FREEZING POINTS, OSMOTICCOEFFICIENTS, AND ACTIVITY COEFFICIENTS OF SALTS
2313
Freezing Points, Osmotic Coefficients, and Activity Coefficients of Salts in N-Methylacetamide. I.
Alkali Halides and Nitrates1
by R. H. Wood,* R. K. Wicker, 11, and R. W. Kreis University of Delaware, Newark, Delaware
19711
(Received J u n e 8, 1970)
Publication costs assisted by the National Science Foundatian
The freezing points of N-methylacetamide solutions of most of the chlorides, bromides, iodides, and nitrates of lithium, sodium, potassium, and cesium have been measured at concentrationsfrom 0.1 to 0.8 m. Osmotic and activity coefficients were calculated from the freezing points and the molal freezing point depression constant. The osmotic and activity coefficients are much higher than in water due to the high dielectric constant of N-methylacetamide. The order of osmotic coefficients is similar to the order in water. This can be explained by the relative ease with which oppositely charged ions can displace solvent from the immediate neighborhood of an ion. The strong structure making action of alkali metal halides in N-methylacetamide indicates that solvent structure breaking is not a necessary condition for observing the order of osmotic coefficient found in both water and N-methylacetamide.
Introduction Water is the only solvent with a high dielectric constant that has been thoroughly investigated. However, aqueous solutions show many anomalous properties such as the expansion of water on freezing, the formation of clathrate compounds, and the formation of “icebergs” around nonpolar solutes. It is recognized that the unique structure of water may have an important influence on the thermodynamic properties of aqueous electrolytes. 3--7 I n particular, structural ideas have been used to explain the relative positions of the activity coefficients of aqueous alkali halidesS6 I n order to understand properties of electrolytes in high dielectric solvents and how they are influenced by the properties of the solvent, it will be necessary to have extensive series of measurements on more than one solvent. The present results represent a step towards the collectian of this kind of data. The solvent N-methylacetamide (NAJA) was chosen for the present measurements because it has a very high dielectric constant ( 6 = 178 a t 30.5’)* which should ensure that effects of ion pairing will be minimized. I n addition, X N A has a chain like hydrogen-bonded structures-16 rather than the branched three-dimensional hydrogen-bonded structure of water. Gas solubility measurements have shomn16 that small amounts of nonpolar solutes in NMA dissolve without appreciably disrupting the linear hydrogen-bonded chains of solvent molecules. The solutes are in contact with the nonpolar groups of the K M A polymer and this explains why NMA is capable of dissolving large amounts of nonpolar solutes. The present paper presents the results of the measurements of the freezing points and thus activity coefficients of some alkali halides and nitrates in S N A .
Dawson and co~orkers,~~-20 and also French and Glover,21 have measured the conductances of many strong electrolytes in NRfA and have shown that there is a negligible association of normal strong electrolytes.
(1) Presented in part at the 154th Kational Meeting of the American Chemical Society, Chicago, Sept 1967. This study was aided by a grant from the National Science Foundation, Grant No. GP5239. (2) For a review see J. L. Kavenau, “Water and Solute-Water Interaction,” Holden-Day, San Francisco, Calif., 1964. (3) H. S. Frank and A . L. Robinson, J . Chem. Phys., 8 , 933 (1940). (4) H. S.Frank and M, W. Evans, ibid., 13, 507 (1945). (5) H. S. Frank and W.-Y. 117en, Discuss. Faraday SOC.,24, 133 (1957). (6) R. W. Gurney, “Ionic Processes in Solutions,” McGraw-Hill, New York, N.Y., 1953, p 256. (7) H. 5. Frank, 2. P h y s . Chem. (Leipeig), 228, 364 (1965). (8) S.J. Bass, W. I. Nathan, R. M. Meighan, and R . H. Cole, J . P h y s . Chem., 68, 509 (1964). (9) J. L. Kats and B. Post, Acta Crgstallogr., 13, 624 (1960). (10) S. Mizushima, T. Simanouti, S. Nagakura, K. Kuratani, 14. Tsuboi, H. Baba, and 0. Fujioka, J. A m e r . Chem. Soc., 72, 3490 (1950). (11) G. R. Leader and J. F. Gormley, ibid., 73, 5731 (1951). (12) R. Linn and W. Dannhauser, J . P h y s . Chem., 67, 1805 (1963). (15) L. A. Planche, H. B. Thompson, and M. T. Rogers, (bid., 69, 1482 (1965). (14) M . Davies and D. K. Thomas, ibid., 60, 767 (1956). (15) I. M. Klots and J. S. Fraeen, J . A m e r . Chem. Soc., 84, 3461 (1962). (16) R. H. Wood and D. DeLaney, J . P h y s . Chem., 72,4651 (1968). (17) L. R. Dawson, P. G. Sears, and R. H. Graves, J . A m e r . Chem. Soe., 77, 1986 (1955). (18) L . R. Dawson, E. D. Wilhoit, and P. G. Sears, ( b i d . , 78, 1569 (1956). (19) L. R. Dawson, E. D . Wilhoit, R. R. Holmes, and P. G. Sears, ibid., 79, 3004 (1957). (20) L. R. Dawson, G. R. Lester, and P. G. Sears, ibid., 80, 4233 (1958). (21) C. M. French and K . H. Glover, Trans. Faradag Soc., 51, 1427 (1955).
