554
ZEEV Luz AND GADYAGIL
Water "0Nuclear Magnetic Resonance Shift in Aqueous Solutions
of 1:l Electrolytes
by Zeev Luz and Gad Yagi1 The Isotope Department and T h Radioisotope Training Center, The Weimnann Institute of Science, Rehovoth, Israel (Received September 19, 1866)
The frequency shift of the water 170signal in aqueous solutions containing varying amounts of diamagnetic electrolytes at 25" was measured. Both high- and low-field shifts were observed. It is shown that the molal shifts in dilute solutions can be divided into the contributions of the constituting ions. Most univalent cations give rise to about the same ionic shift, while the shifts due to anions differ considerably from each other. 170shift of dilute solutions of water in several organic solvents and of water at various temperatures is also reported. From these measurements it is estimated that the breaking of a hydrogen bond shifts the 170resonance 16 ppm toward high field. Analysis of the experimental data shows that the main cause of the shifts in electrolyte solut,ionsis not the modification of the structure of water, but rather direct interaction between the ions and adjacent water molecules.
Introduction The proton magnetic resonance line of water undergoes considerable shift when electrolytes are dissolved in water.' This effect offers a method to study ionsolvent and solvent-solvent interaction on the molecular level. Several investigations have been carried out on the proton shift in aqueous solutions,2+ and the dependence of the shift on concentration for a large number of simple diamagnetic 1:1 electrolytes was determined. It has been found that it is possible to express the observed shift as a sum of individual shifts due to the cations and anions. These ionic shifts were interpreted as being caused not only by specific ion-solvent interactions, but also by modifications of solvent-solvent interactions such as changes in the number of hydrogen bonds in bulk water caused by the presence of the salt. These findings are in line with a series of investigations dealing with the structure of aqueous electrolyte solutions which emphasize the role of hydrogen bond breaking in explaining the properties of these sol~tions.7-~On the other hand, a t least some work on the thermodynamic properties of electrolyte solutions1° has been done which does not invoke changes in the bulk water to explain the properties of these solutions. The Journal of Physical Chemistry
I n this paper we report measurements on the 1 7 0 shift of water in aqueous electrolyte solutions. The question posed was to what extent do the 1 7 0 shifts obey a pattern similar to that of the proton shifts, and whether additional information on ion-solvent ' shifts were found interactions can be gained. The 10 to be of the same order or magnitude (in ppm) as the proton shifts and could again be broken down into contributions of individual ions. Some notable differences were observed, however. I n particular, the I7O ionic shift of the halide and halate ions follows an inverse order with respect to their ionic size to that found with protons, an order which is also inverse to that expected if hydrogen bond breaking between solvent (1) J. N. Shoolery and B. J. Alder, J. Chem. Phys., 23, 805 (1955). (2) H. S. Gutowsky and A. Saika, ibid., 21, 1688 (1953). (3) H. G. Hertz and W. Spalthoff, Z . Elektrochem.. 63, 1096 (1959). (4) M. S. Bergquist and E. Forslind, Acta Chem. Scand., 16, 2069 (1962). ( 5 ) J. C. Hindman, J. C h m . Phys., 36, 1000 (1962). (6) B. P. Fabricand and S. Goldberg, ibid., 34, 1624 (1961). (7) H. S. Frank and W.-Y. Wen, Discussions Faraday SOC.,24, 133 (1957). (8) M. Kaminsky, ibid., 24, 171 (1957). (9) G. R. Choppin and K. Buijs, J. Chem. Phys., 39, 2042 (1963). (10) E. Glueckhauf, Trans. Faraday SOC.,51, 1235 (1955).
I7O
SHIFTSIN 1: 1 ELECTROLYTES
555
molecules was the dominant cause of the shift. AIthough we were not able to give a quantitative interpretation of the results, it, is shown by comparing the shifts with similar shifts in auxiliary systems that direct ion-solvent interaction is the main cause of the 170 shifts in aqueous solutions.
by measuring the position of the water resonance at the desired temperature between two reference water samples a t room temperature. The sample holder used in these measurements was a dewar-like test tube, to keep the temperature constant during the measurements.
