NATUREOF HzO SPECIESIN Hz0-NH3 SOLUTIONS
863
The Nature of Water Species in Water-Ammonia Solutions as Inferred from Proton and Oxygen-17 Nuclear Magnetic Resonance Observations by Mohammed Alei, Jr., and Alan E. Florin University of California, Los Alamos Scientific Laboratory, Los Atamos, New Mexico
87544
(Received July 2 9 , 1 9 6 8 )
H20 proton shifts as a function of H20 concentration in liquid NH3 a t 29.6 f 0.2" are well interpreted in terms of an equilibrium between monomeric and dimeric forms of H20. The intrinsic shift for the average HzO proton in the monomeric species is -268.2 & 0.2 HZrelative t o TMS while for the dimeric species it is -416 rt 12 Hz. The equilibrium constant for formation of dimer is KZ = (1.26 rt 0.15) X l70shifts in H2170SH3 liquid mixtures at room temperature when compared with 170shifts in H2l70-acetone and Hz1iO-(CH3)3N suggest that the shielding at the l70nucleus may be principally determined by the degree to which oxygen lone-pair electrons act as donors to solvent protons in hydrogen bond formation. Such interactions may be more important than heretofore considered. The temperature coefficient of the H20 proton shift in very dilute ppm/deg and appears to be a property of a proton in the Y---H-0 solutions of HzO in liquid X " 3 is 13.5 X hydrogen bond in this system. The temperature coefficient for the proton shift in liquid NH3 is 2.7 X ppm/deg.
Introduction
Experimental Section
Proton shifts were measured against tetramethylsilane (TMS) or tetramethylmethane (TMM) as internal references. TMM was a convenient reference in H20-KH3 solutions -2 M in HzO where the TMS spectral position was matched (and hence the reference obscured) by the upfield member of the 14?X-split, NHI proton triplet. The position of the TMM resonance is 54.3 Hz downfield of the TMS resonance, and this separation is independent of H2O molarity over the range studied. All shifts could, therefore, be easily referred to TMS and could be determined to A0.6 Hz. 170shifts were measured with a magnetometer making use of the control probe and associated electronics which normally supply a proton reference signal for nmr field-lock in the "external reference" version of the Varian field-synchronized recorder. The magnetometer protons were located less than 1 in. from the nuclei under study hence varying field gradients did not introduce appreciable error. A variable radiofrequency oscillator, a frequency counter, and an amplifier for the modulation-frequency phase detector completed the magnetometer. With this technique l7O shifts could be measured with an accuracy of f l ppm. Water enriched to -3% in 170was used for measurement of the 170shifts at high mol % H20while at low percentages, HzO enriched to -30% in 170was used. The NH3 was Matheson anhydrous material, acetone was Baker Analyzed reagent, and trimethylamine was Eastman White Label. Samples were prepared and examined in a plastic-valve, glass nmr cell described elsewhere.2 Water and acetone were pipetted into the
Proton and 1 7 0 shift data mere obtained with the Varian DA-GO instrument. All measurements were made at the equilibrium temperature in the nmr probe which was found to be very stable at 29.6 f 0.2".
(1) Work done Under the auspices of the U. S. Atomic Energy Commission. (2) M.Alei, Jr., and A. E. Florin, J. Phys. Chem., 73, 857 (1969). (3) N . Muller and R . 0.Reiter, J . Chem. Phys., 4 2 , 3265 (1965).
In a recent study2 of H20-NH3 proton exchange in Hz0-NH3 solutions containing OH-, the kinetic data were well fit by an expression which suggested the possibility of more than one HzO species in the systems studied. In particular, it was pointed out that the rate expression would be consistent with an equilibrium between monomeric and dimeric forms of HzO. In this paper we report results of a study at 29.6 f0.2" of the shift, as a function of HzO concentration, of the nrnr peak due to HzO protons in H20-KHs solutions up to 14 M in HzO. The data are well interpreted in terms of an equilibrium between monomeric and dimeric forms of HzO. In an effort to better understand the types of interactions in which the various species might be involved, we have also studied the 170shift as a function of composition in H20-NH3 liquid system at 29.6 =t 0.2". When compared with similar measurements in the HzO-acetone and H2O-( CHB)IN systems, the I7O shifts suggest that the water oxygen may act as an electron-pair donor to solvent protons. Finally, measured temperature dependences of the nmr shift of H20 protons and NH, protons in the HgO-NHa liquid system a t low HzO concentrations support the proposal by Muller and Reiter3 that the temperature dependence of shift may be a characteristic property of a proton in a given type of hydrogen bond rather than mainly a reflection of changes in relative amounts of various hydrogen-bonded species in equilibrium.
