Hydration of tri-n-hexyl amine hydrochloride, tri-n-octyl amine hydrochloride, tri-n-decyl amine hydrochloride and tri-n-dodecyl amine hydrochloride benzene solutions

] inorg, nucl. (?hem.. I976, VoL 38, pp. 104%1048.

Pergamon Press.

Printed in Great Britain

HYDRATION OF TRI-n-HEXYL AMINE; HYDROCHLORIDE, TRI-n-OCTYL AMINE HYDROCHLORIDE, TRI-n-DECYL AMINE HYDROCHLORIDE AND TRI-n-DODECYL AMINE HYDROCHLORIDE BENZENE SOLUTIONS C. KLOFUTAR. S. PAlJK and M. OSTANEK University of Ljubljana, Institute ",1. Stefan", Llubljana, Yugoslavia

I Re¢eired 19 March 19751 Abstraet--lsopiestic equilibrations of THxA.HCI, 7OA-HCI. TDA.HCI and TLA.HCI benzene solutions have been made in the concentration range from 0.02 to 0.13 mole/kg at different water activities and 25°C. The mean hydration number of the chloride ion at unit water activity was found to be I.(]2,-1+0.015 for the systems investigated. To determine the effect of water activity on the amine salt activity the results were treated on the basis of the Gibbs-Duhem equation and a simple linear relation obtained between salt and water activities. IR spectrophotometric measurements were performed to identify the nature of the hydrated species.

INTRODUCTION TERTIARY alkyl amine salts are important extractants for simple and complexed ions. The organic phase processes in solvent extraction are far from simple because of the competition of aggregation, hydration and solvation reactions. Therefore for these s y s t e m s the a s s u m p t i o n of ideal behaviour is not an acceptable approximation[l]. In order to visualize the effects of hydration on the changes of salt activity in ternary s y s t e m s and to elucidate the effects of the size and shape of the cation on the hydration of the anion in the ion pair or higher aggregates, a hydration study of some higher tertiary alkyl amine hydrochlorides in benzene solution in the concentration range where these salts are associated [2-4] to form higher aggregates has been carried out and is presented in this paper. Owing to the hydrolytic decomposition of amine salts during the partition method an isopiestic procedure was used.

Table I. The melting points of the amine hydroehlorides investigated :om,-ou:/d

m.~.

m.p.

(°C] 7H~/,.]](71

53

i Dr, . H C I

vr?

~:)t.t!Cl

">

1:J~'.!ic'_

-

-

', iJ t , ~ r a t u r e } (°C)

51~ r ]

71

""

]

"

"r,~

"!

[' [

~.

5,y i:]

cb

-

lsopiesti{ method. The glass equilibrator, shown in Fig. 1, was constructed according to the idea suggested by Christian et al. [ 11]. For isopiestic measurements 10ml of water or sulphuric acid solutions of known water activity were introduced into the outer part of the equilibrator and 10 ml of pure benzene or benzene solutions of tertiary alkyl amine hydrochlorides into the inner part. Equilibration was allowed to take place for 2 days with the entire apparatus immersed in a water bath held at a constant temperature of 25,00°+ 0.02°C, Preliminary experiments showed

EXPERIMENTAL

Tri-n-hexyl amine hydmehloride (THxA.HCI), tri-n-octyl amine hydrochloride (TOA.HCI), tri-n-deeyl amine hydrochloride ITDA.HCI) and tri-n-dodecyl amine hydrochloride (TLA.HCI). From the commercially available amines tri-n -hexyl amine (pract., Fluka. A.G.), tri-tl-octyl arraine (pract., Fluka. A,G.). tri-n-decyl amine (prac';., Eastman Kodak) and tri-n-dodecyl amine (purum, Rhbne-Poulenc) the amine hydrochlorides were prepared by the method proposed by Kertes[5]. The salts were dried itl cact~o. Their uncorrected melting points are given in Table I. Benzene, Riedel de H~ien, was purified by the method given in Ref.[6] and stored over 'a 4 ~, molecular sieve in well-closed container. The density value of benzene, d> 0.8734+_4xl0 ~g/ml. ~a~ close to the literature value of 0.87370[7]. Density measurements. The density, d (g/mb, was measured at 25.00°=0.02°C with a digital densimeter of type DMA 10 following the manual instructions[g]. In no case was a significant density difference between the dry and corresponding wet systems detected. To conxert molarities to molalities, therefore. the density values of dry benzene and benzene solutions[9] were used. Analysis of water. The stoichiometric concentrations of water in the systems investigated were determined with an automatic Karl Fischer titrator, Aquatest l ltl0], with an accuracy of + 111/,gwater per ml of sample. JIN(' VOL. 38 NO. ~ H

B ~0

i

~.

k

¢6o

Fig. I. Glass equilibrator for isopiestic measurements (the dimensions are in mm).

