Nuclear magnetic resonance determination of hydration numbers of

HYDRATION NUMBERSOF ELECTROLYTES

1563

From Table 111it is reasonably certain that the major, initial indications obtained for the mono- and dimethif not the entire change is accounted for by the differylamines.I5b ence, &AH, which may be set equal to the zero point The agreement of the values of ACp for the protium energy difference. This conclusion is the same as was and deuterium analogs confirms the absence of signifireached from the temperature dependence of k ~ / k ~cant solvation differences accompanying the secondary for the solvolysis of l-butyl chloride but differs from deuterium isotope effect in this system (ref 1, p 158).

Nuclear Magnetic Resonance Determination of Hydration

Downloaded by UNIV OF PENNSYLVANIA on August 22, 2015 | http://pubs.acs.org Publication Date: May 1, 1969 | doi: 10.1021/j100725a064

Numbers of Electrolytes in Concentrated Aqueous Solutions by R. W. Creekmore and C. N. Reilley Department of Chemistry, University of North Carolina,, Chapel Hill, North Carolina

27514

(Received November 1.2, 1968)

Total effective hydration numbers have been determined for a number of alkali metal salts and alkaline earth metal salts by studying the temperature-dependent proton chemical shifts of their aqueous solutions. The hydration numbers obtained for the alkali metal halides were rationalized by assuming that the cation was coordinated to four waters whereas the anions contributed little if any to the hydration number obtained. The hydration numbers for the alkaline earth metal halides were found to include both primary and secondary waters which was attributed to their larger charge-size ratio. Also studied was the effect of several anions (Cl-, Br-, C104-, and p-toluenesulfonate-) on the effective hydration number of their sodium salts. Both NaC104 and Na p-toluenesulfonate gave hydration numbers lower than NaCl and NlaBr. The decrease in the effective hydration number was explained in terms of ion pairing, which was further supported by 23Na relaxation times.

taoduction I n the past decade nuclear magnetic resonance (nmr) has proved to be valuable in studying electrolyte solutions; an excellent review has been given by Hinton and Amis.’ Proton chemical shifts of electrolyte solutions have revealed information about solvent structuring and hydrogen bonding in these solutions. Of recent interest, however, have been the studies of solvation numbers by nmr. For ions which form rather strong coordination with water such 8s Ala+, ]Be2+,and Ciaa+,the exchange of solvation water with bulk water is sufficiently slow so that separate resonances for both types of water may be observed using either “0 r e ~ o n a n c e ~ or, - ~ a t low temperature, using proton resonan~e.~ By ~ ~the addition of acetone to aqueous solutions of Mg (Clod) and Mg (NO3)2, Matwiyoff and Taube? were able to decrease the exchange rate of solvation waters, thus distinguishing the proton signals of Mg (H2O)2’+ ion from those of the bulk water st temperatures below - 70°. The exchange of solvation water for bulk waters is much faster in the case of alkali metals, and their study by the above direct methods is not feasible. However, Malinowski, et al., have shown $ha$ the temperature

effects on the proton shifts of aqueous electrolyte solutions are related to the degree of “effective” hydration.8J‘ Their method is based on the assumption that the hydration number and the chemical shift for the hydration waters do not change with temperature; hence, the observed chemical shift is a weighted average of the chemical shift of the hydration waters and the temperature-dependent shift of the bulk water. They have studied two systems [NaCl and Al(NQa)s] and report a total “effective” hydration number of 4.5 for NaCP and 13.4 for A ~ ( N Q ~ ) Since s . ~ the hydration number for the aluminum ion has been shown by direct (1) J. F. Hinton and E, 8. Amis, Chem. Reo., 67, 367 (1967),and references therein. (2) J. A. Jackson, J. B. Lemons, and H. Taube, J . Chem. Phys., 3 2 , 563 (1960). (3) R. E. Connick and D. N. Fiat, $bid., 39, 1349 (1963). (4) D. N. Fiat and 3%. E. Connick, J. Amer. Chem. Soc., 88, 4754 (1966). (5) R. E. Shuster and A. Fratiello, J . Chem. Phys., 47, 1554 (1967). (6) A. Fratiello, R. E, Lee, V. M. Nishida, and R. E. Shuster, (bid., 48, 3705 (1968). (7) N. A. Matwiyoff and H. Taube. 9.Amer. Chem. Soc.. 90, 2796 (1968). (8) E. R. Malinowski, P. S. Knapp, and B. Feuer. J. Chem. Phys., 45, 4274 (1966). (9) E,R. Malinowski ana P. 8 . Knagp, (Bid., 48, 4989 (1968).

