HARRIET SNELL AND JEROME GREYSON
2148
Water Structure in Solutions of the Sodium Salts of Some Aliphatic Acids192
by Harriet SnelI and Jerome Greyson Rocketdyne, A Division of North American Rockwell Corporation, Canoga Park, California (Received December SO,1969)
Entropies of transfer between heavy and normal water for the sodium salts of the group of aliphatic acids from the formate to the caproate have been determined from combination of heats of transfer obtained from calorimetric measurements with free energies of transfer obtained from cell measurements. The standard heat of transfer for the process salt in heavy water to normal water was negative for sodium formate, near zero for sodium propionate and butyrate, and positive for sodium caproate. Solvent-structureinfluence, established from the sign and magnitude of the transfer entropies, indicated a transition from structure breaking to structure making in passing from sodium formate to sodium caproate. Increases in solute concentration lead to decreases in the absolute value of the entropy of transfer. The structure influence and the concentration dependence of the entropies of transfer of the salts have been explained in terms of the Gurney cosphere model.
Introduction A continuing program has been under way in this laboratory in which the nature of the structure of water and the influence of electrolytes thereon have been investigated. The approach has been thermodynamic in that i t has been directed toward determination of the entropy effects associated with the transfer of ionic salts between heavy and normal water. Justification for relating transfer entropies to water structure has been presented in detail elsewhere as have the details of the experimental technique^.^-^ Suffice it t o say here that the properties of heavy and normal water are similar except to the extent to which they are structured, and the sign and magnitude of the transfer entropy can be related to structure influence. Systems that have been studied and reported in earlier publications have included solutions of alkali halide^^,^ and alkaline earth chlorides.616 I n this paper, we report the results of measurements of solutions of the sodium salts of several low carbon number aliphatic acids.
Experimental Section Materials. Heavy water (99.75% DzO) was obtained from the U. S. Atomic Energy Commission, Savannah River Operations Office, Aiken, S. C. All normal water solution measurements were made with COz-free conductivity water (3.6 X ohm-' cm-l at 25"). For the heat of dilution measurements and for the cell measurements, all starting solutions were prepared by weight in duplicate. All starting chemicals, except sodium caproate and sodium valerate, were reagent grade and were used without further treatment. Sodium formate, sodium acetate, sodium propionate, sodium butyrate, and sodium isobutyrate were oven dried to constant weight at 105". Anhydrous sodium caproate was prepared from 99% n-caproic acid and a 51.2% aqueous solution of sodium T h e Journal of Physical Chemistry
hydroxide in a fashion recommended by VoldS8 The preparation was carried out in a nitrogen-filled glove bag. Caproic acid was dissolved in an equal volume of 70 :30 ethanol-water, and a stoichiometric quantity of 51.2% aqueous sodium hydroxide was added. The solution was then neutralized with 0.05% aqueous sodium hydroxide to the phenolphthalein end point with the aid of filter paper, which had been treated with the indicator and dried. The latter step served to repress hydrolysis and substantially reduce the free acid content of the product. The solvent mas evaporated on a hot plate. Final drying was carried out in a vacuum oven at 155". Anhydrous sodium valerate was prepared similarly from material that chromatographic analysis showed to be 95.1% n-valeric acid. Calorimeter il! easwements. Integral heat of dilution and/or integral heat of solution measurements in normal and heavy water were made for each of the salts. Experimental procedures and the methods for calculating the heats of transfer from the resulting data have been described previously for both the integral heat of solution5 and dilution heate measurements. For the dilution measurements, the starting solutions were 2.5 aquamolal, i.e., moles of salt/55.5 moles of solvent, except for the sodium caproate solution, which was 2.0 aquamolal. A single dilution heat was measured for each process and for each solution by diluting it approximately 100-fold. Dilution volumes were selected such (1) This research was supported by the Research Division of the Office of Saline Water, U. S. Department of the Interior, under Contract No. 14-01-0001-1701. (2) Presented in part a t the 158th National ,Meeting of the American Chemical Society, New York, N. Y., Sept 1969. (3) J. Greyson, J . Phys. Chem., 66, 2218 (1962). (4) J. Greyson, ibid., 71, 2210 (1967). (5) J. Greyson and H. Snell, ibid., 73, 3208 (1969). (6) J. Greyson and H. Snell, {bid., 73, 4423 (1969). (7) R. E. Kerwin, Ph.D. Thesis, University of Pittsburgh, 1964. (8) M. J. Vold, private communication, Department of Chemistry, University of Southern California.
