Hydrogen Bonding of Water in
Organic Solvents
very small (perhaps 2 or 3%) from the relative areas of the three bands The band widths at half-intensity are larger than those of water in dilute solutions as one would expect since there are more possible configurations and hydrogenbond angles in pure water. The large area under the third band indicates that although most of the water is doubly bonded, there Rre a large fraction of the molecules in which the proton-oxygen bonds are of unequal strength. This suggests that the most appropriate model for water might be that proposed by Poplel* some years ago in which the majority of the hydrogen bonds are regarded as distorted rather than broken. This would result in a rather broad spectrum of enorgies for the bonds. Finally data for water in one alcohol are presented (Figure 3). The VOH band of the alcohol lies above 2000 nm and is much less intense than the water hand so that there is no interference. Ail threla water bands are observed for dilute solutions of water in 1-pentanol. The free UOH water band could result from water molecules which are adjacent to the hydrocarbon portion of the molecule but more probably results from the hydrogen bonding of the alcohol proton and the water oxygtm. The large area of band 3 relative to that of band 2 indicates that many of the water molecules are not symmetrically bonded. This could result from the steric effects of a fivta carbon alcohol. The band half-widths and band positions 2re similar to those of pure water.
1727
+
In conclusion, the u2 u3 water band appears to be better suited for the study of water structure than the fundamental modes for at least three reasons. (1)The data are easier to interpret as there is only one band when water is unbonded or symmetrically bonded rather than the three bands ( V I , v3, and 2 ~ which ) interfere with one another. ( 2 ) Interference of other 0-H and N-H bands is avoided. (3) Semiquantitative measurements may be made as shown by the agreement in the values of the integrated intensities which were obtained.
References and Notes R. M. Badger and S. H. Bauer, J. Chem. Phys., 4, 469 (1936). G. Herzberg and H. Verieger, Phys L,37, 444 (1936). A. R. H. Cole and A. J. Mitchell, Aust. J. Chem., 18, 46 (1965). G. Habermahl,,Angew. Chem., 76, 271 (1964). G. Herzberg, Molecular Spectra and Molecular Structure II. Infrared and Raman Spectra of Polyatomic Molecules,” Van Nostrand, Princeton, N. J., 1959, p 281. E. Greinacher, W. Luttke, and R. Mecke, Z. Eleektrochem., 59, 23 (1955). D. P. Stevenson, J. Phys. Chem., 69, 2145 (1965). W. C. McCabe, S. Subramanian, and H. F. Flsher, J. Phys. Chem., 74, 4360 (1970). A. Burneau and J. Corset, J. Phys. Chem., 76, 449 (1972). L. B. Magnusson, J. Phys. Chem., 74, 4221 (1970). K. Buijs and G.R. Choppin, J. Chem. Phys., 39, 2035 (1963). S. C. Mohr, W. D. Wilk, and G. M. Barrow, J. ,4mer. Chem. SOC.,07, 3048 (1965). M. R . Thomas, H. A. Scheraga, and E. E. Schrier, J, Phys. Chem., 89, 3722 (1965). J. A. Pople, froc. Roy. SOC., Ser. A, 205, 163 (1951).
Hydrogen Bonding of Water in Organic Solvents. II. The Change of Water Structure with Composition
. Banner* and Y. S. Choi Department of Chemistry, University of South Carolina, Columbia, South Carolina 29208 (Received February 1 I , 11174)
+
The u2 ug combination band of water has been recorded for tetrahydrofuran-water, acetonitrile-water, arid dimethyl sulfoxide-water solutions over the entire range of water concentrations. Absorbance of the weak bands of the organic component of the solution has been compensated. I t was found that in all solutions the fraction of nonbonded or very weakly bonded OH groups was greater in solutions dilute in water and in pure water than in solutions of intermediate composition. It was also observed that a large fraction of the water molecules bonded with the two protons forming bonds of near equal strength in dilute water solutions but that the bonding became more “unsymmetric” as the mole fraction of water in the solution was increased. Evidence was obtained for the phenomenon of “hydrophobic bonding” in the case of tetrahydrofuran-water solutions. The computer resolution of the bands in the organic solvent-water solutions of various concentrations was substantiated by measurement of the derivative d(absorbance)/d(water concentration) for several solutions over the wavelength region 1850-2050 nm. It was found for these solutions, just as for the dilute solutions reported previously, that the total water absorbance is independent of the cosolvent and dependent only on the amount of water which is present.