Volume '73, flumber 4 April 1969
MOHAMMED ALEI,JR., AND ALAN E. FLORIN
864
triplet with splitting of -80 Hz at HzO concentrations below -10 mol 70. It thus appears that there is no rapid, direct proton exchange between H20 species under conditions where HZO-NHS proton exchange is slow. Because the shift of the HzO proton resonance in Figure 1 approaches a limiting value at low HzO concentrations, we were afforded an opportunity to observe the effect of temperature on the shift of this resonance over a range of HzO concentrations in which changes in HzO concentration produced a negligible effect on the shift at room temperature. The results of the temperature study are shown in Figure 3 where
$ -290 a -295
Figure 1. HSO and NH3 proton shifts in HZO-NH~solutions. -260
cell, and NH, and (CHS)8Nmere added by condensation of the gas from calibrated volumes on a vacuum system.
Results The shift of the nmr peak for HzO protons is shown in Figure 1 where it is plotted against HzO molarity in liquid NH, at 29.6 f 0.2'. The HzO molarity is arbitrarily plotted on a logarithmic scale so that all the data, covering nearly three orders of magnitude variation in HzO molarity, can be legibly displayed on a single plot. The shift of the center component of the NHB proton triplet is also shown in Figure 1. The highest concentration of 14 M HzO represents a practical limit beyond which increasingly rapid averaging of HzO and NH3 protons begins to significantly distort and shift the individual proton resonances. Under these circumstances, accurate shifts cannot be measured without resorting to a detailed analysis of line shapes. 1 7 0 nmr shifts in the Hz0-NH3 liquid system over the entire range of composition at 29.6 f 0.2" are shown in Figure 2. Measured 1 7 0 shifts in the HzOacetone and RzO-trimethylamine systems are also shown. A fact not apparent in these data is that the 1 7 0 resonance is a rather sharp singlet at H20 concentrations greater than -30 mol % and a well-defined
-101 IO
20
30
40 50 60 MOL % H20
70
80
90 'I 0
Figure 2. H2"O shifts in mixed liquid systems: 0 , HzO-NH~; 0 , €120-acetone; (>,HzO-(CH~)~W. The Journal of Physical Chemistry
- Eg
-40
270
-50
-280
-60
1
Y
-290
In z
z
2
-300
0
a a
Ow
-310
z -320 -30
-20
-10
0
IO
20
30
T,OC
Figure 3. Temperature dependences of proton nmr shifts: @,0.048 mol 70 HzO; 0,0.074 mol % HzO; A, 0.44 mol % HzO.
the shift of temperature The shift of triplet over shown.
the HzO proton resonance is plotted vs. at three different HzO concentrations. the center component of the NH3 proton the same temperature interval is also
Discussion The shift of the resonance for the water protons with HzO concentration (Figure 1) provides definite indication of at least two HzO species in this system. Thus, if we assume that in very dilute solutions of HzO in liquid "3, the water is present as a monomeric species (probably associated with one or more NH, molecules through N - - - H-O hydrogen bonds as the extrapolated shift of -268 Hz at infinite dilution is far downfield of the proton shift in gaseous HzO), the dependence of shift on HzO concentration could be explained in terms of an equilibrium between this monomer and at least one other HzO species. If we assume only one other species in equilibrium with monomer, the linear dependence of the shift on water concentration at low [HzO] suggests a monomer-dimer equilibrium. Following the treatment of Saunders and
865
NATUREOF HzO SPECIES IN H20-NH3 SOLUTIONS H ~ n ethe , ~ pertinent equations would then be v = (VlM1
+ 2vzKzM12)/ (M1+ 2K21M12)
(1)
+ 2KzMi2
(2)
of organic solvents have assumed a water monomer complexed t o solvent through hydrogen bonding between water protons and basic solvent sites, i.e.