The Journal of Physical Chemistry, Vol. 76, S o . 16, 1971
2314 I n another series of papers, Dawson and coworkers have measured the activity coefficients of HC1 in NMA up to 0.1 M and have shown that the HC1 behaves as a strong electrolyte in agreement with Debye-Huckel theory. 2 2 , 2 3 Bonner and c o ~ o r k e r s ~have ~ - ~reported ~ some activity coefficients which change very rapidly at low concentrations. However, the freezing point depression constant used by these workers has been shown to be off by 15%.27 Recent resultsz6 on NaI and KI when corrected for the new freezing point depression constant are consistent with the present results ( = k 2 % ) from 0.1 to 0.005 m. Apparently there was an error in the early measuring technique of these authors at low concentrations. Holleck, Cogley, and Butlerz8have found that large errors are possible with the warming curve technique used by Bonner, et al.
Experimental Section The lithium bromide and the cesium salts used in this work were analyzed for sodium and potassium with a flame photometer. Impurities were negligible except for the lithium bromide which contained 0.2% sodium. All other salts were of reagent grade. After being dried in an appropriate manner, the salts were stored in a drybox. In addition, the dryness of those salts of a particularly hygroscopic nature was analyzed by a Karl Fischer titration. N-Rfethylacetamide, obtained commercially, was vacuum distilled after being dried over calcium hydride. The final purification of the solvent is described below. The experimental procedure has been described previously in detail and, therefore, is only briefly summarized here.29p30 Two different freezing point cells were used. The first set of measurements used a dewar flask fitted with a rubber stopper containing openings for the salt inlet tube, thermistor well, sample port, and nitrogen inlet. The freezing point cell used in the later experiments consisted of a solvent well which was sealed at the top by a specially constructed cap fitted with three ground-glass joints and a stopcock which served, respectively, as a salt inlet, thermistor well inlet, sample port, and nitrogen inlet. Stirring was achieved by means of a magnetic stirring bar placed into the solvent well. During an experiment the solvent well was inserted into a strip silvered dewar around which was wound a copper coil through which thermostated water was circulated to minimize heat exchanges with the surroundings. In the first experiments the final purification of the solvent was achieved by slow recrystallization in a drybox. The solvent was then transferred to the freezing point cell with a dry syringe. The major impurity was water and initiaI freezing- points and Karl Fischer titra_ tiOnS show the water content varied from 0.03 to 0.001 m. The freezing point of pure NMA was determined to be 30.56". In later experiments the h-1\/IA was refined by zone The Journal of Physical Chemistry, Val. 7 6 , N o . 16, 1971