1045

C. KLOFUTARel al.

1046

a~ = I (m~° = 0-0385 ---0.0018 mole/kg) and r7 = (Om~/Om,),, is the average number of moles of bonded water per mole of salt at the water activity in question. Table 3 gives the values of slope and intercept of eqn (2) obtained by linear regression analysis from the data listed in Table 2. The calculated mfla~ values are, within experimental error, close to the solubility of water in pure benzene at the water activities in question (see Table 2). Since h is nearly constant at each water activity it may be concluded that the chain length does not affect the r~ value. In Fig. 3 the average value of (Om,~/Om~),,, ap, as a function of water activity is presented. As can be seen, rip = haw, where h is the mean hydration number at aw = 1 and amounts to 1.024-+-0.015 for the amine salts investigated.

that within this period the organic phase came to equilibrium. The same isopiestic time was used by Christian [11] and Frolov [12]. By acidimetric titration it was established that the concentrations of amine salt and sulphuric acid changed less than 0.1% during equilibration. The accuracy of the isopiestic procedure was checked by the determination of water solubility in benzene at various water activities. The value obtained for the solubility of water in benzene, at 25°C and water activity a~ = 1, of 0.0348±6x10-~mole/l, is close to the values of 0.0347±2x 10 "[13], 0.0349111] and 0.0363114]. The solubility data for water in benzene at various water activities fit the relation c, = (0.0343 ±0.0011)a ....

(1)

A value of 0.0345 for the coefficient of a~ was found by Masterton [13]. The water activities of sulphuric acid solutions in the concentration range 0-10mole/kg were obtained from the literature [15]. IR spectrophotometric measurements. The IR spectra of the water-saturated benzene phase were recorded on a Karl Zeiss UR-20 double beam IR spectrophotometer with lrtran 1I. cells in the range from 3000 to 4000 cm-L

i

I

I / /

f

12oo

/ •

/

RESULTS AND DISCUSSION

The primary isopiestic data of THxA.HC1, TOA-HCI, TDA.HC1 and TLA.HCI benzene solutions, listed in Table 2, indicate a linear dependence of water content, m~, on solute concentration, m, at each water activity, in spite of the fact that with the increasing salt concentration increasing aggregation takes place [2-4]. A typical plot of rn~ vs m~ for THxA.HCl--benzene solutions at different water activities is given in Fig. 2, where the lines were obtained by linear regression analysis. The water content can therefore be represented by

j

o

o/ /'/°/ /

1 500

I

300

I 9 O0

1 12O0

a

m~ dO2[mate/kg]

m~, = m,,.°a,, + ~m~

(2)

Fig. 2. m~ vs m, for THxA.HCI benzene solution at various a~ and 25°C (a~ = 0.7740 (11); aw = 0.5906 ([3]); a~ = 0.3938 (O) and a~ = 0.2155 (@)).

where mw° is the solubility of water in pure benzene at

Table 2. The values of the isopiestic equilibrium water concentration, mw (mole/kg), the standard deviation, s, and the number of estimates, n, at amine hydrochloride concentrations, m, (mole/kg) in benzene, at 25°C aw ms

O.0OO0

= 0.7740

mw

aw = 0.5906

s. 10 ~ n

0.0298

8

mw

8

s. 104 n

0.O21, I

7

9

a w = 0.3938 mW

s. 104 n

O.0160

~

8

aw = 0,2155 mw

s. 104 n

0.0085

I

4

THxA. HC1

0.0230

0.0512

I

40.