Volume 7.9,Number 6 May 1069

R. W. CREEKMORE AND C. N. REILLEY

1564


methods to be 6, the hydration number obtained by this technique apparently sees more than the primary hydration shell. I n determining what the measured hydration number means we have considered the following types of interaction. (a) Cation-Solvent. This interaction accounts for the cation effects on bound waters. I n the case of divalent and trivalent ions the cation effect may extend to waters which lie in the secondary hydration layer. (b) Anion-Solvent. One would expect the anion to bind water less effectively than the cations; therefore, this contribution to the total effective hydration number would be small, but not necessarily insignificant. (c) Cation-Anion. This interaction is expected to be of particular importance in salts which form ion pairs and, where operative, is expected to lower the observed hydration number. The purpose of this study is to gain further understanding of the ion-solvent and ion-ion interactions conveyed by the experimentally determined total effective hydration number, using the temperature variation method described by Malinowski, et al. I n the case of aluminum nitrate the large effective hydration number would indicate that the secondary waters are measured, as Malinowski suggested. To substantiate this idea, we report here the effective hydration numbers for both MgClz and CaClz as determined by this method. Hindmanlo has studied the concentration dependency of the chemical shift for a number of electrolyte solutions a t constant temperature, and in theoretically accounting for the various effects giving rise to the observed shifts he obtains hydration numbers for alkali metal ions which decrease systematically in the series from lithium (to which he assigns a value of 4.0) to cesium (to which he assigns a value of 1.0). To determine if such a trend in hydration exists we have studied several alkali metal chlorides by the temperature dependence technique of Malinowski. I n addition, since alkali metals are not as strongly coordinated to water as are divalent and trivalent ions, ion-ion interactions may be expected to have significant influence on the effective hydration numbers obtained. T o evaluate more fully the effects of ion association on the total “effective” hydration number, several different sodium salts were studied by the temperature dependence method and by W a resonance measurements.

Experimental Section The proton magnetic resonance measurements were made with a Varian HA-100 high-resolution spectrometer operating a t 100 MHz. The sodium resonance measurements were also made on the HA-100 operating The Journal

of

Physical Chemistry

HR mode a t 26.452 MHz. The proton chemical shifts of all aqueous solutions were measured relative to external cyclohexane with a precision of f 0 . 2 Hz. The reference, which was sealed in a capillary, was held coaxially in the sample tubes by precision Teflon spacers. The same reference capillary was used throughout the study. The temperature of the samples was controlled by a Varian V-6040 temperature controller to +lo,and the sample temperature was determined by measuring the chemical shifts of methanol or ethylene glycol samples (both samples were calibrated beforehand using a copper-constantan thermocouple). All shifts were corrected for bulk-susceptibility effects according to the equation“ &or

= Bobs

+ +n(xr - xa)