Experimental Section
Results The shift of the 170resonance line of water in solutions of 30 1:l electrolytes as function of concentration is shown in Figures 1-5. The ordinate represents the observed shift of the I7O line as defined in the Experimental part, and the abscissa gives the concentration in moles of electrolyte/1000 g of solvent water. Some of the electrolytes gave straight-line plots, while others deviated from linearity a t high concentrations. Within our experimental accuracy all the lines passed smoothly through the origin; it should, however, be noted that the lowest concentrations of electrolyte that gave measkrable shifts were around 0.5-1 m. A change in the slope of the curve a t more
Preparation of Solutions. The solutions were prepared by weighing commercial analytical grade electrolytes into a weighed amount of water. Water was obtained from the Institute's plant and contained 0.45 atom % I7O. The water also contained about 20 atom % deuterium. The concentrations of the solutions were determined as follows: potassium fluoride by thorium nitrate titration, salts of the other halides by argentometric titration, and perchlorates and nitrates by gravimetric determination of the cation. The presence of metal ion impurities was checked by spectrographic analysis." I n a few cases impurities of paramagnetic ions were detected, and indeed appreciable broadening of the 170line was observed; in purified substances the excess broadening disappeared. For a number of salts the shift was measured both in neutral and in acidified solutions. The amount by which the lines were shifted was not affected by the addition of acid, but very often the lines were broadened in neutral solution due to slow proton exchange.12 The reported measurements were all performed in 0.01 M HCI, except for KF, KN02, and NaNOz, which are salts of weak acids. Nuclear magnetic resonance measurements were performed on a Varian DP-60 spectrometer, operating at 8.13 Ilk. Water of the same isotopic content and acidity as those of the samples was used as a reference. The shift was measured as follows: first the reference sample was run and the position of the line was recorded in terms of field dial units; next, the sample was run, and its position was recorded; finally, the reference was rerun. The field dial units were calibrated by the side-band technique, and the shift was calculated as the difference between the sample and the mean of the two reference readings. The accuracy of the shifts determination is estimated at 0.5 ppm which is about 10% of the water line width. We denote a shift as positive when the observed resonance of the electrolyte was at a higher field than the reference water at fixed radiofrequency. To eliminate contributions to the shift from differences in bulk susceptibilities, spherical sample holders were employed. All of the measurements were conducted a t room temperature, which was 25 A 2". The temperature dependence of the 170shift was determined
8.0
I
I
I
I
I
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Concentration of electrolyte in moles per IOOOgr wter
Figure 1. I7O shift, 6,of water in electrolyte solutions, relative to pure water, as a function of the electrolyte concentration, for nitrates and some halides. Shifts are positive when the resonance of the electrolyte solution is a t a higher field than that of the reference water.
(11) We are indebted to Mrs. I. Schoenfeld and LMrs.S. Held of the Spectrochemical Laboratory of the IAEC, Soreq. for carrying out these analyses. (12) 5. Meiboom, J. Chem. Phys., 34, 375 (1961).
V o l u m e 70,N u m b e r 8 February 1966
ZEEVLuz AND GADYAGIL
556
I
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t\Me,NCI
-“F
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> O . 80 0.58
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Dielectric constant Infrared Model to fit thermodynamic data Infrared up to T, Raman
Luck! Walraf en0
a The numbers give the fraction of hydrogen bonds formed. The fraction is unity when all water molecules participate in four hydrogen bonds, so that there is an average of two bonds per molecule. The fraction of hydrogen bonds broken upon heating from 4 to 72’. c G. H. Haggis, J. B. Hasted, and T. J. Buchanan, J . Chem. Phys., 20, 1452 (1952). K. Buijs and G. R. Choppin, ibid., 39, 2035 (1963). e G. NBmethy and H. A. Scheraga, ibid., 36,3382 (1962). f Figures 50 and 58 in W. Luck, Fortsch. Chem. Forsch., G. E. Walrafen, J . Chem. Phys., 40, 3249 (1964). 4, 43 (1965).
tional to the number of ions present. This already indicates that direct ion-water interaction makes an important contribution to the observed shift. Referring to Figure 6, one observes that most cations are grouped together a t approximately 0.8 ppm/m. The same, though less pronounced, is observed in the case of the proton shifts. It seems that the effect of the univalent cations on the 170shift is small compared to the effect of the anions. This is surprising because oxygen is the part of the water molecule that is expected to be closest to the cation and thus be subjected to the strongest influence of the ion, while the anions are more likely to approach the hydrogen end of the water molecule. It can only be added that a similar lack of effect of cations has been observed before, e.g., on the Raman spectra of electrolytes in water,19 on the infrared of methanol OH stretching frequency discussed below,20 and on the solubility of some nonelectrolytes in water.21 On the other hand, preliminary measurements on salts of the alkaline earth metals indicate a somewhat larger spread of the molal shifts. I n addition, the shift of the solvation shell of the AI3+ion, which can be observed separately from the solvent peak, was foundz2to be -11 ppm, corresponding to a molal shift of -1.1 ppm/m. (Susceptibility correction will lead to an even larger figure.) With anions, a different picture is observed. Roughly, in the two series halides and halates, the ionic shift becomes more negative with increasing ionic size. An attempt to correlate the ionic shift with the ionic volume is shown in Figure 7. I n this figure the ionic molal shift (based on &E,+ = 0) is plotted against the cube of the ionic crystal radii. The values for the ionic radii were taken from ref 23; those for the oxyanions were calculated from the covalent radii of the constituting atoms.23 It is seen that within each The Journal of Phgsical Chemistry