SODIUM SALTSOF ALIPHATIC ACIDS
2149
Table I : Integral Heats of Solution for Sodium Formate, Sodium Acetate, Sodium Propionate, and Sodium Butyrate in Normal and Heavy Water a t 25.0" Final concn, aquamolality
Temp,
Solvent
OC
Measd AH," kcal/mol
Temp,
Solvent
DaO
24.99 25.00 24.99 25.00 25.00
0.1189 0.1217 0.1308 0,1336 0.1440
25.00 25.00 25.00 25.00 25.00
0.1314 0,1410 0.1446 0.1449 0.1486
0.266 0.267 0.268 0.271 0.272 AR = 0.269 f 0.003 0.465 0.467 0.470 0.458 0.459 A n = 0.464 i 0.005
Ha0
25.01 25.01 25.01 25.01 25.01
5
0.01410 0.05610 0.05668 0.07693
25.01 25.00 25.01 25.01
0.02414 0.04729 0.05689 0.07530
0.01240 0.01246 0.01619 0.02706 0.03600
AR
-3.044 -3.060 -2.990 -3.002 -2.979 = -3.015 f 0 016 -3.022 -3.044 -2.976 -2.950 -3.007 = -2.999 f 0.037 I
DaO
25.01 25.01 25-01 25.00 25.01
0.01267 0.01740 0.01939 0.02836 0.03540
An
Sodium Butyrate
Sodium Acetate 25.01 25.00 25.01 25.00
Meaad A H , Q kcal/mol
Sodium Propionate
Sodium Formate
Ha0
OC
Final concn, aquamolality
AR
An
-4.012 -3.979 -3.967 -3.933 = -3.973 f 0.038 -3.915 -3.903 -3.918 -3.885 = -3.905 & 0.015
Ha0
DaO
25.00 25.01 25.01 25.01
0.01942 0.03235 0.03855 0,05334
25.01 25.01 25.01 25.01
0.02358 0.03430 0.05052 0.05719
-3.393 -3.398 -3.397 -3.353 = 3.385 f 0.022 -3.394 -3.429 -3.386 -3.397 A n = -3.402 f 0.019
An
The mean values and the root-mean-square deviations are also listed.
that the final concentration was 0.02 aquamolal for all salts. As reported before,6 the value 31 cal/mol was used to correct the heats measured for the heat effect contributed by the isotopic exchange reaction in the integral heat of dilution experiments. Emf-Transfer Free Energy Measurements. The experimental apparatus and general procedure for obtaining the free energies of transfer of salts between heavy ,and normal water have been described elsewhere. 8 , 4 Potential measurements were made with the same equipment as that described in a previous publ i ~ a t i o n . ~For this work, however, a Leeds and Northrup K-4 potentiometer was used. Measurements were carried out at 25.0 f 0.2'. The emf measurements were made in cells of the type AgjAgCIINaCl ( a d ]+INaR ( a J -1 HzO "20 NaR (a&I +INaCl (4[ AgC1IAg D20 KO where the configuration
AglAgCIINaCl
(&)I +I
H2O is a cation-reversible membrane electrode reversible to sodium ion. The symbol R refers to an aliphatic acid permits anion, and the anion-exchange membrane, the diffusion of anions between the solvents. It has been shown that the emf's of such cells are related to the free energy of the transfer of 1 equiv of salt from DtO a t activity ae to H2O a t activity a,.a,4
I -1,
Results and Discussion The calorimetric data are presented in Tables I and 11. Table I summarizes the integral heat of solution data, obtained for sodium formate, acetate, propionate, and butyrate. Table I1 includes the integral heats of dilution and the heats of transfer calculated from the measurements. The mean values of the heats and the root-mean square deviations are also given in the tables. As noted p r e v i ~ u s l ythe , ~ standard heat of transfer for salts between heavy and normal water is defined as the difference between the values of the heats of solution in the isotopic solvents at inVolume 74,Number 10 M a y 14,1070
2150
HARRIET SNELLAND JEROME GREYSON
Table 11: Integral Heats of Dilution and Heats of Transfer for Sodium Salts of Some Aliphatic Acids in Normal and Heavy Water a t 25.0" Y A H d i l ,
Salt
YConcn, aquamolalityInitial Final
AHdi1=2~,
cal/mol
Sodium formate
2.500
0,02500
0
Sodium acetate
2.500
0,02400
- 387 -408 ...