Introduction In a previous paper1 spectra were reported for small amounts of water In 13 aprotic solvents and one alcohol. The u2 “3 Combination band of water involving the bending mode wa3 chosen so as to avoid interference from OH
+
and NH groups in alcohols, glycols, and amides since it was hoped to extend these studies to systems of biological interest. It is the purpose of this paper to report spectra for some of the ogranic solvent-water systems which have been recorded over the entire concentration range from very diThe Journal of Physical Chemistry. Vol. 78. N o . 17. 1974
0. D. Bonner and Y.S . Choi
172
TABLE I: Computer-Resolved Band Positions, Half-Widths, and Areas for the Tetrahydrofuram-Water Solutions
Mole fraction of water Band 1 Position, nm Half-width, nm Area, nm absorbance X Band 2 Position, nm Half-width, nm Area, nnn absorbance X Band 3 Position nm Half-widlth, nm Area, iini absorbance x Total area. nm absorbance X
v2
f- v 3 W a t e r Band in
0.08
0.21
0.32
0.66
0.82
0.91
0.98
1 .oo
1891
1891
1892
1894
1897
1897
1900
1901
21.7
24.5
24.3
22.9
23.6
22.2
21.6
22.4
4.0
3.3
2.5
0.8
0.8
0.8
0.6
0.8
1920 37.5
1921 42.8
1923 46.6
1924 55.5
1925 56.7
1924 56.7
1923 57.7
1923 57.9
22.3
21 .o
20.1
18.7
17.7
17.5
17.2
17.2
1958 36.6
1963 48.3
1966 53.6
1971 65.8
1975 70.8
1975 69.1
1977 70.0
1977 70.1
10J
6.8
10.0
11.7
12.9
14.2
14.9
14.5
14.6
LOd
33.1
34.3
34.3
32.4
32.7
33.2
32.3
32.6
103
lo3
lute water soiutionrj to pure water and to note the water structure in these solutions.
I
I
/
~ x ~ e r iSection ~ e ~ t ~ ~ ~ Epectroph0torrxeti.i~ or Reagent Grade solvents were used in all instances without further purification. For solutions containing only a small amount of water or for solutions in which the ,water concentration was varied only a small amount a Cary Model 14 spectrophotometer equipped with a 0-0.1- and 0.1-0.2 range slidewire was used. The same insilrarment equipped with the normal 0-1.0 and 1.0-2.0 range slidewire was used for more concentrated water solutions. In order to be consistent with the data reported in the previous paper, all absorbances were normalized t o 1 fil cf E l 2 0 per 5 ml of solution and a cell path length of 0.5 rim. The concentrated solution data may thus be directly cumpared with that for the very dilute solueAOnS. The two slide wires and the cells of varying path length were, oC course, calibrated for internal consistency.
It is not possible because of solubility limitations to study aqueous solutions of all of the nonelectrolytes reported in the previous paper1 over the entire concentration range. Three systems with different organic functional groups were, therefoye, chosen and examined in detail. ( A ) I'etrahydro~ifilran-Water. The tetrahydrofuranwater system is interesting since the very dilute solution of water in T H F exhibited the largest fraction of nonbonded water OH groups of all of the systems reported. Solutions containing various mole fractions of water were prepared and their spectra were recorded in the 1850-2050-nm region with the slight s.bsorbance due to T H F being compensated in the reference beam. The absorbances were then normalized to 1 fil of H20 per 5 ml of solution and subjected to computer resolution in the same fashion as the dilute solutions. The resu1i:ant data are reported in Table 1. It may be noted 1 hat the position of band L shifts in a regular fashion with iiicreasi ng water concentration from 1891 to 1901 nm. The fraction of nonbonded OH groups, however, appears to dewease to a minimum value at a n intermediate concentration and then increase slightly in pure water. The area of band 2 does not change drastically with composition but the positior? of the band appears to reach a maximum wavelength a t i3n intermediate concentration. If this The Journal of F'hyzical Chemisfry. Vol. 78. No. 17. 1974
1850 1879 1908 1937 19% 19951850 1879 1908 :937 1966 1995
Wavelength (nm)
Figure 1. Spectra of the derivative d(absorbance)/d(H20concentration) for tetrahydrofuran-water solutions of various compositions: (A) 0.21 mole fraction water, relative areas band Piband 3 = 1.8; (6) 0.66 mole fraction water, relative areas band 2/band 3 = 0.9; (C) 0.91 mole fraction water; (D) 0.99 mole fraction water.
