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Partial Volume Expansibility of Simple Organic Solutes (18) M. irie, M. Shimizu, and H. Yoshida, Chem. Phys. Lett., 25, 102 (1974). (19) Badger and Brocklehurst" did not measure the spectra at elevated temperature before refreezing, and the viscosity of the 1:l n-butyl chioride-isopentane matrix is not known relative to our 3MP and MCH matrices. However the slight warming mentioned in ref 11 is large enough to equilibrate the monomer-dimer equilibrium (Figure 2 for toluene in ref 11). in such conditions our analogous systems with benzene and tertbutyibenzene always showed a charge resonance band shift. (20) B. Badger and B. Brocklehurst, Trans. Faraday SOC.,66, 2939 (1970). (21) J. M. Bossy, R. E. Buhier, and M. Ebert, J. Am. Chem. Soc., 02, 1099 (1970).
(22) J. M. Bossy and R. E. Buhier, lnt. J. Radiat. Phys. Chem., 6, 85 (1974). (23) I?. E. Buhier, Radiat. Res. Rev., 4, 233-258 (1972). (24) S. Noda, K. Fueki, and Zen-ichiro Kuri, Bull. Chem. Soc. Jpn., 41, 2882 (1968). (25) C. L. Gardner, J. Chem. Phys., 45, 572 (1966). (26) Thanks are due to the referee who pointed out that this controversy stili persists in the view of many chemists. (27) B. Brocklehurst, J. Phys. Chem., 75, 1177 (1971): A. Ekstrom, ibid., 75, 1178(1971). (28) T. Shida and W. H. Hamili, J. Am. Chem. Soc., 68, 3689 (1966). (29) S.Wexler and L. G. Pobo, J. Phys. Chem., 74, 257 (1970).
Partial Volume Expansibility of Simple Organic Solutes from the Temperature of Maximum Density of Aqueous Solutions James R. Kuppers Debrtment of Chemistry, The university of North Carolina at Charlotte, Charlotte, North Carolina 28223 (Received May 9, 1975) Publicationcosts assistedby the University of North Carolina at Charlotte
Partial volume expansibility of some representative organic solu;Ces, in aqueous solution a t 3.98' and infinite dilution, were computed from shifts in the temperature of maximum density. Interplay of electrostrictive disruption of solvent structure by polar groups and the solvent ordering effects associated with the hydrocarbon moiety of solute molecules are thereby brought into focus.
Introduction An acceptable model of solution structure must be consonant with a set of fundamental thermodynamic properties which includes solute partial volumes and partial thermal expansibilities. There is a unique experimental method for acquiring thermal expansibility data on dilute aqueous solutions near 3.98O which is attributable to the existence of maximum solution density near this temperature. Earlier work on this topic was reviewed by Franks.' The shift in the temperature of maximum density under the influence of solute in dilute aqueous solution was expressed by Wada and Umeda2 as A0 = -110
- X ) ~ ~ V ~ * [ X-F aAVM/aT] (UV~~
(1)
where x is solute mole fraction, a is the thermal coefficient of expansion of pure solute, p is the coefficient in the parabolic relation to temperature of the molar volume of water in the vicinity of 3.98', Vzo is the molar volume of pure solute at O", VI* is the molar volume of water a t 3.98', and AVM is the excess volume of mixing. Frank3 identified an approximate relationship between the apparent molal expansibility, a, and A0 = -A.8/m[a2V/aT2]
(2)
where m is the molality and where the second derivative is to be evaluated a t 3.98' and infinite dilution. Subsequently it has been shown4 that eq 1 can be cast into a form which provides a direct link between experiment and a value for the partial molal expansibility of solute a t 3.98' and infinite dilution aVZ*/aT = -2pV1*[A0/x]~
(3)
or the thermal coefficient of partial volume expansion of solute d*
= -~/~V~*/P~*[AO/X]L
(4)
where [AO!X]L is the limiting value of AO/x a t infinite dilution and V2* is the partial molal volume of solute a t 3.98O and infinite dilution, The reliability of such solute expansibility data will be limited by the precision of extrapolation from experimental measurements. AOlx is quite insensitive t o solute concentration for a large number of electrolytes4 at high dilution. However, nonlinear extrapolation functions may be needed for other cases in which this is not true. This is a report of some experimental measurements of AO/x and the calculation of partial expansibility data from these, and from other experimental measurements reported in the literature, for some simple organic solutes in aqueous solution.