...
AHdiiDz0,
Sodium propionate Sodium butyrate Sodium isobutyrate Sodium valerate
...
...
...
2.509
0.02260
2.500
0.02260
Sodium caproate
2.000
0.01728
-1093 - 1104 -1711 - 1693 -2087 -2058
cal/mol
HzO into D20
+83 +85 -368
Cal/mOl-----. DzO into H2O
- 121 - 107
+209 +214 -316 -322
- 348 ...
-408 -412
...
... - 1098 - 1094 - 1757 - 1732
AHtP,b
oal/mol
-121 f 9 -26 i 18
...
...
...
...
...
-1134 -1150 - 1747 - 1741 -2153 -2176
-2186 -2172
AHtpP
cal/mol
-205 f 9 - 195 f 60 -65 f 18 -68 4 42C -16 f 40"
- 1064 - 1071 - 1679 - 1678
$39 zk 11
+17 f 2gC +36 f 11
+12 4 19
+54 f 17
-2070 - 2075
-7 f 17
+99 4 17
a Concentrations equal initial concentrations. Standard enthalpies of transfer have been assumed equal to the transfer values at the final concentrations tabulated. 0 These data were calculated from integral heat of solution measurements.
Table 111: Heat, Emf, Free Energy, and Entropy of Transfer from DzO to HzO for the Sodium Salts of Some Aliphatic Acids at 25.0' 7 -
AHtr',*
Salt
Sodium formate Sodium acetate Sodium propionate Sodium butyrate Sodium isobutyrate Sodium valerate Sodium caproate
cal/mol
-205 -65 -16 +17 +36 $54 +99
f 9 f 18
f 40" f 2ga f 11 f 17 f 17
--Dilute Emf, mV
-0.8 -1.5 -1.9 -2.0
aonl-AFtr',
-0.75 -0.33 -0.20 -0.097
-121 f 9 -26 f 18 $39 rt 11 +12 =k 19 -7 f 17
...
...
+39 +39
f0.050 +0.20
finite dilution. The heats of transfer listed in Table 11, calculated for dilute solutions ranging in concentration from 0.01 to 0.1 aquamolal, were assumed equal to the standard heats of transfer. Justification for doing so has been given previously.s The cell data are summarized in Table I11 along with the calorimetric data and calculations of the associated free energy and entropy changes for the transfer process. The measured emf values are the cell potentials with equal concentrations in each half-cell. The standard emf values and the associated free energies were calculated by assuming equal activity coefficients for the salts in heavy and normal water. The sign of the emf is defined such that positive values indicate a spontaneous transfer of salt from heavy to normal water. The precision of the measurements was approximately *0.2 mV. For the purpose of calculating entropy values, contributions to the cell emf, which have been shown to result from solvent transport through the cell memT h e Journal of Physical Chemistry
19 $34 +45 +46
...
c
Conod solno AHt,, cal/mol
cal/mol deg
-1.7 i:0 . 2 -1.7 f 0.2
a These data were calculated from integral heat of solution measurements. equal to the transfer values at the final concentrations tabulated in Table 11.
-
cal/mol
+
f0.2 f0.2 zk 0 . 2 f0.2
AStr',
...
...
