NOTES
1458
lower activation energies than the slower reactions with primary alcohols. Recently, studies of exchange between decaborane and deuterium oxide2 indicate that hydrogen at the bridge positions exchanges very rapidly, and that terminally-bonded hydrogen exchanges slowly in dioxane as solvent. Corroboration of the exchange experiments has been obtained by the following techniques: 0.5-1.0 gram samples, a, of decaborane (B10H14)were added to flasks, equipped with stirrers and water-cooled condensers, immersed in constant temperature baths. At the start of the experiment, ca. 450 ml. of distilled water, pre-equilibrated thermally, was added to each flask. Aliquot samples were removed as a function of time, quenched in ice-water,
Vol. 62 TABLE I HYDROLYSIS RAT^ DATA Temp.,
80
100
OC.
ti/,, hr. k , (sec.)-l
4.5 4.27
x
10-5
13.5 1.43 X
E,,kcal. mole-' 14.5 measured conductimetrically by Guter and Schaeffer for initial dissocciation of decaborane in aqueous dioxane.3 Therefore, of the two mechanisms offered for dissociation of decaborane in aqueous media, a ion-dipole hydroxylation of boronhydrogen bonds appears more plausible. (3) G. A. Guter and G. W. Schaeffer, ibid., 78, 3546 (1958).
TITRATION OF DECABORANE IN NITROGENEOUS SOLVENTS
1-
0.8 I -
BY R. W. ATTEBERRY High Energy Fuels Division, Olin Mathieson Chemical Corp., Niaoara FalEs, N . Y . Received May $1, 1968
Guter and Schaefferl have described the pK titration of decaborane, Bl0HI4,as a strong monoprotic acid in water-dioxane mixtures. The procedure suffers the disadvantage of simultaneous hydrolysis of the decaborane, such that precision of the method is somewhat dependent upon the rate of addition of titrant. Schaeffer2has described the complex formed between decaborane and acetonitrile. Bridge hydrogen vibrations appearing in the infrared spectrum of decaborane a t 5.3 and 6.4 p disappear in the infrared spectrum of the complex, suggesting that the nitrogen has donated electrons to the electrondeficient bridges of the decaborane. Heating of the complex liberates hydrogen to form bisacetonitrilodecaborane, (CH3CN)2B1oH12. It was found that whenever decaborane was dissolved in an excess of acetonitrile, such that the complex had been formed, the resulting solution I could be titrated quantitatively with aqueous I I I I alkali with a precision of three parts per thousand. 10 20 30 Figure 1 shows the typical pH titration curve, emHours. ploying glass and calomel electrodes. Fig. 1.-Hydrolysis of decaborane. As evident, the pH titration curve of decaborane in acetonitrile is typical of a weak monoprotic acid and titrated with standard base to a phenol- with pKa = 3.5. This titration serves as an admirphthalein end-point. Excess mannitol, ca. 4 g., was able assay for decaborane, with an inherent color then added, and the sample was retitrated to change a t the end-point from lemon-yellow to phenolphthalein end-point to measure X, the con- straw. centration of boric acid Although acetonitrile has an appreciable dielectric constant of 36, the effect of a nitrogeneous solB~oHlr+ 30Hz0 +10B(OH)3 + 22H2 vent with the same dielectric constant offering Figure 1 is a first-order plot of the data, indicating hydrogen bonding opportunities was sought. Figsimultaneous, or parallel hydrolytic reactions. ure 2 shows a titration curve of decaborane in The more rapid reactions are evidently hydrolyses N,N'-dimethylformamide as such an example. a t the bridge positions, with subtracted half-times Although the precision of titrations of decaborane of ca. 0.25 hour (dotted portions). The slower in N,N'-dimethylformamide is only ca. 4%, it is reactions are dependent upon temperature. strikingly significant that decaborane becomes a This activation energy is similar to the values stronger diprotic acid in this solvent. Moreover, for secondary alcohols1; it equals exactly the value the first inflection at pH 1.2 occurs at exactly one (1) H. C. Beachell and T. R. Meeker, J . A m . Chem. Boc., 7 8 , 1796 (1956). (2) M. F. Hawthorne and J. J. Miller, ibid., SO, 754 (1958).