is related to the strength of the bonding in the symmetrical complex as was inferred from the earlier work, it indicates that the bonding is stronger in solutions of intermediate concentration than in pure water or in very dilute solutions. This indication will be confirmed by further data. Band 3, which like band 1 is due to the unsymmetrically bonded complex, increases in area and in wavelength as the fraction of water increases. The constancy of the total area of the water band over the entire range of composition is reassuring. The average residual between the calculated sum of the absorbances of the three bands and that observed experimentally was again less than the experimental error in reading the spectra. The data in Table I give the overall picture for all of the water molecules in a solution of a given water concentration. Because of the unexpected minimum in the fraction of nonbonded OH groups and the maximum in the wavelength of the symmetrically bonded water complex, it was desired to further substantiate the data in Table I. Consequently, a series of experiments was designed in which values of the derivative d(absorbance)/d(concentration) could be measured at various water concentrations. This was done by preparing a solution of a given water concentration and dividing it into two parts. One part was placed
Hydrogen Bonding of Water in Organic Solvents
1729
TABLE 11: Computer-Resolved Band Positions, Half-Widths, and Areas for the Acetonitrile-Water Solutions Mole Eraction of water Band 1 Position, nm Half-width, nm Area, nm absorbance X 10; Band 2 Position, nm Half-width, nrn Area, nm absorbance X 1 O J Band 3 Position, nm Half-width, nrn Area, nm absorbance x 10Total area nm absorbance X lo3
0.56
0.74
0.87
0.96
1.oo
1888 16 “ 5
1889 20.5
1894 21.1
1898 21.6
1900 20.7
1900 20.9
1900 22.1
1901 22.4
0.9
0.9
0.9
0.7
0.4
0.6
0 .7
0.8
1905 27.6
1908 31.9
1909 36.6
1914 43.3
1916 47 . O
1919 50.0
1921 54.2
1923 57.9
22.6
21.4
20.5
18.6
17 .O
17.1
16.5
17.2
1937 44.9
1944 49.9
1947 53.3
1957 55.7
1961 63.6
1967 63.4
1.971 64.5
1977 70.1
9.3
10.7
11.5
12.3
14.2
14.5
14.2
14.6
32.8
33.0
32.9
31.6
31.6
32.2
31.4
32.6
-_A
e2
Water B a n d in
0.24
P
t
vg
0.15
/
oe
-+-
0.05
AI
8
~2
0.4 06 0.8 Mole Fraction (Water)
ai
x
i
1.0
Figure 2. Viscositicis of tetrahydrofuran-water solutions at 25’.