0386

2

4

0.0260

2

4

0.0135

I

4

0,0464

0.0690

2

4

0.0533

4

4

0.0364

4

k

0.0198

I

4

0,0701

0.0885

5

4

0.0667

2

4

0,O459

3

~

0.0253

3

4

0.0940

O. 1097

[,

[, 0.0823

4

[, 0,O566

3

4

0.0308

2

4

0.118/4

0.1316

4

4

0.0984

2

[, 0.0662

4

[, 0.0358

2

0.0231

0.0[,97

7

8 0,0371

4

7

0.0258

3

8 0.0160

[,

0.0[,66

0.0708

6

8 0.05[,9

7

8 0.0371

8

8 0.0210

4

8

0.0705

0.0897

~

7

5

7

6

7

5

7

TOA. HCI

0.0691

0.0[,71

0.0269

7

0.0949

O. 1080

7

7

0.0816

5

8

0.0558

5

8

0.0323

8

7

O. 1197

O. 1302

7

7

O. 1005

7

7

0.0673

7

6

0.0379

3

7

TDA.HC1

0.0231

0.0493

6

8

0.0380

3

8

0.0265

3

8

0.0156

4

8

0.0468

0.0677

5

7

0.0525

7

7

0.0368

[,

8

0.0212

6

7

0.0710

0.0871

5

6

0.0663

6

6

0.O465

6

7

0.0264

6

8

0,0958

O. 10[,5

5

6

0.0792

4

6

0.05[,I

6

6

0.0317

8

6

O. 1211

O. 1250

3

6

O.O965

[,

6

0.0672

9

6

0.0398

6

6

O.O232

O.O472

4

7

O.O372

3

8

0.024O

4

8

0.0136

4

7

0,0470

0.0690

5

6

0.0527

3

7

0-0345

3

6

0.0206

2

4

0o071)4

0.0873

5

6

0.0661

4

6

0.0471

5

6

0.0259

5

6

0,0965

O. I068

5

8

0.0830

3

8

O,O583

3

8

0.0329

6

8

O, 1223

O. 1281

6

8

O.O964

?

8

0.0662

6

8

0.0381

[,

8

TDA. HCl

Hydration of trim-hexyl amine h~drochloride I

/

7-

1047

from which, for a, = 0, it follows that In a*la, °= In y*lTJ >= h

(61

O

where a* and "y* are the activity and the activity coefficient of the salt in binary solution. Therefore, the activity or activity coefficient of the investigated amine salts in benzene solutions at any water activity can be evaluated from the relation In a, = In a* - ha..

Fig. 3. fi,, vs a, for amine salts in benzene solution at 25°C. Table 3. The values of the slope, & and intercept, m~°a~, of eqn (2) at 25°C ~

= "!.v740

*~

= C.5906

a

= '),.303~

0.2155

Tibo% HCI

mlja ;

c
"). 6~

0,42

!.:~3

).0~ 5

0.32:140

?.OlrJ ,

3.0h8(

TOA.HC 1

g

0.(32

O.6L

0.42

).23

m~a w

0.:)}14

3.32!35

0.0!67

3.013L

r]

O. 7 v

3.59

0.40

0,2h

;:[~a

0.0318

0.02k 5

0.0171 .

3,0~,9~

0.~1

O.6F)

o. Lk

0.25

0. 0206

O. 0238

0.0146

O. 0'28~

~bA.ffCl

fLA.ECI

m°a,,

Roddy and Coleman[16] found a value of 0.973+0.004 for the mean hydration number of TOA.HCl-benzene solutions up to 0.1 mole/l and at 0.Smole/1 a value of 1.22, while Frolov [17] at 0.1 mole/1 found a value of 0.70. A comparison of the values of the mean hydration number for the chloride ion in water (h =3)[18], nitrobenzene (h = 3.3)[19], dichloroethane (h ~ 2.5)[19], chloroform (h = 1)[19] and benzene (h ~ 1) shows that the mean hydration number depends on the dielectric constant of the solvent as well as its ability to solvate the chloride ion. Some thermodynamic information about the systems investigated can be obtained regarding these systems as lernary ones. The following cross differentiation relation [20] holds

7)