and
C = Mi
H---S
where v = observed shift, VI = average shift for water protons in monomeric HzO species, v z = average shift for water protons in dimeric H 2 0 species, K Z = equilibrium constant for formation of dimeric HzO species, M1 = concentration of monomeric H20 species, and C = total H20 concentration, [H20]. If a monomer-dimer equilibrium is to adequately interpret our observed shift data, we must be able to find values of the parameters vl, UZ, and K z which, when substituted in the above equations, will produce calculated shifts in agreement with the observed values. For this purpose, we utilized a nonlinear least-squares program written for the CDC 6600 computer to select parameter values which minimized the sum of the squares of the deviations of calculated from experimental shifts. On the basis of the 14 experimental shifts, best values for the parameters were found t o be: V I = -268.2 f 0.2 Ha; v2 = -416 f 12 Hz, and K2 = (1.26 f 0.15) x 10-2 where the error limits in all cases represent standard deviations. With these values, the unweighted sum of the squares of the deviations of the 14 observed shifts from the 14 calculated shifts was 1.5 Hz2 and in no case did an experimental shift differ from a calculated shift by more than 0.6 Hz, Thus the experimental data are fit extremely well by a monomer-dimer equilibrium. The data were tested in similar fashion against a model assuming an equilibrium betiyeen a monomeric and a trimeric water species. For the best fit in this case, the unweighted sum of the squares of the deviations was 5.7 Hz2, and several experimental points differed from calculated points by 0.8 to 1.0 Hz which lies clearly outside our experimental uncertainty of f0.5 Hz. It thus appears that the assumption of an equilibrium between a monomeric and dimeric H2O species is the simplest hypothesis which adequately accounts for the observed shift of the water proton resonance in liquid KH3 solutions up to 14 M in HzO. As noted earlier, such an equilibrium would also be consistent with the observed kinetics for proton exchange in H20-nTH3 solutions containing added [OH-], Though the foregoing discussion makes it seem plausible that water exists primarily in the form of monomeric and dimeric species when dissolved in liquid 3" at room temperature, the proton nmr shift data, per se, give little insight into the specific structures of these species. It seems clear that there must be a t least two different environments for water protons but the exact nature of these environments is not revealed by our experiments. Holmes, Kivelson, and Drinkard6 in their study of dilute solutions of water in a number
/ 0 \
H---S
where S represents a solvent molecule. They interpret their nmr shift data in terms of an equilibrium beheen this monomeric H 2 0 species and a dimeric species with structure S---H
\
/H---S
P---H-o
S---H
These seem to us to be reasonable structures insofar as they indicate likely environments for water protons. Such species could, moreover, certainly account for our reported shifts for the water proton in the HZO-NH, system. On the other hand, the above structures suggest that solvent protons should experience an upfield shift due to some breaking of solvent structure by the HzO molecules. In fact (cf. Figure 1) the NHB proton resonance experiences a very significant downfield shift on addition of increasing amounts of H2O to liquid NH3. Moreover, the proposed structures imply no interaction between water oxygen and solvent protons. In calculating the energy of a single water-solvent hydrogen bond from heat of mixing data, Holmes, Kivelson, and Drinkard do in fact ignore possible hydrogen bonds involving water oxygen as an electron donor and a solvent proton as acceptor. In a more recent publication dealing with water species in organic solvents, Johnson, Christian, and Affsprungs also assume that amines and ketones are unable to interact favorably with the oxygen of the mater molecule. The 170nmr shift data in Figure 2 seem to us to indicate a significant interaction between water oxygen and solvent protons in the systems under discussion. we showed that 170nmr shifts In a recent p~blication,~ in liquid and gaseous water are consistent with a decrease in diamagnetic shielding at the oxygen atom with increasing hydrogen bonding. Whether oxygen lone-pair donation or 0-H bond elongation or both are important in determining the shielding at the oxygen was not evident from the work cited. However, if we examine the data in Figure 2 with the assumption that the observed shifts are due to differences in hydrogen ( 4 ) M.Saunders and J. B. Hyne, J. Chem. Phys., 29, 1319 (1958). (5) J. R. Holmes, D. Kivelson, and W. C. Drinkard, J. Amer. Chem. Soc., 84, 4677 (1962). (6) J. R. Johnson, S. D. Christian, and H . E. Affsprung, J. Chem. Soc., A , 1924 (1967). ( 7 ) A. E. Florin and M. Alei, Jr., J. Chem. P h y s . , 47, 4268 (1967). Volume 73, Number 4
April 196B
866
MOHAMMED ALEI, JR., AND ALAN E. FLORIN