in the reference beam of the spectrophotometer and a very small amount of additiional water (ca. 1%)was added to the other sample which w,as placed in the sample beam of the spectrophotometer. The observed spectrum was that resulting from the increment of water which was added to the solution. This procedure is similar to that employed by Magnusson2 to study the dimerization of water in CC14 and by Bellamy and Pace3 to investigate the polymerization of alcohols. It difftw, however, in that the concentrations of water or alcohol in thleir compensation technique differed by factors of 5-50 whide we changed the water concentration only slightly. These spectral data are thus somewhat analogous to partial molal quantities which may be calculated for a solution a t a given (constant) composition. The derivatives for some of‘the solutions are shown in Figure 1. When a small increment of water is added to solutions containing mole fractions of water in the range 0.15-0.82 the appearance of the derivative, d(absorbance)/d(water concentration), is similar to that of the spectrum of pure water except that the wavelength of maximum absorbance gradually shifts from 1925 to 1948 nm. The wavelength of maximum absorbance of pure water is 1928 nm. A computer resolution of the derivatives indicated that the additional increment of water is a t first primarily symmetrically bonded with the bonds being rather strong as is indicated by the increasing wavelength o f both the observed band maximum and also the maximum1 of resolved band 2 due to the symmetrically bonded water. It is necessary that most of the new hydrogen bonds be water-water rather than the weak-
er water-THF bonds since a t mole fraction 0.33 the availability of T H F oxygen atoms for bonding is exhausted. The strength of the bonds appears to be experimental confirmation of the “cooperative phenomenon” praposed by Frank.* It is also probable that the structure of the water is enhanced by the T H F methylene groups. This phenomenon has been referred to as “hydrophobic bonding” by some persons. As the water concentration becomes greater, it becomes increasingly difficult for the water to hydrogen bond in a symmetrical fashion, primarily due to steric effects. In solutions having mole fractions of water greater than 0.9, the added increment of water exhibits completely unsymmetrical hydrogen bonding and this contributes to the larger fraction of nonbonded OH groups present in pure water than in a solution containing 0.8-0.9 mole fraction water. The derivatives thus tend to validate the computer resolutions for the spectra of the various solutions. The spectral data for the THF-water solutions also correlate with certain physical properties of the solutions. Both the adiabatic c~mpressiblity~ and the water activity6 curves show minima between water mole fractions of 0.9 and 1.0. the viscosity, which is probably a better indicator of structure, shows a maximum a t a mole fraction of water of approximately 0.9 indicating that solutions of intermediate concentrations indeed have stronger forces of attraction than are present in the pure liquids. This behavior is similar to that observed in solution^^-^ of water with dioxane and alcohols. ( B )Acetonitrile-Water. This system seemed to be of interest since the water-solvent hydrogen blonds appeared to be among the weaker of those reported1 for dilute water solutions as indicated by the wavelengths of both the symmetrically and unsymmetrically bonded water. The spectral data for solutions of various water Concentrations are presented in Table 11. It may be noted that the positions of the band maxima for the three bands and also the band areas vary in a more regular fashion with water concentration than was the case with the THF-H2O system except that there again appears to be slightly fewer nonbonded OH groups in solutions of intermediate composition than in either of the pure liquids. There is, however, no solution of intermediate composition in which the wavelength of the symmetrically bonded water band is greater than that for pure water. This confirms the suspicion that the occurrence of this phenomenon in THF-water systems may he due to the enhancement of the water structure (and strength of The Journal of Physical Chemistry, Vol. 78. No. 17, 1974
0 . D. Bonner and Y . S. Choi
1730
TABLE 111: Computer-Resolved Band Positions, Half-Widths, and Areas for the v p Dimethyl Sulfoxide-Water Solutions
Mole fraction of water Band 1 Position, nin Half-width, nm Area, nm absorbance X 103 Band 2 Position, nin Half-width! nm Area, nm absorbance X :LO3
+ vj Water Band in
0.07
0.19
0.31
0.63
0.80
0.90
0.97
1.oo
1896 16.6
1897 17.5
1898 18.2
1900 18.5
1901 19.2
1901 19.7
1901
21.2
1901 22.4
1.4
1.1
1.o
0.7
0.6
0.7
0.8
0.8
1941 47.3
1940 51 .O
1939 51.3
1934 57.3
1930 57.7
1927 57.7
1924 57.7
1923 67.9
28.6
27.4
23.9
20.0
17.8
17 .O
16.3
17.2
1983 35.9
1982 42.3
1980 54.1
1978 62.4
1978 66.0
1978 67.7
1977 69.4
1977 70.1
:1.03
3.9
5.4
8.8
11.7
12.5
12.9
13.8
14.6
.Lo"
33.9
33.9
33.7
32.4
30.9
30.6
30.9
32.6
Band 3
Position, nin Half-wid t h , nm Area, nni absorbance x Total area, nin absorhance x
Wavelength ( nrn)
Wavelength ( n m )
Figure 3. Spectra of the derivative d(absorbance)/d(H20 concentration) for acetonitrile-water solutions of various compositions: (A) 0.10 mole fraction waier, relative areas band P/band 3 = 1.8; (B) 0.24 mole fraction waier, relative areas band Plband 3 = 1.0; (C) 0.74 mole fraction water; (D) 0.96 mole fraction water.