The volume susceptibilities (x) of cyclohexane and pure HzO a t the various temperatures were calculated from specific su~ceptibilities’~J~ and den~ities.’~JsFor the electrolyte solutions the specific bulk susceptibility was computed using the ionic susceptibilities compiled by Selwood.l6 The volume susceptibilities of the electrolyte solutions were computed by the interpolation of density-temperature data,” except in the case of (CH,).,NCl and Na p-toluenesulfonate where such data are not available. I n these two cases the change in density with temperature was assumed to be the same as those in pure water and sodium chloride solutions, respectively. Since the absolute magnitude of the susceptibilities is not important in determining the hydration number, such assumptions are justifiable. All salts used were analytical reagent grade of 99+% purity. With the exception of RbCl and Na p-toluenesulfonate all salts reported were used without further purification. I n the case of RbCl and Na p-toluenesulfonate, it was necessary to repurify by recrystallization from water. Solutions were prepared gravimetrically if the salt was anhydrous. Otherwise the solutions were analyzed to determine their concentration and densities taken to obtain their molalities. Deionized, distilled water was used in the preparation of all solutions. (10) J. C . Hlndman. J . Chem. Phzrs., 3 6 , 1000 (1962). (11) J. A. Pople, W. G . Schneider, and H . J. Bernstein, “High Resolution Nuclear Magnetic Resonance,” McGraw-Hill Book Co.. Inc., New York, N. Y., 1959, p 81. (12) J. W. Emsley, J. Feeney, and L. H. Sutcliffe, “High Resolution Nuclear Magnetic Resonance Spectroscopy,” Pergamon Press, New York, N. Y., 1965, Appendix 0. (13) T. Tiets, J . Chem. Phys., 31, 274 (1969). (14) J. Timmermans, “Physico-Chemical Constants of Pure Organic Compounds,” Elsevier Publishing Co., Inc., New York, N. Y . . 1950.

(15) Handbook of Chemistry and Physics, 43rd ed. (16) P. W. Selwood, “Magnetochemistry,” Interscience Publishers, Inc.. New York, N. Y., 1956, pp 78, 86. (17) “International Critical Tables.”

HYDRATION NUMBERS OF ELECTROLYTES

1565

Table I: Effect of Various Electrolytes on the Temperature Dependence of the Proton Resonance of Water Solution


Hi0 LiCl LiCl NaCl NaBr NaClO4 Na p-toluenesulfonate KCI RbCl CSCl MgClz CaC12 (CHs)aNCl (CHa)aNCP (CHs)aNBP a

Concentration, m

Temp range,

Pure 3.08 4.62 3.08 3.01 2.98 2.73 2.99 3.20 3.02 2.92 2.37 2.66 2.66 2.92

6.8-80.5 (14) 6.8-74.0 (9) 11.2-64.0 (16) 12.2-69.5 (8) 12.2-74.0 (9) 11.2-74.0 (9) 12.7-75.3 (7) 6.8-74.0 (9) 11.7-04.0 (7) 6.8-74.0 (9) 6.8-74.0 (10) 6.8-74.0 (10) 11.7-64.0 (6) 11.2-74.0 (8) 11.7-74.0 (7)

Referenced to the methyl group of the electrolyte.

b

d&or/dt, ppm/deg

- B ppm

0.00958 f0.00009 0.00778 f0.00012 0.00709 f0.00010 0.00713 f0.00006 0.00729 f0.00005 0.00808 z!= 0.00009 0.00829 f0.00009 0.00719 f0.00014 0.00737 f0.00012 0.00756 f0.00010 0.00545 f0.00009 0.00575 f0.00007 0.00929 f0.00014 0.00913 f0.00006 0.00906 f0,00009

3.652 3.569 3.528 3.379 3.311 3.262 3.329 3.376 3.393 3.468 3.795 3.534 3.603 1.736 1.636

The numbers in parentheses represent the number of data points.

by the temperature changes. Thus

Results I n the electrolyte solutions studied a single resonance signal is observed whose chemical shift is a weighted average of those for the different environments in which a water molecule can exist. The chemical shift may be written as &or

Xb6b

+ Xd,

(1)

where &or is the observed chemical shift corrected for bulk susceptibility effects, xb and X, are the mole fractions of bulk and solvation water, respectively, and bb and 6, are the chemical shift of the bulk and solvation water, respectively. The temperature variation method of Malinowski, et al., for determining total effective hydration depends on the assumption that only the chemical shift of the bulk water is affected

ds,,/dt

=

x b (d&/dt)

+ X, (dS,/dt)

(2)

Because (dS,/dt)

dSco,/dt

0,

Xb(d&,/dt)

(3)

Since bulk water should behave as pure water, then d&.,/dt = d6a,o/dt

(4)

where BHzO is the susceptibility corrected shift for pure water. Hence dScor/dt = ( 1

