. M. Fuoss, J. Phys. Chem., 67, L343, (1963).

23. R. A. Robinson and R. H. Stokes, Electrolyte Solutions 3rd ed. (Butterworths:

25

London 1959).

T. Kurucsev, A. M. Sargeson , and B. O. West, J. Phys. Chem.,6l, L567,

Q2Ð.'--'J. T. Edward, J. Chem. Educ.,47,261, (1970).

Nucleør Magnelic Resonance 100

In this chaprer the relaxation times of some simple nitrate salts were initially

examined and the results compared with those of Nicholas and V/asylishen (8). The

measurements were then extended for some salts, to cover a range of concentrations, so

that an estimate of the equilibrium constant could be obtained and compared with literature

values.

Previous studies of the conductance of the monomeric benzyltrialkylammonium

nitrates had shown that the association constants in aqueous solution were small (6). Here

nmr relaxation times of these same salts were used to confirm this result. It was shown

that nmr could be used as a useful analytical tool to probe molecular rotation and ion

association over a range of concentrations and that all results could be interpreted almost

exclusively in terms of ion association.

The relaxation times of nitrate in the presence of the corresponding linear chloride

polymers were then examined and the results compared with a concurrent study using

emf measurements (10). Finally, the effect of the degree of cross-linking on nitrate

rotation was investigated.

5.2 Theory.

5.2.1 Nuclei within a Magnetic Field.

The nuclei of many atoms are known to be spinning about an axis and are said to

possess "spin", symbolised by the spin quantum number, I. Such nuclei will therefore

possess angular momentum and may adopt 2I+1 orientations within a magnetic field; each

characterized by an angular momentum quantum number, DI = I, I-1, ..... , -f as

illlustrated in Fig. 5.1. A charged particle spinning about an axis constitutes a circular

electric cuffent which generates a magnetic dipole. All nuclei with nuclear spin (I > l/2)

behave as such a charged spinning particle and the magnitude of the dipole, called the

magnetic moment p is given by,

p=sBNüïTÐ-Jr-1 (1)

Ho zz

+1

mr=41 ml=0 mr=-l

Fig. 5.1. The permitted orientations of angular momentum for I = I in an applied

magnetic field, Hs.

Ho

m=-1 m=+1

AE

m=0 m=0

m=+1 m=-1

z

,'+//l.T

zeÍo magneticfield

magnetic fieldapplied

Fig. 5.2. 14N nucleus for I = I showing the three energy levels of the nucleus and the

three corresponding magnetic moments in the presence of an applied magnetic field.

Nuclear Magnetic Resonance 102

where g is an experimentally determined value, characteristic of each nucleus, called the

nuclea¡ g factor and B¡ is the nuclear magneton; equal to 5.050 x 10-27 J T-1.

In the absence of an applied magnetic field all of the spin states are degenerate.

However, when the nucleus is subjected to an external magnetic field, H6, there is an

interaction between the applied field and the magnetic dipole of the nucleus which causes

the degeneracy to be lifted.

The nuclear spin axis is usually taken to be the z axis and p has 2I+1 orientations

with respect to the nuclear spin axis corresponding to 2I+L energy states. The energy

difference, for spectroscopically allowed transitions, between two adjacent spin states

when Àm1= t I is given bY,

AE=ISBNHOI¡ = I gBNHOI g' (2)

For the 14N nucleus having I=1 the situation is fully illustrated in Fig. 5.2.

Thus a change of spin is associated with the emission or absorption of energy at a

given frequency which is proportional to the applied magnetic field. This frequency is

specific for a given nucleus and is known as the Larmor frequency, olo.

t! Hoú)^ = ------r-----i-_ HZ = AE" 2n r/t (t + t)

(3)

It is no accident that the Larmor frequency is identical to the energy required to induce a

transition in spin state. The interaction of p and Hs induces a torque in the magnetic

moment of the nucleus given by,

Torque = f = F x Ho= It Ho sin 0

The direction of the torque, as given by the right hand rule, is perpendicular to the plane

of p and Ho, and causes the magnetic moment of the nucleus to rotate, or precess,

about Ho, Fig. 5.3. Thus o16 is known as the Larmor precessional frequency and a beam

of energy having a frequency equal to the Larmor frequency will interact coherently with

(4)

Z

Ho

u

X

v

Fig. 5.3. The interaction of Ho and p induces a torque T in the direction of the y axis

lþ"tof+r€ry-hich causes the magnetic moment to precess about the z axis at

frequency <os.

^-2L+

^-

Fig. 5.4. Schematic representation of ca¡bon monoxide showing back-to-back dipole

moments.


the nucleus and cause transitions from one energy level to another, while a beam of any

orher frequency will cause no such transition. This is the basic principle of nmr. In

practice the nuclei are bathed in a radiation of a fixed frequency of say 300 MHz and the

radiation beam is swept over a range of frequencies until resonance occurs.

5.2.2 Relaxation and Quadrupole Effects.

In an applied magnetic field Hs, the population of the different spin states is

governed by the Boltzmann distribution.

(s)

In the case of nuclei AE is very small and this ratio is very close to unity. That is the

population of the spin states are almost identical. The fact that the lower energy state will

always be slightly more populated makes nmr possible, because if at any time the

population states do become equal, then the probability of upward and downward

known as 2-

Nffi=exP( *-)

transitions also and no more absorbance can occur.

saturation.{owever, resonance spectra do not vanish as they are observed and this

implies that there must be a mechanism for returning nuclei to a lower energy spin state.

The mechanism by which excess spin energy is shared between the surroundings

and other nuclei is known as a relaxation process and the time taken for a fraction

(Ue = 0.37) of the excess spin energy to be dissipated is known as the relaxation time.

For nuclei two different relaxation processes are possible. The excess energy can be

shared with the surroundings; the lattice, by spin-lattice or longitudinal relaxation

characterized by the relaxation time, T1. Alternatively, the excess energy can be shared

with other nuclei by spin-spin or transverse relaxation cha¡acterized by the relaxation

time, T2.

In liquids Tt=TzandT1 can have alarge range of values, 10-4 - 10 seconds

(11). In general T1 reflects the lifetime of a particular spin state and the smaller the value

the more effrcient is the relaxation process.


One of the most efficient means of relaxation is quadrupole relaxation. All nuclei

with spin ) 1, such as 14N, possess an electric quadrupole moment which arises because

the nuclear charge is not spherically distributed. The electric quadrupole moment is a

measure of the departure from sphericity, being positive for egg shaped nuclei, negative

for tangerine shaped nuclei and zero when the distribution is spherical (I= l/2,0). The

electric quadrupole moment may be considered as two dipoles placed back to back. Take

for example carbon dioxide Fig. 5.4., which has no perrnanent dipole moment but does

possess a perrnanent quadrupole moment. The nuclear electric quadrupole moment, Q,

interacts with the fluctuating electric field gradient, eq, at the nucleus and effects the

orientation of p providing an additional relaxation pathway for nuclear spin. For many

nuclei this will be the dominant form of relaxation.

5.2.3. Relaxation and Rotational Correlation Times.

In the extreme narrowing limit, (2ærofs<<l), the spin lattice relaxation rate is

given by (12,13),

I3T1 =10 "'(iã\) (") ('.?)" (6)

where X = #, is the nuclea¡ quadrupole coupling constant expressed in Hz, and ( is

an asymmetry factor which describes the extent of the nuclear quadrupole coupling

constant's (NQCC) departure from cylind¡ical symmetry. tc is the rotational correlation

time. Although not strictly correct, Tc cÍur be defined as the time taken for a molecule to

rotate through an angle of 1 rad (57o).

The nitrate ion exhibits anisotropic rotation due to its oblate shape. In aqueous

solution the ion can be considered to have D3¡ slmmetry and the rotational motion can be

charactenzed by the times for the perpendicular and parallel rotations (t1and t¿) about

the principle axis (Fig. 5.5). In the xy plane the nitrogen nucleus is symmetrically bound

to three oxygen atoms and the largest component of the electric field gradient is

perpendicular to the ion plane, correspondlng to the largest component of the inertial

Z

tI

I Xtt

I0

oXI>X+

I

I

I

I

V

0/

v

Fig 5.5. Coordinate axis system for the nitrate anion.


tensor (14). Under these conditions the asymmetry factor becomes zero and equation 6

reduces to,

+=t "z e+e)r,, (7)

where tl is the correlation time for rotation about the C2 axis.

A rotating molecule will have an inherent rotational correlation time to, due only

to its moment of inertia. In solution the liquid phase has viscosity, which slows rotation,

and causes the observed t to be greater than that of the free ion. The molecule therefore

possesses a rotational correlation time which has components resulting from both inertial

and diffusional effects.

1ç=f6*Îy (8)

For most molecules fe is zero, or small and positive, and rotation in a liquid can be

considered to be almost exclusively viscosity dependent.

The Stokes-Einstein-Debye equation relates the rotational correlation time to

viscosity and is given by,

, =! T! ^t (e)uc -3 kr'

where a is the molecular radius and k is the Boltzmann constant. By comparison of

equations 6 and 9 it has been concluded that the quantity nTl should be concentration

independent for a non-associating solute (15). Any departure from linearity has thus been

attributed to ion association. ,AA u, Cr., Lr,.-r,

In the simplest mode on time is considered to be a,Ðrylon

otf'!futnerelaxation #; "nl'Crl.

There are two limiting cases

dependent upon the rate of chemical exchange relative to the rate of molecular rotation.

When chemical exchange is relatively slow:

N uclear Magnetic Rcsonance 108

1_4,Tt - T1 bound '

(10)

(12)

and whenever chemical exchange is relatively fast:

T1 = T1 bound A + T1 free (1 - A) (1 t)

where A is the mole fraction of nitrate bound.

5.2.4. Equilibrium Constants.

From the concentration dependence of the relaxation times it should be possible to

obtain an estimate of the equilibrium conStânt, K4.

For a 1:1 electrolyte, MX, the equilibrium is,

M++No¡-SrurNo,

characterized by the thermodynamic equilibrium constant

where A is the fraction of nitrate bound, p is the mean ionic activity and C is the

concentration on the molar scale. The equation assumes that the activity of the uncharged

species does not differ greatly from unity.

For a 2:1 electrolyte, MX2, the complex associates in two stages and the

equilibrium is slightly more complicated due to the formation of a charged mononitrate

species in solution.

vr2+ +No¡- I [MNo¡]+ I vr(Nor)t

Nuclear Magnelic Resonance 109

The equilibrium can be simplifîed if K1 is taken to be so small that the salt dissociates

complerely in solution. That is if the species M(NOl)z does not exist in solution. Under

these conditions it follows that

(13a)

Where K5 is the stoichiometric association constant. Multiplying both the numerator and

the denominator by fNO¡-) converts the individual ionic activity coefficients to the mean

ionic activity coefficients of the 2:1 electrolyte M(NO¡)z and the 1:1 electrolyte

(MNO¡+)NO:-.

(t3b)

(14b)

(1s)

(L4a)

=Ks xz 1t:1)

y*3 qz:t¡

While rhe mean ionic activity coefficients of the species M(NO:)z are readily available in

most rexts the mean ionic activity coefficients of the 1:1 species (MNO¡+)NO3- is not

known, although it is expected to deviate substantially from unity.

Assuming unit activity an estimate of the thermodynamic association constant can

be obtained by extrapolation to infinite dilution of

vlMNOa+)KA: Ks x yoþ¡lñ; x

Klt (2:1) = Ksy+3 (2:l)

y(No¡-)

r(No¡-)

An alternative is to make some specific assumption regarding the activity of the

(MNO¡+)NO3- species. It is reasonable to suspect that the mean ionic activity coefficients

will be similar to those of other simple 1:1 electrolytes. The ammonium nitrate salt was


chosen as it is non associating and of similar size to that expected for MNO3+.

Substitution gives

3 (NH¿No¡)K¡2 (2:1) = Ks x

y+3 (zir)

Again extrapolation to infinite dilution allows estimation of the thermodynamic

association constant.

A further complication may a¡ise if there is signif,rcant ion association in solution

(16). This seems more likely to occur in 2:I electrolytes due to an increased charge

density relative to 1:1 electrolytes. Such ion association will cause a substantial decrease

in the observed activity coefficients. Similarly, the ionic strength, I, will be reduced

below 3C due to ion pair formation.

(16)

(17)I=(3 -2¡^)C

The decrease in the activity coefficients relative to the "nomal" activity coefficients in the

absence of any association is expected to be directly proportional to the fraction of nitrate

bound. It follows rhat multiplication of equation 16 by the fraction of nitrate free will

correct the activity factor. This gives a third estimate of the association constant for a2:I

electrolyte.

(18)

The use of this correction removed the concentration dependence of the apparent

association constants, but did not substantially alter the extrapolated estimates of the

thermodynamic association constants.

Analogous equations to those described above can be written for the molal scale

by substitution of molal concentrations and molal activity coefficients. Where possible the

molar concentration scale was adopted in this work for ready comparison with literature

KA3 (2:1) = Ks. vt3 !NIì:¡l x (1 - e¡3y*s (2:l)


values. In theory, at infinite dilution both concentration scales should yield identical

results.

5.3 Experimental Methods.

5.3.1 Solution Preparation.

Solutions suitable for nmr measurements were prepared by weight from vacuum

dried salts dissolved in Milti-Q water. Where a series of concentrations were required the

dilutions were made by weight from a common stock solution. All solutions contained

10Vo D2O w/w, which was necessary to provide an internal deuterium lock for the nmr

spectrometer. Immediately prior to running the spectra the solutions were filtered through

a simple wool plug, placed in a pipette, directly into a clean Wilmald type 513pp 10mm

nmr tube. All molalities were converted to molarities using the previously determined

density coefficients (Appendix A5), or available literature values in the case of simple

nitate salts.

5.3.2 Density and Viscosity.

The experimental methods for the determination of viscosity and density have

been fully described in Chapter 4. In order to correct the relaxation times of the monomer

analogues it was necessary to extend these measurements to the higher concentration

ranges covered by the nmr measurements. The density and viscosity of the

benzyltrialkylammonium nitrate salts were largely available from Guarilloff (10) and the

benzyltrimethylammonium nitrate was the only salt requiring the viscosity measurements

to be extended to higher concentrations and was determined here (Appendix A5). All

viscosity data were fit to the extended Jones and Dole equation and the resultant

coeffrcients appeil in Appendix 45. The density coefhcients of the nitrate salts previously

described in Chapter 4 and Appendix A5 were considered sufficient for the correction of

the viscosity.

Nuclear Magnetic Resonance ll2

The main concern with the available viscosity and density coefficients was that

they related to salts dissolved in pure water, whereas the nm¡ solutions contained 107o

D2O. Since the viscosity and density of water and deuterium oxide are fairly similar this

was not believed to introduce a significant elror. For example, consider the density of

water and deuterium oxide, p(HzO) =0'997048 and p(D2O) = 1'10445 I cm-3 (17)' The

percentage difference between the two is around ll%o,but since solutions were 9ÙVoHZO

this reduces to only about a l7o difference. Even under optimum conditions errors in T1

are greater than3To.It was found that for simple nitrate salts it made very little difference

if literature values of density and viscosity in pure water were used in place of the

experimentally determined density and viscosity of lOVo DZO solutions on which

experiments were actually performed (10).

The viscosity and density measurements of the linea¡ polymers were performed

on the actual solutions used in the nmr experiment, using the standard methods

previously described in Chapter 4.

5.3.3 Design of the Spin-lattice Relaxation Experiment.

The determination of spin-lattice relaxation times, T1, consuûle a considerable

amount of spectrometer time. It was therefore of interest to design the experiment in such

a manner to reduce the total time of determination while retaining a high degree of

precision. The two main areas where optimization was possible were in the choice of

pulse sequence and the choice, number and disfibution of the delay times, r¡.

Choice of Pulse sequence.

Several different pulse sequences have been used successfully in the

determination of relaxation times and the most conìmon a.re summarized in Table 5.1. The

total time for each experiment has also been calculated, ignoring any time involved in the

Table 5.1. Methods for the determination of spinJattice relaxation times.

Method Pulse Sequence Total Time

Inversion Recovery(rRFT) (20,2r).

Freeman-Hill modificationof Inversion Recovery(FH-rRFT) (79,2r).

Saturation Recovery(SRFT) (20,21).

(180o-t¡90o-Tuq-Tdn T=n^ ( Tr,+Tuo+T¿)i=0

(, T;' + 2(r^r+ rd)(I¿- 1 80o-t1-90o-Taq-T¿-90o-T¿q)¡

r.h

>2HSP+ti+=0

(90o-HSP-t¡90o-T¿q-HSP)¡ T=n x ( ruq)

T=nX

Progressive Saturation(PSFr) (21).

(90o-t¡HSP)n

In the symbolic notation adopted in this table, 18@ and 90o represent the respective flip angles, ti represents the delay times, T¿q

represents the aquisition time, T¿ represents the delay or recylce time necessary to allow full recovery between scans and HSPdenotes a homospoiling pulse. Nl is the number of delay times and n is the number of scans at each delay.

1= (n+3) * ( Tnsp *., )i=0

Nucleør Magnetic Resonance ll4

storing of data to computer, which would be constant for a given instrumental setup and

independent of the pulse sequence chosen.

It can be seen that the FHIRFT sequence is the most time consuming of all the

pulse sequences. It is also less sensitive than the IRFT sequence (18). For long T1's or

where time is not important it is thought to be the most accurate method of determination

(le).

In contrast the quickest pulse sequence is PSFT. It is also less sensitive than the

IRFT sequence and is restricted to fairly long relaxation times, because the sequence

requires ri > Taq. The sequence is also restricted to samples for which T2 << T1. Its main

weakness, however, is that in order to obtain accurate values for T1, stringent attention

must be made to the precise setting of the 90o flip angle.

The SRFT sequence is of similar speed to the PSFT sequence, and is less

sensitive to an inaccurate setting of the 90o flip angle. The SRFT sequence is generally

faster than the IRFT sequence and its main drawback involves the use of homospoiling

pulses which are necessary to remove residual x,y magnitization and reduce non linea¡

phase erors. Their use may preclude the use of Autoshimming on some instruments.

The most widely reported method for the determination of T1 relaxation times is

the IRFT sequence. This is most likely due to the method's general applicability to a wide

range of T1's. The IRFT sequence is twice as sensitive as the SRFT or PSFT methods

and like the SRFT sequence, pulse angle settings are less critical. Its main drawback is

that a maximum T1 must be known prior to running the experiment since the sequence

requires Taq * T¿ > 5Tt. However, when T1 is small this condition can easily be fulfilled

by choosing a moderately large delay time T6, around 1 s for nirogen nuclei.

In summary, the choice of pulse sequence was made after consideration of the

magnitude of the relaxation time, the precision required, spectrometer availability and the

technical capabilities of the instrument in both operation and data acquisition. The IRFT

sequence was found to be the most appropriate and accurate method of determination and

was adopted here.

N uclear Magnetic Resonance

Choice of Delay Times.

115

One obvious area for optimization is in the choice of delay times t¡; the shorter the

delay times the shorter the whole experiment. While some researchers have reported delay

times in the range 3Tt < T¿ < 7Tt it is now accepted that f¡¡¿¡¡ = 3.5 Tt is usually

sufficient to ensure an accurate determination of Tt (20,2I).

When all other conditions are equal the accuracy of T1 is proportional to the signal

to noise ratio and hence {n where n is the number of repeat scans at each delay. The

question which arises is whether it is better to have a few t1 values with a high S/l.{; or a

large number of t1 with a lower S/N, In general provided the S/N ratio is greater than

about 10 - 15, precision is improved by increasing the number of t¡ at the expense of

some S¡r¡ (20,21) where the distribution of t¡ does not necessarily have to be uniform

(1e).