Figure 4. Spectra of the derivative d(absorbance)/d(H,O concentration) for dimethyl sulfoxide-water solutions of various compositions: (A) 0.19 mole fraction water, relative areas band 2/band 3 = 3.2; (B) 0.31 mole fraction water, relative areas band 2/band 3 = 1.3; (C) 0.63 mole fraction water, relative areas band 2/band 3 = 1.0; (D) 0.97 mole fraction water.
the hydrogen bonds) by the methylene groups of THF. In this connection it may be observed that at comparable water mole fractions the wavelength of the symmetrically bonded water band is always greater in THF-H20 solutions than in acetonitrile-water solutions. The areas of bands 2 anti 3 indicate that the hydrogen bonding becomes less symmetrical as the concentration of water is increased. The spectra of the derivatives for this system (Figure 3) again confirm the data given by the computer resolution. The fairly shilrp band at 1909 nm for a mole fraction of water of 0.1 indicates that the added increment of water is still bonding ]primarily in a symmetrical fashion but with stronger bonds. At a mole fraction of 0.24 the band is much broader anti i i can he resolved into both a symmetrical and an unsymmetrical bonded portion. The double peaks occurring in the derivative for the 0.74 mole fraction solution confirm that the added water is bonding unsymmetrically. For a mole fraction of water of 0.96 the derivative exhibits a maximum at. 1892 nm and a minimum a t 1918 nm. This substantiates the increase in free OH group concentration and decreaee in the fraction of symmetrically bonded water as one approaches pure water. ( C ) DMSO--Wnter This system was chosen for study since the DMSO-water hydrogen bond strength appears to
be greater than water-water hydrogen hond strength from the dilute solution data. This system also affords a direct comparison of the data obtained from near-infrared spectra with those obtained from Raman spectralo of the fundamental vibrational region. The DMSO-water spectral data, which are summarized in Table 111, are about those which would be expected in the light of the experience gained with the previous systems. It appears that in all aprotic solvent-water systems that the fraction of nonbonded or very weakly bonded water protons is greater in pure water or in very dilute solutions than in those of intermediate composition. Likewise, the fraction of symmetrically bonded water decreases as the mole fraction of water in any of the solutions increases. The band maxima for the three bands again changes wavelength in a regular manner with water composition although there is a shift of only 2 nm in the position of the symmetrically bonded water band between mole fractions of water of 0.07 and 0.31. This may be taken as further evidence that the added water continues to bond primarily to DMSO molecules rather than to other water molecules. The spectra of the derivatives for this system again confirm the computer resolutions of Table 111. Computer resolution of the derivatives which have been made for solutions in which water is the minor
The JOiJrnal of Ph,isicai Cliemisiry. Voi. 78. No. 17. 1974
Coordination and
Ionic Solvation
component of the three cosolvent systems reveal that, for solutions of comparable mole fractions, the relative areas of band I/band 3 are in the order DMSO-H2O > THF-H20 > CH&!N-H20. This is the same order as the relative strengths of thc aprotic solvent-water hydrogen bond. Although this 8tud.y yields data which are more quantitative concerning the state of water in various DMSO-water solutions it is interesting that the model proposed by Scherer, et nl.,lQapplies rather well. Both studies conclude that the water exists almost entirely dihydrogen bonded with a portion of it having one stronger and one weaker bond and another portion having hydrogen bonds of near equal strength A very small fraction of the protons are very weakly bonded.