- X,)d&,o/dt

(5)

Since X , = hm/55.55, where h is the total effective hydration number and m is the molality

[

h = 55.55/m d b o g t -/d&o,ldt] (6) ~~0 dt This expression is equivalent to that of Malinowski but is cast directly in terms of the temperature-dependent quantities measured here. It is apparent from eq 6 that the absolute magnitude of the susceptibility used to correct the observed chemical shifts is not important but only the variation of the susceptibilities with temperature must be known; this can be computed from density-temperature data. Figure 1 is a plot of 6,, vs. temperature for LiCl a t two different concentrations and includes the curve for pure water. All three lines should intersect at the same point if the hydration number is independent of concentration, because a t the intersection point the chemical shift of the solvation water is equal to the shift of the bulk water. I n the case of LiCl solutions the intersection points coincide within experimental error. As will be shown later this is not necessarily true for all electrolytes. Table I gives the slope and intercept values for the

I

20

40 60 TEMPERATURE ("C1

80

Figure 1. Proton chemical shifts (relative to cyclohexane) of aqueous LiCl solutions at various temperatures and concentrations: e, pure water; A, 3.08 m LiCI; m,4.62 m LiCI.

Volume 78,Number 6 May 1969

R. W. CREEKMORE AND C. N. REILLEY


1566

electrolytes studied along with the standard deviations in the slopes; the experimental data, after correction for susceptibility effects, were fitted to least-squares lines, The slope of the pure water line, 0.00958 ppm/deg, agrees very well with that of other investigators who report values of 0.009518and 0.00956 p ~ m / d e g . ~I n the case of the (CH3)4NC1, both internal and external referencing were used to check the validity of the susceptibility corrections. The two slopes were found to be identical within experimental error. Concentrations around 3 na were selected in order to obtain the most accurate data with the least amount of ion pairing. I n Table I1 we have compared our interpolated molal shift values a t 25’ with those of Hindman,lo taking into account his “B” term. The agreement is satisfactory within experimental error. Table 111 lists the total effective hydration numbers calculated from eq 6 for all the electrolytes studied. The value 4.6 for NaCl agrees, within experimental error, with the value of 4.5 determined by Malinowski, et al. I n general, the alkali metal salts were found to give total effective hydration numbers around 4 while the corresponding numbers for the alkaline earth metal salts are about double this value.

Discussion Hydration of Alkali Metals. It is quite difficult to assign a primary coordination number to the alkali metal ions on the basis of the total effective hydration number obtained experimentally for alkali metal salts, since it is not possible to separate rigorously each of the previously mentioned contributions to the total effective hydration number. However, the data suggest that the alkali metals are hydrated and from the magnitude of the total effective hydration numbers we can assume that only the primary hydration shell contributes significantly to the values determined. The fact that the total effective hydration numbers for the alkali metals are about 4 is consistent with the viewpoint that these ions have a coordination number of 4. Bernal and Fowler,19 for example, proposed

Table 11: Effect of Electrolytes on the Proton Resonance of Water at 25” Solution HzO LiCl NaCl XaBr KCl

RbCl CSCl NaC104 0

Concentration, m

,,a

Pure

3.08 3.08 3.01 2.99 3.20 3.02 2.98

Hindman, ref 10.

The Journal of Physical Chemistry

ppm

-3.412 -3.375 -3.201 -3.129 -3.196 -3.209 -3.279 -3.060

6”. ppm

... -0.012 -0.069 -0.094 -0.072 -0.064 -0.044 -0.118

-0.010 -0.067 -0.091 -0.067 -0.056 -0.045 -0.120

Table 111: Hydration Numbers of the Electrolytes Studied Electrolyte

LiCl LiCl NaCl NaBr NaClO4 Na p-toluenesulfonate

KC1 RbCl CSCl MgClz CaCll

(CHa)4NCl (CHd4NBr

Concentration, m

h (f0.2)

