COMMUNICATIONS TO THE EDITOR
116 BizHm(C0)z f HzO
Vol. 86
+
many powers of ten is simply unfeasible by direct experimental determination. Resort has been made, therefore, to such devices as the use of acidity functions, salts. Acidity function studies indicate carboxyl for example, in proton transfer equilibria, to estimate proton ionization constants of the effects of changing acid medium. N o direct experiThe chemistry of the dicarbonyls is of extremely mental data have been available to test the accuracy broad scope. Some of their reactions are outlined of the acidity function method over wide ranges (corbelow for the Blo system; analogous reactions occur with responding to many powers of ten in equilibrium conthe Blz dicarbonyl. stant), and indeed recent evidence leads to certain rBCH3OH e ( H + ) Z B ~ ~ H ~ ( C O O C H ~ ) ~ - ~ serious questions of this usage.3 We report herein the determination of reversible 4NH3 +( N H I ' ) ~ B ~ O H COK"z)2-z ~( electrode reduction potentials in acetonitrile solution BloHa(CO)z+ LiA1H4--+ BloHs(CH3)2-2 a t 25' which provide quantitative measures of the NaN3 -+ B I O H ~ ( N C O ) Z - ~ (large) substituent effects on the standard free energy H20 + HzNOS03H +BioHs( XH3)z ZCOz change of a which gives relative cation stabilii.e. zation energies,4 The BloHs(NH3)2prepared in this fashion is identical In with that prepared from BloHa(N2)z and "3. R' e l / K-R ~ (1) water the following equilibria exist. In this process, the widely delocalized charge of R + becomes essentially localized (for most practical purBloHs(NH3)z S H+BioHsNHz.NHa- e ( H +)zBloHs(NHz)z-' poses) in the nonpolar C-C bond. For general R + , Dehydration of CSzBioHa(CONH2):, is effected by important steric effects on formation of R-R may be heating to about 300". We have not encountered a involved. In the present report, however, we have restricted R+ to p-substituted triphenylmethyl (trityl) CSZBioH8(CONH2 jz --+ CSzBloHs(CN)z cations so that steric interactions in the hexaphenylchange in stereochemistry in any of the reactions of ethane derivatives are essentially constant. BioI-Ii(C0)z. The cell, Pt 1 Rz(cH~cP*;), R+BF~-(cH*cN)I Ag+Halogenated carbonyls (e.g . , BloCls(CO)2 , BloBrs.( CO)2)N ~ ~ - ( c H ~ 1cAg, N ) is prepared on a vacuum h e with a have been prepared by aqueous halogenation of weighed amount of R+BF4- E, into which highly purified BloHs(CO)2and B12Hlo(CO)zfollowed by dehydration, acetonitrile6 is distilled. The e.m.f. cell contains a a t elevated temperatures, of the resulting halogenated volumetric compartment enabling a measured volume carboxylic acids. These halogenated carbonyls are of known concentration of the R+BF*- acetonitrile even more thermally stable than the perhydro carbonyls, solution to be withdrawn, reduced (to the equilibrium e.g., there was no evidence of decomposition after heatmixture of Rzand R.) by metallic zinc, and titrated ing BloCls(CO)2a t 300" in an evacuated sealed tube into the solution in contact with the Pt electrode. for 1 hr. A diverse chemistry based on halogenated Contact between the two electrode compartments is carbonyls has been demonstrated. For example, obtained by use of an asbestos fiber. The e.m.f. BioCl~(C0)zreacts with sodium borohydride to give measurements are made with an L. and N. K-3 potenthe bishydroxymethyl species B ~ ~ C ~ E ( C H ~ O H and ) Z - ~ tiometer and are reproducible to 0.002 v. The e.m.f. couples with dimethylaniline to give BloCls(COC6H4values are stable over time intervals of a few minutes NMe2)2-2 BloCls(C0)2is converted to B10Cl~(NH3)2 up to a t least 2 hr. to the same precision. Variation by reaction with aqueous hydroxylamine-o-sulfonic M to 2.6 X M and of of (R+)from 0.6 X acid. Diazotization of this chlorinated diamine readily (R2) from 0.10 X ;M to 0.80 X M gives gives B1oClE(Nz)z which, in turn, can be converted to e.m.f. values which follow the equation BloClB(Na)2-2by reaction with sodium azide or to BloClB(SH2)2 by reaction with hydrogen sulfide. H+BizHioCOOH.CO-
