Cecil
1000
M .Criss
tion of Ionic Entropies in Various Solvents Cecil M, Criss Department of Chemistry, University of Miami, Coral Gables, Florida 33724
(Received October 29, 7973)
Pubiication costs assisted by The National Science Foundation
The correspondence principle has been employed to divide recently published partial molal entropies for electrolytes in acetone and acetonitrile into their ionic components, The data are expressed in terms of the equation Sz0(X) = a bSz"(H2O) where a and b are constants characteristic of the solvent and So (X) and S2"(H20) are the entropies of the corresponding ions in the nonaqueous solution and water, respectively. Coefficients in the above equation for these solvents are compared with the coefficients for eight other nonaqueous solvents, deuterium oxide, and water in relation to the structural characteristics arid basicities of the solvents. It is shown that solvent structure, rather than basicity, plays the predominant role in determining the value of ionic entropies. A general equation is presented which enables one to estimate ionic entropies in nonaqueous solutions for which there are no data.
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A reliable means of estimating thermodynamic properties of electrolytes in nonaqueous solutions would be of considerable value to chemists and engineers. Some degree of success has been realized in the estimation of standard partial molal. ionic entropies, &".Through the use of correspondence plots, several investigator~l-~ have divided the entropies of electrolytes in nonaqueous solutions into their ionic components and have shown that these entropies are related to the entropies of ions in water by where a and b are constants characteristic of the solvent. Criss, Held and Luksha,3 and Criss and Salomon4 have employed the correspondence plot method and the "absolute" entropies of ions in water5 in an effort to evaluate "absolute" entropies of ions in eight nonaqueous solvents and D20. Within the limits of experimental error, the entropies of both cations and anions (and hence any pair of ions) in the various solvents are observed to increase in the order NH3 < DMF 2 EtOH < MeOH < PC = DMSO < NMF < F < H20 < D20. This order was rationalized in ternis of the inherent structure within the solvent resulting from hydrogen bonding or strong dipoledipole interactions. The highly ordered solvents a t the right end of the series presumably cause the entropies of the ions to be relatively positive because of the large enQopy increase from solvent disruption, while low entropy values are exhibited by ions in solvents at the left end of the series since very little solvent disruption occurs. It was mggesteds that entropies in various solvents could be represented by the equation
how many are restricted, Justification for this assumption was made on the basis that the interactions are electrostatic in nature and that solvent molecules of large dipoles will be strongly oriented near the ion but loosely oriented away from the ion because of the rapid decrease in field strength. On the other hand, for solvent molecules of low dipole moment, the nearest molecules are less severely restricted in their motion, but more molecules are affected since the field surrounding the ion does not diminish as rapidly. Consequently, the first three terms in eq 2 may be replaced by a constant, and since the entropy of disruption of the solvent, AS,', should be directly proportional to the degree of structure in the solvent, it was proposed3 that eq 2 be replaced by
s2"= kS,,,
f &
(3)
where S,,, is the "structural" entropy of the pure solvent and k and C are constants characteristic of each ion. It was further proposed that Sstrwas proportional to the difference in the actual boiling point of the solvent and the boiling point that the solvent would have if there were no strong interactions present, and eq 3 was modified to
S2" = k'ATb,
+&
(4) While this equation is only approximately valid, it can be employed for estimating entropies of ions in solvents for which no data exist, if solvent structure is the major factor affecting the entropy and if one can estimate a reliable value for ATbp for a solvent. In contrast to the above argument, one can make a case for correlating the ionic entropies with the base strengths of the solvents. Within any solvent type (alcohols, amides, Szo S o ASFo ASoo ASDO (2) etc.), ionic entropies become more negative with the genwhere So is the inherent entropy of the ion or pair of erally accepted solvent basicities. Support for this view ions,G AS," is the entropy change because of loss of decomes from studies of enthalpies of solution or of transfer grees of translational freedom of the gaseous ion, AS,' is between solvents. In nearly all of the solvents that have the entropy of orientation of solvent in the electric field been investigated enthalpies of solution or transfer bearound a n ion (always