bonding experienced by a water molecule in going from pure liquid HzO to a solution of water in a given solvent, some interesting inferences may be drawn. In the first place, the downfield shift of 170resonance for HzO in liquid XH3 would indicate generally stronger hydrogen bonding between H2O and NH3 than between water molecules in pure liquid HzO. If the principal factor determining the shielding at the 170nucleus in a water molecule is the length of the 0-H bond and the degree to which it is altered by hydrogen bonds involving the HZO proton as an acceptor for a pair of electrons from a solvent donor atom, then one would expect the 1 7 0 shift in a dilute solution of HzO in trimethylamine to be comparable with that for a dilute solution of HzO in NH3 since in both cases the solvent donor is an amine nitrogen. If, on the other hand, the shielding at the water oxygen is largely determined by the degree to which the oxygen lone-pair electrons participate in hydrogen bond formation with solvent protons as acceptors, then one might expect the shift of the "0 resonance for a dilute solution of water in (CH3)aN to be more comparable with that for a dilute solution of water in acetone since both solvents offer similar types of protons as acceptors for lone-pair electrons of the water oxygen. The fact that the l7O shift for water in (CH3)3Nis observed to be comparable with that for water in acetone suggests that the shielding at the oxygen nucleus of the water molecule is primarily determined by the degree to which oxygen lone-pair electrons are involved in hydrogen bond formation. Thus the downfield shift, relative to pure liquid water, of the l7O resonance in dilute solutions of suggests that a water oxygen donates water in 3" electrons better t o protons in liquid KH3 than it does, on the average, to protons in liquid water. Whether or not this further implies that the 0 - - - E - N hydrogen bond is stronger than the 0 - - - H-0 hydrogen bond in liquid water depends on the degree to which pure liquid water is considered to be hydrogen bonded. On the other hand, in acetone or (CH3)3Kas solvent, the upfield shift of the 1 7 0 resonance would indicate an 0 - - - H-C hydrogen bond weaker than the 0 - - H-0 hydrogen bond in liquid water. What is most significant, however, is the fact that even in dilute solutions in acetone or (CH3)3N the 1 7 0 resonance of the water molecule is still some 21 ppm downfield of its position in gaseous water.? If our proposed model is correct, the 0 - - H-C hydrogen bond could thus conceivably be of the order of QQ or -0.6 as strong as the 0 - - H-0 hydrogen bond in liquid H2O assuming complete hydrogen bonding in pure liquid HzO and a linear relationship between hydrogen bond strength and 1 7 0 shifts for the water molecules in these systems. Of course the uncertainties in these assumptions preclude any reliable quantitative assignment of relative bond strengths. However, in view of the quite large downfield shift relative to Hz170vapor of the Hz170resonance
-
-
The Journal of Physical Chemistry
for dilute solutions of water in acetone it seems highly unlikely that there is no significant interaction between the oxygen of the mater molecule and the protons of acetone. We therefore suggest that ignoring possible interaction between water oxygen and solvent protons in dilute solutions of water in organic solvents may not be generally valid. One ot,her aspect of the H2170shift data which should be pointed out is the qualitative agreement between the 170shifts and water proton shifts in the Hz0-NH3 system. Though the minimum in the Hz170shift at ~ 3 mol 0 % (-12 M ) H 2 0 is not very pronounced, it appears to be real and would be consistent with the presence of at least two water species in dilute solutions of HzO in S H 3 . In view of the preceding discussion, it seems likely that the species are monomeric and dimeric forms of HzO associated with solvent KH3 not only via hydrogen bonds involving water protons and NH3 nitrogen, but also through bonding of NH3 protons by water oxygen. If this is so and if as we have suggested earlier the 0 - - - H - S hydrogen bond is a rather strong one, one should expect addition of HzO to NH3 to produce a downfield shift of the resonance for the average NH3 proton. As pointed out earlier this is indeed observed (cf. Figure 3 ) . The temperature dependence of the nmr shift of the HzO proton in very dilute solutions of HzOin NHa (Figure 3) seems clearly to be a property of the HzO proton in the monomeric HzO species. As is evident in Figure 3, the temperature coefficient of 0.81 Hz/deg or 13.5 X low3ppm/deg does not vary by more than -5%, if at all, over an approximately tenfold variation in HzO concentration between 0.176 and 0.019 M . Over this same concentration range, the change in shift with varying H 2 0 concentration is not experimentally detectable indicating negligible concentrations of dimeric HzO species at room temperature. Under these conditions, if the large observed temperature coefficient of shift were due to changes with temperature in the equilibrium constant for formation of dimeric from monomeric HzO species, one would expect the temperature coefficient t o be rather strongly dependent on total HzO concentration. Since this is clearly not so, the change in shift with temperature is probably related to changes in properties of hydrogen bonds involving water protons in monomeric HzO species. Muller and Reiter3 have proposed that variations with temperature of nmr chemical shifts for protons involved in hydrogen bonds may be attributed to changes in the degree of excitation of the hydrogen bond stretching vibrational mode in a given hydrogenbonded species. These authors suggest that the temperature coefficient of 9.5 X 10-~ppm/deg8 for the proton shift in pure liquid water between 25 and 110' may be substantially due to such distortion rather than ( 8 ) W. G. Schneider, H. J. Bernstein, and J. A. Pople, J. Chem. P h y s . , 28, 601 (1958).