5.3.4 NMR.

The laN nrnr spectra were obtained on a Bruker CXP300 spectrometer using a

broad band variable temperature 10 mm probe (10mm VSP BB probe) operating at2l.66

MHz. The temperarure was maintained at 298.2 + 0.3 K using a nitrogen flow. All spin

lattice relaxation times were determined using the standard inversion recovery pulse

sequence (IRFT) described above (18). Typical acquisition conditions included a 90o

pulse width of 80 ps with a 5000 Hz spectral width, 8K data points, and a recycle time,

Td = 1s. All Tr values were calculated from peak heights obtained at 12 or more T delays,

using a non-linear least squares fitting programme (22).Depending on the concentration

between 12 - 1500 scans at each delay were required to ensure a signal to noise ratio for

Tæ grearer than 50:1.2M Ca(NO¡)Z was used to assign the reference frequency, set

arbitrarity to zero, and for initial tuning, all other chemical shifts were measured relative

to this.

N uclear Magnetic Resonance ll6

5.4 Results and Discussion.

5.4.1 Relaxat¡on in Simple Nitrate Salts.

In order to rest that the method reported by Nicholas and Wasylishen (8,23) was

suitable for use with the available nmr hardwa¡e and softwa¡e, the relaxation times of a

number of simple nirates were investigated. The results appearing in Table 5.2. were in

good agreement with those of Nicholas and Wasylishen (8). The salts chosen represent a

wide range of simple nitrate salts. Ammonium and potassium nitrate a¡e believed to be

non-associating while sodium nitrate is slightly associating at high concentrations.

Calcium and lead nitrate have been included as examples of two salts having higher

degrees of association in aqueous solutions even at lower concentrations.

For these simple salts the number of scans were controlled to give a SA{ ratio in

most instances greater than 125:1. This was thought to be suff,rcient to ensure accurate T1

determinations and no significant improvement was gained by increasing the number of

scans for the same number of delay times. The values of T1 appearing in Table 5.2. were

determined from a non-linear least squares fit of delay times and peak intensities to

equation 19.

rx = r- (r - ze *p (+)) (1e)

where Ix is the peak intensity corresponding to a given tau value f¡, I- is the peak

intensity at infinite time, corresponding to the equilibrium magnetization and A is a

scaling factor approximately equal to unity (18). The final values of T1, A and f-, were

always found to be independent of the initial estimates.

As previously mentioned the solutions were preparedas lÙVoDzO/HzO mixtures.

Gua¡illof.f (10) had shown that there was no signifrcant difference between T1's corrected

using experimental viscosities and densities or T1's corrected using literature values for

ninate salts in pure water. Guarilloff used the viscosity parameters reported by Jones and

Talley (25) to correct his experimental relaxation times. These parameters were from a fit

N uclear Magnetic Resonance

Table 5.2.Effect of counterion on spin-lattice relaxation times of some simple nitrate

salts in water at 25 oC.

Cation c (m) Tlr T1 corr. (ms) a Tr (ps)

tr7

NH¿+

K+

K+

Na+

Ca2+

Pb2+

0.09800.9771

0.09310.8831

0.10731.0082

0.10440.9924

0.09911.0183

0.1 1050.9568

0.994sb0.9653 b

0.997 ac0.980r c

0.9960 d

0.917 oa

1.0066 e

l.O626e

1.0193 fI.298¿,r

117.0 + 6.699.3 + 4.0

t21.5 + 10.689.2 + 5.6

t09.r + 7.386.5 + 4.2

104.5 + 3.1

83.6 + 2.9

77.9 + 5.337.9 + 1.0

55.7 + 1.8

29.0 + 0.5

1.04 + 0.051.24 + 0.O5

1.00 + 0.141.36 t 0.09

r.tz + 0.071.41 + 0.06

1.16 + 0.031.46 + 0.05

1.56 + 0.103.21t 0.09

2.18 + 0.074.20 + 0.07

1.015181.2407I

a-

b-

c-

d-

e-

f

In calculating the error in T1 corr. = 1'ìr x T1 expt. a standa¡d error in 1¡ of 0.002

was assumed. The standard error in T1 expt. was determined from a least

squares fit of intensity and tau values.

The viscosi e

polynomial e

wai treated the

molarity to the molality scale. The temperature was taken to be 25 oC.

From reference 10. The solutions were only 57o DZOIEI2O and the viscosityappearing in the table was determined experimentally.

Calculated using viscosity parameters reported by Jones and Talley (25).

Interpolated from data in reference 26.

Interpolated from data in reference2T.

Calculated from a fit of available data (28) to an appropriate polynomial.(v-Þ


to the Jones and Dole equation and as such are only valid up to 0.1 M. At high

concentrations the slight departure between experimental and literature corrected

relaxation time resulted from this, rather than any real difference between the relaxation

tlmes.

The experimentally determined relaxation times were corrected for viscosity using

the available literature values for salts in pure water. The magnitude of the correction was

directly proportional to the viscosity of the solution and provided the concentration was

fairly small, such that the relative viscosity was close to unity, the correction also

remained small and much less than the inherent elror in Tt. Thus the correction can safely

be ignored if only qualitative inferences are to be made. However, it should be possible

to use relaxation data to obtain a quantitative estimate of association. In this case the

corection is required in order to obtain true thermodynamic estimates of association

constants and was therefore applied throughout this work.

Although within the expected experimental error, the relaxation times determined

here were consistently smaller than those obtained by Nicholas and V/asylishen (8). This

suggests some systematic difference between the two determinations. The measurements

of Nichotas and Wasylishen were performed at 23 + 2 oC, while in this study the

remperarure was slightly higher, 25.O + 0.3 oC. The higher temperature would be

expected to increase rotation and decrease ion pairing; both of which should result in

larger observed relaxation times. This is directly opposite to the observed difference. The

work of Nicholas and Wasylishen differs from this study in two more important aspects.

Their values correspond to the molarity scale and make no allowance for the viscosity of

the solution. In a recent study relaxation times determined by Nakahara et al. (29) have

also been observed to be consistently smaller than those of Nicholas and'Wasylishen.

V/here comparisons were possible the relaxation times determined here are virtually

identical to those reported by Nakahara et aI. (29).

Even given that the data of Nicholas and Wasylishen (8) are not strictly

comparable, relaxation times agree within the expected experimental error and follow

exactly the same trends. For a given cation the relaxation time of the nitrate ion decreased

with an increase in concentration. This was consistent with greater association at the

Nuclear Magnetic Resonance I 19

higher concentration. In general, for a given concentration, the relaxation time also

decreased with the suspected complexing power of the cation. In discussions of

relaxation times and the complexing powers of cations, it is important to remember that

the observed T1 represent the average relaxation time due to all species in solution,

making interpretation difficult if speciation is unusually complex.

If there is significant association and the lifetimes of the distinct species are

sufficiently long, then it should be possible to obtain two distinct signals corresponding

ro the bound and bulk nitrate. Large chemical shifts have been observed for bound and

bulk nitrate ions using 15N nmr, but only when they were coordinated to paramagnetic

ions such as the trivalent lanthanide ions (30,31,32,33). The shifts are due almost

exclusively to the presence of unpaired electrons in these metal ions and parallel their

magnetic moments. None of the cations in this work were paramagnetic so resolution of

distinct peaks on this ground was unexpected.

Another consideration is temperature. Fratiello et al. (33) used fairly low

temperatures C80 oC to -90 oC ) where the rate of exchange would be lower. All spectra

in this work were recorded at ambient temperatures (25.0 oC + 0.3 oC), where the rate of

exchange would be much faster. Indeed only single peaks were observed for all samples

considered here. Even spectra recorded at a greatly expanded sweep width failed to

indicate the presence of any other signals. It is unknown if lowering the temperature

would resolve the observed signal into distinguishable bound and bulk signals. Lowering

the temperature to such an extent would however pose technical problems for aqueous

solutions, which would freeze at these temperatures. Nicholas and Wasylishen (23)

recorded the spectra of 1 M NaNO3, in the absence of paramagnetic ions, at -4 oC

without mention of resolution.

The rotational corelation times were calculated from equation 7 using NQCC

(1¿Ð = 745 k]Frz (34). This choice assumes firstly that NQCC in aqueous solution is

identical to the value measured in the solid state and secondly that NQCC remains

unaffected by association. The first statement has been shown to be correct by Gourdji et

al. (34).


The rotational correlation times of some of the more concentrated solutions can be

compared with values obtained using Raman spectroscopy. Spectra of 1.0 M solutions of

sodium and ammonium nitrate have been determined at 29 + 1.5 oC by Kato et al. (35).

For ammonium nitrate they obtained t1 = 1.33 ps comparedto 1.24 ps determined here.

For sodium nitrate there is again excellent correlation. They obtain Î-¡ = 1.44 ps

compared to the value of 1.46 ps determined here and for potassium nitrate 1.40 ps

compared to 1.40 ps. Similarly, Whittle and Clarke (4) obtained a value of t1 = 3.8 ps

for 0.9 M calcium nitrate, in good agreement with the value of 3.21ps for the 1.0 molal

solution determined here. The agreement not only confirms the accuracy of the nmr \

method but also suggests that the assumption made previously, that NQCC remains þ.lt'

unchanged due to association, is substantially correct. ¡'\ \)

\The variation in the observed relaxation time with 90o pulse angle was

investigated using calcium nitrate. As can be seen from Table 5.3. the relaxation time was

uneffected by small changes in the magnitude of the pulse angle. This confirmed the

insensitivity of the IRFT sequence to an incorrectly set pulse width.

Table 5.3. Variation of spin-lattice relaxation time with 90o pulse for 0.7884 molal

calcium nitrate at25 oC.

Dl = 900 (us) Tr (ms) a

rt7107127

33.2 + 0.932.9 + 0.633.7 + 2.2

a,- the relaxation times are uncorected for viscosity

Concentration Dependence

The concentration dependence of a number of simple nitrate salts had previously

been reported (8,9,10,23). V/hile some support exists for the interpretation of this data in

teÍns of a hydrodynamic interaction model (36), in this work the concentration

dependence of the relaxation times was attributed predominantly to ion association in

Nuclear Magnetic Resonance l2l

solution. Estimates of thermodynamic association constants were obtained by

extrapolation of apparent association constants to infinite dilution. The definition of the

association constants and any assumptions made regarding the activity of the solutions

were discussed in section 5.2.4.In order to minimize any error resulting from the

inherent assumptions made regarding these activities, thermodynamic association

constanrs were estimated by linea¡ extrapolation to infinite dilution of the logarithm of the

apparent association constants as a function of ionic strength.

Although both the fast and slow exchange mechanisms were considered, the fast

exchange mechanism invariably gave association constants of an order of magnitude

larger than expected and the slow exchange mechanism was used in all calculations. This

choice reinforces the conjecture of Nicholas and Wasylishen (8) that the rate of chemical

exchange of the nitrate ion in aqueous solution is slow relative to its rotational correlation

time.

As potassium nitrate was used as a source of nitrate ions in the later investigations

of polymer-nitrate binding, it was worthwhile to explore the relaxation times of this salt in

more detail. The concentration dependence of the viscosity corrected spin-lattice

relaxation times had previously been determined by Guarilloff (10). The relaxation times

varied linearly with concentration (Fig. 5.6) and appeared to support the conjectures of

Ibuki et al. (36) regarding the applicability of a hydrodynamic interaction model.

However, nitrate ions do not readily associate in solution because of their low

charge density. For example, a 0.1 M KNO3 solution would have only 3Vo of the nitrate

ion associated (37). Since the concentration dependence of the relaxation times is

expected to be linear when the fraction of ion pairs is much less than unity, it is not

surprising that a linear relationship is observed in this case.

An estimate of the association constant was obtained by assuming a simple 1:1

equilibrium. Activity coefficients for potassium nitrate in pure water were interpolated

from reported literature values (37,38), while density and viscosity were experimentally

by Guarilloff (10) on the actual solutions upon which the relaxation experiments were

of the a association

rf1c\J

performed. Linear regression of the variation

constants with the square of ionic strength KA = 0.063 + This is to be

of the

zu--'( f otw ,'(e. /

oo

o

b

o

(t)

E

F

140

130

r20

110

100

90

800 0.20 0.40 0.60 0.80 1.00

m (mol kg-l¡

Fig. 5.6. Viscosity corrected longitudinal relaxation times of potassium ninate versus

molal concentrarion (10). The dark line corresponds to the line of best ht obtained by

linear regression of the experimental data.

r40

r20

1m

80(/)

E

F60

40

20

00.2 0.4 0.6 0.8 1.0 t.2

C (M)

Fig. 5.7. Longitudinal relaxation times of simple nitrate salts in water at25 oC. Triangles

(A): silver nitrate; diamonds (0): calcium nitrate; circles (o): barium nitrate; squares @):

lead nitrate. The da¡k lines correspond to calculated relaxation times.

ô ôAô

ô

A

o

tr


results are qualitatively as expected based on ion associative grounds. For a given

concentration the observed relaxation times decrease such that Pb2+ < Ba2+ < CP+ <

Ag+. This is in agreement with the suspected decrease in "complexing power" of the

cations along this series. At higher concentrations the 2: I electrolytes exhibit significantly

lower relaxation times than the 1:1 electrolyte indicating that they participate in

significantly more ion pair formation. The relaxation times of the 2:1 salts decrease with

increasing charge density of the cation. This was again attributed to increased ion pair

formation.

However, in contrast to the 1:1 electrolytes, the evaluation of association

constants for 2:l electrolytes is complicated and necessarily involves some assumptions

regarding the activity of the singly positive charged species, MNO3+.

In the absence of any activity assumptions the stoichiometric association

constants, K5, were found to be concentration dependent and were extrapolated to C = 0

to give reasonable estimates of the association constants, but much lower than expected

(Table 5.4.).

Table 5.4. Extrapolated thermodynamic association constants foraqueous solutions of simple 2:1 nitrate salts at 25 oC.a

Cation Kg Kfl Kf3Kf2 ¡1''tr /' '' '' I

Ca2+

Ba2+Pb2+

0.38 + 0.010.63 + 0.021.01 + 0.04

1.96 t 0.041.68 + 0.113.67 + 0.10

1.50 + 0.031.34 + 0.082.81, + 0.07

1.64 + 0.031.56 + 0.072.88 + 0.07

a - the enor refers to one standard deviation of the mean.

The second set of extrapolated association constants K¡01, were derived by

assuming unit activity for the mononitrate species, which is perhaps the weakest of all the

assumptions and consequently these estimates give the highest association constants.

The f,rnal set of extrapolated association constants K$, assume that not only that

the activity of the mononitrate species could be approximated specifically by ammonium

nitrate, but also that the activity of the electrolyte was already reduced somewhat due to


ion pair formation and attempted to allow for this. Although this assumption did remove

some of the concentration dependence of the apparent association constants, the added

complication did not significantly alter the estimates of the thermodynamic association

constants relative to those obtained by assuming only that the activity of ammonium

nitrate approximated the activity of the mononitrate species, K$.

The apparent stoichiometric association constants were concentration dependent

and plots of log K5 versus the square of the ionic strength had a negative slope. The

extrapolated stoichiometric association constants were reasonable estimates of the

thermodynamic association constants even though they made no allowance for solution

activities.

In comparison, when allowance for activity was made, such as for K602, the

apparent association constants still exhibited some concentration dependence. However,

in this case plots of log K4 versus the square of the ionic strength had a positive slope. It

was clear that such activity assumptions overestimated the correction, although when

such plots were extrapolated to infinite dilution reasonable estimates of the

thermodynamic association constants were again obtained.

Since both the extrapolated thermodynamic association constants K! and Ka02

were prone to some error a better estimate of the "Eue" association constant was obtained

from their arithmetic mean. It is this value that appears in Table 5.5. This seemed justifred

since the logarithmic plots which enabled their calculation were shown to converge from

opposite directions, Fig. 5. 8.

Thermodynamic ion association constants for some simple nitrate salts have

previously been determined using a variety of methods and some are summarized in

Table 5.5. for comparison with the association constants determined in this work using

the nmr technique. Conductance measurements are widely used to study ion association

and consequently such measurements dominate this comparison. What i v/ø A4/.-

clear is that tance measurements association constants ofJ

magnitude larger tha¡r those obtained using alternate techniques.

Some of the conductance measurements were reported in the older literature and it

was initially tempting to simply assign the larger reported association constants to the

2

1

vÞo

0

-1

0.2 0.4 0.6 0.8 1.0 r.2 1.4 r.6

Fig. 5.8. Logarithm of apparent association constants versus the square root of ionic

strength of lead nitrate in water af 25 oC. Open circles (o): Ks ; filled circles 1o¡: Kez .

The dark lines correspond to lines of best fit obtained by linear regression of the relevant

data series.

o

o

Nuclear Magnetic Resonance

Table 5.5. Thermodynamic ion pair association constants of simple

nitrate salts in water at25 oC.

Cation. rf, (nmr) Kf, (other)

t28

Ag*

K+ 0.063 + 0.004 a

0.31 + 0.04

c*+ 0.94 + 0.02

Bp+ 0.99 + 0.05

Pbz+ 1.91 + 0.06

conductance b

conductance c

conductance d

conductance b

conductance c

conductance d

conductance e

Raman fconductance e

Raman Ispectroscopy h

polarography h

potentiomeüry h

nm¡ shift I

conductance j

0.73 + 0.010.7r0.58

1.06 + 0.010.630.512.04t.678.712.84

1.6 t 1.02.8 + 0.053.3 + 0.5

t.23315.1

a-

b-

determined using data of Gua¡illoff (10).

determined by fitting the data of McKenzie and Fuoss (39) tothe Lee-Wheaton model.

reference 40.

reference 37.

reference 42.

reference 43,0.5 M solution.

reference 43,O.4 M solution.

reference 44, constant ionic strength I = 2.0 M.

reference 45.

reference 46.

c-

d-

e-

f-û-

h-

i-j-


undeveloped state of conductance theory at that time. However, when the same data were

refitted using the latest Lee-Wheaton model (6,47-53), even larger association constants

were obtained. Notwithstanding that some of the association constants were determined at

consrant ionic strengrh (44) while others were extrapolated to infinite dilution (46), it was

clea¡ that association constants derived from conductance measurements were consistently

higher than expected.

Jones and Frost (43,54) have suggested that this discrepancy might result from

the existence of a two step equilibria in solution.

(Mn*)uq + (NO3-)uo <+ (M, H2O, NO3)uo(n-t)* <+ (MNO¡) "o(n-t)+

where the first species formed corresponded to a solvent-separated ion pair and the

second species formed is a contact ion pair. They suggested that association constants

derived from conductance measurements related to the solvent-sepa¡ated ion pair rather

than the contact ion pair. Consequently, the smaller association constants derived from

alternate methods, such as relaxation measurements, may relate to the formation of

contact ion pairs.

Irrespective of their relative magnitudes or method of determination the average

extrapolated association constants would be expected to exhibit identical trends.

Conductance and Raman measurements have shown the trend in association constants to

be

K+ < Ag+ <Ca2+ <B*+ <Pb2+

The association constants obtained in this work from relaxation measurements are in full

agreement with this trend.

The simple ion association model provided a good description of the concentration

dependence of the spin-lattice relaxation times of the nitrate nucleus in aqueous solution.

It was shown that such measurements could be used to determine association constants.

The most reliable results were obtained when allowance for activities were made and the


concentration range explored was not too high. A range having an upper limit near I M

seems to be most appropriate. Rotational correlation times calculated from the viscosity

corrected relaxation times determined here were in good agreement with literature values.

The slow exchange mechanism was found to fit the data slightly better and yielded more

reasonable estimates of the association constants. The association constants so derived

were in good agreement with literature values, with the exception of conductance

measurements. In this case the discrepancy was attributed to the existence of solvent-

separated and contact ion pairs in solution.

