NOTES
1076
in D20, log K , = 8.5. The acid dissociation constants of histidylhistidine in H20 and DzO were those of Table I. In the complexation of histidine and the peptides to Cd++, the amino group probably is involved in binding. Since the amino hydrogens are easily exchanged in DzO medium, we would expect a slight deuterium isotope effect on the formation constants, as has been found to be the case. Acknowledgment.-We wish to thank Mr. E. Tucci for preparing the fully deuterated acetic acid. KINETICS OF BOROHYDRIDE HYDROLYSIS BY 171. H. STOCKMAYER, ROYR. MILLERAND ROBERT J. ZETO Deportment
0.i
Chemistry, Massachusetts Instatute of Technology, Cambridge 69, Mossachusetts Recewed December 0, 1060
Recently Davis and Swain' observed that the hydrolysis of borohydride ion EiHI-
+ 4Hn
+ 3H20
was subject to general acid catalysis, according to the expression
- d In (BHI-)/dt
kl = Cki(HAi) i
From measurements of the rate over a wide range of pH with several buffer systems at 25", they determined an accurate value of kHao+and reported approximate values for kHaBos,kHco, - and
-.
kHaP04
Our earlier studies2J of the rate of this reaction supplement the results of Davis and Swain. Confining our experiments to the neighborhood of pH 9, but employing a range of buffer concentrations, we determined kHsO + and kHaBosat 0.0 and 25.0' at an ionic strength of 0.16, and we also evaluated k",+ at 25". Our value for kHaO+ at 25' is in good agreement with that of Davis and Swain, but ~ higher than our value for ~ H ~ BisOconsiderably theirs. Experimental Research grade potassium borohydride (Metal Hydrides, Inc., Beverly, Mass., over 96% KBH,) was used a t an initial concentration of approximately 0.01 M in most runs. The iodate method4J was used to determine borohydride concentrations. The buffers were prepared from HaBO3NaOH or NI-14C1-NaOH,. with addition of NaCl where necessary to bring the ionic strength to 0.16 M . All measurements of pH were made on a Reckman Model G meter a t 25.0'. The values of pH in the borate solutions a t 0" were evaluated from those measured a t 25" from the known change in the ionization constant of boric acid6 (5.79 X 10-10at250to3.09 X 10-laatOo).
-___
(1) R. E. Davis and C. G. Swain, J. Am. Chem. SOC., 82,5949 (1960) ; earlier literature is quoted there. (2) R. R. Miller, S.B.Thesis, M.I.T., May. 1958. (3) R. J. Zeta. S.M. Thesis, M.I.T.. January, 1959. (4) D. A. Lyttle, E. H. Jensen and W. A. Struck, Anal. Chem., 24, 1843 (19.5.2). (5) E. H. Jensen, "A Study on Sodium Borohydride," Nyt Nordisk Forlag, Arnold 13usch, Copenhagen, 1954. (6) H. S. Hsrned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," Reinhold Publ. Carp., New York, N. Y., Third Edition, 1958, p. 755.
Vol. 65
7 t 0
0.05
0.10 0.15 (HzBOa), M . Fig. 1.-Pseudo-first-order rate constant kt, for hydrolysis of borohydride at O", ionic strength 0.16 M and H 9.27 f 0.02, as a function of the molar concentration ofundissociated boric acid.
Results and Discussion Results for the borate buffer system a t 0.0" and pH 9.27 f 0.02 are shown in Fig. 1, where the pseudo-first-order rate constant kl is plotted against the molar concentration of un-ionized boric acid. The slope of the line gives directly the value of kHsBOa, while that of kHaO+ is found from the intercept and the pH with an assumed value of 0.72 for the activity coefEcient of any univalent ion in the solution. These constants and those found in similar fashion a t 25" are collected in Table I. The energies of activation calculated from these results are 9 f 1 kcal. mole-' for the reaction with H30+and 14 f 2 kcal. mole-' for that with Hac),. Our value for kHsO+ at 25" compares well with that of Davis and Swain, (10.0 f 0.4) X lo5 M-' set.-' a t p = 0.10 M ; with a reasonable estimate of the difference in salt effect, their figure and ours agree to better than 10%. However, their estimate of ICH~BO, was (1 f 5 ) X M-' see.-', and even their upper limit is well below any figure that could accommodate our results (Table I). TABLE I REACTIONS OF BOROHYDRIDE IONWITH ACIDS Reaction
BH4-
+ HaO+, 0" H30+, 25" &Boa, 0' H3B03, 25' "a+,
a
25"
ki
(2
(M-1
sec.-V
o f 0.2) x
105
(8 f 1) X lo6 (2 3 0.2) f 10-4 ( 2 o f o 3) x 10-3 (1 5 f 0.4) x 10-3
*
All a t ionic strength 0.16 M .
From the rates for H 3 0 +and NH,+, we estimate an exponent of well over 0.9 in the Bronsted catalysis law for acids of this charge type. Clearly the H-A bond of the attacking acid is very largely broken in the transition state. This conclusion is also supported by the magnitude of the observed difference in activation energy for the H30+ and H3B03 reactions, which within the experimental error equals the enthalpy of ionization of H3B03. Since the enthalpy of ionization of NH4+ is about 12 kca1.-1 mole-', we might expect the activation energy for the NH4+ reaction with BH4- to lie near 20 kcal. mole-', as is in fact borne out by very rough results2for this reaction a t 0".