1731
of the solvent-water hydrogen bond relative to that of the water-water bond from the wavelength at which the band occurs. The technique of observing the derivatrve, d(absorbance)/ d(water concentration), should be extremely valuable in that i t enables one to observe changes in water-water and water-cosolvent interactions as a function of solution concentration. In cases where water miscibility is not achieved over the entire range of concentrations such as for biological membrane systems and solutions of electrolytes one may observe at what point the additional water molecules begin to bond to other water molecules instead of to the co. solvent molecules. A study of biological systems by successive dehydration might furnish data for distinguishing between so called “bound” and “free” water in these systems.
Conclusions
+-
The v 2 1’3 combination water band is useful for studying the nli ture of organic solvent-water solutions over the entire range of compositions from pure water to solutions containing very small quantities of water. There is no interference of neighboring bands as is the case in the v1, v3, 2112 region and no need to considerlo the effects of Fermi resonance. The water band of any solution can be resolved into its Gaussian components and from these data one can estimate (1)t,he relative amounts of symmetrically and unsymmetrically bonded water molecules, (2) the fraction of protons which are very weakly bonded, and (3) the strength
References and Notes (1) 0. D.Bonnerand Y. S . Choi, J. Phys. Chem., 78, 1723(1974). (2) L. B. Magnusson, J. Phys. Chem., 74, 4221 (1970). (3) L. J. Bellamy and R. J. Pace, Spectrochim. Acta, 22, 525 (1966). (4) H. S. Frank and W. Y. Wen, Discuss. faraday Soc., 24, 133 (1957). (5) E. K. Baumgartner and G. Atkinson, J. Phys. Chem., 75, 2336 (1971). (6) K. L. Pinder, J. Chem. Eng. Data, 18, 275 (1973). (7) R. L. Kay and T. L. Broadwater, Necfrochem. Acta, 16, 667 (1971). (8) T. L. Broadwater and R. L. Kay, J. Phys. Chem., 74, 3802 (1970). (9) R. L. Kay, C. Cunningham, and D. F. Evans, “Hydrogen Bonded Solvent Systems. A. K. Covington and P. Jones, Ed., Taylor and Francis, London, 1968. (IO) J. R. Scherer, M. K. Go, and S. Kint, J. Phys. Chem., 77, 2108 (1973).
4;
Coordinatiam and Ionic Solvation
. G . Cox, A. d. Parker, and MI. E. Waghorne”’ The Research School of Chemistry. Australian National Unwersity. Canberra, A.C. T. Australia
(Received November 19. 7973:
Revised Manuscript Received April 15, 1974)
Free energies of transfer of a variety of ions from pure solvents to binary solvent mixtures have been estimated by the application of an extrathermodynamic assumption to free-energy data from emf and solubility measurements. In addition, the composition of the inner-coordination spheres of several cations in solvent mixtures have been estimated from pmr spectroscopic results. These results have been compared to those calculated from the measured equilibrium constants for the coordination of the cation by the solvents, using a simple coordination model of ionic solvation in which variations in solvation energy are considered to arise solely from differences in the energy of complex formation between the ions and molecules of the component solvents, and in which interactions outside the first coordination sphere are considered to be constant. For ions in a variety of nonaqueous solvent mixtures, the agreement between measured free energies and those predicted by the model is excellent. In aqueous mixtures, deviations from the behavior predicted by the simple model are attributed to the effects of nonideality of the solvent components and interactions between the complexed ion and solvent molecules outside the first coordination sphere.
Introduction A number of‘ recent studies2 have pointed out the inability of the 13orn equation to account satisfactorily for the changes in solvation energy of electrolytes on transfer from one medium to another. Indeed in many cases the sign of the free.energ!y change is incorrectly predicted. For
example, the transfer of silver halides from methanol ( E 32.6) to propylene carbonate (c = 65) is energetically unfavorable in all cases,3 while the opposite is predicted by the Born equation. The unfavorable change in the free energy of anions from protic to dipolar aprotic media has been explained in The Journal of Physical Chemistry. Vol. 78. No. 17. 1974