3.08 4.62 3.08 3.01 2.98 2.73 2.99 3.20 3.02 2.92 2.37 2.66 2.92

3.4 3.2 4.6 4.4 3.0 2.8 4.6 4.0 3.9 8.2 9.4 0.6 1.0

(from a theoretical standpoint) that the alkali metals would have a coordination number of 4 consistent with the average coordination of water (which has been shown from X-ray studies20 to be 4.6) and the small volume alterations of solutions of alkali metal salts. Samoilov21also points out the strong influence that the existing water structure (tetrahedral) has on the coordination of alkali metal ions. It is reasonable to assume that these ions will assume a coordination which will be the most energetically stable, and that in the case of alkali metals where the electrostatic interactions with water is small (comparatively speaking) the most stable coordination would be that which would have the smallest effect on the water structure. Therefore one could envision the hydrating of an alkali metal ion as a substitutip of the ion for a water molecule, which is tetrahedrally coordinated. Several X-ray studies have been conducted which offer support to the above conclusions. For example, X-ray studies22 on aqueous solutions of LiC1, LiBr, and RbCl have found that the diffraction maximum of pure water is retained in these solutions, indicating that the water has not changed except for substitution of a metal ion for a water molecule. Brady and K r a u ~ on e ~the ~ ~ ~ ~ basis of their X-ray studies have assigned to both K+ and Li+ a coordination number of 4. There will not, however, be any attempt to compare the hydration numbers obtained by this technique with those obtained by the more classical techniques, such as mobilities, density, compression, and activity measurements, since there is little agreement among these various methods. (18) W. G. Schneider, H . J. Bernstein, and J. A . Pople, J . Chem. Phys., 2 8 , 601 (1958). (19) J. D.Bernal and R. H. Fowler, ibid., 1, 515 (1933). (20) J. Morgan and B. E. Warren, i b i d . , 6, 666 (1938). (21) 0. Ya. Samoilov, “Structure of Aqueous Electrolyte Solutions and the Hydration of Ions,” translated by D. J. G. Ives (Consultants Bureau Enterprises, Inc., New York, N. Y., 1965,pp 132, 133. (22) I. Beck, Phys. Z . , 40, 474 (1939). (23) G.W.Brady and J. T. Krause, J . Chem. Phys., 2 7 , 304 (1957). (24) G.W. Brady and J. T. Krause, i b i d . , 28, 464 (1958).


HYDRATION NUMBERSOF ELECTROLYTES

If the C1- ion is assigned a hydration number of 0.6 (a value which is consistent with the total hydration number determined for (CH3) 4NC1, assuming that the (CH3)4N+ion is neither hydrated nor associated to any detectable degree), a value of 4.0 would then be obtained for Na+, 4.0 for K+, 3.4 for Rb+, and 3.3 for Cs+. The regular decrease in the hydration number in going from sodium to cesium suggests an increasing degree of ion association if one is to hold to the idea that the coordination for these ions is the same. It has been shown that RbCl and CsCl do indeed form ion pairs.26 Assuming that the coordination number of these ions is 4 in dilute solution and that the ion pairs formed are contact ion pairs (which we have assumed for the sake of simplicity to have a hydration number of 0), the degree of ion pairing in 3 rn solutions of RbCl and CsCl was estimated as 13 and 14%, respectively. On the other hand, LiCI, NaC1, and KC1 are not expected to form ion pairs a t the concentrations employed. For example, Hindman,lo from his studies of metal and halide nuclear magnetic resonance, could not find any evidence which would indicate direct ion association for LiC1, NaC1, and KCI. However, both metal and halide resonance shifts for RbCl and CsCl solutions could be interpreted as indicating ion association. Since ion association may be reflected in the total hydration number obtained for any given cation, the total hydration number value will in this sense depend indirectly on the anion. To determine this influence of the anion, we chose to study several sodium salts since the above effects could be further substantiated by 2aNanmr. Due to the quadrupole moment of sodium, the symmetry and correlation time of the solvent around sodium is important in determining its relaxation time. For a coordination number of 4 or 6 the field gradient around sodium would become 0. Any deviations from a symmetrical field environment would be reflected in a decrease in TI resulting from the quadrupole interaction with the fluctuating electric field gradientnZ6Therefore, a measurement of T I or line width (since TI = T z ) for sodium should reflect

Table IV: Anion Effect of W a Resonance of Aqueous Solutions of Several Sodium Salts Sodium compound NaCl NaCl NaBr NaI NaC104 Na p-toluenesulfonate

Concentration, m 2.86 3.08 3.00 3.00 2.62 2.73

l/Ti, sec-1

20.8

...