series the correlation is quite linear, but the shift is less negative for the oxyanions. The ionic shift thus cannot be explained solely by the size of the anion. It appears that the oxygen atoms of the oxyanion have a specific effect on the water molecules causing an upfield shift. The points for nitrate, nitrite, and perchlorate in Figure 7 support this idea. Nitrate, having three oxygen atoms, falls on the line of the halates, while perchlorate and nitrate are displaced above and below the halate line in accordance with the number of oxygens they contain. There are a t least two possible explanations to the effect of ionic volume (size) on the anionic shifts. The first, suggested, e.g., by Bergquist and F ~ r s l i n d ,is~ that the larger the anion, the more efficiently it breaks the water structure (“the steric effect”). However, as discussed above, this effect should shift the 170resonance line in a direction opposite to the observed one, and can therefore be ruled out. AlternativeIy, what might be involved is not the ionic volume itself but a related property such as the ionic polarizability (which indeed gives similar patterns to Figure 7). The larger the polarizability of an anion in a series, the more likely are interactions between it and adjacent water molecules which underlie the chemical shift. However, the theory of the 1 7 0 shift is not sufficiently advanced to permit further examination of this possibility. Recent infrared and ultraviolet measurements by several a ~ t h o r sbring ~ ~ , evidence ~~ for the existence of solute-solvent interaction between halide ions and (19) G. E. Walrafen, J. Chem. Phgs., 36, 1035 (1963). (20) A. Allerhand and P. von R. Schleyer, J . Am. Chem. soc., 85, 1233 (1963). (21) P. R. Robinson and W. P. Jencks, (bid., 87, 2470 (1965). (22) R. E. Connick and D. N. Fiat, J . Chem. Phys., 39, 1349 (1963). (23) M. W. Sidgwiok, “The Chemical Elements and Their Compounds,” Vol. I, Oxford University Press, London, 1950, Table V.
170
56 1
SHIFTSIN 1 : 1 ELECTROLYTES
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Figure 7. Ionic molal shift of anions as function of the cube of their crystal radii.
methanol. Thus, Allerhand and SchleyerZ0 measured the shift of the OH stretching frequency, (around 3600 cm-l) of methanol highly diluted in solvents such as methylene chloride. Upon the introduction of various tetraalkylammonium halides, shifts of this line up to 340 cm-' were observed. The infrared shift was almost independent of the cation, but strongly dependent on the anion, as in the 170shift, leading to the c o n c l u ~ i o nthat ~ ~ ~anion-methanol ~~ interactions are responsible for the shift. Since methanol is very similar to water, this system is a suitable one to examine whether direct interaction of anion with hydroxyl group can lead to shifts similar to those observed in the aqueous electrolyte solutions. For that purpose a series of solutions of methanol (0.2 M ) and each of the four tetrabutylammonium halides (0.6 M ) was
prepared in methylene chloride. The "0 resonance of methanol in these solutions was measured relative to a methylene chloride solution containing only methanol at the same concentration.26 The results are: tetrabutylammonium fluoride, -0.9 ppm; tetrabutylammonium chloride, -7.9 ppm; tetrabutylammonium bromide, - 5.7 ppm; tetrabutylammonium iodide, -8.0 ppm. These figures clearly demonstrate that anion -OH interaction causes a considerable lowfield shift of the 170resonance line. Moreover, except for the chloride, the shift increases with ionic size, iodide causing the largest shift. These results cannot be compared quantitatively with the ionic molal shift of the halides because it is not known to what extent association between methanol and the anions is complete and because the solvation number of the halide anions in water is not known; also, the effect of an anion on methanol probably differs somewhat from the effect on water. More recently, it has been shown that similar infrared shifts occur when HzO interacts with anions in CC14.27 The 1 7 0 shifts described above are, however, large enough to support the conclusion that direct anionwater interaction is the dominant cause for the I7O shift in aqueous solutions. This conclusion suggests that more attention be paid in future examination of electrolyte solutions to the interactions of anions with their neighboring solvent molecules. (24) (a) H. Lund, Acta Chem. Scand., 12, 298 1958); (b) J. Buffalini and K. H. Stern, J . Am. Chem. SOC.,83,4362 (1961); (e, J. B. Hyne, ibid., 85, 304 (1963). (25) M. S. Blandamer, T. E. Gough, and M. C. R. Symons, Trans. Faraday Soc., 60, 423 (1964). (26) Tetraalkylammonium salts were recrystallized from benzene or
chloroform. Tetrabutylammonium fluoride was prepared as suggested by Allerhand and Schleyer.20 (27) S . C. Mohr, W. D. Wilk, and G. M. Barrow, J . Am. Chem. SOC., 8 7 , 3048 (1965).
Volume 70,Number 9 February 1966