NOTES
June, 1961 Finally, we record in Table I1 some observed values3 of the rate constant for the second-order reaction between acetone and borohydride in water, at such high pH that the hydrolysis is completely negligible. These results were preliminary to an intended study of the competition between H30+ and acetone for borohydride at lower pH, in the hope of obtaining some information on the steps beyond the first hydride transfer. Competitive systems involving ferricyanide would also be specially interesting, since the rate-determining step of the reduction of ferricyanide by borohydride is knownlJ to be the same as that of the hydrolysis. Since these projects have been abandoned, and since we are unaware of other observations of the temperature coefficient of the acetone reaction in water, we give the figures in Table 11. Our rate at 24' is about 50% higher than that reported by Jensen.6 The activation energy is found to be 7.6 ==! 0.5 kcal. mole-'. TABLE I1 REDUCTION OF ACETONE BY BOROHYDRIDE Temp., '(2.
k X 101 (M-1 sec.-l)
15 20 25 30
1.45 1.83 2.27,2.33 2.73
(7) B. Lowry, S.B. Thesis, M.I.T., May, 1958. (8) T. Freund, J . Inoru. Nuclear Chem., 9, 246 (1959).
1077 Experimental
Materials.-The preparation and purification of the palkyl compounds have been described.' The solvents used were reagent grade materials, dried and distilled before use. The mixed pyridine-carbon tetrachloride solvent way made up to contain 16% pyridine by volume. Procedure.-A Perkin-Elmer Model 21 spectrometer with sodium chloride optics was employed. The mechanical slit width was set at 0.022 mm., corresponding to a calculated spectral slit width of 7.0 cm.-l. A cell of 1.00 mm. thickness was employed; solution concentrations ranged from 0.01-0.08 M. At least four samples of varying concentration were examined for each compound. For each sample a value of apparent intensity was calculated; these were then graphed vs. the absorbance at band maximum. The best line was extrapolated to zero absorbance to obtain the limiting value of intensity. The curves were integrated over an interval of 65 em.-' for the CCb and CCb-pyridine solutions, and over 78 cm.-' for the CHCla solutions. No wing corrections were employed. Results.-The results of the intensity measurements arc shown in Table I. The frequencies of band maxima were the same for all of the compounds in any one solvent. These were 2231, 2230 and 2232 em.-' for carbon tetrachloride, chloroform and the mixed solvent, res ectively. The halfintensity widths, A Y I / ~in, the three soyvents were 11, 13.5 and 11.5 em.-', respectively, exce t that A Y ~ / *for the butyl compound waa about 1 cm. -1 farger than those values in each case. Every effort was made in this study to obtain a high degree of precision, since only small differences in the intensities were expected. The values listed in Table I are thought to be relatively accurate to within 0.01 intensity unit. Where comparison with other data is possible (only the methyl compound has been reported previously)2 the agreement is good. The slightly higher value obtained in chloroform solution a~ compared with the previous one for p-CH32 is the result of integrating over a larger interval.
TABLE I
THE ELECTRONIC PROPERTIES OF ALKYL GROUPS. 111. THE INTENSITY OF T H E INFRARED NITRILE ABSORPTION I N paraALKYLBENZONITRILES BYTHEODORE L. BROWN
Integrated Intensities of the C=N Absorption in p Alkyl-Benzonitriles in Various Solvents Intensities are in units of 1 X 104mole-11iter cm.-2. Solvent
pMethyl
pEthyl
P Isopropyl
Butyl
CCl, CHCla CsHsN-CCL
0.280 ,553
0.300 ,560 .402
0.312 ,573 .402
0.322 ,589 .413
Noyes Chemical Laboratory. University of Illinois, Urbana, Illinois Received December 17, 1060
,408
p-t-
Discussion
The data in Table I show that in chloroform and It was concluded from a study of the dipole carbon tetrachloride solvents the alkyl groups moments of the p-alkylbenzonitriles in cyclohexane cause an increase in the C=N band intensity in that the electron distribution in these molecules in the inductive order. This order is essentially the ground state does not reflect a preferential absent in the mixed solvent containing 16% pyrielectron release from the alkyl groups via hyper- dine. These results thus parallel the dipole conjugation. There is some evidence, however, moment data in showing a change in the relative for a solvent effect in the dipole moment results. order in basic solvents. In dioxane, a basic solvent, the order of dipole On the basis of a simple molecular orbital model moments, after correction for electrostatic effects, for intensities in substituted aromatic compounds3 approaches the Baker-Nathan order: methyl > it has been shown that the square root of the C s N ethyl > isopropyl > t-butyl. intensity in substituted benzonitriles should relate In this note the results of a study of the inte- linearly to the electrophilic substituent constants, grated intensity of the infrared C=N band in the a+.' When the U + values for the p-alkyl groups p-alkylbenzonitriles are reported. The intensity are applied to this correlation it is found that the of this band in substituted benzonitriles is quite predicted intensities vary in the order: methyl > sensitive to the electronic properties of the meta or ethyl > isopropyl > t-butyl. This is the inverse para substituent.2 It was felt, therefore, that a of the order actually found in carbon tetrachloride, determination of the band intensity in the p-alkyl although in the more basic mixed solvent the compounds would provide data which would be of changes in relative values are in the direction of the interest in evaluating the properties of the alkyl Baker-Nathan order. One can speculate that if groups. the intensities were obtained in 90% aqueous T. L. Brown, J. Am. Chsm. Soc., 81, 3232 (1959). (2) P. J. Krueger and H. W . Thompson, Proc. Bog. SOC.(London), 8960, 22 (1958). (1)
(3) T. L. Brown, J . Phys. Chem., 64, 1798 (1960). (4) H. C. Brown and Y. Okamoto, J. A m . Chem. SOC.,80, 4979
(1958).
