J. Chem. Thermodynamics 1971, 3, 837-841

The osmotic and activity coefficients of some solutes in deuterium oxide

O. D. BONNER

Department of Chemistry, University of South Carolina, Columbia, South Carolina 29208, U.S.A.

(Received 4 May 1971)

Osmotic and activity coefficients are reported for 11 solutes in D20 at 25°C. Activity coefficients of lithium and sodium salts are somewhat larger in D20 than in H20 while those of caesium and silver salts are smaller. The activity coefficients of dextrose and sym- dimcthylurea are surprisingly large in D20. A comparison is made of the excess Gibbs free energies of the solutes in H20 and DzO, and the relation between excess thermodynamic quantities and the corresponding thermodynamic quantities for the transfer of a solute from one solvent to the other is noted.

I. Introduction

Water is the solvent most commonly used for the study of solutions as is evidenced by the preponderance of publications on the properties of aqueous solutions. I t has become evident in recent years that the colligative and other properties of aqueous solutions are influenced by the structure of water. A complete bibliography on this subject is scarcely possible. Several excellent monographs are available (~-3) which cite many literature references. One method of detecting structural effects is to study the properties of solutions of the same solutes in other solvents. Aprotic solvents such as dimethylsulfoxide or ethylene carbonate and solvents having high dielectric constants such as the N-substituted amides have been used. Another alternative is to compare the solutes in H 2 0 and D20. The two liquids have very similar physical properties (4~ such as dielectric constant, molar volume, and dipole moment. The properties of their solutions are not identical, however, as has been shown (s-12) by the Gibbs free energy, enthalpy, and entropy of transfer of certain solutes from one solvent to the other or by the calculation of excess thermodynamic properties in the two solvents. The data for solutions in D 2 0 are not nearly so extensive as those in H 2 0 and these newly deter- mined osmotic and activity coefficients should supplement those determined by Kerwin (~3) thus permitting calculations to be made on a larger number of systems.

2. Experimental

Commercially available reagent grade salts were used except for sodium toluene sulfonate and sodium mesitylene sulfonate. Mesitylene sulfonic acid was prepared by sulfonation of the purified hydrocarbon. I t was then neutralized with sodium hydroxide to obtain the sodium salt. The salt at this time is slightly contaminated with sodium

838 O.D. BONNER

sulfate as is the commercially available sodium toluene sulfonate. Both salts were then

recrystallized at least three times from water and w a t e r + m e t h a n o l mixtures and dried in a vacuum oven to constant mass. The salts were checked for sulfate ion and for molecular weight by conversion to the acid in an ion exchange column and subsequent t i tration. The samples of urea and sym-dimethylurea also were repeatedly re-

crystall ized before use. Activity and osmotic coefficients of the solutes were determined by isopiestic

compar ison of their solutions with solutions of sodium chloride (or l i thium chloride in the case of concentrated solutions) according to the method of Rob inson and Sinclair. (~4) The values reported by Kerwin (la) for NaC1 and LiC1 coefficients were used in the calculations according to the equat ion:

m r e f

In ~, = In Tref+ln(mref]m)+ ~ {(mref[m)--1} d ln(Yre f mref). 0

TABLE 1. Osmotic coefficients ff of DaO in solutions containing "aquamolalities" m~ of various solutes at 25 °C