22.5 24.0 50.5

...

Peak width, HZ 6.70 7.60 7.2b 7.6‘~ 16.10 11-50

Eisenstadt and Friedman, ref 27, spin-echo measurements. Speight and Armstrong, ref 26, spin-echo measurements. c This laboratory, line-width measurements.

1567 the interaction of the anion with sodium. Table IV gives the relaxation times for several sodium salts, The data suggest that C1-, Br-, and I- exhibit little if any interaction with sodium. However, the lower TI values for the (3104- and p-toluenesulfonate- solutions suggest that these anions interact with the sodium ion. The low effective hydration numbers for NaC104 and Na p-toluenesulfonate given in Table 111, therefore, appear to be due to ion association. Lending support to this conclusion, it has been shown from conductivity measurements2Sthat ion association exists in KC104 solutions (which one would assume behaves somewhat analogously to NaC104) and also from 2aNa relaxation measurement^,^^ where ion association in NaC104 solution was suggested as a possible explanation for the faster relaxation rates found for NaC104 as compared to NaCl solutions. LiCl yields a total effective hydration which is lower than expected on the assumption that Li+ is coordinated to four waters. Significant ion pairing would tend to lower the hydration number obtained, but ion pairing seems improbable since ’Li nmr26has indicated a very symmetric environment for lithium in LiCl solutions and as previously mentioned the concentration-dependent proton chemical shifts for LiCl solutions show no deviation from linearity. Hydration of Alkaline Earth Metals. For the divalent cation salts, MgCL and CaCl2, the total effective hydration numbers (8.2 and 9.4) obtained are greater than the primary hydration number of 6.0 which has been established for Mg2+.6~7The high values may be rationalized in two ways. First, it was suggested earlier that a hydration number of 0.6 could be assigned to the C1- ion. By taking 0.6 for the hydration of each C1- ion and 6 for the hydration of the divalent cation, one arrives a t a total effective hydration number of 7.2 (MgCI2). This is still lower than the experimental value of 8.2 A value of 1.1 for the hydration number of the C1- ion would, however, account for the deviation but would not explain the large hydration number determined for CaC12. Malinowski and Knapp suggested that the hydration number for the NOs- ion might be 2 or 3 as a possible explanation for the large hydration number obtained for A1 (NO$)3. However, such large hydration contributions from anions seem improbable from our studies. Second, the high charge to size ratio causes the secondary hydration layer to be held sufficiently tight so that this layer makes a positive contribution to the hydration number measured by this technique. This seems reasonable, since each primary water could provide two binding sites. If secondary hydration (25) C . W. Davies, “Ion Association,” Butterworth, and 00. Ltd., London, 1962, Appendix, p 169. (26) P. A. Speight and R . L. Armstrong, Can. J. Phys., 4 5 , 2493

(1967). (27) M. Eisenstadt (1966).

and H. L. Friedman, J . Chem. Phys.,

4 4 , 1407

Volume 7.9, Number 6 May 1968


1568

D. B. PETERSON, J. HOLIAN,AND W. M. GARRISON

effects are reflected in the measured hydration numbers, they would be expected to be even larger in the case of the trivalent metal ions such as AI3+. A larger hydration number was obtained for CaCl2 than for MgC12, and this would not have been expected unless the primary hydration number for Ca2+ were larger than 6.0, or unless solvent-separated ion pairing is more predominant in MgClz solution than CaClt solutions (assuming that if ion pairing does occur in these solutions it would affect only the secondary hydration layer). It is interesting to note that Swift and SayreZ8using a quite different technique also obtained a larger hydration number for Ca2+ than for Mg2+. Despite the problems associated with this technique, we feel that it can be useful in providing further understanding not only of hydration in the alkali metal

series but also of ion-ion interactions. Some of the interpretive difficulties may lie in the basic assumption that the total effective hydration number and 6, do not change over the temperature range measured and in this sense, the technique will require more investigation before more definitive assignments are possible.