Standard enthalpies of transfer have been assumed Concentrations equal t o the initial values in Table 11.
branes,S have been neglected. Justification for neglecting them results from a study which indicated that for anion-exchange membrane cells containing halide ions, differences in solvent contributions arising among members of the same family were second order relative to the total contributions characteristic of the family.'O It is assumed that similar characteristics would be exhibited by the aliphatic acid anions and that solventtransport corrections would not change the sequence of emf values shown in Table 111or the resultant sequence of calculated free energies of transfer. Thus, even if the solvent correction characteristic of the family of aliphatic acid anions were large relative to the measured emf values in Table 111,the net effect would be to shift the calculated entropies of transfer up or down scale but not to change the sequence as shown. However, the nonpolar moieties of the aliphatic acid anions do not (9) J. Greyson, J . Phys. Chem., 71, 259 (1967). (10) J. Greyson, ibid., 71, 4549 (1967).
SODIUM SALTB OF ALIPHATIC ACIDS solvate to any large extent,ll and it is not expected that solvent-transport corrections would be large. Thus we conclude that the entropies as shown are correct in sequence and probably correct in value within the general limits of our experimental reproducibility. It is to be noted that the emf values and the resulting transfer free energies for all of the salts indicate that they pass spontaneously from normal to heavy water. However, the standard heat of transfer for the process salt in heavy water into normal water is negative for sodium formate, near zero for sodium propionate and sodium butyrate, and positive for sodium caproate. Calculated entropy values follow the same general pattern as the heats, decreasing in the transfer process for the formate and increasing for the caproate. It is also to be noted that the heat of transfer diminishes in absolute value with increased concentration. Lack of activity coefficient data prevents formal calculation of the concentration dependence of the entropy of transfer. However, a characteristic of transfer measurements between heavy and normal water solutions has been that the value of the entropy is dominated by the heat of transfer, even in concentrated solutions.12 Dominance of the heats can be seen in Table I11 for the dilute solutions of formate and caproate. It seems, therefore, that it is not unreasonable to expect dominance of the heats for these salts in concentrated solutions also. Thus, the diminishing heat of transfer can be used as an indication of diminishing entropy of transfer with increasing solution concentration. Considerable attention has recently been given to the Gurney cosphere model13 for structural interactions in aqueous system^.'^^^^^'^ As will be seen, the model provides an almost ideal explanation for the results of the work reported here. The Gurney model in its simplest terms states that the solute particles in aqueous systems are surrounded by spherical shells of solvent modified in structure from that in the pure solvent. The concentration dependence of properties attributable to structural interactions results from the increasing overlap of the spherical shells with increasing concentration i.e.,the concentration of cosphere solvent per mole of solute decreases with increasing solute concentration. In the process of solution of solute to infinite dilution, the entropy per mole of salt associated with the formation of cosphere solvent could be negative or positive as determined by whether the modified structure of the shell were more or less ordered than the pure solvent. Similarly, because of the changing concentration of cosphere solvent per mole of solute, the process of dilution will give rise to an entropy change dependent on the relative ordering of pure and cosphere solvent. A structure-breaking solute would, therefore, exhibit a positive entropy change in the process of solution to infinite dilution because of the formation of disordered
2151 SALT + H20 OR D20
\,,
CONCENTRATION I N H20 c
SALT I N I N F I N I T E L Y D I L U T E H20 SOLUTION
C O N C ~ ~ ~ ! l cO N
SALT I N I N F I N I T E L Y D I L U T E D 2 0 SOLUTION
Figure 1. Process for transfer of salt from
DzO t o HpO.
