Thermodynamic Behavior of Electrolytes in Mixed Solvents

18 Heats of Solution and Dilution of Lithium Perchlorate in Aqueous Acetonitrile

R. P. T. TOMKINS, G. M. GERHARDT, L. M. LICHTENSTEIN, and P. J.

Downloaded by CORNELL UNIV on July 18, 2016 | http://pubs.acs.org Publication Date: June 1, 1976 | doi: 10.1021/ba-1976-0155.ch018

TURNER Department of Chemistry, Rensselaer Polytechnic Institute, Troy, N. Y. 12181

Heats of solution of lithium perchlorate in 20, 40, 60, 80, 90, and 100 wt % acetronitrile-water mixtures at 298.16°K

are re

ported. The heats of dilution were measured for lithium per chlorate in the mixed solvent containing 90 wt

%

CH CN. 3

The heats of transfer (∆H ) of lithium perchlorate from water tr

to aqueous acetonitrile were calculated. The results are dis cussed in terms of the structure of the solvent system and se lective solvation properties of the lithium ion.

T

he

heats of solution a n d d i l u t i o n

of electrolytes

i n nonaqueous-

aqueous solvent mixtures have been l i m i t e d mostly to alcohol-water systems

and a few measurements i n dimethylsulfoxide-water and dioxane-water mixtures (J, 2).

T h e structural m a x i m u m i n aqueous-organic solvents at h i g h water

content has been well established b y a variety of techniques (8,4), but few systems have been explored over the w h o l e composition range. T h e purpose of this study is to examine the structural features of acetonitrile-water mixtures over the whole composition range using the heats of solution a n d d i l u t i o n of l i t h i u m perchlorate as a probe.

T h e effect of water on thermo-

d y n a m i c properties such as heats of solution is also of interest. Acetonitrile has been selected as the solvent i n this study since it is a possible candidate for a nonaqueous electrolyte battery (5).

F r o m this v i e w p o i n t , ace-

tonitrile has several attractive p h y s i c a l properties, as s h o w n i n T a b l e I.

It has

a useful l i q u i d state range a n d a reasonably l o w vapor pressure a n d viscosity at ambient temperature. acetonitrile.

I n a d d i t i o n , m a n y c o m m o n electrolytes are soluble i n

A c e t o n i t r i l e is a good m o d e l solvent for solvation studies, as the

molecule is a linear aprotic dipole. 297

Furter; Thermodynamic Behavior of Electrolytes in Mixed Solvents Advances in Chemistry; American Chemical Society: Washington, DC, 1976.

298

THERMODYNAMIC

T a b l e I.

BEHAVIOR O F ELECTROLYTES

P h y s i c a l P r o p e r t i e s o f A c e t o n i t r i l e (1)

Boiling point Freezing point Vapor pressure Viscosity Dielectric constant

81.60° C —43.84° C 88.81 Torr (at 25°C) 0.375 cP (at 15°C) 37.5 (at 2 0 ° C ) Nonaqueous Electrolytes H a n d b o o k


Experimental L i t h i u m perchlorate was p u r i f i e d as described earlier (6). Acetonitrile was Fisher A C S reagent grade a n d was used without further purification. W a t e r was distilled twice, the second t i m e i n a C o r n i n g A G - 2 a l l glass still, a n d h a d a specific conductivity at 2 5 ° C of 8 X 10~ o h m c m . T h e m i x e d solvents were prepared b y weight as shortly as possible before heat measurements were made. 7

- 1

- 1

Heats of solution were measured at 298.16° K using an L K B 8700-1 precision calorimeter. Details of the procedure have been described earlier (6). T h e p r i n c i p a l difference i n procedure i n this case was that the ampoules were to be f i l l e d w i t h solid instead of solution. T h e ampoules were f i l l e d w i t h crushed l i t h i u m perchlorate using a o n e - m m copper w i r e as a r a m r o d , a n d their contents were weighed by difference i n a nitrogen-filled drybox. T h e y were placed about four at a time i n a small desiccator a n d transferred to the sealing apparatus a n d sealed as q u i c k l y as possible. Solid sticking to the f i l l i n g stem was w e i g h e d b y difference on washing a n d d r y i n g the stem after sealing. Heats of dilution were measured at 2 9 8 . 1 6 ° K w i t h an L K B 11700 batch microcalorimeter. T h e heats of solution of l i t h i u m perchlorate i n aqueous acetonitrile were measured at concentrations between 0.01 a n d 0.1m. T h e concentration dependence was small c o m p a r e d w i t h the experimental scatter of about 0.1-0.2 kcal m o l e . A H s values are g i v e n i n T a b l e II. T h e heats of solution i n anhydrous acetonitrile were corrected to i n f i n i t e d i l u t i o n using measured heats of dilution (6), and the corrected values were averaged. T h e heats of dilution were measured for l i t h i u m perchlorate i n the m i x e d solvent c o n t a i n i n g 90% M e C N . - 1

Table II.