OH -
( H +)zBizHio(COOH)z-2+BizHio(COO)z-4
1
+
+
*
CONTRIBUTION S o 930 W H KNOTH J C SAUER to a precision of b0.003 v. CENTRAL RESEARCH DEPARTMENT A N D COMPANY H C MILLER E I D U PONTDE NEMOURS The standard free energy change, A F z 0 ,for the reacE L MUETTERTIES tion EXPERIMENTAL STATION WILMINGTOU, DELAWARE R+ '/z(Ro-Roj = Ro' '/z(R-R) (2) RECEIVED NOVEMBER 4, 1963
+
+
An Electromotive Force Cell for Carbonium Ion StRbilization Energies'
Sir: Substituent effects on thermodynamic properties afford one of the principal building blocks of theoretical organic chemistry. The accepted standard, both for application of theoretical analysis and for analysis of reaction rates, is the substituent effect on the equilibrium process.2 However, there has been an inherent dilemma in the application of such standards. For convincing tests of applications of basic theory to essentially complex organic molecules, substituent effects of substantial (optimum) magnitude are required. Variation of equilibrium constants through ( I ) Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for partial support of this research. (2) L . P. Hammett, "Physical Organic Chemistry," McGraw-Hill Book C o . , New York, N. Y., 1940, Chapters3, 4, 7 , and 9,
where Ro designates unsubstituted trityl cation and R the p-substituted derivative, is obtained from Eo corresponding to this cell reaction (which is negligibly affected by the liquid junction potential). Table I lists values of EO,K e q u i l , and AF2" for'reaction 2. Also listed in Table I are AF30 values for the reaction RC
+ Ro-OH
=
Ro+
+ R-OH
(3)
obtained through the application of the HR acidity function. (3) E. M. Arnett, in "Progress in Physical Organic Chemistry," Vol. I , Editors, S. G. Cohen, A. Streitwieser, Jr., and R. W. T a f t , interscience Publishers, Inc., 1963, p. 223. (3a) J. B. Conant, et ol., J . A m . Ckem. SOC.,47, 1959 (1925); 66, 3752 (1833). (4) Cf.N. Muller and R. S. Mulliken. i b i d . , 8 0 , 3489 (1958). ( 5 ) H. J . Dauben, Jr., L. R . Honner, and K. M. Harmon, J . Org. Chcm., 26, 1442 (1960). (6) J. F. Coetzee, G. P. Cunningham, D. R . McGuire, and G. R. Padmanabhan, Anal. Ckem., 84, 1139 (1962). (7) (a) N. C. Dino, J. J. Jaruzelski, and A. Schriesheim, J. A m . Chem. S O L , 77, 3044 (1955); (h) N. C. Deno and A. Schriesheim, i b i d . , 7 7 , 3051 (1955); ( c ) N. C. Deno and W. L. Evans, i b i d . , 79, 5804 (1957).
Jan. ti, 1964
COMMUNICATIONS TO THE EDITOR
Table I shows striking agreement between our directly measured AFZovalues for reaction 2 in acetonitrile solution and corresponding AF30 values for reaction 3 in aqueous solution. This result then offers compelling contLrmation of the HR acidity function (up to about .50 wt. % acid). The agreement further implies that the free energies to a very major extent measure internal energy effects,8 i e . , relative cation stabilization energies. The substituent effect on the free energy of solvation of the trityl cation is apparently substantially the same in the aqueous as the nonaqueous solvent (in spite of their great difference in dielectric and other properties), a result which must be attributed to the chemical inertness of the R+, R-R, and substituent toward the solvent.3,9810 TABLE I RELATIVESTABILIZATION ENERGIES OF p-SUBSTITUTED TRITYL CATIONS IN ACETONITRILE SOLUTION p-Subst.
Kequil
- E o , v.
AFx', kcal.
AFn', kcal.
H, H, Ha (1.00) (0.000) ( 0 . 0 ) ( 0 . 0 ) F, H, Ha 0.98 ,006 0.1 O.lb ,086 2.0 1.9 CHI, H, H" 3 . 6 X lo-* 2 . 0 x 10-4 ,219 5.1 4.4 OC&, H, Ha 8.8 7.4 3.7 X ,581 OCHz, OCH8, Ha OCH3, OCH3, OCH3" 1.7 X ,518 12.0 1 0 . 2 ,660 1 5 . 2 1 5 . 5 6 . 5 X 10-l2 N(CHa)2, H , H" N(CH312, N(CHa)g, K(CH3)2" 3 . 1 X lo-'* 1.037 23.9 21.8 See ref. 7. * Estimated from u + and p values listed in ref. 7 c . 'R. J . Goldacre and J . N. Phillips, J . Chem. SOC.,1724 (1949).