R. H. WOOD,R. K. WICKER,11, AND R. W. KREIS melting in a side tube of the freezing cell. Prior to each experiment the entire apparatus was dried at 120" and assembled while hot. The entire cell and zone purification tube were evacuated through the stopcock and flamed with a Bunsen burner. Upon cooling, the cell was put under a slight positive pressure of dry nitrogen and this was maintained until the end of the experiment. Approximately 135 ml of NMA was then transferred from the drybox in a syringe and loaded into the zone melting tube. After purification by repeated zone melting the top 25 ml of NMA was melted and poured into the freezing point cell. In these experiments the concentration of impurities varied from 0.001 to about 0.0002 m as determined by fractional freezing experiments. Karl Fischer titrations showed that the water content was less than 0.001 m. Calorimetric measurements on a sample prepared in the same way showed a total impurity of 0.0005 m.27 Temperatures were measured by a thermistor connected to a resistance bridge. After temperature equilibration, the freezing point of the solvent was taken and a small amount of salt was added from the salt pistol. After equilibration at the new temperature, a sample was removed using a heated syringe with a long needle and more salt was added. The process was repeated until saturation was reached. Karl Fischer titrations indicated that no appreciable water was introduced during an experiment. The actual freezing point of the solvent in each run was used to calculate the freezing point depression. In this way any water present in the solvent would affect the initial and final freezing points by the same amount provided the water did not preferentially solvate the salts and the amount of solid in the equilibrium mixture stayed constant. The halide concentrations of the samples were determined by titrating with standard AgNO3 from a micropipet, using an absorption indicator. Synthetic standards showed results accurate to 0.1%. The nitrate solutions were analyzed by evaporating to dryness a t 140". Synthetic samples showed that this procedure was accurate to 1%. (22) L. R. Dawson, R. C. Sheridan, and H. C. Eckstrom, J . Phys. Chem., 65, 1829 (1961). (23) L. R. Dawson, W. H. Zuber, Jr., and J. C. Eckstrom, ibid., 69, 1335 (1965). (24) 0. D. Bonner, C. F. Jordan, and K. W. Bunzl, ibid., 68, 2450 (1964). (25) 0. D. Bonner, K. W. Bunal, and G. B. Woolsey, ibid., 70, 778 (1966). (26) 0. D . Bonner, S. J. Kim, and A . L. Torres, {bid., 73, 1968 (1969). (27) R . \V. Kreis and R. H. Wood, J . Chem. Thermodyn., 1, 523 (1969). (28) G. Holleck, D. R. Cogley, and J. N. Butler, J . Electrochem. Sac., 116, 952 (1969). (29) R. K. Wicker, Ph.D. Thesis, University of Delaware, 1966. (30) R. W. Kreis, Ph.D. Thesis, University of Delaware, 1969. (Note that the molecular weight used for CsBr is incorrect.)
2315
FREEZING POINTS, OSMOTIC COEFFICIENTS, AND ACTIVITYCOEFFICIENTS OF SALTS and upon rearrangement
Results The activity coefficients were represented by an extended form of Guggenheim's3I equation log
y* =
+ ml/') + Bm + (?ma'/" (3)
-Aml/'/(l
The Debye-Huckel coefficient, A , was calculated using the dielectric constant of Bass, Nathan, Meighan, and Cole8 and the density of Dawson and Griffith.32 The result was A = 0.14128 for NMA. Integration of eq 3 in the normal manner yields the following expression for osmotic coefficients, 4
where ff(m'l') =
+ ml'/" 1/(1 + m1j2)- 2 In (1 + m1'2)](5)
(3/m8'/")[l
The practical osmotic coefficient, (6, of a solvent as a function of temperature depression, Bi is given by33
(A-
2ACp0f 3TrAH"t
+ ..