5.4.2 Relaxation in Benzyltrialkylammonium Salts.

Having confirmed the utility of the nmr method for simple nitrate salts it was

extended with some confidence to the relaxation times of the benzyltrialkylammonium

nitrates. These salts are of interest because they can be thought of as monomeric

analogues of the polymeric species. It is the polymeric species which are believed to be

nitrate selective, but this needed to be confi¡med. Conductance measurements on

benzyltriatkylammonium nitrates (6) showed no significant association and the relaxation

method provided an independent method of confirming this.

The concentration dependence of the relaxation times may expressed as a function

of four parameters,

Tt = Tt (C;Tr,Tu,Kn) (20)

where C is the molar concentration and Tf and T5 are the free and bound relaxation times

of the nitrate ion taken to be 130 and 9 ms respectively. In fitting the relaxation data T¡

and T6 were taken to be constants. The definition of the parameter K4 may change

depending on rhe nature of the equilibrium. Here the only equilibria which were

considered corresponded to an association constants resulting from a 1:1 as previously

described in section 5.2.4. Although it is possible to obtain estimates of stoichiometric

association constants, it is more correct to include activity coefficients to obtain estimates


of thermodynamic association constants. Thermodynamic association constant, K¡, may

be estimated from a weighted non-linear least squares f,rt of the concentration dependence

of the relaxation times using equation 20. Weighting was assigned by using the standard

deviation of Tr obtained from fitting of equation 19.

Alternatively an estimate of the association constant may be obtained by

extrapolation of apparent association constants to infinite dilution. This technique was

applied to simple ninate salts in the previous section in an attempt to minimizÊ any activity

contribution to the thermodynamic association constants for 2:1 electrolytes. When

applied to 1:1 electrolytes the least squares and extrapolation techniques gave identical

results within experimental error and consequently only the least squares results will be

presented here.

Experimental relaxation times were corrected for viscosity and are gathered in

Appendix A5. The relaxation times were found to fit equation 10 better than equation 11.

This was also found to be true by Nicholas and Wasylishen (8) for simple nitrate salts.

The implication being that the rate of chemical exchange is slow relative to the rate of

molecular rotation. In fitting the data the relaxation times of the free and bound nitrate ion

were taken to be 130 and 9 ms respectively. Two resultant fits are illustrated in Fig. 5.9.

and Fig. 5.10. The fits were extremely good given the inherent elrors in the determination

of relaxation times. It is also clear from these figures that the relationship between

relaxation time and concentration is not a linear one. A linear relationship ìù/as proposed

by Adachi et al. (9) for simple nitrate salts.

As with the simple nitrate salts, it was possible to obtain estimates of K4.

However, in the case of the simple salts the concentrations could be corrected using

literature activity coefficients even up to the fairly high concentrations employed. No such

activity data on benzyltrialkylammonium nitrates were available making such a correction

difficult. The activity coefficients were estimated using the extended Debye-Hückel law,

. _A{Tlos v+ = _____*+ b.I

1+Baril(2r)

o

oo

o

o

o ooo

140

120

100

60

40

20 0 0.20 0.40 0.60 0.80 1.00 r.20

C (M)

Fig. 5.9. Longitudinal relaxation times of benzyltrimethylammonium nitrate in water at

25 oC. The dark line corresponds to the calculated line of best fit of the data.

vtÉ..:, g0

F

140

120

100

60

o

40

0 0.10 0.20 0.30 0.40 0.50 0.60

C (M)

Fig. 5.10. Longitudinal relaxation times of benzyltripropylammonium nitrate in water at

25 oC. The dark line corresponds to the calculated line of best fit of the data.

20


Where A and B are constants which have been tabulated as a function of temperature

(37), ais the ion size parameter in cm as previously described in Chapter 3, and I is the

ionic strength of the solution. The b'coefficient is usually assigned a value so as to give a

best fit of the experimental determined activity coefficients. Lacking any experimental data

such a treatment was not possible. The b'coefficient was chosen, for each salt, so as to

generate activity coefficients that gave a best fit of the relaxation data. The use of activity

coefficients generally resulted in a better fit of the relaxation times when compared to the

fit obtained with no allowance for activity. The butyl and propyl salts exhibited a

marginally poorer fit when allowance for activity was made. The poorer fit almost

certainly resulted from the use of calculated activities in place of experimentally

determined values.

It seems clear that the calculated activity coefficients of the higher alkyl derivatives

may not adequately represent the actual activity coefficients of these solutions.q

Unfortunately, only the activity coefficients of the benzyltrimethylammonifum chloride

and bromide salts have been reported in the literature (55), and these differ significantly

from the activity coeff,rcients of benzyltrimethylammonium nitrate calculated here.

The tetraalkylammonium salts are the closest available analogues for which

activity coefficients have been reported for a variety of anions (56). The non polar

quaternary ammonium cations are believed to enforce the structure of water in their

vicinity (56). As the size of the alkyl substituent increases, so does the hydrophobic

nature of the cation. This results in a greater free energy of the ion and a corespondingly

larger activity coefficient. It might be expected that the activity coefficients of such salts

would be dominated by such cation effects. However, as can be seen in Fig. 5.11. the

activity coefficients of both the tetramethylammonium and benzyltrimethylammonium

salts change significantly with the anion present.

The lower activity coeff,rcients of the bromide salts relative to the chlorides has

been attribured to hydrophobic bonding. Hydrophobic cations and anions will tend to

combine with each other and minimize their interaction with bulk water; decreasing the

free energy of the system. This results in an increase in water structure in the vicinity of

the ions and a drastic lowering of the activity coefficients. Bromide is believed to

'(1 ,

l(r"'ltl^,Á

'1,

0.0

-o.2

+l>.; -0.4o

BzMe3N-NO3

o Me¿N-Cl

I

o

a

Me4N-Br

BzMe3N-Cl

BzMe3N-Br

-0.6

-0.8 o.2 0.4 0.6 0.8 1.0 1.2 r.4

{;

Fig. 5.1 l. Concentration dependence of the logarithm of the molal activity coefficients

of several methyl derivatives. The da¡k line corresponds to coefficients calculated here

using the extended Debye-Hückel relationship.


pafücipare in hydrophobic bonding; while chloride does not; presumably because it is too

hydrated (56). Tetramethylammonium iodide would also seem to participate in

hydrophobic bonding as it has activity coefficients lower than either the corresponding

chloride or bromide (56). Unfortunately, due to its low solubility in water it cannot be

directly compared with the other halides above 0.25 molal and does not appear in

Fig. 5.11.

It might have been expected that the replacement of a benzyl for a methyl group

would increase the hydrophobicity of the salt and result in a higher observed activity

coefficient. However, the benzyltrimethylammonium halides have activity coefficients

significantly lower than the corresponding tetraalkylammonium halide. In the case of the

bromide salt the decrease may be attributed to extensive hydrophobic bonding, but if

chloride cannot participate in such bonding, this does not explain its much lower activity

coefficients. Boyd et al. (55) offered no explanation for the lower activity coefficients of

this salt.

The calculated activity coefficients of the benzyltrimethylammonium nitrate

therefore, at first glance, seem unusually high. However, this may simply be due to the

specific nature of the anion. No activity coefficients of similar nitrate salts are available

for comparison. Indeed, in order to generate activity coefficients similar to those of the

benzyltrimethylammonium halides the b'coefficient must be negative and contrary to

theory. Coefficients so generated, for b' = -0.2 for example, also result in a poorer fit of

the experimental relaxation times. The increase in y+ with concennation is also consistent

with conductance data which suggested little association (6)'

The main reason for choosing calculated activity coefficients in preference to

another method was simply that the method of generating the activity coefficients needed

to be applicable to all four of the alkyl derivatives studied here. Lacking further

experimental data, the use of the Debye-Hückel extended equation provided such a

consistent method. The activity coefficients so generated do differ significantly from

those of the corresponding halides, but this may merely be due to some significant anion

effect in these salts. The calculated activity coefficients are not unreasonable and generally

result in a better fit of the experimental relaxation times when compared to their absence.

Nuclear Møgnelic Resonance I37

Regardless of the choice of activity coefficients, since all samples were dealt wiîh

identically, and being mainly rht relative differences between salts,

comparisons between samPles w possl ble without serious error. The calculated

relaxation times for four of the nitrate salts obtained from a best fit of the data in

Appendix A5 to equation 20 appear in Fig. 5.12. The results are qualitatively correct. For

a salt of a given concentration the longitudinal relaxation time decreased with the size òf

the alkyl substituent. This could be interpreted as an increase in nitrate binding with alkylt9

length. The increase in bindin gretndtßthe motion of the nitrate ion and rcsuIyf{ínthe

observed lower relaxation times. Also, for a given salt the relaxation time decreased with

concentration, consistent with an increase in ion pair formation at higher concentrations.

The association constants of the hve nitrate salts considered here are presented in

Table 5.6. The benzyltripentylammonium salt was fairly insoluble, limiting the

concentration range which could be studied by nmr. The two determinations were too few

to enable an accurate fit of the data. The reported K4 waS the average of the two K¡'s

determined for each corresponding relaxation time by manual substitution into equation

12. Since the concentrations were below 0.1 M, the activity coefficients were calculated

using the Debye-Hückel law, where the ion size parameter was taken from Steel et al. (6).

The thermodynamic association constants were always larger than the corresponding

stoichiometric constants and were of the same order as the values previously obtained

using conductance (6). The trends are identical regardless of which group of association

constants were considered, that is an increase in association with an increase in size of the

alkyl substituent, Fig. 5.13.

As noted for the simple nitrate salts the association constants derived from

conductance data were always larger than those obtained using nmr. This was again

attributed to the formation of solvent-separated and contact ion pairs in solution. Both

conductance and nmr methods yielded association constants of simila¡ magnitude and

supported the major conclusion of the previous conductance study (6). That is, that there

was no signif,rcant association of nitrate with these monomer analogues.

140

120

100

60

40

20 0 0.20 0.40 0.60

C (M)

0.80 1.00 1.20

Fig. 5.12. Calculated longitudinal relaxation times for benzyltrialkylammonium nitrates in

water at 25 oC using the best fit parameters for equation 20. Experimental points have

been excluded to avoid crowding. Examples of goodness of fit can be found in Fig. 5.9.

and 5.10.

U)Ée80F

@-N -R

3

1

3 24

1. R = Methyl

2. R = Ethyl

3. R = Propyl

4. R = Butyl

Table 5. 6. A s sociation con stants for benzyltrialkylammonium nitrate

salts in water at 25 oC.

Atkyl Group Kr aKlK1

Methyl

Etiryl b

Propyl

Butyl b

Pentyl c

0.24 + 0.01

0.70 r 0.09

0.91 + 0.02

1.88 t 0.11

3.08 + 0.17

0.41 + 0.01

1.08 + 0.12

1.30 + 0.03

2.29 + 0.r4

3.83 + 0.13

2.02 + 0.r8

2.45 + 0.rl3.28 r 0.10

3.80 + 0.09

4.46 + 0.14

a-

b-

c-

determin ed from cond uctance measurements (6).

calculated using data from reference 10.

The reported K4's were not detemrined by fitting, but werethe average of the K4's calculated from two measurements.The error refers to one standard deviation of the mean.

5

K¡ (nmr)

Kf; (nmr)

K! (cond)

1 ') 3 4 5 6

4

3

2

!4

I

00

n

Fig. 5.13. Variation of association constant, K4, with n , the number of carbon atoms ln

the alkyl group. The shaded symbols denote thermodynamic association constants while

the unshaded symbols assume unit activity.

N uclear Magnetic Reso nonce

5.4.3 Relaxation in Linear Polymer-Nitrate Solutions.

140

The longitudinal relaxation times of the nitrate ion were determined in mixtures of

linea¡ chloride polymer and potassium nitrate at various concentration ratios. In the case

of the monomer analogues and the simple nitrate salts, where the viscosity of the samples

was not very different from that of pure water, it was acceptable to use literature or

experimental viscosities for the samples in pure water, rather than the viscosity of the

sample inHzOlDzO. However, for polymer samples, where the viscosity effects become

more important, it was decided to determine the density and viscosity of exactly the same

solution as was used to determine the longitudinal relaxation times.

Densities and viscosities were determined using the methods described in

Chapter 4. Densities determined here were consistently larger than those of the

corresponding monomer analogue solutions at the equivalent molarity due to the addition

of both D2O and KNO3 cosolute. For this reason variations in both density and viscosity

were independent of polymer concentration.

Since this study is primarily concerned with ion association, which is believed to

decrease T1 from the free nitrate value of 130 ms (8), potassium nitrate provided a useful

source of nitrate ions because it has a negative viscosity coeff,rcient. This means that any

decrease in T1 is more likely to be due to association rather than microviscosity effects.

Although the relaxation times are corrected using the bulk viscosity it is not unknown for

microviscosity effects to be significant. The results are summarizedín Appendix A5. In

the case of the poly pentyl derivative the solubility of the polymer-nitrate complex was too

low to permit any relaxation measurements. When samples were rerun at a later date the

experimental T1's were reproducible within 3Vo, well within the acceptable enor (18).

This difference was most likely due to either changes in field homogeneity or slight

variations in temperature between determi nations.

In all cases the presence of linear polymer caused a decrease in the observed spin-

lattice relaxation time of the nitrate ion relative to that expected for pure potassium nitrate

in the absence of polymer. This could most readily be attribute to nitrate association of

some kind which hindered molecular rotation. The exact nature of this association was

1.0

0.8

0.6

o.2

0.0

0.4

0 0.002 0.004 0.006 0.008 0.010

lPolyl C(M)

Fig.5.14. Fraction of nitrate bound, A, versus polymer concentration for the

poly(vinylbenzyltrialkyammonium chloride) - potassium nitrate system. Circles (O):

methyl; triangles (A): ethyl; diamonds (0): propyl; squa.res (o): butyl.

cA

ô

o

o

ô

o oBo C

o oOO o oOO oo o

ooo o o oo

oo

Fig. 5.15. Graphic representation of the coiled structure of a linear polymer illustrating

three types of possible binding. A) site binding, B) localized binding and C) hindered

binding. The positive charges denote fixed polymer exchange sites and the negative

charges denote free cosolute anions.

o

oo

o

oo

Ao

OO

o

oo

oo


counterion binding may involve rhree distinct types of binding :

(a) Site binding - specific binding between counterions and a specific charged

exchange site along the resin backbone.

(b) Local binding - nonspecific electrostatic binding where counterions are

localized in the vicinity of the resin backbone by the potential of the polyion.

(c) Hindered binding - the coiled structure of the polyion occludes counterions

and hinders movement.

These three types of binding are illustrated in Fig. 5.15. While site binding could

conceivably be examined by some association constant it was not possible to

unambiguously distinguish between this and the other types of binding occuring in

solution.

As a first approximation it was assumed that the moles of polymer would

correspond to the maximum number of exchange sites available for site binding. The

concentration ratio of polymer to nitrate therefore corresponded to the maximum fraction

of nitrate bound if only site binding were in operation. It can be seen from the average

ratios of A to Ama* shown in Table 5.7. that all of the polymers considered here had

fractions of nitrate lar due to slte This ratio called the

localized binding rario A¡, indicated that all the polymers participated in local or hindered

binding of nitrate. This trend increased with the size of the alkyl functional group

indicating significant localized binding by the larger polyelectrolytes.

Table 5.7. Variation of localized binding ratio with alkyl length foraqueous chloride polymer - nitrate solutions at25 oC-

Alkyl Group A¡ = Au/A¡¡¿¡

MethylEthylPropylButyl

1.6 + 0.22.8 + 0.45.4 + 0.68.5 + 1.1

In practise it is unlikely that all binding sites would be exclusively occupied by the nitrate

ion and consequently A6¿¡ would be reduced. Emf measurements (10) suggested that at

nitrate to polynìer mole ratios near 2 over 80Vo of the nitrate was bound. So the above


assumption appears to be reasonable. This study also found the same order in the

retention of nitrate when the size of the alkyl group was varied.

Me<Et<Pr<Bu

For a given concentration of polymer the fraction of nitrate bound increased with

nitrate concentration. When the KNO3 concentration was sufficiently large, so as to

ensure a vast excess of free nitrate, the determined T1 values were identical within

experimental error (Table 5.8.). This occurred when the mole ratio of nitrate to polymer,

R, was greater than approximately 100.

Table 5.8. Variation of spin-lattice relaxation times for aqueous solutions of

poly(vinylbenzyltriethylammonium chloride) in the presence of a vast excess

of potassium nitrate at25 oC.

polymer

C (M)

mtrate

C (M)

Density

G cm-:¡

T1 corr b

(ms)

Ra Tlr

0.001164

0.001016

0.001061

0.001036

0.302496

0.208758

0.095505

0.04933s

267.7

209.7

90.9

47.8

1.026855

t.020494

1.013553

1.010790

1.023585

t.026079

r.044456

1.055577

93.4 + O.9

90.7 + 0.8

91.6 + 0.6

67.9 + 2.r

a- R is the mole ratio of nirate to polymer.

the error refers to one standard deviation of the mean.b-

5.4.4 Relaxation in Cross-linked Polymer-Nitrate Solutions.

The relaxation times of the nitrate ion were determined in mixtures of cross-linked

chloride polymer and potassium nitrate at various degrees of cross-linking. All of the

samples appearing in Table 5.9. were prepared using divinylbenzene, DVB, as a

N uclear Magne tic Resonance 146

Table 5.9. Variation of spin-lattice relaxation time with percentage crosslinking ofpoly(vinylbenzyltriethylammonium chloride)-divinylbenzene in potassium nitrate at

25 0C.

A-Vo a masspolymer (g)

mass nirate(e)

KNO3C (molal)

B b T1 expt. c

KNO3

0 (linear)

1.5

1.5

3.8

5.7

7.9

7.9

9.4

2t.6

2r.6

0

0.00117

0.000960.00096

0.0r400.01400.0125

0.00133

0.00117

0.00116

0.0093

0.00158

0.001480.00148

0.0115

0.0212

0.02220

0.020600.02060

0.04160.04160.0436

0.02126

0.02127

0.02126

0.0212

0.02124

0.020550.02055

0.0213

0.04682

0.04905

0.050240.05024

0.091930.091930.09770

0.04501

0.o4784

0.04606

0.0521

0.04655

0.050150.05015

0.0518

t29.1 + 9.9

67.9 + 2.1

oo

47.8

53.9s3.9

40.1

45.6

46.0

5.7

33.7

34.934.9

4.6

46.648.165.9

r12.6

104.5

111.8

108.9

1t6.7

113.4r16.4

T13.7

+ I.7+ 1.9+ 2.0

+ 8.4

+ 7.6

+ 8.4

+ 8.3

+ 9.7

+ 9.9+ 9.7

+ 8.3

8+7.28+7.0

97.00.1

10.010.08.8

a-

b-

c-

The percentage cross-linking was estimated from the mole ratio ofdivinylbenzene to polymer.

R was the mole ratio of nitrate to polymer assuming a base My¿ for the cross-linked polymer identical to that of the linear polymer.

The T1 values were uncorrected for viscosity and the elrors referred to onestandard deviation of the mean.


cross-linking agent. This resulted in the formation of an insoluble polymer gel which

made any discussion of polymer concentration in solution nonsensical. The addition of

DVB also meant that the base molecular weight was not know with any cert;inty.

However, because the amount of DVB added was relatively small the base molecular

weight was assumed to be identical to that of the linea¡ polymer.

The spin-lattice relaxation times were not decreased significantly from that of free

nitrate and consequently the fraction of nitrate bound was extremely low. This suggested

that the large decrease in T1 observed for the linear polymers was predominantly due to

an increase in hindered binding. The more rigid structure envisioned to occur due to

cross-linking in these polymer gels, being far less flexible than the linear counte{part, was

unable to trap as much nitrate. The fraction of nitrate bound decreased with the amount of

cross-linking, Fig. 5.16, in agreement with an inability to hinder nitrate rotation.