AQUEOUSCHEMISTRY OF INORGANIC FREERADICALS rupture of hydrogen bonds. If we accept this interpretation, our observed temperature dependence for the shift of water protons in very dilute solutions of HzO in YH3 could be due to changes in the degree of excitation of the stretching mode of ;I’- H-0 hydrogen bonds involving both water protons hydrogen bonded by NH3 nitrogens in a single, monomeric HzO species. It is interesting to compare the observed coefficient of ppm/deg for a proton in such an environ13.5 X ment with the value of 9.5 X ppm/deg for the ppm/deg proton resonance in liquid H20and 2.7 X for the proton resonance in liquid NH3 (cf. Figure 3 ) . If we assume that the shift of nonhydrogen-bonded protons is temperature independent, and that liquid HzOand liquid “3 are as completely hydrogen bonded as possible, then we might consider that the temperature coefficient characteristic of an - - H-K proton is 3 X 2.7 X low3= 8.1 X ppm/deg compared ppm/dcg for an 0 - - H-0 proton and with 9.5 X ppm/deg for an N - - - H-0 proton. The 13.5 X magnitudes of the temperature coefficients of proton shift would thus seem to fall in the order one might expect for the energies of the hydrogen bonds involving the protons in question. Whether the temperature
-
I
~
-
867 coefficient for proton shift might be generally useful in assigning relative hydrogen bond strengths seems, however, highly questionable. Even assuming the temperature coefficient to be entirely determined by excitation of the hydrogen bond stretching vibration, the magnitude of the temperature coefficient mould depend on the anharmonicity of the vibration as well as upon the bond strength. Moreover, Hindman, et U L . , ~ have demonstrated that dispersion interactions can account for rather large temperature coefficients for proton shifts in liquid HzS, HzSe, and H2Te. Although such interactions might be expected to be less important in liquid HzO or liquid NH3, they cannot, in general, be ignored.
Acknowledgment. The authors are indebted to Dr. B. B. McInteer and Mr. R. M. Potter of this laboratory for the preparation of the enriched 170used in this study. We are also indebted to Dr. T. W. Newton of this laboratory for assistance with the nonlinear least-squares computer program used in some of the calculations, (9) J. C. Hindman, A. Svirmickas, and W. B. Dixon, J. Chem. Phys., 47, 4658 (1967).
Aqueous Chemistry of Inorganic Free Radicals.1 VI. The Effect of Oxygen on the Rate of Photolysis of Hydrogen Peroxide in Aqueous Solutions Containing Carbon Monoxide by F. P. Laming, George Buxton, and W. K. Wilmarth Department of Chemistry, University of Southern California, Loa Angeles, California 90007
(Received Julg 39, 1968)
When aqueous R202is photolyzed in solutions containing both CO and 02, the rate of photolysis is decreased to a value less than that observed in the absence of CO. I n the interpretation presented for these results, the conclusion is reached that in their reactions with the OH radical, the rate constant for CO is 72 i: 3,6 times larger than that of HzOz. To supplement our earlier kinetic studies in the 02-freesystem, tracer studies using 14C0 in the photolysis have been made in an effort to identify the free radical reactive intermediate involved in chain termination at low Hz02 concentrations.
Introduction An earlier paper in this series deals with the use of CO as a scavenger for the OH radicals generated by the photolysis of Hz02.2 Most of the studies were carried out in solutions in which the 0 2 concentration was thought t o be low enough not to influence the kinetic results. One of the objectives of these experiments was the evaluation of the relative scavenger efficiencies of CO and HzOz in their reactions with the OH radical.
I n the present studies, the photolysis of aqueous has been carried out in solutions containing both CO and 0 2 . The work was carried out because the results provide an alternate method of evaluating the relative scavenger efficiencies mentioned above. The present evaluation of the scavenger efficiencies leads to the conclusion that in their reactions with the hyH202
(1) This work was supported by the U. S. Atomic Energy Commission. (2) G. Buxton and W. K. Wilmarth, J . Phys. Chem., 67,2835 (1963). Volume 75, Number 4 April 1969