where the heat of fusion (AHf = 31.78 f 0.16 cal/g), change in heat capacity (ACpor = 0.12 =t 0.03 cal/g OK), and freezing point constant (K = 5.77 f 0.02"K/m) are taken from Kreis and Wood.27 The experimental dataZgsa0 were fit to eq 9 by the method of least squares to calculate values of B and C. I n the case of NaCI, KCl, and CsCl where the solubility was low, the C term was set equal to zero since it was not needed for a good fit. The results of the calculations are given in Tables 1-111 and Figure 1. The osmotic and activity coefficients were calculated by eq 4 and 3, respectively. The results in Table I and examination of the data indicate that the great majority of the differences between the least-squares fit and the experimental osmotic coefficients are less than 0.01 indicating an accuracy better than 1%in (6. The data for S a 1 are not as good as the rest, and errors as high as 2% would be possible. The first measurements of NaCl, KC1, and LiBr and LiN03 showed errors much higher than this so the osmotic coefficient was remeasured using the redesigned apparatus. The new measurements changed the
Standarda Salt I
(6)
11,the relative partial molal enthalpy, and J1, the relative partial molal heat capacity, are not available for NMA solutions.34 The term in A b takes into account the change in heat capacity, ACpglr with temperature and is negligible relative to other uncertainties. Ignoring terms contributing less than 0.1% of 4 at 8 = 10" or less and omitting those terms which presently are not available, eq 6 reduces to
LiCl NaClb KC1b CsCl LiBrb
NaBr KBr CsBr NaI KI
CSI LiN08 NaN03
KNOi
Substituting eq 4 in eq 7 yields vKm[l
- 2.303 --Am"'g(m1~')]
B
0.5555 0,2101 0.1406 0.0835 0.2872 0.3087 0,1415 0.3284 0.3776 0.1084 0.1616 0.2127 - 0.0481 -0.0498
C
error of e
-0.4089
0.017
Standard error of d
0.006 0.011 0.017 0.021 0.003
0.0034 -0,0685 -0.0070 -0,4363
0.015 0.028 0.031
0.008
0.003
-0,0498
0,071 0.046 0.012 0.048
0,024 0.004 0.004 0.014 0.006 0.010
0.0859 -0.0828 0.0474 0.0069 -0.1051
0.023 0.024
0.005 0.007
This is the standard deviation of the left-hand side of eq 9. Remeasured using zone melted NMA.
+
3 2.303KBm2
+ 2.303K(6/s)CmK//"(9)
Table I : Results of the Least-Squares Fit
Ab L +-6AH0r +--Tr2AHor 2J1 )83 3TrAH" f
2.303KBm2
+ 2.303(6/s)KCm6//9=
(31) E. A. Guggenheim, Phil. Mag., 19, 588 (1935). (32) L. R. Dawson and E. J. Griffith, J . Phys. Chem., 56,281 (1952). (33) G. N. Lewis and M. Randall, "Thermodynamics," K. S. Pitzer and L. Brewer, Ed., McGraw-Hill, New York, N. Y . , 1961, p 406. (34) By analogy with aqueous solutions, the inclusion of ZI and 71 in eq 6 would result in a correction to + of less than 1% over the
concentration range studied.
The Journal of Physieal Chemistry, Vol. 76, N o . 16, 1971
R. H. WOOD,R. K. WICKER,11, AND R. W. KREIS
2316 Table I1 : Osmotic Coefficients, 6 m
LiCl
NaCl
KC1
CsCl
LiBr
NaBr
KBr
CsBr
NaI
KI
CsI
LiNOa
NaNOa
KNOa
0.01
0.996 1.008 1.023
0.993 0.994 1.001 1.020 1.042
0.992 0.990 0.994
0.992 0.987 0.987
0.994 0.999 1.011 1.039 1.070 1.101 1.133 1.166
0.994 0,999 1.010 1.035 1.061 1.086
0.992 0.990 0.993 1.004 1.017 1.030
0.994 0.994 0.996 0.994
0.995 1.003 1.019 1.053 1,089 1.124 1.159 1.194 1.228 1.262
0.992 0,990 0.994 1,008 1.027 1,048 1.071 1.095 1.122
0.992 0.990 0.991 0.997 1.003 1.009
0,993 0.995 1.004 1.027 1.054 1.082 1.112 1.143
0,990 0.979 0.972 0.962 0.955 0.948 0.942 0.937 0.931 0.926
0,990 0.978 0.967 0.948
0.05 0.10 0.20 0.30 0.40 0.50 0.60 0.70
1.050 1.068
1.111 1.135
1.158
0.80
Table 111: Activity Coefficients,
y5
m
LiCl
NaCl
KC1
CeCl
LiBr
NaBr
KBr
CsBr
NaI
KI
CsI
LiNOa
NaNOs
KNOa
0.01 0.05
0.982 0.993 1.020 1,073 1.120
0.976 0.966 0.971 0,996 1.031
0.974 0.958 0.955
0.973 0.952 0.943
0.977 0.974 0.988 1.033 1.088 1,151 1.220 1.295
0.977 0.974 0.988 1.028 1.075 1.126 1,179 1,235 1.293
0.974 0.957 0.955 0.964 0.980 1.000
0.977 0.968 0.966 0.962
0.979 0.983 1.005 1.065 1,135 1.213 1.296 1.386 1.482 1.584
0.973 0,956 0.954 0.968 0.993 1.025 1.062 1.105 1.154
0.974 0.956 0.952 0.953 0,959 0.966
0.975 0.966 0.974
0.970 0.937 0.915 0.886 0.864 0.847 0.832
0.969 0.934 0.907 0.864
0.10 0.20 0.30 0.40 0.50 0.60 0.70
0.80
I
I
I
.4
.6
I 1.1
0 1.0
0.9
-
‘Cs Br
.2
m -
Figure 1. Osmotic coefficients in AT-methylacetamide us. m.