Experimental Section Measurements of A0 were made by the method of Wada and Umeda2 with the following modifications. Matched dilatometers, one containing the aqueous solution and the other water, were immersed side by side in a controlled temperature water bath. This type of differential measurement eliminated the need for glass expansion corrections and rendered less critical errors in thermal equilibration or temperature measurement. The temperature of maximum density, TMMD, was calculated from a t least eight data points (height, h , of liquid in dilatometer arms) taken over a range of TDMf 2 O . Multilinear regression was used to find the best fit to the equaThe Journal of Physical Chemistry, Vol. 79, No. 20, 1975
James R. Kuppers
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TABLE I: Partial Specific Volumes, Shifts in Temperature of Maximum Density Vs. Concentration, and Thermal Coefficients of Partial Volume Expansion for Some Simple Organic Solutes in Water at Infinite Dilution and 3.98" [ A O/XlL,
u*,
Met hano 1 Form'c acid Ethanol E thanediol
CH,O CHz02 C2H60 C2H602
Acetic acid Glycine 2- Propanol 1- Propanol 1,2- Propanediol Glyc e r ol 2- Propanone Propanoic acid cy- Alanine p- Alanine 2-Methyl- 2-propanol 2-ButanOl 2-Methyl- 1-propanol 1-Butanol Diethyl ether Tetrahydrofuran
C2H402 C2H502N
C3H80 C3H80?
C3H803 C3H60 C3H60?
c3 HT02
f
C4H100
ml g-'
1.186a 0 .760d
1.187a 0.864
0.833d 0.530 1.188" 1.l65" 0.926 0.768 1.121
0.877 0.657 0.640 1.185b 1.160" 1.159" 1.156" 1.093 1.O5Oc
deg mole fraction-'
1o3Z*,
42e -418 141e -1 8ge -172 -413 -548 140e 11P 4 6" -164e -120e -407 -399 -501 393e 140e
-0.32 3.44 -0.74 1.02 0.92 2.38 3.99 -0.56 -0.4 8 -0.19 0.67 0.53 1.80 1.96 2.53 -1.29 -0.47 -0.24 -0.24 0.16 0.85 0.91 1.46
7 le
?le -44 c 4 HSO -223e -238 C4H802 2-M,ethylpropanoic acid 0,917 -408 1,4-Dioxane 0.895c -410e -414 1,3-Dioxane 0.724 -451 C5H1002 3-Methylbutanoic acid 0.957 -43 8 gH1 l0ZN Valine 0.756 -368 C5H1302N Leucine 0.760 -580 Isoleucine 0.794 -485 C5H1003N2 Alanylglycine 0.610 -618 C6H1203N? Al any1a1anine 0.670 -628 C8H1603N? Glycylleucine 0.714 -894 a Reference 6. * Reference 7 . Reference 8. Reference 5. e Reference 2. f Amino acids are all in the dl form
tion, h - hMD = c(T - T M D )where ~ , c is a constant. This interpolation equation is based upon the empirical parabolic relation between temperature and volume of water and dilute aqueous solutions near 3.98O. A corresponding value of TMDOfor water in the matched dilatometer was found and AO/x = ( T M D- TMDO)/xwas recorded for each solution. Values of AO/x for a number of solutes were taken from the literature, subject to the condition that they could be extrapolated to infinite dilution with the desired precision. I t was possible to obtain a limiting value of A0/x at infinite dilution, [ A ~ / X ] L for , most solutes by use of a linear extrapolation function, A0/x = ax b, where a and b are constants. Better extrapolation for alcohols, as a class, was achieved with A0/x = axil2 + b. In either case the reliability of the extrapolated value was within f5%. Standard pycnometric methods were employed to measure densities used in the calculation of partial specific volumes within error limits no greater than &I%, but where possible, values were interpolated from density measurements reported in the literature.