,n~ molkg- 1 NaC104 KNO3 AgNO8 KI CsI NaTolaNaMesbNa2SO4Dextrose Urea DMU o

0.1 0.930 0.906 0.903 0.932 0.916 0.929 0.924 0.793 1.002 0.998 0.990 0.2 0.920 0.873 0.870 0.922 0.895 0.918 0.912 0.753 1.006 0.997 0.981 0.3 0 .915 0.851 0.847 0.918 0.880 0.913 0.900 0.725 1.012 0.995 0.973 0.4 0 .912 0.833 0.827 0.917 0.870 0.908 0.883 0.705 1,017 0.994 0.965 0.5 0 .911 0.817 0.811 0.917 0.860 0.905 0.862 0.690 1.023 0.992 0.956 0.6 0.912 0.802 0.795 0.918 0.850 0.898 0.832 0.676 1.029 0.990 0.948 0.7 0 .914 0.790 0.779 0.919 0.840 0.890 0.804 0.662 1.035 0.988 0.940 0.8 0 .916 0.778 0.765 0.922 0.834 0.883 0.775 0.650 1.041 0.986 0.933 0.9 0 .918 0.766 0.750 0.924 0.829 0.876 0.748 0.641 1.047 0.985 0.926 1.0 0 .921 0.752 0.735 0.927 0.825 0.869 0.720 0.635 1.051 0.983 0.919 1.2 0.924 0.729 0.710 0.931 0.814 0.840 0.617 1.069 0.975 0.896 1.4 0 .929 0.709 0.690 0.933 0.804 0.826 0.615 t.078 0.972 0.886 1.6 0 .932 0.692 0.671 0.934 0.794 0.812 0.612 1.083 0.969 0.880 1.8 0 .937 0.675 0.652 0.936 0.780 0.795 0.614 1.089 0.965 0.875 2.0 0.940 0.660 0.633 0.937 0.760 0.628 1.106 0.956 0.863 2.5 0.954 0.627 0.589 0.944 0.741 0.656 1.123 0.944 0.853 3.0 0.970 0.592 0.557 0.958 1.139 0.936 0.847 3.5 0.986 0.528 0.971 1.156 0.927 0.842 4.0 1.002 0.502 0.990 1.174 0.920 0.837 4.5 1.016 0.478 1.005 0.912 0.832 5.0 1.031 0.461 0.906 0.830 5.5 1.050 0.446 0.900 0.830 6.0 0.431 0.894 0.830 6.5 0.416 0.891 0.829 7.0 0.402 0.888 0.829 7.5 0.394 0.885 0.827 8.0 0.389 0.882 0.825

0.880 0.823 0.878 0.821 0.876 0.820

a Sodium p-toluenesulfonate. b Sodium mesitylenesulfonate. c Sym-Dimethylurea.

DEUTERIUM OXIDE SOLUTIONS 839

In the overlapping molal i ty ranges where both NaC1 and LiC1 were used as references

there was always perfect agreement. Individual isopiestic rat ios differed f rom the

smoothed curve o f (mref/m) as a funct ion of molal i ty by no more than three parts in

one thousand and the average deviat ion was less than one par t per thousand.

The isotopic puri ty o f the D 2 0 solvent was 99.77 moles per cent.

3. Results

Osmotic and activity coefficients o f 11 solutes are given in tables 1 and 2. Composi t ions

are expressed as "aquamola l i t i e s " rn~ defined by

m~ = n ( s o l u t e ) / n ( D 2 0 ) M ( H 2 0 ),

where n denotes amount o f substance and M molar mass, so that these results are

TABLE 2. Activity coefficients 7 of solutes having "aquamolalities" ma in D20 as solvent at 25 °C

m~ 7 molkg-1NaCIO4 KNOa AgNOa KI CsI NaTolaNaMes~Na2SO4Dextrose Urea DMU c

0.1 0 .775 0.740 0.732 0.775 0.753 0.773 0.769 0 .452 1.002 0.992 0.974 0.2 0 .729 0.662 0.656 0.726 0.692 0.727 0.719 0.372 1.009 0.984 0.952 0.3 0 .701 0.612 0.605 0.702 0.651 0.697 0.682 0 .325 1.017 0.976 0.931 0.4 0 .683 0.574 0.566 0.686 0.622 0.677 0.651 0 .294 1.027 0.969 0.911 0.5 0 .669 0.542 0.535 0.672 0.597 0.660 0.618 0.271 1.038 0.957 0.890 0.6 0 .659 0.517 0.508 0.663 0 .575 0.645 0,585 0.252 1.049 0.954 0.871 0.7 0 .652 0.495 0.484 0,655 0.556 0.629 0.553 0 .236 1.061 0.945 0.852 0.8 0 ,645 0.475 0.463 0.651 0.541 0.616 0.522 0 .222 1.073 0,938 0.835 0.9 0 ,640 0.456 0,443 0.645 0.528 0.601 0 .494 0.212 1.085 0,931 0.818 1.0 0 ,637 0,439 0.424 0,642 0.516 0.590 0,467 0.202 1.094 0.962 0.803 1.2 0,630 0,409 0.393 0.636 0.494 0.567 0.420 0.•868 1.117 0.913 0.773 1.4 0 ,626 0,384 0,368 0,630 0 .475 0.545 0.175 1.137 0.901 0.746 1.6 0 .622 0.362 0.346 0.626 0.457 0.526 0.1662 1.158 0.890 0.720 1.8 0,619 0.342 0,327 0.622 0,440 0,508 0.1583 1.176 0.880 0,701 2.0 0 .618 0.326 0.308 0.618 0.490 0.1523 1.192 0.871 0.682 2.5 0 .619 0.292 0.271 0.613 0.450 0.1416 1.240 0.847 0.641 3.0 0 .625 0,263 0.243 0.616 0.421 0.1363 1.289 0.825 0.608 3.5 0.632 0.220 0.620 1.336 0.807 0.580 4.0 0.641 0.200 0.631 1.386 0.789 0,556 4.5 0.650 0.184 0.641 1.439 0.772 0,545 5.0 0.664 0.171 0.756 0.516 5.5 0.678 0.160 0.743 0.501 6.0 0.150 0.728 0.488 6.5 0.141 0.715 0.476 7.0 0.133 0.704 0.465 7.5 0.127 0.693 0.456 8.0 0.121 0.682 0.446