solvent shells about each particle. Dilution of a concentrated solution of the same solute to infinite dilution would also exhibit an increase in entropy because ordered solvent would be used to complete the disordered solvent shells, which in the concentrated solution had been completed by overlapping cospheres. The contrary, of course, would be true of structuremaking solutes and, in general, the sign of the entropy change associated with both solution and dilution processes for a given solution will be the same. Application of the Gurney model to transfer of solute between heavy and normal water is best illustrated by reference to the processes shown in Figure 1. The entropy of transfer between the infinitely dilute solutions is given by the equation
AS,:
= AS,oiHSo
-
ASsolD20
(1) for the process of transfer of salt from heavy into normal water. For the same process in concentrated solution, the entropy of transfer can be expressed by the equation AStr’ = ASdilDZ0
+
AStr’
-
ASdilHZ0
(2) Since heavy water is more structured than normal water, the entropy change associated with the solution of structure-breaking salts (Le., the formation of structure-broken cospheres) in the DzO would be more positive than the entropy associated with solution of the same species in normal water, if one can assume that cospheres for each of the solvents are equally disordered. (11) T. L. Kavanau, “Water and Solute Water Interactions,” Holden-Day, Inc., San Francisco, Calif., 1964. (12) Y. C.Wu and H. L. Friedman, J. Phys. Chem., 7 0 , 166 (1966). (13) R. W. Gurney, “Ionic Processes in Solution,” Dover Publications, Inc., New York, N. Y., 1962. (14) R. H. Wood, R. A. Rooney, and J. N. Braddock, J. Phys. Chem., 7 3 , 1673 (1969). (15) J. E.Desnoyers, M.Arel, G. Perron, and C. Jolicoeur, ibid., 7 3 , 3346 (1969). Volume 74, Number 10
M a y 14, 1970
2152 Thus, from eq 1, the entropy of transfer, A&$, for a structure-breaking solute is seen to be negative in passing from heavy into normal water. I n a dilution process, in which a structure breaker is in solution, the resulting increased concentration of disordered shells per mole of solute will result in a more positive entropy change in the more highly structured D20than it will in H2O. Thus, values of AXdilDZo will also be more positive than A&'d,lHZ0. The net result, as examination of eq 2 reveals, is that AStrCis less negative than ASt:. Consideration of solution and dilution processes for structure-making solutes leads t:, similar conclusions except that the signs are reversed. That is, st>ructuremaking solutes will give rise to shells with more structure than pure solvent. Thus, solution of a structure maker in D20will lead to a larger reduction in entropy (more negative AS) than solution of the same species in normal water. From eq 1, it can be seen, therefore, that ASt$ will be positive. Dilution processes will also lead to reduction in entropy because of the formation of more shells per mole of salt. However, the value of ASd,lDzo will be more negative than ASd,lH20 because more structure will be made in the cospheres of heavy water than in the cospheres of normal water. Thus, as can be seen from eq 2, the value of AStrowill become less positive as concentration increases. I n general, then, application of the Gurney model to processes under way in transferring solutes between heavy and normal water leads to the conclusion that increases in solute concentration lead to decreases in the absolute value of the entropy of transfer, regardless of whether the solute behaves as a structure maker or breaker. That the characteristics as described in the foregoing discussion are displayed by the aliphatic acid salts can be seen by examination of the data in Table 111. I n passing from sodium formate to sodium caproate, the transfer entropies in the dilute solutions pass from negative t o positive values, indicating a transition from structure-breaking to structure-making influence.
T h e JOUTTW~ of Physical Chemistry
HARRIET SNELLAND JEROME GREYSON Such results are in accord with the prevailing view that, in proportion to their carbon number, saturated hydrocarbon moieties increase the degree of hydrogen bonding among neighboring water molecules and the apparent structuring in water. The data are also in accord with the proposed formation of shells about the solutes with structural properties progressing from less ordered to more ordered than solvent. An noted before, the heats of transfer for the concentrated solutions, compared to those for the dilute solutions, decrease in absolute value for all the salts, except perhaps for the sodium isobutyrate. The isobutyrate displays no difference within experimental error; however, it may represent a transition entity, which is neither structure breaking nor making. The extremes, sodium formate and caproate, do display a clear decrease in absolute value of the heat of transfer with increasing concentration. If the assumption that the heats of transfer are representative of the entropy values is valid, the data are in agreement with the model presented; i e . , decreases in transfer entropy are exhibited with increased concentration regardless of the structure-influencing properties of the solute. I n conclusion, it is worth noting that the concentrations of sodium valerate and caproate in the solutions used for the dilution experiments were in excess of their critical micelle concentrations.16 Rlicelle formation can be looked upon as a process in which solute collects into aggregates of individual particles with solvent shells coalescing about the individual aggregates. Upon dilution, demicellization would result in the formation of shells about the individual members of the aggregates, with a corresponding reduct,ion in ent,ropy per mole of solute. Thus, the presence of micelles in the concentrated solutions strengthens the fit of the cosphere model to the data since the coalescence of solvent about an aggregate of individual particles is really an example of extreme cosphcre overlap. (16) E. R. B. Smith and R. 70 (1942).
A. Robinson,
Trans. Faraday SOC.,38,