Heats o f S o l u t i o n o f L i t h i u m Perchlorate i n Aqueous Acetonitrile

MeCN (wt %) 0 20 40 60 80 90 100

—&H (kcal mot ) S

1

6.35 7.5 10.0 12.0 14.2 15.8 13.6

Mean deviation (kcal mot ) 1

No. of measurements

0.05 0.09 0.16 0.26 0.28 0.20 0.15


10 14 11 9 10 5 4

18.

TOMKINS E T A L . T a b l e III.

Lithium

Perchlorate

299

in Aqueous Acetonitrile

Heats of D i l u t i o n of L i t h i u m Perchlorate i n 1 0 % Aqueous Acetonitrile -AH b (cal mol~ )

m (final) a

Da

(cal

l

0.0180 0.0448 0.1124 0.3049 1.0172

0.134 0.212 0.335 0.552 1.009

0L

mol'

69 109 172 341 885

816 776.2 712.3 543.7 0

Initial concentration for each dilution is 1.0172m. b Estimated uncertainty of measurement: ±30 cal mol" .

a


1

T h e data were analyzed b y the m e t h o d of H a r n e d a n d O w e n (7).

T h e r e was

little difference between a linear a n d a square root extrapolation at the lowest measured molalities. A square root extrapolation was used.

Heats of d i l u t i o n

and 0 L values are listed i n Table III. Figure 1 shows the dependence of 0 L versus

Discussion Heats of solution for l i t h i u m perchlorate i n pure water have been reported previously (8).

T h e heats of transfer to the aqueous mixtures a n d anhydrous

acetonitrile are given i n T a b l e I V a n d F i g u r e 2. 900

T

800 700 600 •g E

500

3

400 300 200 100 0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1.0

l/

m 2(mole'£kg~£) Figure 1. perchlorate

Apparent relative partial molal enthalpies of lithium in 9:1 (w/w) acetonitrile-water as a function of m /


1

2

THERMODYNAMIC

300

BEHAVIOR O F ELECTROLYTES

Table I V . Heats o f Transfer o f L i t h i u m Perchlorate f r o m Water to


Aqueous Acetonitrile MeCN (wt %)

Mole Fraction of MeCN x 2

-AH (W-+S) (kcal mol' )

20 40 60 80 90 100

0.0989 0.2264 0.3971 0.6372 0.7982 1.000

1.2 3.7 5.7 7.9 9.5 7.26

t

1

T h e free energies of transfer, A G , to the aqueous mixtures have not been t

measured, but f r o m the s u m m a t i o n of ionic free energies (9), the free energy of transfer of L i C l 0 4 f r o m water to anhydrous acetonitrile is f o u n d to be +7.6 kcal mol . - 1

Thus the entropy of transfer (as T A S ) is —14.9 kcal m o l . - 1

T h e fact that

the solubility (on a m o l e fraction basis) of LiClC>4 is very nearly the same i n acetonitrile as i n water is the result of differences i n the free energy of the crystalline state i n e q u i l i b r i u m w i t h solution (10), i.e., L i C l 0 4 « 4 M e C N or L i C l 0 - 3 H 4

2

0

a n d of i o n association, a n d does not reflect the true situation of the ions at h i g h dilution. T h e general tendency i n aqueous organic solvents (I J) is for the free energies of transfer to be monotonic functions of solvent composition. O f t e n the enthalpy of transfer shows a m a x i m u m i n the water-rich region, a n d for salts of small ions there m a y be a further reversal i n the o r g a n i c - r i c h region. T h e heats of d i l u t i o n are nearly linear i n m / b e l o w 0 . 1 m , w i t h a slope of 1

about 500 cal k g / m o l e / . 1

2

- 3

2

2

This is very close to the theoretical and measured


18.

TOMKINS ET AL.

Lithium

Perchlorate

values for 1:1 salts i n pure water (12).

301

in Aqueous Acetonitrile

T h e temperature dependences of the

density a n d dielectric constant of the m i x e d solvent are not k n o w n , so it is i m possible to verify that this is the expected l i m i t i n g slope.

If we assume that these

missing values are close to those for pure acetonitrile, i.e., a ~ 0.0014 K d l n D / 6 T ~ 0.005 K cal k g / m o l ~ / . 1

2

3

2

_ 1

-

1

and

, then the p r e d i c t e d l i m i t i n g slope is of the order of 1700

It is possible that the low value observed is caused by the work

r e q u i r e d to separate i o n pairs on d i l u t i o n . A b o v e 0.1m the 0 L values of aqueous salts flatten out a n d m a y become negative except where small ions are involved.

In 90% acetonitrile, the 6

becomes steeper w i t h increasing concentration.

curve

L

T h i s is p r o b a b l y caused b y


depletion of the free water i n the solvent, so that i o n solvation b y acetonitrile becomes significant.