The small differences (AFz0 is generally about 10% larger than AF3O) may result from the fact that there is greater localization of the positive charge of R + in the C-C than the polar C-0 bond. There are also probably some small effects on the A F o values due to symmetry considerations and on AF," due to solvation of the OH group. The ionization potential of organic radicals to cations also has been a quantity of much theoretical interest.I4 The applicability of gas-phase values to chemical reactions in solution is uncertain, however, for .reasons of solvation energy as well as of possible excess energy content of the gaseous ions. The great contrast between gaseous ionization potentials of the metallic elements and the corresponding aqueous e.m.f. series serves to illustrate how different in character ionization potentials in the gas and condensed phases can be. Literature values15 for the equilibrium constants for ( 8 ) L. G. Hepler, J. Am. Chem. Soc., 86, 3089 (1963). (9) (a) R. W. Taft, ibid., 81, 2965 (1960); (b) N. C. Deno, "Equilibria in Copcentrated Mineral Acid Solutions," in A. F . Scott's "Survey of Progress in Chemistry," Academic Press, New York, N. Y., 1963; (c) R. H . Boyd, J . Am. Chem. SOL, 86, 1555 (1963); (d) E. M. Arnett and R. D. Bushick, ibid., in press. (10) This result may be contrasted with proton transfer equilibria of mand p-substituted benzoic acids in water11 and with 1,3-diphenylguanidine in benzene.12 For m-substituents there is precise fit t o the Hammett u p relationship, but substituent effects are 230% greater in benzene than water. Further, for -R-P-substituents, enhanced resonance effects are observed in water.18 Evidently, the nonequivalence of substituent effects is t o be attributed a t least in part t o the chemical interactions between solvent and the C02H and CO2- functional groups as well as t o different dielectric properties of water and benzene, cf., also, C. D. Ritchie and W. F. Sager, "An Examination of Structure-Reactivity Relationships," in "Progress in Physical Organic Chemistry," 1864, in press. (11) Ref. 2, Chapter 7. (12) M. M. Davis and H . B. Hetzer, J. Res. Nall. BUY.S l d . , 6 0 , 569 (1958). (13) R . W. T a f t , S . Ehrenson, I. C. Lewis, and R. E. Glick, J. Am. Chem. Soc., 81, 5352 (1959). (14) A. Streitwieser, Jr., "Progress in Physical Organic Chemistry," ref. 3, p . 1. (15) (a) Cf. G. W. Wheland, "Advanced Organic Chemistry," 2nd Ed., John Wiley and Sons, Inc., New York, N. Y., 1949, p. 694; (b) J. B. Conant and N. M. Bigelow, J. Am. Chem. Soc., 63, 676 (1931).
+
117
-
the reaction, R. '/s(Ro-Ro) Ro.+ '/z(R-R), in combination with Eo values of Table I, provide for the first time quantitative measures of +substituents on the ionization potentials of trityl radicals in acetonitrile solution. Table I1 lists the ionization potential data so derived. Unfortunately, gas-phase values (such as those recently reported for benzyl radicals) are not presently available for comparison. TABLE I1 RELATIVEIONIZATION P O T E N T I A L S O F P - S U B S T I T U T E D TRITYL RADICALS I N ACETONITRILE SOLUTION p-Subst.
( I - Io), v.
H, H, H F, H, H CH, H, H OCH3, H, H h'(CH3)2, H , H
(0.000) - ,000 - ,043 - ,187 - ,587
A subsequent detailed report will be made of our e.m.f. cell, as well as the results of studies of the temperature coefficients of the e.m.f. values. (16) A G. Harrison, P. Kebarle, and F . P. Losing, i b r d . , 89, 777 (196 1). (17) National Science Foundation Graduate Fellow, 1961-1964.
DEPARTMENT OF CHEMISTRY EVAND. J E N S O N n THEPENNSYLVANIA STATE UNIVERSITY ROBERTW. TAFT UNIVERSITYPARK, PENNSYLVANIA RECEIVED OCTOBER30, 1963
Electronic Interactions in Biradicals. Triplet Equilibrium',
11. A Singlet-
Sir : We recently reported a study of the radicals derived 2,2-bis(3,5-di-t-butyl-4-hydroxyphenyl)propane from (I). The removal of both phenolic hydrogen atoms of I gave an equilibrium mixture of singlet (bonded) and triplet species. In this case the equilibrium favored the singlet tautomer almost completely. We have modified I in order to make formation of the internal bond more difficult, shifting the equilibrium toward the triplet state. The system we chose is formed by replacing the gem-dimethyl group in I with a 2,2'biphenylene group, yielding 11. The ordinary (aluminum chloride-carbon disulfide) FriedelLCrafts reaction of 2,6-di-t-butylphenol with 9,9-dichlorofluorene did not appear successful. However, the addition of a solution of silver oxide in acetic acid containing excess perchloric acid to a solution of the phenol and the dichloride in acetic acid produced I1 in ca. 50y0 yield, m.p. 272-273' from benzenehexane. Anal. Calcd. for C41H5C02: C, 85.66; H, 8.77. Found : C, 85.94, 85.76 ; H , 8.57, 8.77. The infrared spectrum of the product was consistent with the assigned structure, Am,, 2.7 p , as was the ultraviolet spectrum, A$Ela299 mp (log e 2.6), 310 (2.8). An ethereal solution of the bisphenol (11) was readily oxidized by aqueous alkaline ferricyanide in the usual manner3 and gave a dark yellow-green solution. Evaporation of the solvent' with nitrogen gave a dark oil, which crystallized when treated with pentane. The infrared spectrum (carbon tetrachloride) of this product indicated that both biradical (111) and spirodienone (1) Paper I : E. A. Chandross and R . Kreilick. J . Am. Chem. SOL.,86, 2530 (1963). (2) Partial support (at Washington University) of this work by the U. S. Air Force under contract, t h e National Science Foundation, the Petroleum Research Fund of the American Chemical Society, and an equipment loan contract with the office of Naval Research is gratefully acknowledged. (3) E. Muller and K. Ley, Chem. BEY..87, 922 (1954).