Although there was a some decrease in the experimentally determined T1, by all

cross-linked samples compared to that of free nitrate, it is unclear if any significant

conclusions can be made. This was because the samples were not completely soluble in

water and the resultant gel made the determination of viscosity impossible. In the absence

of viscosity measurements it was not possible to unambiguously attribute the observed

decrease in longitudinal relaxation time to nitrate association or to an increase in solution

viscosity. Because of the limitation of the nmr technique when applied to cross-linked

gels only poly(vinylbenzyltriethylammonium chloride) was considered in any detail.

Gel samples may more conveniently be studied using ion-exchange or

electrochemical techniques. Emf measurements (10) in the concentration range where

nitrate was present in mole ratios less than or equal to unity, indicated that nitrate

association tended to increase with the amount of cross-linking and also with the size of

the alkyl group. For poly(vinylbenzyltriethyammonium chloride), 97o cross-linked with

DVB an average selectivity coefficient of nitrate over chloride, K[ = 3.14, was

suggested. This study also indicated that higher degrees of cross-linking resulted in a

lower capacity for nitrate. If the results of the emf study are correct it may indicate that the

gels studied here have a low capacity for nirate. Since the observed spinJattice relaxation

time is a weighted mean of all relaxation times in solution between bound and free

0.03

o.02

0.01

0.0010 30

CL Vo

Fig. 5.16. Fraction of nitrate bound versus the degree of cross-linking for a lightly

cross-linked poly(vinylbenzyltriethylammonium chloride) - potassium nitrate system. The

concentration of potassium nitrate was held constant near 0.05 M and approximately

0.001 g of polymer were added in each case.

200

o

o

o

ooo

oo


extremes. It seems likely that the observed relaxation time will be weighted significantly

towards that of the predominantly free potassium nitrate and the resultant decrease in Tt is

small even though the resin may bind signif,rcant amounts of nitrate relative to its capacity.

5.5 Conclusion.

The measurement of spin-lattice relaxation times using 14N nmr provides a quick

and easy method for the determination of nitrate association. The method is limited only

by the time taken to perform density and viscosity measurements, which are necessary to

co11ect relaxation times. However, when the solutions are sufficiently dilute and only a

comparison between two solutions of similar viscosity is required the corrections can be

ignored without serious error.

Atl the association constants determined in this work gave excellent agreement

with a vêlu,ag whçn--aç-t-i-Yity-.ç9rreç[i94s were made. Even in the absence

of such conections the determined K4's gave agreement of the same order as the

expected values. The results were also found to agree well with a concurrent study using

emf measurements (10).

N uc lear Mag netic Resonance

Literature Cited.

10

G. A. Guter, U.S. Pat. 4,479,877 (30 October 1984).

M. B. Jackson and B. A. Bolto, React. Polym., 12,277, (1990).

R. E. Hester and R. A. Plane, J. Chem. Phys.,40(2),411, (1964).

M. V/hittle and J. H. R. Clarke, Chem. Phys' Len.,98(2), 172, (1983).

R. L. Darskus, D. O. Jordan and T. Kurucsev, J. Chem. Soc. - Faraday Trans'

62(526),2876, (1966).

B. J. Steel, A. S. Kayaalp, T. Kurucsev, A. D. V/a¡d and M. B. Jackson, Aust.

J. Chem., 43, 1.983, (1990).

R. M. Fuoss and K-L. Hsia, Proc. Natl. Acad. Sci. U.S,A., 57, 1550,

(re67).

A. M. P. de Nicholas and R. E. Wasylishen, Can. J. Chem.,65, 951,

(1e87).

A. Adachi, H. Kiyoyama, M. Nakahara, Y. Masuda, H. Yamatera, A. Shimizu

and Y. Taniguichi, J. Chem. Phys.,90(1), 392, (1989).

P. Guarillofl Ph.D. Thesis, The University of Adelaide, to be submitted.

C. N. Banw ell., Fundamentals of Molecular Spectroscopy 3rd edn.,252

(McGraw-Hill: London 1983).

R. T. Boeré and R. G. Kidd, Annu. Rep. NMR Spectrosc. (Ed. G. A. Webb)

L3,319 (Academic Press: London 1982).

M. Witanowski, L. Stefaniak and G. A. Webb, Annu. Rep. NMR Spectrosc.

(Ed. G. A. Webb) 18, 201 (Academic Press: London 1986).

V/. T. Huntress Jr., Adv. Magn. Reson., 4, l, (1970).

W. Egan, J. Amer. Chem. Soc.,98(L4), 4091, (1,976).

B. J. Steel, private communications, (1993).

R. C. Weast, (Ed.), Handbook of Chemistry and Physics 50th edn. (Chemical

Rubber Company: Ohio 1970).

M. L. Martin, J. Delpuech, and G.J. Ma¡tin, Practical NMR Spectroscopy,

(tleyden: London 1980).

11

t2

13.

t4.

15.

16.

17.

150

1

2

3

4

5

6

7

8

9

18.

N uc lear M ag ne t ic R e so nanc e

19. R. K. Harris and R. H. Newman, J. Magn. Res',24,449, (1976).

20. T. Kurucsev and E. H. Williams, J. Crystallogr. Spectrosc. Res., L2(5),397,

(1e82).

21. G. C. Levy and I. R. Peat, J. Magn. Res., 18, 500, (1975).

22. T. Kurucsev, J. Chem. Educ., 55, 128, (1978).

23. A. M. P. de Nicholas and R. E. Wasylishen, J. Phys. Chem.,89(25), 5446,

(1e8s).

24. V. M. M. Lobo, Handbook of Electrolyte Solutions Physical Sciences Data4L,

@lsevier: Amsterdam 1989).

25. G. Jones and S. K. Talley, J. Am. Chem. Soc.,56,624, (1933).

26. T. Isono, J. Chem. Eng. Data.,29,45, (1984).

27. E. W. Washburn (Ed.), International CriticalTables Vol. V. lst edn.,

(McGraw Hill: New York 1929).

28. E. I. Chernen'kaya and O. M. Kuznetsova, Zh. Prikl. Khim.,52(6), 1255,

(1979); J. Appl. Chem. (Engl. Transl.) 1188, (1979).

29. M. Nakahara, A. Adachi, H. Kiyoyama, A. Shimizu, Y. Taniguchi and Y.

Masuda, J. Phys. Chem.,94, 6L79, (1990).

30. A. Fratiello, V. Kubo-Anderson, S. Azimi, E. Marinez, D. Matejka, R.

Perrigan, and B. Yao, J. Sol. Chem', 20(9),393' (1991).

3I. A. Fratiello, V. Kubo-Anderson, T. Bolinger, C. Cardero, B. DeMerit, T.

Flores, D. Matejka and R. D.Perrigan, J. Magn' Res., 83, 358' (1989).

32. A. Fratiello, V. Kubo-Anderson, S. Azimi, T. Flores, E. Marinez, D. Matejka,

R. D. Perrigan, and M. Vigil, J. Sol' Chem.,19(8)' 811, (1990).

33. A. Fratiello, V. Kubo-Anderson, S. Azimi, F.Laghaei, R. D. Perrigan, and F.

Reyes, J. Sol. Chem.,21(10), 1015, (1'992)-

34. M. Gourdji, L. Guibé, and A. Peneau., J. Phys. (Les Ulís, Fr.)'35,477,

G97Ð; Chem. Abstr., 81, 113338 , (1974).

35. T. Kato, J. Umemura, and T. Takenaka, Mol. Phys.,36, 621, (1978).

36. K. Ibuki and M. Nakahara, J. Chem. Phys.,90(1)' 386' (1989).

151


37. R. A. Robinson and R. H. Stokes, Electrolyte Solutions 3rd edn.

(Butterworths :London 1 965).

38. G. W. C. Kaye, and T. H. Laby, Tables of Physical and Chemical Constants

15th edn., 255 (Longman: London 1986).

39. I. D. McKenzie and R. M. Fuoss, J. Phys' Chem.,73(5)' 1501, (1969).

40. R. M. Smith and A. E. Martell, Critícal Stabiliry Constants Vol.4 (Plenum

Press: New York 1975).

4I. G. Jones and J. H. Colvin, J. Am. Chem.\oc.,62,338' (1940).

42. L. G. Sillén and A. E. Martell (Eds.), Stabiliry Constants of Metal-lon

Complexes. Special Publication No. 25., 93 - 1O4 (The Chemical Society:

l,ondon ß7Ð.

43. D. W. James and R. L. Frost, Aust. J. Chem.,35, L793, (1982).

44. H. M. Hershensotr, M.E. Smith, and D. N. Hume, J- Am. Chem. Soc','75,

507, (1953).

45. P. G. Harison, M. A. Healy, and A. T. Steel, J. Chem. Soc. DaltonTrans.,

1845, (1983).

46. G. H. Nancollas, J. Chem. 9oc.,1458, (1955).

47. W. H. Lee and R. J. Wheaton,,/. Chem. Soc. Faraday Trans.II,74(4),743,

(1e78).

48. W. H. Lee and R. J. Wheaton,,I. Chem. Soc. Faraday Trans.11,74(8),1456,

(1e78).

49. V/. H. Lee and R. J. Wheaton, "I. Chem. Soc. Faraday Trans.II,75(8>,1L28,

(1e79).

50. R. J. V/heaton, /. Chem. Soc., Faraday Trans. 11,76,1093 (1980).

51. R. J. Wheaton, ,I. Chem. Soc., Faraday Trans. 11,76, 1599 (1980).

52. R. J. V/heaton,,I. Chem. Soc., Faraday Trans.II,77,631 (1981)'

53. R. J. Wheaton, ./. Chem. Soc., Faraday Trans. II,'77,1343 (1981).

54. R. L. Frost and D. W. James, J. Chem. Soc., Faraday Trans. 1,78,3263,

(re82).


55. G. E. Boyd, A. Schwarz. and S. Lindenbaum, ,/. Phys' Chem.,70(3),821'

(1e66).

56. S. Lindenbaum and G. E. Boyd, J, Phys. Chem.,68(4)' 911' (1964)'

57. P. Molyneux (Ed.), Water-soluble Synthetic Polymers: Properties and Belnviour,

Vol II, Chapter 2, (CRC Press: Boca Raton 1984)-

58. S. A. Rice and M. Nagasawa, Polyelectrolyte solutions, chapter 9,

(Academic Press: New York 1961).

59. J. W. Lyons and L. Kotin, f . Amer. Chem. Soc., 87(8), 1670, (1965).

154

6

Conclusion.

A number of ion-exchange resins have been proposed to remove nitrate from

contaminated drinking water. Resins which contained a quaternary ammonium functional

group were shown to be most selective for nitrate but the reason for this selectivity was

poorly understood. This work was the first systematic study of nitrate association of

benzyltrialkylammonium type resins.

Initially, monomer analogues were prepared to determine the role of the functional

group in influencing selectivity. The benzylrialkylammonium series of salts were chosen

and studied up to and including the pentyl derivative. Solubility measurements were

conducted on saturated solutions of these salts for a range of counter ions including

nitrate. Such high concentrations were thought to most closely mimic the equilibrium at

the resin-ion interface in an ion-exchange resin. It was found that all nitrate salts had

unusual thermod when compared to corresponding sulfates and halides.

In addition, b gt ryJt¡_U-u tfJ44qglon ig q nitrate also had a much lower and free

energy of solution than its nearest neighbours in the alkyl series. This was interpreted as

an increase in hydrophobic bonding which enforced the structure of water in the vicinity

of the cation. Such an effect may be important in ion-exchange resins where the linked

charge density and close proximity of functional groups could lead to cooperative

enhancement of this hydrophobic effect.

Density and viscosity measurements were used in parl to determine the hydration

numbers of these salts. From viscosity B coefficients, the hydrodynamic volumes of the

solutes were calculated and compared to the apparent molar volumes obtained from

density measurements. Hydration numbers of these salts increased with the size of the

alkyl group attached to the nitrogen atom. This was consistent with an increase in

clathrate-like water cage formation as suggested by solubility measurements. A monotone

Conclusion 155

increase in nitrate association with alkyl length was expected from conductance

measurements. This monotone increase in hydration with alkyl length suggested that

nitrate association may in part be influenced by hydrophobic hydration.

Apparent molar volumes determined here from density measurements were in

good agreemenr with published values. B coefficients determined from viscosity

measurements showed similar trends to those observed for teuaalkylammonium salts and

it was concluded that the solution properties of the benzyltrialkylammonium salts

considered here were not signihcantly different from those of tetraalkylammonium salts.

A new method for the determination of nitrate association in solution was

developed which utilized 14N spin-lattice relaxation measurements. The method was

initially tested by studying the association of a number of 1:1 and2:1' simple nitrate salts.

Where comparisons were possible derived rotational correlation times, fç, wero in perfect

agreement with literature values. The concentration dependence of the relaxation times

indicated that the slow exchange model was the preferred model. It was thus concluded

that chemical exchange was slow relative to the rate of molecula¡ rotation of the nitrate

ion. It was a major conclusion of this work, that contrary to some doubts expressed

recently, a simple ion pair formation model could be used to explain the concentration

dependence of spin-lattice relaxation times. Association constants so derived were in

good agreement with with those obtained using alternate methods.

It was observed that association constants derived from conductance

measurements were consistently larger than those obtained using other methods. This

was attributed to the formation of a solvent-separated ion pair in solution while the lower

association constants determined here from relaxation measurements corresponded to a

contact ion pair.

The nm¡ technique was then extended to the monomer analogues. The slow

exchange model was again found to fit the data better than the fast exchange model. The

derived association constants were of the same magnitude as those obtained from

conductance measurements, although again consistently lower due to contact ion pair

formation. Nitrate association was found to increase with the size of the alkyl substituent.

However, aS was found for conductance measurements, this increase was not as dramatic

Conclusion 156

as that observed for ion-exchange resins. This indicated that the polymeric structure of the

resin together with the functional group was essentiai in determining nitrate selectivity.

Some previous studies have suggested that resins containing a tributyl quaternary

ammonium group have the highest preference for nitrate. In contrast all techniques

explored here, with the exception of solubility, indicated that the preference for nitrate

continued to increase as the size of the alkyl group was increased above the butyl

derivative. Any selectivity observed by monomer analogues was enhanced by

polymerization.

When applied to linear polymers the viscosity corrected spin-lattice relaxation

times indicated that the fraction of nitrate bound increased with the alkyl length. The

fraction bound was significantly larger than that expected based on the available exchange

sites. This indicated that, while a large proportion of nitrate was undoubtedly bound to

the polymer, the relaxation times referred mainly to hindered nitrate trapped in the tertiary

coil structure of the polymer and not necessarily to site bound nitrate. Relaxation

measurements on lightly cross-linked resins, which are less flexible, showed a dramatic

decrease in the fraction ofnitrate bound and tended to support this conjecture.

A concurrent study utilizing emf measurements was more suitable for the study of

nitrate association in cross-linked samples and indicated that the preference for nitrate

increased with both alkyl length and amount of cross-linking.

This study has shown that the selectivity of the monomer analogues for nitrate is

enhanced by polymerization, which in turn is further enhanced by the degree of cross-

linking. However, as cross-linking increases the capacity for nitrate decreases. It is hoped

that this work will be the prelude to the design of a commercial ion-exchange resin having

greater affinity for nitrate while maintaining a high capacity. Five membered cyclic

quaternary ammonium compounds based on N,N, dialkyl-3,4, dimethylpyrrolidine

currently seem to offer the best chance of accomplishing this.

A-l

Appendix A0

Selectivity Parameters.

Some of the more extensive data reviewed in ùhe introduction which described

previous studies of nitrate selective resins are presented here.

Table 40.1 Selectivity coefficients of poly(alkylamine) resins.

Table 40.2 Selectivity coeff,rcients of polystyrene-divinylbenzene weak base resins.

Table 40.3 Selectivity coefficients of polystyrene-divinylbenzene resins containing

quaternafy ammonium functional groups.

Table 40.4 Relative retention times of nitrate and sulfate anions on

polystyrene-divinylbenzene trialkylammonium anion exchange resins.

Table 40. 1 . Selec tivity coeffrcients of poly (alkylallyamine) resins. a

Resin Type CL(vo) 6- "f "! c'[ c'$ KU NSS Ref.

p¡¡q b

P¡4sPA.,qbPPTDAA b

PTAA b

IrTAAPMeTAA Qo0Vo)PMeTAA Qo22Vo)PMeTAA (Do577o)

PMeTAA Qo\3Vo)QrßxAQrßxAQrßxADADMAC-QHÐ(ADADMAC - QIßXA

ll d

13d

2.55

5

25

3;48

652030

0.940.540.460.420.300.550.870.920.390.49

2.0d3.7 d

6.6 d

1.3 d

2.0d

9.57l

16053047039

165

45012

34

-r.00.10.51.1

1.2

-0.2

0.30.7

-0.s-0.2

4.51.81.41.6

r8.013.08.08.0

r.7r.8t.72.2

0.09 c

0.14 c

0.21c0.28 c

0.7 d

3.0 d

5.6 d

1

III

2,32,32,32,32,32,32,32,32,32,3

r- amine), PPTDAA = Poly(propyldiallylamine), PTAA =amine quaternised X%a withdodecylbromide, DADMAC =

a¡y (Me) FIEXA.

calculated from reference l.reference 2.

b-c-d-

Appendix A0 A-3

Table A0.2. Selectivity coefficients of polystyrene-divinylbenzene weakbase resins.

pl p2C

(meq/ml)

cx,N rlt " NSS b Ref'

MeEtPr

i-PrBuj-BuPei-Pej-Pe

MeEtPr

i-PrBui-BuPei-Pei-Pe

3.8r.020.800.531.1

0.950.551.0

0.92

0.3 2082

148380370

t2001200750750

-r.27-0.100.270.850.501.10r.330.870.91

8.26.2

10-1616-r98-10

422aJ

2,3JJJ2

a- at 20 oC except for reference 4.

b- calculated using NSS = log Kf - log e +1

Table 40. 3. S electivity coeff,rcients of poly styrene-divinylb enzeîe re sin s

containing quaternary ammonium functional groups.

pl B2 p3 K NSS Ref.

(meq/ml)

NSe

Mel\4eN4e

EOHEtEtEtEtPrPrBuBuPei-Pe

MeMe

EOHEOH

EtEtEt

EOHPrPrBuBuPer-Pe

MeEOHEOHEOH

EtEt

EOHEOH

PrPrBuBuPei-Pe

r.4lr.42t.4lr.23

100501010

100040010050

17001 100

230001 1000

16001200

-0.14-0.45- 1.15-1.09+0.92+0.80-0.11-0.4r

+1.4

+2.7

+I.7

1.190.661.291.300.62

0.430.660.320.19

55555255)a6J7

2,33

a- EOH=CH2CH2OH.

Appendix A0 ^4

Table 40.4. Relative retention times of ni[ate and sulfate anions on

polystyrene-divinylbenzene trialkylammonium anion exchange resins.a

Resin pl p2 p3 e tnCI tNO3 tSO¿

(meq/g) (mins)

TMA b

TMA C

TMA d

TMA E

DMEA b

MDEAbMDEAdTEA b

TEIA b

TPA b

TBA b

TBA d

TBA E

TBP E

TT{A b

TOA b

SATMA (n=1) fSATMA (n=2)SATMA (n=3)SATMA (n=4)SATMA (n=6)

N,le

N4e

N4e

N{e

MeN4e

N4e

EOHEtPrBuBuBuBu

HexOctN{e

l\{eN{eN4e

Me

Me

1,1"N{e

IvIe

EOHEOHEOHEOH

EtPrBuBuBuBu

HexOctlvle

N{eN{eMeI\'Io

0.0270.0460.0920.0900.0250.0260.0900.031o.0260.026o.o270.0960.0930.0970.0280.0280.0490.0520.0500.0470.061

8.33.r3.0

0.859.48.42.518.68.0

10.69.4r.66r.431.331.33r.62

1.301.56r.932.49r.281.221.481.38t.321.38r.544.323.684.3rt.63r.723.623.072.462.432.65

5.684.772.09

4.64

4.212.422.32

7.368.278.538.508.s0

ì\4e

l\{eMeN,Ie

lvfe

EOHEOHEOH

EtPrBuBuBuBu

HexOctlyle

N{eN{eN4e

Ivfe

9.58.93.41.11

1.19

b-

a-

f

d-

e-

All resins except SATMA have the strong base structurg giygn by Fig 1.7.lethanolamine;MDEA =

amine; TEtA = triethylamine;TPA =BP = tributylphosphine; TFIA =H=CH2CH2OH. All retention times are

given relative to chloride.

reference 8; eluent = benzoic acid.

reference 9; pH 6.0; eluent = 0.2 mM potassium phthalate.

reference 9; pH 6.0; eluent = 0.4 mM potassium phthalate.

reference 10; pH6.0; eluent =0.22 mM sodium phthalate.

reference 11; SATMA = spacer arm trimethylammonium having the structure inFig. 1.8., where n denotes the number of methylene groups in the spacer alm.pH 6.0; eluent = 0.2 mM sodium phthalate.