It is also possible that, because ionic electric fields are sig-

nificant at greater distances i n a solvent of lower dielectric constant, the overl a p p i n g of ionic co-spheres has a m u c h more drastic effect on the enthalpy than it does i n water for this salt. A p a r t f r o m a small exothermic region at the w a t e r - r i c h e n d , A H

E

for the

acetonitrile-water system is strongly positive at 2 9 8 ° K (J 3), i n d i c a t i n g that the water structure is significantly b r o k e n u p b y the a d d i t i o n of large amounts of acetonitrile. A H

E

passes t h r o u g h a m a x i m u m at about X

2

= 0.65, but A H

t r

ap-

parently does not change d i r e c t i o n u n t i l later than this. It seems l i k e l y that:

(a) the smaller amount of water is still c o m p e t i n g

successfully w i t h the acetonitrile for solvation sites, (b) this effect is accentuated by the c o n t i n u i n g b r e a k u p of water structure so that " n o n a q u e o u s " water is a far more effective solvent than " a q u e o u s " water, a n d (c) o r d e r i n g b y the ions is m u c h more significant i n a w e a k l y associated aprotic solvent.

W h i l e this

general type of behavior has been f o u n d before for small ions i n m i x e d solvents, these results indicate that the reversal of the effect i n a c e t o n i t r i l e - w a t e r occurs at m u c h lower water content than was observed where both components are h i g h l y structured, e.g., for H C l a n d HCIO4 i n aqueous ethylene g l y c o l

(14,15),

for K B r i n aqueous f o r m i c a c i d (16), a n d for H C l i n aqueous ethanol (17).

In

contrast, H B r and H I i n aqueous methanol (18,19) do seem to behave i n a similar manner i n this region, although H C l does not (20).

In each of these ( H B r a n d

HI) cases, it is possible that the m o v e m e n t of A H t r is controlled b y a large a n d sudden increase i n A G . t r

T h e path of A H

t r

T h i s m a y a p p l y to aqueous acetonitrile also.

above X

2

= 0.8 is especially interesting.

In order to gain

a reliable interpretation i n this region, further work is planned to investigate the effects of small quantities of water on the t h e r m o d y n a m i c properties of electrolytes.

Acknowledgment W e wish to thank R. H . W o o d of the U n i v e r s i t y of D e l a w a r e for the use of the L K B 11700 batch m i c r o c a l o r i m e t e r i n his laboratory.


302

T H E R M O D Y N A M I C BEHAVIOR O F E L E C T R O L Y T E S


Literature Cited

1. Janz, G. J., Tomkins, R. P. T., "Nonaqueous Electrolytes Handbook," Vol. I, New York, Academic, 1972; Vol. II, New York, Academic, 1973. 2. Covington, A. K., Dickinson, T., Eds. "Physical Chemistry of Organic Solvent Sys tems," London, Plenum, 1973. 3. Feakins, D., in "Physico-Chemical Processes in Mixed Aqueous Solvents," F. Franks, Ed., New York, American Elsevier, 1967. 4. Franks, F., Ives, D. J. G., Q. Rev. Chem. Soc. (1966) 20, 1. 5. Jasinski, R., "High Energy Batteries," New York, Plenum, 1967. 6. Tomkins, R. P. T., Turner, P. J., J. Chem. Eng. Data (1976) 21, 153. 7. Harned, H. S., Owen, B. B., "The Physical Chemistry of Electrolytic Solutions," New York, Reinhold, 1958, p 334. 8. Parker, V. B., "Thermal Properties of Aqueous Uni-Univalent Electrolytes," NSRDS-NBS 2, Washington, D.C., 1965. 9. Abraham, M. H., J. Chem. Soc., Faraday I (1973) 1375. 10. Tomkins, R. P. T., Turner, P. J., J. Chem. Eng. Data (1975) 20, 50. 11. Bennetto, H. P., Willmott, A. R., Q. Rev. Chem. Soc. (1971) 25, 501. 12. Young, T. F., Seligmann, P., J. Am. Chem. Soc. (1938) 60, 2379. 13. Morcom, K. W., Smith, R. W., J. Chem. Thermo. (1969) 1, 503. 14. Stern, J. H., Nobilione, J. M., J. Phys. Chem. (1968) 72, 3937. 15. Stern, J. H., Nobilione, J. M., J. Phys. Chem. (1969) 73, 928. 16. Ivanova, E. F., Fesenko, V. N., Zh. Fiz. Khim (1969) 43, 1006. 17. Butler, J. A. V., Robertson, C. M., Proc. R. Soc. (1929) 125 A, 694. 18. Schwabe, K., Ferse, E., Ber. Bunsenges. Phys. Chem. (1966) 70, 849. 19. McIntyre, J. M., Ph.D. Thesis, University of Arkansas, 1968. 20. Slansky, C. M., J. Am. Chem. Soc. (1940) 62, 2430. RECEIVED September 4, 1975.


Thermodynamic Behavior of Electrolytes in Mixed Solvents

Recommend Documents