values of 9 by less than 1% for NaCl, 1.5% for KCl, 1.5Yc for LiN03, and 4Yc for LiBr. Points which differed from the final least-squares fit more than about 2 u were rejected. This resulted in the dropping of one point from the KBr data and one from the CsBr data. The points above about 0.8 m tended to have larger random errors so they were not included in the leastsquares fit. T h e Journal of Physical Chemistry, Vol. 16, N o . 16, 1971
1.007 1.051 1.102 1.160 1.224
0.818 0.805 0,794
Measurements on LiI solutions had very high random errors in both sets of measurements for unknown r e a s o n ~ . These ~ ~ data are not reported.35 At low concentrations the osmotic coefficients are much higher than in water, and this is due to the high dielectric constant of NMA which results in a DebyeHuckel slope that is one-fourth as large as the DebyeHuckel slope for water. At the highest concentrations where the most accuracy is obtained, the order of osmotic coefficients is NaI > LiBr > LiCl > LiK08 > KaBr > NaCl > KI > KBr > CsI > CsBr > CsCl > Na;‘U’03> KNO, with KC1 close to KI and KBr. This order is consistent with Li+ > Na+ > K + > Cs+ and I- > Br- > C1- > xo3-. The lowconcentration data for LiCl and CsBr do not follow this order but this could easily be due to the greater experimental error at low concentrations. The unusually large values of C (Table I and eq 3) for these two salts also suggests this kind of error. The solubility of KC1 is not high enough t o see if it fits this order. The order observed a t high concentrations in NMA is the same as the order in water except that in water the cesium salts show the reverse anion order. It should be noted that although the cesium salts do show a reversal (35) The measurements show LiI to have the highest value of 4 measured and this is consistent with the interpretation presented.
FREEZING POINTS, OSMOTIC COEFFICIENTS, AND ACTIVITY COEFFICIENTS OF SALTS in order in water and KRIA, the osmotic coefficients of the cesium salts in both solvents are very close to each other. This similarity in trends is brought out by a plot of + N M A us. +H%O which shows that there is a rough correlation. Although the data are limited, it looks as if a similar rough correlation will hold for alkali halides in formamide. 37 These regularities indicate that the same factors are important in both the water and the NMA solutions (and possibly in formamide) and that the effects depend in roughly the same way on the sizes of the ions. It may be that this similarity will hold for any strongly hydrogen-bonded solvent with a high dielectric constant. At moderate concentrations the osmotic coefficients of the alkali metal halides and the nitrates will depend mostly on the interactions of the two oppositely charged ions although like-charged ion interactions will also c ~ n t r i b u t e . ~The * ~ ~interactions ~ between ions are influenced by the solvation of the the relative ease with which ions of opposite charge displace solvent from the coordination sphere of an ion,44-46the structure of the solvent in the solvation spheres of the ions,3j6,7147 and the “localized solvolysis” induced by weak a ~ i d s . ~All~ ,of~ these ~ effects and others effects not listed may be important, and the problem is that it is impossible to distinguish between effects which predict the observed experimental results. For this reason the present discussion will be limited to showing how the competition of oppositely charged ions and solvent for the immediate company of an ion can account for the data and how the evidence indicates that solvent structure breaking is not an important factor. P r ~ ehas ~ proposed ~ t ~ ~ that the competition of solvent and oppositely charged ions for the immediate company of an ion is an important effect of the solvent. A crude model predicts that ions that are both small or both large can more easily come into contact. This can be understood by comparing the interaction energy of a small ion with a coordinated solvent molecule and with an oppositely charged ion. If