Appendix A0

11.

A-5

t

2

J

Literature Cited.

B. A. Bolto, M. B. Jackson and R. V. Siudak, Isr. J. Clrcm., 26, 17 , (1985).

M. B. Jackson and B. A. Bolto, React. Polym., 12,277 , (1990).

M. B. Jackson and L. J. Vickers, Effect of structure on nitrate selectiviry. Part 2

Resins with either sryrenelDVB, PECH or allyamine backbones., CSIRO Report

No. DB-127 (April 1987).

D. Clifford and W. J. Weber, Jr., React. Polym., L,'77, (1989).

G. A. Guter, Removal of Nitate from ContaminatedWater Supplíes for Public

U s e., EP A-600/2-82-042 (March 1982).

G. A. Guter; quoted by.M. B. Jackson, ín Effect of structure on nitrate

selectívity. Part I. Resins with undecanoic groups., CSIRO Report No.

DA/B-I11, pg 5 (August 1986).

G. A. Guter, U. S. Patent 4,419,877 (30 October 1984).

R. E. Barron and J. S. Fritz, J. Chromatogr.,284,13, (1984).

R. E. Baron and J. S. Fritz, J. Chromatogr.,316,20I, (1984).

L. M. Warth, R. S. Cooper and J. S. Fritz, J. Chromatogr., 479, 40L,

(1e8e).

L. M. Warth and J. S. Fritz, J. Chromatogr. 9ci.,26, 630, (1988).

4

5

6

7

8

9

10.

A-6

Appendix A1

Preliminary Synthesis' of Improved Capacity Resins.

Although the polymers of vinylbenzyltrialkylammonium salts have shown some

utility as nitrate selective resins (1) they suffer from a low capacity. A more commercially

viable polymer, with a higher capacity for nitrate, could possibly be obtained by siting the

nitrogen atom within a five membered ring. The salts of the N,N dialkyl, 3,4, dimethyl

pyrrolidine fulfill these conditions and would be suitable monomeric analogues

(Fig.41.1.).

H3n

N

R R

monomeric analogue polymer

Fig. A1.1. Monomeric and polymeric compounds based on N,N dialky-3,4,

dimethylpyrrolidine.

Preliminary reactions have shown that of such monomeric model compounds

could be obtained from maleic anhydride by following Scheme 41. Some preliminary

steps in this synthesis are outlined below. The synthesis of such monomeric analogues

was abandoned when it was found that no significant selectivity for nitrate was exhibited

by benzyltrialkylammonium salts (2). It appears that selectivity is not largely influenced

by functionality, but results from specific interactions in the polymeric material.

Consequently, a study of these monomer analogues seemed unwarranted.

{

o o

o

Q N

2 + N + 2CO2 + (1)

NH2 I

NH-R

H

o o

HH

N-RLiAlHa<_

HH

IV

HHR

o N xO vrR

-RG

HH

(1)

H2lcatalyst

+

o

NHzl+nY

oMe MeMe

Hzo

H

N-R

v

+RX

oII

oIII

H

Scheme A1

Appendix AI

Preparation of Dimethymaleic Anhydride (I).

A-8

The preparation of dimethylmaleic anhydride closely followed the method

described by Baumann et al. (3). The mechanism of this unusual reaction has also been

reported (4). A typical preparation is given below.

2-aminopyridine (20.12 9,214 mmol) was dissolved in glacial acetic acid (46 ml)

in a round bottom three necked flask and the solution brought to the boil. Maleic

anhydride (41.36 g,426 mmol) was dissolved with heating on a steam bath in glacial

acetic acid (66 ml) and added dropwise to the boiling 2-aminopyridine solution. During

addition the solution darkened to the colour of port. The mixture was refluxed for t hr.

Acetic acid was removed by distillation at room temperature (b.p. = 115 oC). 4 M

HzSO¿ (106 ml) was added to the remaining solution and the combined solution was

refluxed for t hr. The resultant solution was poured into a preheated beaker and allowed

to slowly cool to room temperature and crystallize out overnight. The solid was collected

by buchner f,rltration and washed twice with ice cold 4M HzSO¿ to give 17.55 g (65Vo) of

white product. m.p. = 92-94 oC, lit. 93-94 oC (3,4). TLC: Rp (CHzClz) 0.60. lH NMR:

õ(60 MHz, CDCI3) 2.03 (s, 6H, CH3). 13C Ntr¡R: ô(CDCI:) 165 (s, C=O); 140.5 (s,

C=C); 9.102 (s, CH¡). IR (NaCl): 1870; 1.825:1800; 1755; 1695 cm-1. Mass spectrum

mlz 127 (70, CoHeO¡H+); 82Ø2);54(100); 39(55).

Preparation of N-alkyl-3,4-dimethyl maleimides.

A large number of N-alkylmaleimides have been prepared from the corresponding

maleamic acids (5-19). Two of the latest methods of synthesis were considered here for

their suitability in the preparation of N-alkyl-3,4-dimethylmaleimides (III) from

dimethylmaleic anhydride (II).

In a method proposed by Mehta et al. (9), maleamic acids were initially prepared

by reaction of a primary amine with maleic anhydride. Subsequently the isolated

maleamic acids were cyclized using the mild conditions described by Searle et al. (5).

Appendix Al A-9

Mild conditions were required to avoid significant amounts of polymeric by-product

which occured at elevated temperatures (8).

Alternatively, Kita et al. (19) proposed a "one pot method" where the produced

maleamic acids were immediately converted to the maleimide by ring closure imidation in

an organic solvent by azeotropic removal of water in the presence of an acid catalyst. In

general, the latter method was found to give products of greater purity and in higher

yield.

Preparation of N-methyl-3,4-dimethylmaleimide. "one pot method"

Dimethylmaleic anhydride (10.07 g, 80 mmol) was dissolved in toluene (75m1) in

a two-necked round bottom flask fitted with a Dean-Stark apparatus and a pressure

equalising dropping funnel. Aqueous methylamine solution 40Vo wlw (6.48 g, 83 mmol)

was added dropwise to the rapidly stirred solution via the funnel. Stirring was continued

after addition for a further 30 min.

Phosphoric acid (1.59 g, 16 mmol) was added and the solution azeotropically

distilled until the evolution of water in the Dean-Stark apparatus ceased

(b.p.95-114.5 oC). The solution was cooled and the upper organic layer washed with

water (2 x 25 ml) and dried over sodium sulphate. Toluene (60 ml) was removed by

distillation at reduced pressure (b.p.29.2-36.5 oC at25.26 mmHg) to yield 8.39 g (72Vo)

of clear liquid product. TLC: Rp (CHzClù 0.39. lH NMR: ô(60 MHz, CDCI¡) 2.93 (s,

3H, NCH3):1.92 (s, 3H, CH3C=C). IR (NaCl): 2860 (w, NCH:); 1775 (s, Q=Q); 1715

(s br, C=O); 1685 (m, C=C) cm-1.

Preparation of N-propyl-3r4-dimethylmaleamic Acid.

Dimethylmaleic anhydride (2.01 g, 16 mmol) was dissloved in anhydrous ether

(80 ml) and the solution was cooled to 2 oC using an ice bath. Propanamine (1.0 g,

17 mmol) in ether (16 ml) was added via a pressure equalising dropping funnel over 1.5

hrs. While stirring constantly the solution was slowly raised to room temperature over

Appendix AI A-10

2.5 hrs. The resultant clay-like product was collected by Buchner filtration, washed with

cold ether and vacuum dried to yield 2.12 g (727o) of white solid. TLC: Rp (CHzClz)

0.59. lH NMR: ô(60 MHz, CDCI3) 6.67 (s, lH, COzH); 6.43 (br s, 1H, NH); 2.89 (m,

2H, NCH2); 1.89 (s, 6H, CH3-C=C); 1.60 (m, 2H, CHz); 0.90 (t, 3H, CH3). IR

(NaCl): 3650-2300 (br, COzH); 3080 and 3250 (br m, -CONH-); 1770 (w, C=O

monomer); 1715 (m, -C=C-C=O); 1630 (br m, -CONH- Amide I); 1550 (br m, -CONH-

Amide II) cr¡-l.

Preparation of N-propyl-3,4-dimethylmaleimide.

N-propyl-3,4-dimethylmaleamic acid (0.70 g, 3.8 mmol) was added to a 50 ml

two necked round bottomed flask fitted with a condensor and a thermometer. Anhydrous

sodium acetate (0.19 g, 2.3 mmol) and acetic anhydride (12 mI) were added and the

solution heated at 100 oC for t hr. The solution was cooled to 60 oC before addition to

water (12 mt). This mixture was stirred continually for a further 4.5 hrs. The cooled

mixture was extracted with dichloromethane (1 x 20 ml, 2 x l0 ml) and the lower organic

layer concentrated and f,rltered to yield 0.26 g (477o) of a pale yellow liquid. TLC: Rp

(cHzclù 0.51. lH NMR: ô(60 MHz, CDCI3) 3.56 (t,2H, NCHz); 1.95 (s,6H, CH¡-

C=C); 1.52 (m,2}l,CHz);0.92(t,3H, CH¡). IR (NaCl):1775 and 1715 (s, -C=C-

C=O); 1660 (br s, 30Amide -N-C=O) cm-l.

Preparation of N-propyl-3,4-dimethylmaleimide. "one pot method"

Dimethylmaleic anhydride (2.10 g, L6.6 mmol) was dissolved in toluene (15 ml)

and a solution of propylamine (2.00 g, 33.8 mmol) in toluene (5 rnl) was added dropwise

via a pressure equalising dropping funnel over 30 min. Phosphoric acid (0.39 g, 5 mmol)

was added and the solution was refluxed under azeotropic distillation until the evolution

of water ceased after approximately 2 hrs. The cooled solution was washed with \vater

(2 x 10 ml) and dried over magnesium sulfate. Toluene was removed by distillation

under reduced pressure (b.p. = 26-30 oC at24-27 mm Hg), to yreld2.l5 g (76Vo) of an

Appendix AI A-11

orange liquid. TLC: Rn (CHzClù 0.52. lH NMR: ô(60 MHz, CDCI3) 3.38 (t, 2H,

NCHz); 1.83 (s,6H, CH3-C=C); 1.45 (m,2}l,CHz);0.85 (t,3H, CH3). IR (NaCl):

1775 (w) and 1715 (s, -C=C-C=O) cm-I.

The main disadvantage of the synthetic route outlined in Scheme 41. was the

number of steps required. It is also possible to obtain five membered cyclic polymers by

polymerization of diallyamine monomers (20,21). Although six-membered rings might be

expected on thermodynamic grounds, as initially suggested by Butler et al. (22,23), 13C

nmr studies clearly show that five-membered rings predominate for kinetic reasons

(24,25). As discussed in Chapter 1, the possibility of five-membered rings as nitrate

selective resins has not been fully explored.

1

2

J

4

5

6

7

8

9

Appendix AI

I l.

t2.

t3.

14.

15.

t6.

t7.

A-12

Literature Cited.

M. B. Jackson and B. A. Bolto, React. Polym., 12,277, (1990).

B. J. Steel, A. S. Kayaalp, T. Kurucsev, A. D. Ward, and M. B. Jackson, Aust.

J. Chem., 43, 1983, (1990).

M. E. Baumann and H. Bosshard, Helv. Chím. Acta., 6I,275I, (1978).

M. E. Baumann, H. Bossha¡d,W. Breitenstein, G. Rihs and T. Winkler, Helv.

Chím. Acta., 67, 1897, (1984).

N. E. Searle, U. S. Pat. 2,444,536 (6 July 1948).

E. Bellack and J. B. DeWitt, J. Agric. Food. Chem.,2(23), L176, (1954).

Y. Liwschitz,Y. Edlitz-Pfefferrnann and Y. Lapidoth, J. Amer. Chem.|oc.,76,

3069, (1956).

L. E. Coleman, J. F. Bork and H. Dunn, J. Org. Chem.,24, 135, (1959).

N. B.Mehta, A. P. Philips, F.F. Lui and R. E. Brooks, J. Org. Chem.,25,

LOLZ, (1960).

Y. Kanaoka, T. Sekine, M. Machida, Y. Sôma, K. Tanizawa and Y. Ban, Chem.

Pharm. Bull., l2(2), 127, (1964).

Y. Kanaoka, M. Machida, Y. Ban and T. Sekine, Chem. Pharm. BuIl.,15(11),

t738, (1967).

R. Istratoiu. M. Farcasiu and CL. Nicolau, Rev. Roum. Chim.,12,1429,

(re67).

J. R. Heitz, C. D. Anderson and B. M. Anderson, Arch. Bíochem. Bíophys.,

127, 627, (1968).

M. K. Hargreaves, J. G. Pritchard and H. R. Dave, Chem. Rev.,70(4),439,

(1e70).

T. Miyadera, E. M. Kosower and N. S. Kosower, 'I. Med. Chem., 15(5), 534,

(1e70).

T. Miyadera and E. M. Kosower, J. Med. Chem.,15(5), 534, (1972).

B. Rubin, O. Kirino and J. E. Casida, I. Agric. Food. Chem.,33,489, (1985).

10.

Appendix AI A-13

18. K. Takatori, T. Hasegawa, S. Nakano, J. Kitamura and N. Kato, Microbíol.

I mmunol., 29(12), 1237, (1985).

19. Y. Kita, K. Sakamoto, M. Baba and A. Okubo, Eur. Pat. 0,165,574 (14 June

198s).

20. Y. Chang and C. L. McCormick, Polymer,35 (16), 3503, (1994).

21. B. A. Bolto, M. B. Jackson and R. V. Siudak, Isr. J. Chem.,26, 17 , (1985).

22. G. B. Butler änd R.J. Angelo, J. Am. Chem.ioc.,79,3128, (1957).

23. G. B. Butler, Acc. Chem. Res., 15,370, (1982).

24. S. R. Johns, R. I. Willing, S. Middleton and A. K. Ong, "I. Macromol. Sci.,

Chem. A., I0,875, (1976).

25. J. Lancaster, L. Vaccei and H. J. Panzer, J. Polym. Sci., Polym. Lett. Edn., 14,

549, (1976).

A-t4

Appendtx AZ

Accuracy of the Spectrophotometer.

There had been some concern with the accuracy and reproducibility of the Cary

2200, which had; at the time of use; been recently installed. It was decided that the

accuacy of the machine should be verified by using an ultraviolet standard.

Potassium dichromate in a weakly acidic medium is one of the more common

standards used for the evaluation of UV spectrophotometers (1,2) and was also cited as

the manufacturers reference material and is probably most suitable. Any error arising

from either small temperature variations or small discrepancies in cell positioning between

runs was considered to be negligible (3).

Potassium dichromate was available from Univar as an analytical grade reagent

and was used without further purificication after drying at 110oC for several days.

Solutions were prepared by weighted dilutions from a 2.1 glkg stock solution in 0.001M

perchloric acid. The results appearing in Table A2. agree closely with values previously

determined (4) and thus confirm the reliability of the machine. A more recent study of

potassium dichromate absorptivities in sulphuric acid (2) has been included for

comparison as these give slightly better agreement. The absorptivity of K2Cr2O7

solutions in perchloric and sulphuric acid differ only from 0.0lVo to 0.06Vo and so

comparison should be possible without any serious error.

Appendix A2

Table A2. Molar Absorptivities of 0.0670 g kg- 1 Potassium Dichromate Solutions.

ì. (nm) e (kg g-1"--1)u e (kg g-l "--1)b Vo dtff e (kg g-1"--1)" 7o diff

A-15

350

346

345

323

322

313

257

235

10.763

10.696

10.663

6.1 15

s.893

4.885

t4.603

t2.536

r0.692

10.629

10.599

6.058

5.837

4.8t314.390

12.357

0.65

0.63

0.60

0.94

0.96

1.50

1.48

1.45

r0.704

t0.643

r0.624

6.t645.938

4.837

14.456

12.445

0.55

0.5

o.37

-0.79

-0.76

0.99

r.02

0.73

a-b-

This work.

Interpolated values of potassium dichromate in 0.001 M perchloric acid from

¡eference 4.

Interpolated values of potassium dichromate in sulphuric acid from reference2.

Literature Cited.

G. Rohle, H. Schlebusch, R. Kruse and W.J. Geilenkeuser, J. Clin. Chem.

Clín. B iochem., 27 (5), 323, (1989).

M. Gil, D. S. Escolar, N. Iza and J. L. Montero, Appl. Spectrosc.,40(8),

(1e86).

S. Caroli and N. Violante, Spectosc. Lett., L2(9),67I, (1979).

R. W. Bourke and R. Mavrodineanu, -/. Res. Nat. Bur. Stand., Sect. 4.,

80A(4), 63t, (r976).

c-

I

2

3

4

A-16

Appendix A3

Enthalpy and Entropy of Tetraalkylammonium Halides.

The enthalpies of solution of the tetraalkyammonium halides were reported in the

literature, being determined calorimetrically. In comparison, there is little information

available on the corresponding entropies of solution.

Where literature values were unavailable, the entropy of solution was determined

using,

ASosol -AHosol + 2RT ln Sy+

T

where AH was the average of the available enthalpies of solution in Table 43. and S was

the solubility in mol kg-l at25oC. R was the gas constant taken to be 8.31441 J K-l

mol-1 and y+ was the activity coefficient.

When required literature values were converted to SI units using 1 cal = 4.184 J.

The standa¡d state was taken to be the hypothetical ideal solution of unit molality.

The solubility data of the tetraalkylammonium salts are scanty, and although

Lindenbaum (15) provides a convenient source of solubility data, his values were only

intended to be approximate and where possible more reliable solubilities were used. His

activity coefficient data is without question the most substantial available, but suffers

since it does not usually extend to high enough concentrations requiring slight

extrapolations. The quoted values in Table A3. corresponded to his obtained activity

coeffr.cient at the highest measurable concentration and was therefore, not necessarily the

value used in entropy calculations.

Tetramethylammonium chloride:- Four values of enthalpy of solution have been reported

(1,2,3,4) giving an average of AHosel = 4.4! 0.2 kJ mol-1. The activity coefficient and

estimate of the solubility was available from Lindenbaum (15) giving ASosol = 71.6 J K-l

mol-1.