0.673 0.437 0.663 0.429 0.655 0.420 0.647 0.413

a Sodium p-toluenesulfonate. b Sodium mesitylenesulfonate. c Sym-Dimethylurea.

840 O. D. BONNER

1.15

0 1.10

1.05

~" 1.00

0.95

I [ I I 7 - - -

M N

I I _ _ _ _ 1 [ I

2 4 6 8 10 vm/tool kg-1

FIGURE 1. Activity coefficient ratios of solutes in H20 and D20 plotted against vm where v is the stoichiornetric number of ions for the salt or 1 for non-electrolytes and m is molality.

A, Dextrose; B, Sym-Dimethylurea; C, LiC1; D, Na toluenesulfonate; E, ZnSO~; F, Na mesitylene- sulfonate; G, NaBr; H, NaI; I, NaC104; J, Me4NCI; K, Urea; L, NaCI; M, KCI; N, CsC1; O, NaaSO4; P, KNOa; Q, AgNOa; R; KI; S, CsL

directly comparable with those of the same solutes in water. When these results are combined with those determined earlier by Kerwin, (~3) there are available Gibbs free energies for 19 solutes in D20. The activity coefficient ratios of these solutes in the two solvents at various molalities are represented graphically in figure 1. Data for the aqueous solutions are taken from various sources. °5-17)

4. Discussion

A comparison of the thermodynamic properties of solutions in the two solvents H 2 0 and DzO has usually been made in one of two ways. Partial molar Gibbs free energies of transfer of electrolytes from one solvent to the other have been calculated from e.m.f, measurements <1°' ~ 2) of electrolytes in the two solvents. These measurements are normally made in dilute solutions and the Gibbs free energy difference which is very small is taken to be the standard Gibbs free energy of transfer. Standard enthalpies of transfer are obtained from enthalpies of dilution in the two solvents to infinite dilution and the standard partial molar entropy of transfer is thence calculated. This type of calculation affords a comparison of the standard thermodynamic quantities in each solvent but gives no information about the composition dependence since Gibbs free energies of transfer are limited to a single composition. I t is also possible to calculate ~o called "excess" or relative non-ideal partial molar thermodynamic quantities.

DEUTERIUM OXIDE SOLUTIONS 841

Excess Gibbs free energy, defined as G2 E = v R T l n ?±, was first proposed by Scatchard and discussed (18) by Hildebrand and Scott. The excess partial molar enthalpy and entropy are defined as H2 E and by TSE2 = H ~ - G ~ , respectively. These quantities are available for many aqueous solutions as a function of composition and for a limited number of solutions °3~ in D 2 0 as a solvent. The results reported in this paper now permit a comparison of the excess Gibbs free energies of 19 solutes in H 2 0 and D 2 0 and figure 1 presents this comparison graphically. A scarcity of H2 E values as a function of composition prevents the comparison of excess entropies in solutes other than those already reported. (5' s) It has been noted earlier °a) that excess Gibbs free energies and Gibbs free energies of transfer may be related if the solubility of the solute in both solvents is known and if the solid phase in equilibrium with the saturated solution in each solvent is a pure salt and not a hydrate. I t is unfortunate that solubilities (4~ in D 2 0 are unavailable for any of the presently reported solutes.

An examination of figure 1 permits a few generalizations to be made about the behavior of electrolytes in HzO and D20. The activity coefficients of lithium and sodium salts are somewhat larger in D 2 0 while those of caesium and silver salts are smaller. The activity coefficient ratio of chlorides in D 2 0 and HzO is larger than that of iodides but it is possible for this ratio to be greater or less than unity depending upon the cation of the salt. The most surprising aspect of a comparison of solutes in the two solvents is the large increase in the activity coefficients of the non-electrolytes dextrose and sym-dimethylurea upon their transfer from H 2 0 to D20. An explanation of this effect is complicated by the fact that in either solvent the activity coefficients of dextrose are appreciably larger than unity while those of sym-dimethylurea are considerably less than unity. They appear thus to differ in their degree of solvation or in their effect upon the structure of H 2 0 or D20.

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