the oppositely charged ion is muchlarger than the solvent molecule, the interaction energy of the two ions is less than the interaction energy of the ion with the solvent dipole. Because of this, a large ion cannot easily enter the coordination sphere of a small ion and direct ion-ion contacts do not often occur. I n the case of two small ions or two large ions the ion-ion interaction is larger than the ion-solvent dipole interaction so that in these cases ion-ion contact is more f r e q ~ e n t . ~ ~ , ~ ~ The cation order in the osmotic coefficients of NMA solutions, Li+ > IYa+ > I Csf, and the anion order, I- > Br- > C1- > SOa-, are explained by this model with the assumption that the alkali metal cations are small strongly solvated ions and the anions are relatively weakly solvated. As the cation gets smaller and the anion gets larger, it is more difficult for the solvent displacement to occur, ion-ion repulsion increases, and 36t
2317
the osmotic coefficient goes up. The nitrate ion acts as an ion that is smaller than the chloride ion because its coordinating groups, the oxygen atoms, are about the same size as a fluoride ion. The hydration association model of Frank44 is an early attempt to put some of these effects on a more quantitative basis. A calculation of the energy of displacing two solvent molecules and forming a contact ion pair in an aqueous solution predicts the correct trends in activity coefficients. Specifically, there is evidence for the reversals that occur in water as the anion gets smaller and as the cation gets larger. The model takes into account dielectric saturation, ion-ion, ion-dipole, and Van der Waals forces in calculating the energy of the displacement reaction. The following paper shows that with an ion that is able to displace solvent from the coordination sphere of the alkali metals, the order of osmotic coefficients is reversed in NMA. This is just what would be predicted by the theory. Recent calculations on a charged square-well mode150 show that very small changes in the energy of the ions a t contact ( E < IcT) are adequate to account for the osmotic coefficients of aqueous solutions. These small energy changes do not greatly affect the pair correlation function, and thus do not change the probability of an ion-ion encounter to any great extent. It appears that osmotic coefficients are a very sensitive measure of ionion interactions. Frank and Wen5 picture the water around an ion as being divided roughly into three regions: strongly hydrated water near the ion; a transition region where the normal water structure is broken; and an outer region of relatively undisturbed water. This picture of (36) E. N. Vassenko, Zh. F i z . Khim., 21, 361 (1947); 22, 999 (1948); 23, 959 (1949). (37) The data of E. Luksha and C. M. Criss [ J . Phys. Chem., 70, 1496 (1966) ] in N-methylformamide indicate trends opposite to those in water but in view of the difficulty of cell measurements this is
not certain. (38) R. H. Wood and R . W.Smith, J . Phys. Chem., 69, 2974 (1965). (39) R. H . Wood and H. L. Anderson, ibid., 70, 992 (1966). (40) N. Bjerrum, 2. Anorg. Allg. Chem., 109, 275 (1920). (41) R . H. Stokes and R. A . Robinson, J . Amer. Chem. Soc., 70, 1870 (1948). (42) (43) (44) (45)
E. Glueckauf, Trans. Faraday Soc., 5 1 , 1235 (1955). D. GI. Miller, J . Phys. Chem., 60, 1296 (1956). H. S. Frank, J . Amer. Chem. Soc., 63, 1789 (1941).
(a) J. E. Prue, “Chemical Physics of Ionic Solutions,” B. E . Conway and R. G. Barradas, Ed., Wiley, New York, N. Y., 1966, p 163; (b) J. E. Prue, “The International Encyclopedia of Physical Chemistry and Chemical Physics,” Vol. 3, “Ionic Equilibria,” Pergamon Press, Oxford, 1966, Chapter 10, Topic 15. (46) R. Garnsey and J. E. Prue, Trans. Faraday Soc., 64, 1206 (1968). (47) R . H . Wood and H. L. Anderson, J . Phys. Chem., 71, 1869 (1967). (48) R. A . Robinson and H. S. Harned, Chem. Rev., 28, 419 (1941). (49) R. M . Diamond, J . Amer. Chem. SOC., 80,480 (1958). (50) R . F. Rasaiah and €I. L. Friedman, J . Phys. Chem., 72, 3352 (1968).