Appendix A3 A-17

Table 43. Enthalpy and Enropy of Solution of Teüaalkylammonium Halides

Compound ÂHsol

(J mol-l)

Ave. ÂH5s1

(J mol-l)

Activity^{x

Solubility

S (mol kg-l¡

ASsol

(J mol-l)

Me¿N-CI

Me¿N-Br

Me¿N-I

Et¿N-Cl

EqN-Br

Et¿N-I

Pr¿N-Cl

Pr4N-Br

Pr¿N-I

BuaN-Cl

Bu4N-Br

Bu¿N-I

PeaN-ClPe4N-BrPe¿N-I

HexaN-I

4,079 (t)4,435 (2)4,519 (3)4,602 (4)24,267 (t)24,493 (5)24,769 (6)24,602 (3)24,686 (4)42,070 (t)42,049 (7)41,798 (8)42,t75 (2)42,384 (3)42,258 (4)

-t2,845 (6)-t2,970 (t)-t2,636 (9)6,234 (6)5,774 (9)27,907 (6)28,200 (8)27,907 (t)28,517 (9)28,702 (3)28,033 (4)

-22,175 (t0)

-4,31O-4,250-4,602

11,54811,569

(10)(11)(3)(10)(8)

-30,543 (r2)

-9,205 (t2)-8,661 (5)

-8,580 (13)15,899 (r2)

-38,200 (10)3,222 (t0)t7 400 (14)

38,300 (14)

4,409+ 230

24,563+ 194

42,122+ 201

-L2,817+ 168

6,004+ 32528,22r+ 343

-22,175

-4,397+ 188

11,559+14

-30,543

-8,815+ 339

15,899

-38,2003,222r7,400

38,300

1.s96 (15)

0.349 (1s)0.4620 (7)

0.s40 (r5)

1.98s (ls)

0.619 (1s)

0.199 (1s)

3.74s (1s)

0.439 (1s)

0.332 (rs)

0.688 (15)

0.0e1 (15)

I

1

1

1

1

19.060 (ls)

s.s8s (1s)6.3s3 (7)

0.2314o.2714

0.2624(0.2740 (

(1s)(7)

16a)16b)

9.47r (rs)8.5076 (16c)

12.6s (1s)r4.74 (r7)r.927 (rs)

1.7554 (16c)1.9068 (16b)1.8r7 (18)

18.66 (1s)19.00 (1e)8.964 (1s)10.73 (19)

0.50750.626

0.s9630.61915.6120.5326.392r.48

(15)(1e)(16c)(1 8)(15)(1e)(15)(1e)

0.06s8 (14)

40.28 (1e)0.232 (19)1.8e-3 (14)

4.5e-4 (r4)

1r.6

100.391.s6 (7)

rto.2110.08 (7)

5.6

s5.3

78.0

-3.6

11.1

rt.4

-57.9

- 16.5

8.1

-66.7- 13.5-46.7

0.3

Appendix A3 A-18

Tetramethylammonium Bromide:- The enthalpy of solution of the bromide salt has been

reported in five articles (I,3,4,5,6) giving an average of AHosol = 24.6 + 0.1 kJ mol-I.

The activity coefficient and solubility were available from two sources (7,15) and differ

somewhat. The data of Levine (7) are believed to be more accurate and were used to

obtain ASosol = 100.3 J K-l mol-l, fairly close to her previously reported value of 91.6 J

K-l mol-l zurd the value of 98.1 J K-l mol-l calculated using Lindenbaum's data.

Tetramethylammonium Iodíde:- Six values of the enthalpy of solution (1,2,3,4,'7 ,8) were

averaged to give AHosol = 42.1 + 0.2 kJ mol-I. Th¡ee values of solubility are in close

agreemenr (7,I6a,I6b) giving S = 0.268 + 0.006 mol kg-l. The approximate solubility of

Lindenbaum (15) was somewhat lower and was discarded. His limited activity coefficient

data were linearly fit to obtair y+ = 0.576 corresponding to the average solubility. The

resultant entropy of solution ÂSosol = 110.2 J K-1 mol-l is in good agteement with the

two previously reported values (7,18).

Tetraethylammoníum chloríde:- Three reported enthalpies of solution (1,6,9) give an

average of AHos61 = -12.8 + 0.1 kJ mol-l. The activity coefficients of Lindenbaum (15)

were f,rt to a quadratic. Two available solubilities (15,16c) do not agree. The older value

of Peddle and Turner correspondS to y+ = I.872 and aSossl = 3'0 J K-l mol-l' V/hile the

approximate solubility of Lindenbaumcorresponds to y+ =2.275 andÂSoss¡= 8.1 JK-l

mol-l. Lacking further solubility data it is difficult to prefer one derived entropy over

another, giving an average, ASosol = 5.6 * 3.6 J K-l mol-1.

Tetraethylammonium Bromide:- Two reported enthalpies of solution (6,9) were averaged

to give AHosol = 6.0 * 0.3 kJ moll . The activity coefficient data above 5 molal obtained

by Lindenbaum (15) were fit linearly to obtain estimates of the activity coefficients. The

two reported solubilities are quite different. Using the value of Giacomelli (17) S = 14.74

corresponds to y+ = 0.576 and yields ASosol = 55.7 J K-l mol-l compared to ASorol =

54.9 J K-1 mol-1 using Lindenbaum's solubility (15) S = 12.65 and y+ = 0.638. These

two values average to give ASosol = 55.3 t 0.5 J K-1 mol-1.

Appendix A3 A-19

Tetraethylammonium iodide:- Six values of the enthalpy of solution have been reported

(1,3,4,6,8,9) giving an average of AHosot=28.2 + 0.3 kJ mol-l. Four values of

solubility give an average of S = 1.85 + 0.07 mol kg-l and the activity coefficient of

Lindenbaum (15) was used to obtain ASosol = 78.0 J K-l mol-1 , in close agreement with

the value of Johnson (18).

Tetrapropylammonium chloride:- Krishnan (10) has reported the enthalpy of solution,

ÀHosol = -22.2 kJ mol-l. The activity coefficient was taken from Lindenbaum (15) and

two solubilities agree closely (15,19) giving S = 18.8 + 0.2 mol kg-l , yielding ASos6¡ =

-3.6 J K-1 mol-1.

Tetrapropylammonium bromide:-Three enthalpies of solution have been reported

(3,10,11) averaging to AHo.ot= -4.4 + 0.1 kJ mol-l The activity coefficient was taken

from Lindenbaum (15). The activity coefficients have also been reported by Pepala (20)

and were in good agreement with Lindenbaum but did not extend to high enough

concentrations. Although two values of solubility have been reported (15,19), Wen (21)

has studied the density up to 10 molal; indicating that the value of Lindenbaum must be a

poor estimate of the solubility. The solubility of Nakayama (19) was therefore used

yielding ASosol = 11.1 J K-l mol-I.

Tetrapropylammonium iodide¡ The two reported enthalpies of solution (8,10) agree

closely giving AHosol = 11.56 + 0.1 kJ mol-I. Of the four reported solubilities the value

of Lindenbaum (15) again seems to be somewhat low and was discarded. The remaining

solubilities ',vere averaged to give S = 0.61 + 0.01 mol kg-l. The corresponding activity

coefficient f + = 0.317 was extrapolated from the data of Lindenbaum (15) using a

quadratic f,rt, giving an entropy of solution ASosol = ll.4 J K-l mol-l.

Tetraburylammonium chloride:- The enthalpy of solution was available from Fuchs (12).

The two reported solubilities (14,15) differ substantially. The higher concentrations and

Appendix A3 A-20

the corresponding activity coefficients of Lindenbaum (15) were fit to a quadratic to

obtain exrrapolated activity coefficients. For the solubility of Lindenbaum (15) S = 15.61,

y+ = 0.715 giving ÀSosol = -62.3 J K-l mol-1 and for the solubility of Nakayama (14)

S = 20.53, \+ = 0.927 giving ASosot = -53.5 J K-l moll. There is no reason to reject

either of these values giving an average entropy of solution ASosol = -57.9 + 6 J K-1

mol-l.

Tetraburylammonium bromide:- Three values of enthalpy of solution have been reported

(5,12,13) giving an average of AHo5s1 = -8.8 + 0.3 kJ mol-I. The activity coefficients of

Lindenbaum (15) show little change at high concentrations and the lowest reported value

y+ = 0.091 was used here. Like the previous sample two values of solubility have been

reported which differ substantially and again there is no apparent reason to reject either

value. Using the solubility of Lindenbaum (15) gives ASosol = -14.6 J K-1 mol-l

compared to ASossl = -18.4 J K-l mol-l using Nakayama's (19) reported solubility

giving an average entropy of solution ASosol = -16.5 t2.61K-l mol-l'

Tetrabutylammonium iodíde:- The enthalpy of solution was reported by Fuchs (12) and

was combined with the solubility of Nakayama (14) to yield ASosot = 8.1 J K-l mol-l.

There was no activity coefficient data available in the literature and unity was assumed.

Tetrapenrylammonium chloride:- No activity coefficient data were available in the

literature. The enthalpy of solution of Krishnan (10) was combined with the solubility of

Nakayama (19) to yield the entropy of solution ASosol = -66.7 J K-l mol-l assuming unit

activity.

Tetrapenrylammonium bromide:- The enthalpy of solution from Krishnan (10) was

combined with the solubility of Nakayama (19) to give ASosol = -13.5 J K-l mol-1

assuming unit activity.

Appendix A3 A-21

Tetrapenryammoniuim iodide:- The enthalpy of solution was determined non

calorimetrically by Nakayama (14). Two solubilities have been reported. The older value

of Peddle and Turner (16c) was discarded as it was far larger than the value of Nakayama

(14) which was in fair agreement with a value determined by Franks (22) at 26 oC. Unit

activity was assumed; giving ASosol = -46.7 J K-1 moll.

Tetrahexylammonium lodide¡ This is the only hexyl salt for which an enthalpy of

solution has been reported. The enthalpy of solution was determined non calorimetrically

by Nakayama (14). It was combined with his solubility (14) assuming unit activity to

obtain ASosol = 0.3 J K-l mol-l.

Appendix A3 A-22

I

Literature Cited.

R. C. Weast, (Ed.), Handbook of Chemistry and Physics 50th edn., D-76

(Chemical Rubber Company: Ohio 1970).

C. V. Krishnan and H. L. Freidman, J. Phys. Chem.,74(II),2356, (1970).

R. H. Boyd, J. Chem. Phys.,51(4), 1470, (1969).

M. F. C. Ladd, Z. Phys. Chem. (Frankfurt),72,91, (1970).

M. J. Mastroianni and C. M. Criss, J. Chem. Thermodyn', 4,32I, (1970).

Y. V/u and H. L. Friedman, J. Phys. Chem.,70(6), 2020, (1966).

B. J. Levien, Aust. J. Chem., 18, 1161, (1965).

Om. N. Bhatnager and C. M. Criss, J. Phys. Chem.,73(l), 114, (L969).

E. M. Arnett and D. R. McKelvey, J. Amer. Chem. Soc., 88(1I),2598, (1966)'

C. V. Krishnan and H. L. Friedman, J. Phys. Chem.,73(Il),3934, (1969).

C. de Visser and G. Somsen, J. Phys. Chem.,78(17), 1719, (1974)-

R. Fuchs, J. L. Bear, and R. F. Rodewa\d, J. Amer. Chem. Soc.,9I(21), 5797 '

(1e6e).

C. de Visser and G. Somsen, J. Chem. Thermodyn',5,147, (1973).

H. Nakayama, H. Kuwata, N. Yamamoto, Y. Akagi, and H. Matsui, -Bøl/.

Chem. Soc. Jpn., 62,985, (1989).

S. Lindenbaum and G. E. Boyd, J. Phys. Chem.,68(4),911, (1964).

W. F. Linke and A. Seidell, Solubilities of Inorganic and Metal Organic

Compounds 4th edn., 682-704 (Van Nostrand Company Inc.: New York

1953); a Hill (1917), b Walden (1906), c Peddle and Tumer (1913).

A. Giacomelli and R. Menicagli, Ann. Chim. (Rome),59(10), 860, (1969).

D. A. Johnson and J. F. Martin, J. C. S. DaltonTrans.,15, 1585, (1973).

H. Nakayama, BulI. Chem. Soc. Jpn., 54,37I7, (1981).

C. N. Pepela and P. Dunlop, J. Chem. Thermodyn.,4, ll5, (1972)-

W-Y. Wen and S. Saito, J. Phys. Chem.,68(9), 2639, (1964)-

F. Franks and D. L. Clarke, J. Phys. Chem.,71'(4),1155, (1'967).

1

J.

4.

5.

6.

7.

8.

9.

10.

11.

t2.

r3

t4

15

16

t7.

18.

19.

20.

2t.

22.

A-23

Appendix A4

Calibration of Pycnometers.

Initially the mass of the tare minus the mass of the empty pycnometer was

determined. Then the pycnometers were filled with Milli-Q water and degassed. The mass

difference of the full pycnometer minus the tare, was then recorded as a function of the

height of water from a predetermined scratch ma¡k.

Measurements of height were accomplished using a cathetometer after

equilibration in a 25 oC water bath for at least I/2 hour. The pycnometers were then

removed from the bath, dried, polished and placed near the balance to equilibrate for at

least l/2 hour before weighing.

If we assume a uniform cross-sectional capillary area the radius of each

pycnometer; in cm; can be determined using;

1l- ah *

".p-

where ðm/âh is the slope of a plot in g cm-l and pw is the density of water taken to be

0.991048 g cm-3. The plots also provided the intercept; equal to M1,

Ml = mass pycnometer full of water to mark - mass of tare

From this the mass of water to the mark was determined using:

mass to mark = Ml + (mass tare - mass empty pycnometer)

and the corresponding volume to the mark was determined by division of the water

density.

Appendix A4 A-24

Vma¡k =mass to mark

Pw

All the relevant constants have been tabulated in Table 44. for the four

pycnometers considered and the appropriate calibration plots appear in Fig. A4.I. - A4.4.

Table A4.

Pycnometer 1 Pycnometer 3 Pycnometer 4 Pycnometer 53

slope (g mm-l)

slope error

intercept (M1)

intercept error

radius (cm)

mass tare - mass empty

mass water to mark

Volume to mark (cm3)

0.001168

0.00001

1.8624

0.0004

0.0611

29.1025

30.9649

31.0566

0.001307

0.00004

-0.9856

0.0005

0.0646

31.0036

30.0178

30.1067

0.001172

0.00004

0.6812

0.001

0.0612

3r.3154

31.9960

32.0913

0.001210

0.00002

0.2063

0.0005

0.0622

31.1575

31.3638

31,.4567

Cleaning.

The pycnometers were cleaned by soaking overnight with chromic acid and

rinsing four times with Milli-Q water. The pycnometers were then rinsed with AR

Methanol and oven dried before use.

Sample Measurements.

The height and mass of each pycnometer containing solution was recorded in a

similiar manner to that used for the calibration. Although only one measurement of height

and corresponding mass for each solution was required, in practice the density of solution

was calculated from the average of two or more heights and corresponding masses. The

values were corected to obtain M2, the mass of solution to the mark minus the mass of

tare, and the density was obtained using;

1.88

1.87

!'aà 1.86

1.85

1.84

-o.97

-0.98

-0.99

-1.00

-20

-10

-10 10

Height from mark (mm)

Fig.44.1. Calibraúon plot of Pycnometer 1.

-5 5

Height from Mark (mm)

0 20

10

Fà

0

Fig. A4.2. Calibration plot of Pycnometer 3.

FÀ

0.7r

0.70

0.69

0.68

0.67

0.66

o.23

0.22

FI-à 0.2r

0.20

0.19

-13 -9

-15 -11

-5 1 3 7 11

5 9


Fig.44.3. Calibration plot of Pycnometer 4.

-7 -3 1


I

+

Fig.44.4. Calibration plot of Pycnometer 53

Appendix A4 A-27

lrsol - \, . ^' v marK( 1

Pa

Pm) * pwater

M2-M1

in conjunction with an iteration method. The values pa and Pm a¡e the densities of air and

the balance weights, taken tobe 1.2 x 10-3 and 8.4 g cm-3, respectively.

A-28

Appendix A5

Experimental Data and Fitting Parameters.

Tabulations of extensive data were summanzed in this appendix to avoid any

distraction when reading the main body of the text. All fitting parameters used and

discussed in the main body are also tabulated here. Selected plots of some data have been

included together with a line representing the calculated "best fit" values to give an

indication of the typical goodness of fit. Fitting parameters were estimated by non linea¡

least-squares analysis of data using DATAFIT (1). All calculations were performed on a

Sun 4/280 computer. Some of the nitrate data fitted here were kindly supplied by Philippe

Guarilloff prior to submission of his thesis.

Section I - Density Data.

All the densities discussed in Chapter 2 are tabulated here (Tables 45.1 - 45.10).

The "best fit" parameters for each individual data set are gathered in Table 45.11. Density

measurements were the average of at least four determinations at each concentration. The

method of determination was described in Appendix 44.

Although rhere was very linle difference between fits of density to the Root

equation or to the corresponding Redlich-Meyer equation, the Root equation was used for

the discussions in Chapter 2. It was believed that the concentrations were not sufficiently

high to w¿uïant the use of the derived Redlich-Meyer equation. Additionally, the Root

equation had previously been used to investigate these salts and were the only literature

values available for comparison. Therefore, although the Redlich-Meyer equation was

more recent, the Root equation was simpler and equally well explained the concentration

dependence of density and was adopted here. The extended form of the Root equation

was found to give large errors in the coefficient, by

Appendix A5 A-29

Although htting of the concentration dependence of the apparent molar volumes to

the Masson relation also gave reasonable estimates of Quo , and are included here, they

were again not used for any of the calculations in Chapter 2. At low concentrations when

the measured solution density approached that of the pure solvent the error in the apparent

mola¡ volume increased sharply, such that a 6 ppm error in density could result in a

3 cm3 mol-l difference in the apparent molar volume (2). For this reason fits which

extrapolated on the basis of density, rather than apparent molar volumes, were preferred.

Section II - Viscosity Data.

The relative viscosity of all aqueous solutions studied in this work are gathered in

this section. The B coefficients reported in Chapter 4 were obtained from fits of viscosity

to the Jones and Dole equation for samples of low concentration (strictly < 0.1 M).

However, for the corection of relaxation times reported in Chapter 5 a higher

concentration range was required and in this case the full data were fit to the extended

Jones and Dole equation. All frtting parameters appeü in Table A5.22.In all cases the A

coefficient was considered to be a constant determined from conductance measurements

(3) using the Falkenhagen equation.

Of all the salts considered only benzyltributylammonium chloride was taken to

significantly high concentrations. As can be seen in Fig. 45.9. the Jones and Dole

equation begins to fail for such high concentrations and the extended Jones and Dole

equation (inset in Fig. 45.9.) becomes a more appropriate method of frtting.

Section III - NMR Data.

The concentration dependence of the spin-lattice relaxation time was interpreted as

arising almost exclusively to simple ion pair formation in solution. Estimates of

associations constants were obtained by fitting the concentration dependence of viscosity

corrected relaxation times to the slow exchange model. The free and bound relaxation

Appendix A5 A-30

times of the nitrate anion, T¡ and T6, were taken to be 130 and 9 ms, respectively. A full

description of the methodology is provided in Chapter 5.