+
The Journal of Rhysical Chemistry, Vol. 79, No.20, 1975
deg-'
1.50
1.51 2.04 1.29 1.20 1.67 1.34 1.81 1.69 1.91
Results Three comparisons between values of thermal coefficients of partial volume expansion, h*, computed from the measurements of Wada and Umeda2 and those from my measurements are available in Table I, i.e., ethanediol, tetrahydrofuran, and l,4-dioxane. Agreement is within the confidence limits of f 5 % deduced from my own measurements. The values of h* listed in Table I draw particular attention to monohydric alcohols in this set, all of which are negative, and to the highly symmetrical 2-methyl-2-propanol standing in a class by itself. The dramatic effect of additional hydroxy groups is apparent: compare ethanol with ethanediol and 1-propanol, or 2-propanol with l,2-propanediol and glycerol. Conformational relationships among tetrahydrofuran (THF), 1-butanol (BA), and diethyl ether (EE) suggest the following consideration. If THF is thought of as cyclized EE with concomitant restriction on freedom of reorientation of the hydrocarbon portion of the molecule, then one sees a significant increase in c%*. If, on
Partial Volume Expansibility of Simple Organic Solutes
2107
the other hand, the hydroxy group of BA is replaced by an ether linkage with ring closure to form THF, there is a much greater increase in &*. The series of carboxylic acids included in this set starts with formic acid which has a value of &* comparable with ammonium halide^.^ Increasing size of the hydrocarbon moiety in the progression formic, acetic, propanoic, 2methylpropanoic, and 3-methylbutanoic acid is accompanied by an uninterrupted decline in &* values, ranging from 3.44 x 10-3 to 1.29 x 10-3. The effect of increasing size of the hydrocarbon moiety upon the value of &* for a sequence of amino acids is even more pronounced, ranging from 3.99 X for glycine to 1.20 X for valine. However, the values for the six-carbon isoleucine and leucine are greater than the sequence minimum of valine. The effect of a hydrocarbon tail, unrestricted by hydrophylic substituents or by ring closure, is again illustrated in the comparison of &alanine and a-alanine. The three dipeptides in this set have values of d* less than the average of the respective amino acids from which they are formed.
higher temperatures. Thus a rise in temperature would cause some disruption of the short-range order, and an accommodation of the guest molecule in a somewhat more limited space. Formal electric charge or electric polarity on a guest molecule produces quite a different effect: solvent electrostriction, disruption of the normal hydrogen-bonded structure of liquid water, with the guest molecule occupying a relatively small volume element. This volume element, of course, will expand as the thermal energy of the water molecules increases. When a polar organic molecule with a hydrocarbon moiety is placed in an aqueous environment, each of these two effects of the solvent will be manifest to a greater or less extent, and the partial volume expansibility should reflect the manner in which these two solution-structuring effects are balanced against each other: electrostrictured solvent contributing to greater thermal expansibility and quasiclathrate formation causing a reduction. Moreover, if expansibility data relate to infinite dilution, the complications of solute-solute interactions and their effect on solution structure are eliminated. If the integration of concepts considered in this discusDiscussion sion is accepted, the collection of data in Table I then illustrates the manner in which these contrasting effects on solAlthough some insight into aqueous solution structure was provided by a separatiqn of measured shifts of T M D vent structure balance out in some representative classes of organic solutes. In any event such data should be useful in into and Aflstructural,2further significance of these testing any solution structure model. measurements is revealed by the computation of partial volume expansibilities of the solute. When comparing solutes of vastly different partial molal volumes, it is espeReferences and Notes cially advantageous to remove the volume dimension from expansibility data, as done in defining &*, the thermal coefF. Franks and D. S. Reid, “Water, A Comprehensive Treatise”, Vol. 2, ficient of partial volume expansion. F. Franks, Ed,, Plenum Press, New York, N.Y., 1973, Chapter 1, p 26, Chapter 5, pp 364-365. The concept of a cage-like solvent structure, involving an G. Wada and S. Umeda, Bull. Chem. SOC.,Jpn., 35, 646 (1962); 35, increase in hydrogen bonding in the solvation sphere, is 1797 (1962). H. S. Frank, unpublished work as cited by F. Franks, Ann. N.Y. Aced. supported by a variety of properties of many dilute aqueSci., 125, 287 (1965). ous so1utions.l Especial attention has been focused on the J. R. Kuppers, J. Phys. Chem., 78, 1041 (1974). alcohols since it appears that, among relatively soluble orE. W. Washburn, Ed., “InternationalCritical Tables”, Vol. 111, McQrawHill, New York, N.Y., 1933. pp 122-123. ganic solutes, they have the most profound effect.ly9J0 T o M. E. Friedman and H. A. Scheraga, J. fhys. Chem., 89, 3795 (1965). the extent that such a clathrate-like structure exists, it F. Franks and H. T. Smith, Trans. Faraday SOC.,84, 2962 (1968). F. Franks, M. A. J. Quickenden, D.S. Reid, and B. Watson, trans. Faramust be associated with increased short-range order .in the dav Soc.. 66.582 (1970) solvent near the guest molecule. One can expect that this (9) J. R. Kuppers, J. k g n . keson., 4, 220 (1971); 8, 201 (1972). type of structure would be more favored a t the T M Dthan a t (10) J. R. Kuppers and N. E. Carriker, J. Magn. Reson., 5, 73 (1971).
The Journal of Physical Chemistry, Vol. 79, No. 20, 1975