~~
(1) G. A. Guter and G. W. Schaeffer, J . A m . Chem. (1956). (2) R. Schaeffer, ibid., 79, 1006 (1957).
SOC.. 78, 3546
Nov 1958
NOTES
with active hydrogen, as is the case with N,N'dimethylformamide, then I could have one of its four hydrogen bridges substituted by a nitrogen bridge and a strongly-acidic proton before the tautomeric equilibrium is affected in the manner described above.
8
%
1459
6
THE SOLUBILITY OF SOME CHLOROMETHANES I N WATER' BY JAMES E. BOGGS AND A. ERWIN BUCK,JR.
4
Department of Chemistry, The University of Tezas, Austin id, Texas Received M a y 29. 1968
2 4
I
I
6 8 Meq. of base per gram of sample. 2
4
Fig. 1.-Titration
of decaborane in acetonitrile.
2 4 6 8 Meq. of base per gram of sample. Fig. 2.-Titration
of decaborane in N,N'-dimethylformamide.
In connection with certain kinetics studies we have determined the solubility of CHsCl, CHzFCl and CHFzClin water as a function of temperature. Experimental
A flask, of accurately known volume, was supported in a
thermostated oil-bath and connected to a vacuum line, a manometer and a second smaller flask containing water. After evacuation, the flask was fi11ed with gas to a desired pressure and allowed to stand until temperature equilibrium was established. The connections from the flask were made of capillary tubing so that their volume was negligible. Then an accurately measured volume of boiled distilled water was drawn in (the pressure inside the flask still being less than atmospheric). The flask and contents were then shaken until the indicated pressure became constant. If P I is the pressure of the gas alone in the flask. at .the temperature of the experiment, P2is the final equillbrrum pressure, Pa is the vapor pressure of water at the temperature of the experiment, VI is the volume of the flask and VB is the volume of water admitted corrected for the temperature change, the pressure of gas in equilibrium with the solution is PO - Pa and the molar concentration of the gas dissolved is C - P'Vl - (Pz - Pd(V1 - Vd VzRT It is assumed that the gases obey the ideal gas law, that a t the low concentrations involved the va or pressure of water from the solution is the same aa that $om pure water, and that the volume change in the liquid phase is negligible. The authors wish to express their appreciation to Dr. W. B. McCormack of E. I. du Pont de Nemours and Co. for supplying us with a sample of CHZFC1.
Results fourth the stoichiometry of the second inflection The results of the solubility measurements are at pH 6.0. These facts permit several speculations about the structure of decaborane molecules. shown in Table I, which lists values of K = (molar A tautomeric equilibrium between form I, con- concentration of solution)/(gas pressure in atmostaining four hydrogen bridges, and form 11, con- pheres). From 2 to 5 determinations were made taining two hydrogen bridges, exists in decaborane, a t each temperature, the average deviation of the individual results from the mean being a little over as 1%. At the low concentrations resulting, all of the solutions showed excellent agreement with Henry's law. It will be noted from Table I that, over the temperature range studied, CHzFCl is a little more f - H soluble than CHaCl but CHFzCl is appreciably less soluble than either. The solubility of CFaCl is too low to be measured by our techniques. The thermodynamic quantities AFO, ASo and AHo, for the transference of one mole of organic chloride from the gas phase a t unit pressure to a I I1 111 hypothetical solution of unit activity have been Whenever a nitrogeneous solvent is used, the evaluated and are shown in Table 11. equilibrium is shifted to the right, the two electronFigure 1 shows a plot of A H 0 vs. absolute temperadeficient hydrogen bridges of I1 being replaced by ture. It can be seen that within experimental an electron-donating nitrogen bridge. If the nitro(1) The authors are grateful to the Research Corporation for finangeneous solvent is capable of hydrogen bonding aid support of this investigation.