Appendix A5 A-31

Section I . Density Data 25 oL

Table 45. l. Benzyltrimethylammonium chloride

Conc. (M) Density (e cm-l¡ 0u (cm3 mol-l)

Expt. Calc. a Calc. b Expt. Calc. c

0.00787s0.008336O.0089220.O140470.02647 t0.05390¡0.O824910.1982450.2887360.57230t

0.9971830.997r770.997198o.9972840.9974830.997927o.9984341.000s07L.OO2L79r.007702

0.9971740.997r82o.997r910.9972750.9974800.9979430.9984351.0005011.002180t.007702

0.997t830.997t910.997201o.9972890.997s010.9979710.9984631.000500r.002t521.007709

169.067170.736169.384169.403169.767t69.894169.398168.750168.427167.579

169.846169.838169.829t69.753169.6L4t69.39s169.222t68.725168.436167.746

a-b-c-

P=Po+aC-6ç3/2p=po+aC-6ç3/2+cC2

0u=quo+Su*€

Table A5.2. Benzyltriethylammonium chloride

Conc. (M) Densitv (e cm-3) 0u (cm3 mol-l)

Expt. Calc.a Calc.b Expt. Calc.c

0.00421t0.0045920.00899s0.0186250.O3240¿,0.0488020.0573700.10125t0.1521230.22273t0.2758720.32684e

0.9971810.9971570.9972r20.9973680.9976200.9979200.9980700.9987980.9997591.001067t.0020571.003008

0.9971210.997r270.9972030.9973700.9976120.9979010.9980530.9988390.9997631.001062r.0020521.003009

0.997r230.997t300.9972070.9973770.9976200.9979090.9980600.9988400.9997561.001052r.0020471.003019

196.844204.6892r0.t7821r.227210.753210.5362r0.5872r7.tzl210.580210.357210.2442r0.166

2t0.798210.7562t0.7122r0.6702r0.65t2t0.5722r0.4992r0.4162t0.3612r0.3r4

P=Po+aC-6ç312p=po+aC-6ç3/2+cC2

0u=Ouo+Sv*{e

a-b-c-

clI

Eobo

o

1.008

1.006

1.004

1.002

1.000

0.998

0.9960 0.10 0.20 0.30 0.40 0.50 0.60

C(M)

Fig. 45.1. Density of aqueous solutions of benzyltrimethylammonium chloride at

25 oC. The full line corresponds to the best ht of the data using pammeters obtained

from a non-linear least-squares f,rt of the data to the Root equation.

c.lIH

(,)

bo

o-

1.004

1.003

r.002

1.001

1.000

0.999

0.998

0.997Qr 0.10 0.20

C(M)

0.30 0.40

Fig. 45.2. Density of aqueous solutions of benzyltriethylammonium chloride at25 oC

The full line corresponds to the best fit of the data using pammeters obtained from a

non-linear least-squares fit of the data to the Root equation.

Appendix A5 A-y

Table 45.3. Benzyltripropylammonium chloride

Conc. (M) Densiw (e cm-3) Qu (cm3 mol-l)

Expt. Calc. a Calc. b f*Pt. Cut".'+0.00584e0.01042s0.01952e0.O339320.04563q0.0611920.09171s0.0984670.1164020.1450530.21723e0.3993330.686102

0.997t240.9971740.9972620.9973820.9974980.9976050.9978620.9979720.9982320.9984640.99930s1.00160s1.006134

0.9970920.997r280.9972030.9973280.9974340.9975830.9978890.9979590.9981500.9984650.9993111.0017131.006093

0.9971O3o.997146o.9972310.997367o.9974790.9976290.99793r0.9979990.9981830.9984830.9992841.0015981.006137

257.609258.523259.649260.766260.7s026t.50926r.73826t.228260.437260.848260.2r9259.194257.357

260.374260.334260.275260.206260.t60260.108260.02r260.004259.962259.900259.7672s9.509259.203

a-b-c-

P=Po+aC-bC3/2p=po+aC-bc3l2+cCZ

0u=ouo+su*{c

Table 45.4. Benzyltributylammonium chloride

Conc. (M) Densitv (e cm-3) Quo (cm3 mol-l)


0.0055530.01295t0.02075s0.0396100.0624460.10104r0.15898¿0.25394t0.38111+0.5431500.67131ø

0.9971150.997rr30.9971,690.9973150.9974100.9976740.9981060.9989111.0001311.001802r.003248

0.9970750.9971r40.997r570.9972690.9974r50.9976830.9981230.9989221.0001151.0018021.003250

0.9970850.997r340.9971850.9973r00.9974640.9977320.9981570.9989161.00006s1.0017601.003291

300.756307.833307.015306.105307.050306.650306.189305.506304.750304.085303.601

306.061305.978305.913305.796305.688305.54430s.372305.148304.907304.653304.478

a-b-c-

P=Po+aC-6ç312p=po+aC-bc3l2+cC2

0u=quo+Su*{e

Cî.E

(-)

äo

o-

1.008

1.006

1.004

t.o02

1.000

0.998

0.996Qi 0.20 0.40 0.60 0.80

C(M)

Fig. 45.3. Density of aqueous solutions of benzyltripropylammonium chloride at

25 oC. The full line corresponds to the best fit of the data using parameters obtained

from a non-linear least-squares fit of the data to the Root equation.

CîI

E()äo

o-

1.004

1.002

1.000

0.998

0.9960 0.20 0.40 0.60 0.80

C(M)

Fig. 45.4. Density of aqueous solutions of benzyltributylammonium chloride at

25 oC. The full line corresponds to the best frt of the data using parameters obtained

from a non-linear least-squares fit of the data to the Root equation.

Appendix A5 A-37

Table 45.5. Benzyltripentylammonium chloride

Conc. (M) Density (e cm-3) quo (cm3 mol-l)

Expt. Calc.a Calc.b E*pt. Calc.c

0.0055800.01904¿0.0308800.0392720.0604850.08877s0.1035820.1 174860.1890000.2859450.416372

0.9970500.997058o.9970780.9970740.9970880.9971360.9971810.99726r0.9973840.9976890.998036

0.9970520.9970650.9970780.9970890.9971210.9971690.9971970.9972240.9973850.9976450.9980s9

0.9970560.9970760.9970930.9971050.9971360.9971810.9972060.99723t0.99737 50.9976240.998066

354.689354.52r354.074354.384354.385354.054353.760353.230353.26s352.800352.668

354.7t9354.472354.325354.238354.052353.8s0353.7s6353.67s353.3r7352.928352.497

a-b-c-

P=Po+aC-6ç312p=po+aC-6ç3/2+cC20u=Qvo+Sv*1re

Table 45.6. Benzyltrimethylammonium nitrate

Conc. (M) Density (e cm-l¡ Quo (cm3 mol-l¡


0.00745e0.00974s0.0213800.04834r0.07652+0.0904300.14266e0.2237400.43378t0.702502

0.9972220.9973430.9976900.9984890.9993150.9997221.0012351.0035661.0095021.016854

0.9972730.9973410.9976890.9984890.9993180.9997251.0012421.0035701.0094841.016861

0.9972720.9973400.9976870.99848s0.9993t40.99972r1.001240r.0035721.0094931.016857

t89.482182.526t82.76r182.981183.166r83.223t83.444183.660184.083184.601

t82.639182.766r82.963183.1 151 83.1 80183.386183.641I 84.1 38184.618

P=Po+aC-bc3l2p=po+aC-bC3/2+cC2

0u=0.ro+Sv*1re

a-b-c-

o

o

o

o

0.9982

0.9980

0.9978

0.9976

0.9974

0.9972

0.9970

cnÉ()èo

è

0 0.10 0.20 0.30 0.40 0.50

c(M)

Fig. 45.5. Density of aqueous solutions of benzyltripentylammonium chloride at

25 oC. The full line corresponds to the best f,rt of the data using parameters obtained

from a non-linear least-squares f,rt of the data to the Root equation.

Appendix A5 A-39

Table A5.7. Benzyltriethylammonium nitrate

Conc. (M) Density (e cm-:¡ Quo (cm3 mol-l)


0.0085920.0138450.02529e0.03641e0.0583990.0950600.12983ø0.21526+0.4292320.560652

0.99728r0.9974200.9977320.9980300.9986310.9996221.0005451.0028801.008168r.0t1240

0.9972930.9974410.9977610.9980690.9986710.9996601.0005831.0028031.008131r.ort274

0.9972890.9974360.9977s30.9980600.9986620.9996551.0005831.0028181.0081521.011255

227.900228.r3s227.963228.039227.893227.927228.069227.910229.r00229.695

227.708227.766227.86t227.934228.0s1228.204228.324228.s63229.O03229.2r8

a-b-c-

P=Po+aC-6ç312p=po+aC-6ç3/2+cC20v=qvo+Sn*€

Table 45.8. Benzyltripropylammonium nitrate

Conc. (M) Densitv (e cm-3) quo (cm3 moll)


0.00745s0.0094120.0152470.0322560.06983e0.15567+0.23684s0.36199¿,0.43326g

0.9971810.9972750.9972430.9976150.9982540.99974r1.0010801.003145r.044278

0.9971800.9972t50.9973180.9976150.9982650.9997241.0010801.0030351.004288

0.9971810.9972r5o.9973180.9976160.9982670.9997271.0010821.003134r.004286

279.392279.482284.450279.647279.9s6279.927280.203280.385280.54r

280.592280.585280.569280.533280.479280.396280.337280.264280.228

a-b-c-

P=Po+aC-6ç312p=po+aC-6ç312+cC2

0u=quo+Sv*{e

Appendix A5 440

Table 45.9. Benzyltributylammonium nitrate

Conc. (M) Densitv (e cm-3) Quo (cm3 mol-l)


0.0082540.0110930.02140s0.0395750.0784010.1792460.34944s0.5632180.80s628

0.9972250.9972810.9974830.99786r0.9986151.0006761.0039731.008063r.ot2623

0.9972180.9972770.9974880.9978s8o.9986421.000651t.0039771.008069r.012620

0.9972r9o.9972770.9974890.9978s90.9986441.0006s3r.0039761.008067t.01262l

317.975318.4163 19.100318.8783r9.4363t9.r82319.606319.867320.092

3r8.576318.605318.688318.7943t8.9573t9.245319.s83319.903320.r99

a-b-c-

p=po+aC-6ç312p=po+aC-6ç3/2+cC2Qu=Quo+Su*€

Table A5. 10. Benzyltripentylammonium nitrate

Conc. (M) Density (e cm-l¡ Quo (cm3 mol-l)

E*Pt. Calc.a Calc.b Expt. Calc.c

0.O071250.0098520.01264t0.0323600.039380

0.9971690.9972r50.9972640.9975920.997713

0.9971690.9972150.9972620.9975940.997712

0.9971690.9972r50.9972620.997s940.9977t2

364.654364.694364.545364.826364.7s3

364.623364.644364.662364.760364.787

a-b-c-

P=Po+aC-6ç312p=po+aC-6ç312+cCZ

0u=Quo+Sv*1re

Appendix A5 A-41

Table 45.11. Least squares parameters and derived coefficients from molarity fits of densityto the derived density equations from either the Masson or Redlich-Meyer equations andfrom fits of apparent molar volumes to the Masson equation.

Salt 100 a 1000 b 0uo(cm3 moll)

Su*@m3 Yrl2 ro13/2¡

bv(cm3 L mol-2)

BzMe3NCl

BzEr3NCl

BzPr3NCl

BzBu3NCl

BzPe3NCl

BzMe3NNO3

BzEt3NNO3

BzPr3NNO3

BzBu3NNO3

BzPe3NNO3

1.5706L.t29r

1.6998r.7826

0.685810.94815

0.441000.67524

0.044660.16165

2.9t262.954r

2.89522.8250

2.07892.0868

1.7109r.7r37

-3.8473-4.7984

-2.t699-4.6167

-7.6352-7.7375

-5.8933-6.0683

-3.0701-4.8762

0.99947-0.39s88

4.77872.7031

t.6262-0.16707

r.3036-2.20r4

168.9170. l211.42t0.6210.9263.8

260.5308.4306. l306.2354.6353.43ss.0r82.4

182.62r82.37226.0226.7227.5279.3279.3280.7

318.63318.55

318.4364.5364.5364.5

-3.1 + 0.7-2.2 + 0.6

-1.5 + 1.9-5.9 + 0.1

-3.9 + 0.42.6 + 0.1

2.68 + 0.084.8 + 0.5

_0.6 + 2.61.6 + 0.1

2.0 + 0.41.3 + 1.1

1.4 + 0.8

-4.8 + 0.2

-4.6 + 0.6

-6.r + o.2

0.6 + 0.1

2.7 + O.3

0.0 t 0.5

-0.17 + 0.07

-2.2 + 3.6

170.05 + 0.06 -3.86 + 0.09

0.70.30.1o.70.10.11.10.20.r0.10.10.060.030.20.1

++++t++++t+t+++++t+++++++++++

1.r26

0.10.20.30.1

o.20.30.11.00.080.050.10.20.10.1

1

1

2.3.

0.60.4

-1.0 +-7.7 +

2.50.4

++

2.3 + 0.51.8 + 0.5

-7.8 t 0.1

-4.9 + 0.5

1.7908 1.8190r.7945 0.030689

Appendix A5 A42

Section II . Viscosity Data 25 oC

Table A5.12. Benzylrimethyl ammonium chlorideto = 381'96 s'

Conc. (M) r1, (Expt.) n. (Calc.) a Ir (Calc') b

0.0138530.02057s0.03651s0.0435720.06097s0.07275s0.0854650.09632s0.105766

1.005971.008831.015101.017841.02464t.029071.033571.037381.04072

1.005971.00864t.014871.01761r.024321.02884r.033711.03786

1.006511.008881.015191.017911.024521.028901.033541.037441.04078

a-b-

1'ìr=1+Ar/e+BC1'ì, = I +Ar/C+BC+DC2

Table 45. I 3. Benzyltriethylammonium chlorideto = 381'99 s'

Conc. (M) r¡. (Expt.) r1, (Calc.) a Ir (Calc.) b

0.O0347g0.0104530.0194200.03107¿0.O378220.04842s0.0596820.06966e0.07628s0.08875s0.0988020.1074470.1 15870

t.002421.006551.011531.01865t.02249r.028921.03575t.04215t.04602r.05417r.060261.065931.07160

t.002451.00680r.ot226r.oL927r.02330r.029621.03630r.04221r.046t21.053481.05941

t.00233r.00647r.0tr731.01859t.022591.028951.035601.04r951.046081.053961.06040t.06602r.07157

a-b-

1'lr=1+AIE+BCIr=1+ArE+BC+DC2

1.05

1.04

1.03

Ë r.o2

1.01

1.00

0.990 0.o2 0.04 0.06 0.08 0.10 0.r2

C(M)

Fig. 45.6. Relative viscosity of aqueous solutions of benzyltrimethylammonium

chloride at 25 oC. The full line corresponds to the best fit of the data using parameters

obtained from a non-lineil least-squares fit of the data to the Jones and Dole equation.

1.08

1.06

1.04

ÈF

t.02

1.00

0.98 0 0.o2 0.04 0.06 0.08 0.10 0.12

C(M)

Fig. 45.7. Relative viscosity of aqueous solutions of benzyltriethylammonium chloride

at 25 oC. The full line corresponds to the best ht of the data using parameters obtained

from a non-linear least-squares fit of the data to the Jones and Dole equation.

o

Appendix A5 445

Table 45.14. Benzy chloridero-

Conc. (M) r¡. (Expt.) q. (Calc.) a îr (Calc.¡ b

s.

0.0069400.0136820.0206520.0360100.0407200.0523Ot0.06345a0.07011+0.0871010.t020200.r0202s0.1094910.121917

1.007461.013481.02034t.035371.040071.05133r.062561.068831.086201.101361.101121.10804t.t2rt9

r.00728t.oL3971.020831.035841.040421.051691.062501.068951.08539

t.N7 t4r.013721.02050t.035471.040061.05142t.062421.069001.085951.10098l.100981.108551.12t23

a.-

b-î. = 1 + ArE+BCI, = I +Ar/c+BC+DC2

Table 45. I 5. Benzyltributylammonium chlorideto = 381'69 s'

Conc. (M) '4, (Expt.) r1, (Calc.) a Ir (Calc.) b

0.00219000682e01599s027O1203996+05763¿,07176209125s

20647224097234827t

0.471O5g

1.005311.010251.02130r.03552r.05123r.075651.09486t.122671.16061t.2r07r1.254041.300961.36t54r.56929r.84276

r.003231.00952r.021751.036331.053401.076621.095t41.12068

1.003011.008861.020381.034401.051 171.074661.09398r.t274t1.158701.21055t.254521.30237t.362901.569991.84190

I 1678r15059s17793t

0000000000000

a- Ir = 1+Ar/õ+BC1'ìr = 1+Ar/e+BC+DC2b-

F

T.T4

t.t2

1.10

1.08

1.06

t.o4

r.02

1.00

0.98 0.10 0.12 0.140' o.o2 0.04 0.06 0.08

C(M)

Fig. 45.8. Relative viscosity of aqueous solutions of benzyltripropylammonium

chloride at25 oC. The full line corresponds to the best fit of the data using parameters

obtained from a non-linear least-squares fit of the data to the Jones and Dole equation'

-

2.0

1.8

1.6

t.4

t.2

1.0

2.01.91.8

t.71.61.5

t.41.3

t.21.1

1.00.9

o

0 0.10 0.20 0.30 0.40 0.50

0.10 0.20 0.30 0.40 0.500

C(M)

Fig. A5.9. Relative viscosity of aqueous solutions of benzyltributylammonium chloride

at25 oc.The full line corresponds to the best ht of the data using parameters obtained

from a non-linear least-squares fit of the data to the Extended Jones and Dole equation'

The inset is the relative viscosity and best ht of the data to the Jones and Dole equation'

Appendix A5 A-48

Table 45. 1 6. Benzyltripentylammonium chlorideto = 383'73 s'

Conc. (M) r¡,. (Expt.) r1, (Calc.) a Ir (Calc.) b

0.0038020.00527¿0.00640g0.01506e0.O237470.03382s0.O4482s0.O5629t

1.008331.01156t.012321.03054t.042441.059871.077 5r1.09684

1.007031.009631.011631.027391.040861.0s939r.078431.09826

t.007641.010431.0t259t.02853t.04283r.06089r.078531.0959r

a-b-

rl. = 1+Ar/õ+BCî, = I +ArE+BC+DC2

1.10

1.08

1.06

o

0.980 0.01 0.02 0.03 0.04 0.05 0.06

c@1)

Fig. 45.10. Relative viscosity of aqueous solutions of benzyltripentylammonium

chloride at25 oC. The full line corresponds to the best fit of the data using parameters

obtained from a non-linear least-squares fit of the data to the Jones and Dole equation.

Ë 1.04

t.02

1.00

Appendix A5 A-50

Table 45. I 7. Benzyltrimethylammonium nitrate

Conc. (M) r¡, (Expt.) 11, (Calc.) a 1r (Calc-) b

to = 36.70 s.

0.0081460.0134430.01701¿0.02819r0.0358300.0641860.09075s0.1 100920.L445O7O.l445ot0.17360e0.19386r0.21316t0.24949s0.24949s0.2938950.33344t

0.00289¿0.0069250.01258sO.O234870.03849+0.06105e0.07884e0.098845

1.00361r.005471.006981.01099t.0t3761.02396r.032451.040911.05358t.052741.06438t.071821.082531.093161.095181.11310r.t2754

1.00116r.00257t.004741.008731.01411t.022051.028191.03515

r.003421.00541t.006721.01078r.01353r.02360t.03294

1.00136r.0029rr.00s001.008911.01419r.02204r.028t71.03504

1.003411.005391.006711.01080r.013571.023881.033611.04076t.05364r.05364t.06470t.072501.08000r.09434r.09434t.tt2221.12850

to = 382.51 s.

a-b-

lr=1+ArlC+BC1'1. = 1+Ar/e +BC+DC2

1.15

1.10

Þ 1.05

1.00

0.950 0.10 0.20 0.30 0.40

C(M)

Fig. 45.11. Relative viscosity of aqueous solutions of benzyltrimethylammonium

nitrate at25 oC. The full line corresponds to the best fit of the data using parameters

obtained from a non-linear least-squares fit of the data to the Jones and Dole equation.

Appendix A5 A-52

Table A5. 18. Benzyltriethylammonium nitrateto = 103'64 s'

Conc. (M) q. (Expt.) q, (Calc.) a îr (Calc.) b

0.00168s0.003250O.00649t0.01254g0.0136160.0180500.02595q0.0518750.0816000.12195+0.25913s0.3157720.46818r0.6094130.829223

0.999631.000341.001391.005961.006671.009331.01304r.027711.04678r.06320t.tr6291.13915t.20422t.26123r.35172

1.00120r.o02t41.004001.007391.00798t.0t042r.0t474t.02872r.0446r

1.001081.001901.003541.006491.007011.00912r.012861.02491,1.038521.05675t.tr728t.t4L70r.2059rr.26353r.34971

a.-

b-îr=1+ArE+BC1'ì, = 1+ArE+BC+DC2

Table 45. 19. Benzylripropyto = 382'41

lammonium nitrates.