Radiation Chemistry of t.he a-Amino Acids.

Aclenoutledgment. One of the authors (R. W. C.) wishes to acknowledge the financial assistance provided by the University of North Carolina Materials Research Center Contract SD-100 with the Advance Research Projects Agency. We would also like to acknowledge the National Science Foundation Grant 325-NAT-190 for the purchase of the Varian HA-100 used in these studies. (28) T.J. Swift and W. G. Sayre, J . Chem. Phys., 44, 3567 (1966).

Radiolysis of Solid Cysteine’

by Donald B. Peterson,2 John Holian, and Warren M. Garrison Lawrence Radialion Laboratory, University of California, Berkeley, California

94720

(Received November 1 3 , 1 9 6 8 )

The y radiolysis of solid cysteine is shown to yield hydrogen, hydrogen sulfide, ammonia, and cystine as major products. Radiolysis data are also given for a series of related compounds: cystamine, N-acetyl cysteine, S-methyl cysteine, glycine, and alanine. The proposed reaction schemes provide a basis for correlating the radiation chemistry of these solid-state systems.

Introduction Radiolysis of the simpler a-amino acids such as glycine and alanine leads to deamination as a major chemical consequence both in aqueous solution3 and in the solid state: and there is accumulating evidence that the intermediate processes of deamination in the solid state are closely analogous to those that occur in aqueous solution. For example, in aqueous solution the ionization step6 HzO---aH+, OH, eaq-

(1)

is followed bye eaq-

+ NH3+CH(R)COO- -+

+ CH (R) COO-

NHB

(2)

OH

+ NH,+CH (R) COO-+

HzO

+ NHa+C( R )COO-

(3)

where eaq- represents the hydrated electron. The experimental evidence is that ea,- adds initially to the The Journal of Physical Chemistry

C=O linkage of the carboxyl group and that dissociation of the reduced intermediate then ensues.6blo Subsequent interactions of CH (R) COOand (1) This work has been supported by the U. 8. Atomic Energy Commission. (2) Department of Chemistry, University of Ban Diego, San Diego, Calif. (3) (a) W. M. Dale, J. V. Davies, and C . W. Gilbert, Biochem. J . , 45, 93 (1949); (b) G. Stein and J. Weiss, J . Chem. Soc., 3256 (1949): (c) C. R. Maxwell, D. C. Peterson, and N. E. Sharpless, Radiation Res., 1,530 (1954); (d) B. M. Weeks and W. M. Garrison, ibid., 9, 291 (1958). (4) (a) W. M. Dale, J. V. Davies, and 0.W. Gilbert, ref 3e; (b) B. Rajewski and K. Dose, Z . Naturforsch., 12B, 384 (1957); (c) G. Meshitsuka, K. Shindo, A. Minegishi, H. Suguro, and Y.Shinosaki, BuEl. Chem. SOC.Jap., 37,928 (1964);(d) A. Minegishi, Y. Shinosaki, and G. Meshitsuka, (bid., 40, 1271 (1967):(e) W. 0.Gottschall, Jr., and B. M. Tolbert, J . Phys. Chem.. 72, 922 (1968). (5) For a recent review of the radiation chemistry of water, see M. 9. Matheson, Advances in Chemistry Series, No. 50, American Chemical Society, Washington, D. C., 1965,P 45. (6) (a) W. M. Garrison, Lawrence Radiation Laboratory Report No. UCRL-10827, May 1963;Radiation Res. Suppl.. 4, 1964; (b) B. M. Weeks, 9. A. Cole, and W. M. Garrison, J . Phys. Chem., 69, 4131 (1965); (c) R. L. 8. Willix and W. M. Garrison, ibid., 69, 1579 (1965); (a) R. L. 9. Willix and W. M. Garrison, Radiation Res., 32, 452 (1967).

Nuclear magnetic resonance determination of hydration numbers of

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