For example for 0.1 mole/kg of tri-n-dodecyl amine hydrochloride in benzene at 25°C we have: In a, - -- 3.146 (a, =0); -2,.658 (a, =0.5) and -4.170 (a .... 1). The respective y* value was taken from Ref.[4]. IR spectra of the amine salt-benzene solutions investigated, containing various amounts of water, show characteristic bands in the OH stretching region ascribed to the monohydrate of the halide ionI211. The narrow band located at 3685 _+5 cm ' results from the overlap of two bands: the first of them is due to the z,~ antisymmetric stretching mode of free water (u~ 3678 cm ~[22]), while the second belongs to the stretching vibration of free OH of bonded water. The broad band at 3 4 2 5 + 5 c m c characterises the stretching vibration of the OH group of the water molecule bonded to the halide ion, while the band at 3200 _+5 cm ' corresponds to the overtone of the bending vibration. From the constant and concentration independenl spectral shift, '3u = 2 6 0 c m . for all the amine salts investigated, it may be concluded that the donor-acceptor properties of the chloride ion in ion pairs or higher aggregates are similar and not affected by the size and shape of the cation. The IR spectra do not show any evidence of interactions between the oxygen atom of the water molecule and the alkyl ammonium ion [211. The difference of ~35 cm ' in the position of the broad band in benzene and carbon tetrachloride123] can be ascribed to the effect o1' solvent on the hydration of the chloride ion. In Fig. 4 the: IR spectra of TOA.HCI benzene solutions at two water activities are presented. An attempt to represent the experimental data on the basis of a simple hydration model, which takes into account the formation of the monohydrate of the chloride ion, was unsuccessful, owing to the impossibility of determining the concentrations and activity coefficients for each of the species present.

0-77

v-,,,

80 (O In a,/O In a~) ..... = -(On JOin) ......

-~

(3)

where indexes s, w and b are related to the amine salt, water and benzene, respectively. Since On./On~=am,,/Om,, eqn (3) upon integration within limits In a .... 0 and In a,,. yields

E E

lna

I n a / a , ° = l n y j % °= -

L

(Om,/Om,) .........Olna,

(4)

where a, ° and %" are the activity and the activity coefficient of the salt at unit water activity. Putting 6, = ha,, we obtain In aJaJ' -In %/T.~° = - h(a,, - 1)

(5)

2O I

I

3800

I

I

I

I

3600 3~00 V [ cm-I]

I

~__1

3200

Fig. 4. IR spectra of TOA.HCI benzene solutions equilibrated at two water activities.

1048

C. KLOFUTARet al.

Acknowledgements--Grateful acknowledgement is made to the Slovene Research Community for financial support of this project. The authors also wish to thank Mrs. J. Burger for her skilful technical assistance. 1. 2. 3. 4. 5. 6. 7.

8. 9. 10.

REFERENCES Y. Marcus, J. Phys. Chem. 77, 516 (1973). A. S. Kertes and G. Markovits, J. Phys. Chem. 72, 4202 (1968). M. Aguilar and E. H6gfeldt, Chem. Scripta 3, 107 (1973). M. ~umer, Disertation, University of Ljubljana, Yugoslavia (1972). A. S. Kertes, J. lnorg. Nucl. Chem. 27, 209 (1965). J. H. Richards and S. Walker, Trans. Faraday Soc. 57, 399 (1961). J. A, Riddick and W. B. Bunger, Organic solvents, physical properties and methods of purification. In Techniques of Chemistry (Edited by A. Weisberger), Vol. II, Wiley, New York (1970). A. Paar K. G., DMA 10 Information, Digital Densimeter for Liquids and Gases. S. Paljk, C. Klofutar and M. 7,umer, J. lnorg. Nucl. Chem., in press. Operating Manual, Aquatest 11, Photovolt Corporation 1115 Broadway, New York 10010.

I I. S. D. Christian, H. E. Affsprung and J. R. Johnson, J. Chem. Soc. 1896 (1963). 12. J. G. Frolov and V. V. Sergievsky, Radiochim. 13, (5), 760 (I971). 13. W. L. Masterton and M. C. Gendrano, J. Phys. Chem. 70, 2895 (1966). 14. J. W. Roddy and C. F. Coleman, Talanta 15, 1281 (1968). 15. R. A. Robinson and R. H. Stokes, Electrolyte Solutions, p. 477. Butterworths, London (1970). 16. J. W. Roddy and C. F. Coleman, J. Inorg. Nucl. Chem. 31, 3599 (!969). 17. J. G. Frolov, V. V. Sergievsky, A. V. Ochkin and A. P. Zujev, Radiochim. 14(4), 578 0972). 18. W. C. McCabe and H. F. Fisher, J. Phys. Chem. 74, 2990 (1970). 19. T. Kenjo and R. M. Diamond, J. lnorg. Nucl. Chem. 36, 183 (1974). 20. H. A. McKay, Nature 169, 464 (1952). 21. S. C. Mohr, W. D. Wilk and G. M. Barrow, J. Am. Chem. Soc. 87, 3048 (1965). 22. E. Greinacher, W. Ltittke and R, Mecke, Z. Elektrochem. 59, 23 (1955). 23. J. F. Desreux, Anal. Chim. Acta 52, 207 (1970).