Conc. (M) q, (Expt.) q. (Calc ) a q, (Calc ) b

0.0031120.00718q0.01397a0.02743s0.0487020.0698700.0828620.09249+0.1002370.1393 1o0.1719370.22068t0.29277s0.35875s0.40163s

r.002701.006281.0t3461.026081.04s66r.06544r.077371.08638r.093821.108421.1 3582r.r7694r.24373r.3ro371.35609

1.003281.007221.01365r.02623r.045961.065501.077471.08633r.09345

1.002881.006321.011951.023001.040431.057871.068631.076631.083081.1 1593r.143751.186001.249901.310011.34990

a,- 1'ìr=l+Ar/C+BCl, = 1+ArE+BC+DC2b-

Appendix A5 A-53

Table A5.20. Benzyltributylammonium nitrateto = 381'66 s' a

Conc. (M) r1' (Expt.) q, (Calc.) b r'ìr (Calc ) c

0.0017850.0035560.00455e0.00785¡0.03005r0.0437300.06092t0.0832130.09963s0.17 547 o0.21833s0.28988¿0.3615700.482920

.00295

.003s5

.0051s

.00803

.03667

.05436

.07720

.t0622

.t2929

.24268

.3Lll4

.43734

.57805

.86273

1.002591.004951.006251.010551.038951.056311.07807t.106221.12692

1.002351.004471.005651.009551.036011.05285r.074681.10411rJ2663r.239971.31088t.440341.583981.85896

1

1

II1

1

1

1

1

1

1

III

a-

b-c-

relative viscosities above 0.1 M were determinedusing to = 195.76 s.

1r=1+Ar/õ+BCIr=1+Ar/e+BC+DC2

Table A5.21. Benzyltripentylammonium nitrateto = 383'87 s'

Conc. (M) q. (Expt.) t1, (Calc.) a rlr (Calc.) b

0.00073r0.00238t0.00419t0.00843¿0.01264t0.03235e0.039380

r.005931.007001.010861.01807r.02559t.0562s1.06571

r.00t47r.004451.008701.015021.022281.056071.06807

1.001811.005501.010641.01802r.026041.0s689r.06524

a- rlr=1+Ar/õ+BC"1r = 1+ArE+BC+DC2b-

Appendix A5 A-5/

Table A5.22. Coefficients of mola¡ity fits of viscosity to either theJones and Dole or the extended Jones and Dole equations of aqueoussolutions of benzyltrialkylammonium salts at 25 oC.

^

Compound Afakenhagen DB

BzMe3NCl

Bz.Er3NCl

BzPr3NCl

BzBu3NCI

BzPe3NCl

BzMe3NNO3

BzEt3NNO3

BzPr3NNO3

BzBu3NNO3

BzPe3NNO3

0.00710.00710.00750.007s0.00790.00790.00820.00820.00840.00840.0074 b

0.0074 b

0.0074 c

0.00780.00780.00830.00830.00860.00860.00890.0089

0.370 + 0.0010.386 + 0.0020.630 + 0.0040.54r + 0.0020.954 + 0.0030.931 + 0.004

1.30 + 0.011.190 r 0.004t.77 + 0.021.88 + 0.05

0.331 + 0.0010.336 + 0.0040.338 + 0.0020.52 + O.O20.448 + 0.03

0.906 + 0.0020.699 r 0.006

1.25 + 0.011.11+ 0.011.68 + 0.062.t6 + O.l

-0.2r + 0.02

0.47 + 0.02

0.33 + 0.04

r.24 + 0.0r

-3.83 + 1.1

0.11 + 0.01

-0.04 + 0.01

0.42 + O.0l

1.36 + 0.03

-1.38 + 4.56

a-

b-

c-

Fits to the Jones and Dole equation (rl. = 1 + Ari C +BC)

were restricted to concentrations stictly less than 0.lM, while

fits to the extended Jones and Dole equation

(q, = 1 + Ar/C +BC + DC2) utilize the full concentration

ranges given in Tables A5.12. - 1^5.21.

run 1: to = 36.70 s.

run 2: to = 382.51 s.

Appendix A5

Section III. NMR Data 25 oC

Table A5.23. Variation of spin-lattice relaxation time with concentration for aqueoussolutions of silver nitrate at 25 oC.

A-55

C (m) p G cm-3) u t'1, u C (M) T1 corr (ms) cby+

0.04990.08610.08610.20440.20440.39450.59660.78331.0693

1.0041r.00921.00921.02s81.0258r.05241.08061.10651.1460

00300053005s01250125024s

0.7710.7410.7410.6600.6600.5680.5080.4700.410

0.04970.08570.08570.20260.20260.38910.s8530.7649r.0370

tlt.2107.2t74.4110.7108.0100.8

94.592.3

1

1

1

1

1

1

1

1

1

0373050s073s

+ 10.4+ 8.3+ 7.5+ 7.5+ 7.8+ 6.1+ 5.0+ 4.9+ 3.787.0

a.-

b-

c-

Both density and viscosity were calculated using coefficients determined byJones and Colvin (4).

Activity coefficients were calculated from a fit of tabulated molal activitiycoefficients (5) upto I molal to an appropriate polynomial.

The error refers to one standard deviation of the mean.

Table A5.24. Variation of spin-lattice relaxation time with concentration for aqueoussolutions of calcium nitrate at25 oC.

C (m) p (g cm-3) u q, b Y+c C (M) T1 corr (ms) d

0.05330.09150.19550.37500.6042o.78540.78840.78841.0183

r.00471.0099r.0209r.04t3r.06741.08811.08841.0884t.Lt46

1.01551.024t1.04351.084s1.147 t1.204t1.205t1.205t1.2872

0.5300.4850.4320.3830.354o.3440.3440.3440.339

0.05320.09850.19340.36780.58680.77700.75980.75980.9725

90.7 + 10.378.3 + 5.368.r + 4.755.6 + 2.847.r + r.343.7 + r.640.1 + t.24r.2 + r.237.6 + t.0

a-

b-

c-

Density was calculated from a linear fit of available data (6) upto I molal.

Viscosity was calculated from a fit of available data (7) to quadratic.

Activity coefficients were calculated from a fit of tabulated molal activitiycoefficients (5) upto 1 molal to an appropriate polynomial.

The error refers to one standa¡d deviation of the mean.d-

Appendix A5

Table A5.25. Variation of spin-lattice relaxation time with concentration for aqueous

solutions of barium ninate at25 oC.

A-56

C (m) p (C cm-:¡ . t1. b y+c C (M) T1 corr (ms) d

0.05380.10250.15000.20070.25060.30380.3270

t.00621.0159r.02531.03531.04511.05571.0603

1.01051.0182

0257

0.4770.4280.3860.3470.3160.29r0.282

0.05340.10140.14800.t9740.24580.2971o.3194

76.3 +66.3 +61.1 +49.9 +45.4 +43.8 +40.4 +

033s042¡05 1¿

0552

8.57.35.63.2

I1

1

1

I

2.64.01.5

a-

b-

c-

d-

Density was interpolated linearly from available data (8).

Viscosity was calculated using the coefficients of Doan and Sangster (8).

Activity coeff,rcients were calculated from a fit of tabulated molal activitiycoefficients (5) upto 1 molal to an appropriate polynomial.

The error refers to one standard deviation of the mean.

Table A5.26. Va¡iation of spin-lattice relaxation time with concentration for aqueous

solutions of lead nitrate at25 oC.

c (m) p (g cm-3) u t1, u C (M) Tl corr (ms) cbyr

o.04620.11050.19970.37260.56750.77240.9568

1.0099r.0279r.0526r.09931.1503r.2020r.2469

1.001s1.015rl.O34sL.075¿,I.l25s1.183q1.240t

0.4950.3760.3050.2380.2000.r740.156

0.04590.10960.r9720.36460.54950.73930.9059

69.0 + 1.955.7 + 1.845.7 + 1.439.3 + 1.533.3 + O.730.7 + 0.529.0 + 0.5

a-

b-

Both densitypolynomials.

and viscosity were calculated from fits of available daø (9) to

Molal activity coefficients were calculated using reported parameters (10).

The error refers to one standard deviation of the mean.c-

Appendix A5 A-57

Table A5.27. Va¡iation of spin-lattice rela¡<ation time with concentration foraqueous solutions of benzyltrialkylammonium nitrates at 25 oC.

c (M) p (e cm-r¡ u r'ì, u y+ b T1 corr. (ms) c

Benzyltrimethylammonium nitrate.

b' = 0.045

1.0506o.87t40.78110.68060.60470.46110.46110.37690.27780.27780.19600.t4950.12160.10060.0511

0.99100.99100.86s30.79120.69600.45700.3083o.235r0.16090.04830.02s4

1,.0262r.02141.0r901.01631.0t421.01021.0r021,.00791.00511.00511.00281.00141.00061.00000.9986

r.0226r.02261.01951.01761.01521.00901.00531.00331.00140.99840.9977

1.473t3752328t2790243s178¿,

.043a

.036r1.0186

1.45121.45121.397 61.36551.323sl.2l5s1.14831.11411.079s1.O2461.0133

0.7600.7550.7530.752o.752o.7540.7540.7580.7660.7660.7790.7900.7990.8070.839

0.7810.7810.7790.7780.7770.7800.7870.7950.8070.8530.879

1

I1

I1

1

1

1

1

II1

1

178¿,1427102ø102607lt053s

46.7 + 2.945.4 + 2.542.8 + I.346.7 + t.652.8 + 3.160.5 + 2.963.8 + 1.358.7 + 2.370.8 + 3.975.9 + 3.472.4 + 3.584.6 + 6.4

104.2 + 8.3101.0 + 7.2104.5 + 10.5

Benzyltriethylammonium nitraæ. d

b' = 0.03

23.6 + 0.723.2 + t.629.5 + r.625.1 + 1.829.r + 1.935.r + 2.546.5 + 2.453.4 + t.277.8 + 2.285.7 + 6.691.4 + t.2

Appendix A5 A-58

Table A5.27 . (continued).

c (M) p (s cm-l¡ u rl. u y+ b T1 corr (ms) c

Benzyltripropylammonium niuate.

b' = 0.125

0.57620.47520.39650.31780.r7930.13s00.05660.0396

1.00661.00501.00371.00251.0002o.99940.99810.9978

1.00611.00551.00551.00391.00211.00051.00000.99880.99880.99790.9976

0.9974o.9973

1.545¡1.43351.3490t.2690l.l42s1. l05z1.042g1.0300

1.7253l.66ts1.661q1.494el.34It1.21361.17731.09801.09801.044s1.027e

0.8580.8430.8330.8250.8210.8250.8500.863

0.9520.9410.9410.9140.8900.8730.8700.8680.8680.8810.894

0.8890.903

27.7 + t.428.r32.638.348.2s6.374.393.2

23.2r.2t.23.29.34.36.3 + 1.749.8 + 5.450.3 + 3.485.4 + 2.885.4 + 5.5

74.7 + 4.283.9 + 5.1

+ 2.4+ 1.8+ 1.9+ 3.1+ 3.8+ 9.8+ 6.8

Benzyltributylammonium nitrate. d

b' = 0.21

o.42850.40010.40010.32000.23800.16030.13620.07960.07960.03780.0234

6!O.45+1.23+r.4o!r.25+0.73+2.5

B enzyltripentylammonium nitrate.

b'=0

0.02090.0139

36s24+

1.01.0

a-

b-

both density and viscosity were calculated using coefficientsreported prèviously and resulted from a fit of experimental datato appropriate equations.

the calculation of activity coefficients were described Chapter 5.

the error in Tl refers to one standard deviation.

datafromreference 11.

c-

d-

Appendix A5 A-59

Table A5.28. Density, viscosity and spin-lattice relaxation times for aqueous

polymer-nitrate solutions at 25 oC.

polymer

c (M)

nitrate

C (M)

Density

(s cm-l¡

T1 corr b

(ms)

Ra 1r

Polybenzyltrimethylammonium chloride - Potassium nitrate.

0.000262

0.0004s9

0.000459

0.001004

0.001901

0.002968

o.o07126

0.000113

0.000291

0.000484

0.001036

0.001975

0.003087

0.004739

0.007674

0.008940

0.046963

0.048259

0.048259

0.050486

o.05r757

0.049724

0.046566

0.055483

0.051066

0.049651

0.049335

0.048848

0.051838

0.048100

0.050678

0.050774

180.2 c

105.6 d

105.6 d

50.5 d

273d16.8 c

6.6 c

1.006906

1.0r04t7

l.olo4t71.0r I 131

1.011097

1.006521

1.006660

1.006878

r.006022

1.010828

1.010790

1.010871

t.007t75

1.006756

1.007005

r.006293

1.033734

t.047398

r.047398

1.082539

1.135908

r.203181

r.515612

t73.4 + lr.2108.4 + 2.4

706.L + 2.2

92.7 + 3.0

77.7 + t.351.9 + 6.1

33.2+ r.9

120.3 + 9.7

to2.5 + 4.8

9L.I + r.767.9 + 2.1

51.8 + 1.2

44.3 + 2.8

34.4 + 4.8

20.5 + 3.7

20.5 + 4.0

Polybenzyltriethylammonium chloride - Potassium nirate.

494.8 c

176.s c

103.1d

47.8 d

24.8 d

16.9 c

10.2 c

6.6 c

5.6 c

L.Or54l4

1.026202

1.03777t

r.055577

r.092956

1.140665

1.236579

t.469922

1.554850

Appendix A5 A-60

Table A5.28. (continued).

polymer

c (M)

nrEate

C (M)

Density

(s cm-r¡

T1 corr b

(ms)

Ra flr

Polybenzyltripropylammonium chloride - Potassium nitrate.

0.000203

0.000510

0.000943

0.002160

0.003554

0.009790

0.047933

0.047426

0.048142

0.0481 12

0.044898

0.044489

238.0 c

93.4 d

51.3 d

223 d

r2.7 c

4.6 c

180.9 c

97.0 d

45.7 d

21.9 d

12.6 c

8.3 c

1.006930

1.010345

r.0ro473

1.010479

1.006984

r.007769

1.006796

t.010742

1.010504

r.010842

r.006624

1.007226

1.021547

t.o52t791.110954

r.243816

r.t921831.636858

r.022558

r.036928

1.057563

1.101512

1.189619

r.413169

97.7 + 8.3

69.1 + 2.4

52.9 + L.r

31.3 + 0.9

18.8 + 2.1

9.7 + 0.7

80.1

60.7

36.7

19.7

10.9

9.1

+ 6.7

+ 3.7

+ 2.O

+ 1.3

+ 1.1

+ 3.0

Polybenzyltributylammonium chloride - Potassium nitrate.

0.000252

0.000507

0.001056

o.oo2t57

0.004198

0.007017

0.045330

0.048984

0.048044

0.o47046

0.052720

0.048813

a-

b-

c-

d-

R is the mole ratio of nitrate to polymer.

the error refers to one standard deviation of the mean.

reference 11.

this work.

Appendix A5

11.

A-61

J

1

2

4

5

8

9

Literature Cited.

T. Kurucsev, J. Chem. Educ., 55, 128, (1978).

F. J. Millero, "The Partial Molal Volumes of Electrolytes in Aqueous Solutions",

in W ater and Aqueous Solutions : S tructure, T hernndy namics, and T ransport

Processes. (Ed. R. A. Horne), Chapter 13 (Wiley-Interscience: New York 1972).

B. J. Steel, A. S. Kayaalp, T. Kurucsev, D. Wa¡d and M. B. Jackson, Aust. J.

Chem., 43, 1983, (1990).

G. Jones and J. H. Colvin, J. Am. Chem.9oc.,62,338' (1940).

R. A. Robinson and R. H. Stokes, Electrolyte Solutions 3rd edn.

(Butterworths: London 1965).

L. Gmelin (Ed.), Gmelins Handbuchder Anorganíschenchemie, ca [B],356

(Verlag Chemie: Berlin 1950)

E. V/. Washburn (Ed.), International CriticalTables Vol. V. 1st edn.,

(McGraw Hill: New York 1929).

T. H. Doan and J. Sangster, J. Chem. Eng. Data, 26, l4l, (1981).

E. I. Chernen'kaya and O. M. Kuznetsova, Zh. Prikl' Khim'52(6),1255,

(1979); J. Appl. Chem. (Engl. Transl.) 1188, (1979).

R. N. Goldberg, J. Phys. Chem. Ref. Data., 8(4), 1005, (1979).

P. Guarillofl PhD thesis, The University of Adelaide, (1994).

6

7

10

1

PUBLICATIONS

The following publications have arisen from work covered in this thesis:

14N Nuclear Magnetic Resonance Relaxation of the Nitrate Ion and Ion Pairing in

Aqueous Solution, Gary Owens, Philippe Gua¡illoff, Barry J. Steel and Tomas

Kurucsev, Aust. J. Chem., 48(2),207, (1995)-

2. Nitrate Selectivity of Ion-Exchange Resins and Their Model Compounds. II

Viscosity and Density of Benzyltrialkylammonium Salts in Aqueous Solution and

14N NtrrtR Relaxation of the Nitrate Ion., Gary Owens, Philippe Guarilloff and

Tomas Kurucsev, submitted for publication in Aust. J. Chem.

Owens, G., Guarilloff, P., Steel, B.J., and Kurucsev, T., (1995) 14N nuclear magnetic

resonance relaxation of the nitrate ion and ion pairing in aqueous solution.

Australian Journal of Chemistry, v. 48 (2), pp. 207-215.

NOTE:

This publication is included in the print copy

of the thesis held in the University of Adelaide Library.

Owens, G., Guarilloff, P., and Turucsev, T., (1995) Nitrate selectivity of ion-exchange

resins and of their model compounds. II. viscosity and density of

benzyltrialkylammonium salts in aqueous solution and 14N N.M.R. relaxation of the

nitrate ion.

Australian Journal of Chemistry, v. 48 (8), pp. 1401-1411.

NOTE:

This publication is included in the print copy

of the thesis held in the University of Adelaide Library.

It is also available online to authorised users at:

http://dx.doi.org/10.1071/CH9951401

THE UNIVERSITY OF ADELAIDEBARR SMITH LIBRARY

Th shoul.l be reh,.-'ed no later than the last dar

"el.w l-ì r

ERiL:.,i¡ì. ',r; ',

Page 2,line 23-24:

Page3,line 4,6:

Page L6, line 9:

Page2S,line 4:

Page 46,line 23:

Page 47,line 14:

Page 5L, F,qn.2:

Page94,line 9:

Page 96, line 3:

Page L07,line 20:

Page l07,Ltne2l;

Page l2l,line 30:

Page l26,hne27:

Page I34,line 14:

Page l34,line 16:

Page l37,line 3:

Page l37,line 8:

Page l|l,line 28:

Figure caption 5.3

methaemoglobin

methaemoglobin

... was calculated ...

15 MÇ2 cm.

... low temperature and high pressure.

replace "¿+" and all subsequent appearances with (r +"

aMx

replace "difussion" by "diffusion"

replace "linearity" with "faithful agreement"

replace "function of both" by "linear combination of'

... relaxation times (or inverted relaxation times) of ...

All association constants, K¡, henceforth have the units dm3 mol-l.

... conductance measurements often yield association ...

delete "ou" from "benzyltrimethylammonioum"

... calculated here (Fig. 5.11.).

... comparison between samples was inferred to be possible

without serious error ...

"retards" and "results"

delete "counterion"

delete "(out of page)", the diagram is already in perspective.

Documents

NITRATE SELECTIVE RESINS · 14N nmr spin-lattice relaxation times and a method for interpreting the non-linear concentration dependence of these times, in terms of ion association,