The Journal of Physical Chemistry, Vol. 76, No. 16, 1971
R. H. WOOD,R. K. WICKER,11, AND R. W. KREIS
2318 the structure of water around an ion explains the viscosity, conductance, and heats of interaction of the ions very n i ~ e l y . ~ - ~ Gurneye s ~ ~ proposed that the relative amounts of structure making or structure breaking present were responsible for the trends of the activity coefficients of the alkali halides. The similarity of the trends of activity coefficients in water and XMA raises the question of whether the structural properties of water and NMA are similar. All of the evidence indicates that structure breaking is not important in solutions of alkali halides in any of the N-methylamides. Viscosity measurements of electrolytes in N-me thylf or mamide, N-met hylacetamide, l7 and N methylpr~pionamide~~ indicate that structure making is the predominant effect for alkali halides. The transference numbers of KC1 in NR4A led Gopal and Bhatnagar5a to conclude that the potassium ion has no appreciable structure breaking. Engle and Hertzs4 conclude from nuclear magnetic relaxation studies on alkali halides that there is no structure breaking in N-methylformamide. However, French and G10ver21 and Singh, Rastogi, and GopaIs6 conclude from conductance measurements that tetraalkylammonium ions depolymerize KRlA. Similarly RlillerosZbhas shown that benzene and pyridine depolymerize N-methylpropionamide solutions. Apparently the alkali halides are all net structure makers while some nonelectrolytes and perhaps the tetraalkylammonium ions are structure breakers. This strong structure-making effect of the alkali metal halides is just what might be expected from the properties of the solvent. Each ion in N U A should act as a cross-linking agent for the chains of hydrogenbonded solvent molecules that occur in the pure solvent, 9-16 because each ion mould serve as a starting point for four to six chains, depending on its coordination number. The chains would be fairly long because coordinated X l I A molecules would be strongly polarized and thus have an increased tendency to hydrogen bond,13
r
The Journal of Physical Chemistry, Vol. 76, N o . 16, 1971
The fact that NMA solutions of the alkali halides show no evidence of structure breaking together with the similarity56ain the order of osmotic coefficients in water and NMA indicates that net structure breakingseb is not a necessary condition for observing this order. Frank and K e r ~ i nmeasured ~~ osmotic coefficients of alkali halides in D20 and showed that the osmotic coefficients of the alkali halides were not sensitive to the structural differences between D20 and H20a5*Similarly, Robinsonsgshowed that the osmotic coefficients of a mixture of NaCl and KC1 are very similar in H 2 0 and D20. It may be that osmotic coefficients in both N3IA and water are not sensitive to small changes in the structure of the solvent. I n contrast to this the enthalpies and entropies of dilution60p61appear to be very sensitive to the structural difference between H2O and
D2O.
Acknowledgment. The authors would like to acknowledge the helpful discussions with Dr. John E. Prue and Professor Henry S. Frank. (51) D. Feakins and K. G . Lawrence, J. Chem. Soc., 212 (1966). (52) (a) T . B. Hoover, J. Phvs. Chem., 68, 876 (1964); (b) F. J. Millero, {bid., 72, 3209 (1968). (53) R. Gopal and 0. N. Bhatnagar, ibid., 69, 2382 (1965). (54) G . Engel and H. G. Hertz, Ber. Bunsenges, Phys. Chem., 72, 808 (1968). (55) R . D. Singh, P. P. Rastogi, and R. Gopal, Can. J . Chem., 46, 3525 (1968). (56) (a) The anion order for cesium salts is the only difference between the two solvents (see above). (b) It is probably true that there will always be at least a small transition region between the solvated ion and the solvent which will have less structure than the solvent. The term net structure breaker is used here to denote an ion where the structure-breaker region is small compared with the region where structure-making predominates. (57) R. E. Kerwin, Ph.D. Thesis, University of Pittsburgh, 1964. (58) G . Nemethy and H. A . Scheraga, J. Chem. Phus., 41, 680 (1964). (59) R . A. Robinson,
J. Phys. Chem., 73, 3165 (1969).
(60) Y. C. Wu and H. L. Friedman, ibid., 70, 2020 (1966). (61) R. H. Wood, R. A. Rooney, and J. N. Braddock, 1693 (1969).
ibid., 73,
