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J. Phys. Chem. 1980, 8 4 , 3270-3272
(7) Pardo, M. P.; Fourcroy, P. 10, 665.
H.;Flahaut, J. Mater. Res. Bull. 1975,
(8) Edwards, J. G.; Gates, A. S.;Sobel, E. S.; Baker, N.; Turner, C.; Wagner, W. Ohio J. Scl. 1979, 79, 186. (9) Viswanadham, P.; Edwards, J. G. Mater. Res. Bull. 1975, 10, 651. (10) Viswanadham, P.; Edwards, J. G. J. Chem. Phys. 1975, 62, 3875. (11) Wiedemeier, H.; Gilles, P. W. J. Chem. Phys. 1965, 42, 2765. (12) Uy, 0. M.; Muenow, D. W.; Ficalora, P. J.; Margrave, J. L. Trans. Faraday SOC. 1968, 64, 2998. (13) Kashkooii, I. Y.; Munir, Z. A. High Temp. Sci. 1972, 4 , 82. (14) Roberts, J. A.; Searcy, A. W. Science, 1977, 796, 525. (15) Starzynski, J. MS Thesis, University of Toledo, Toledo, OH, 1980. (16) Stearns, C. A,; Kohl, F. J. NASA Tech. Note 1974, 07613. (17) Edwards, J. G. J. Vac. Scl. Techno/. 1974, 7 1 , 400. (18) Gates, A. S.;Edwards, J. 0. Rev. Sci. Instrum. 1977, 48, 488. (19) Freeman, R. D.; Edwards, J. G. In “The Characterization of High Temperature Vapors”; Margrave, J. L., Ed.; Wiiey: New York, NY,
1967; Appendix C. (20) Edwards, J. 0.; Freeman, R. D. High Temp. Sci., in press. (21) Pelino, M.; Vlswanadham, P.; Edwards, J. G. J . Phys. Chem. 1979, 8 3 , 2964. (22) Freeman, R. D. In “The Characterkation of High Temperature Vapors”; Margrave, J. L., Ed.; Wiley: New York, NY, 1967; Chapter 7. (23) Voimer, M. 2. Phys. Chem. 1931, Bodenstein Festband, 863. (24) Cater, E. D. In “Techniques of Metal Research”; Rapp, R. A., Ed.; Wiley: New York, NY, 1970; Vol. IV, Part 1, Chapter 2A. (25) Cubicclotti, D. J. Phys. Chem. 1966, 70, 2410. (26) Stull, D. R.; Prophet, H. ”JANAF Thermochemical Tables”, 2nd ed.; US. Government Printing Offlce: Washlngton, D.C., 1971. (27) Keliey, K. K. U.S., Bur. Mines, Bull. 1960, No. 584. (28) Kelley, K. K.; King, E. G. U . S . , Bur. Mines, Bull. 1961, No. 592. (29) Mills, R. C. “Thermodynamic Data for Inorganic Sulphldes, Selenkles, and Tellurides”; Butterworth: London, 1974.
Rate and Mechanism of the Ketene Hydrolysis in Aqueous Solution E. Bothe, A. M. Dessouki, and D. Schulte-Frohllnde” Instltut fur Strahlenchemle lm Max-Planck-Instltut fur Kohlenforschung, 0-4330 Mulhelm/Ruhr, West Germany (Received: March 17, 1980; In Flnal Form: June 18, 1980)
The hydration of ketene in aqueous solution has a rate constant of 44 s-l at 25 “C which is independent of pH in the range 4.4-9.85. The reaction shows a solvent kinetic isotope effect of kH20/kDa0 = 1.9. Activation enthalpy and entropy were measured to be 10.3 kcal mol-l and -16 eu, respectively. In dioxane/water and acetonitrile/water mixtures the rate constants were found to be approximately proportional to the water content. It is concluded that the first step in the hydrolysis is the nucleophilic reaction of water with the unsubstituted ketene.
Introduction The widespread use of ketenes in acylations and cycloadditions, and their implication as intermediates in reactions of various acyl derivatives with nucleophiles, led to the recent interest in the reaction mechanism of ketenes.l+ Various ketenes undergo hydrolysis which leads to formation of the corresponding acids, a reaction which can be described in an overall reaction by eq 1. The rate and RiR,C=C=O
+ H2O
-
RiRzCHCOOH
(1)
the mechanism of reaction 1 were investigated for some substituted ketenes previo~sly.~-l~ These studies included an investigation of the kinetics of spontaneous and acid-catalyzed reactions of dimethylketene in ether at 25 “C with water and alcohols by Lillford and Satchels7They reported that the catalytic effect of an acid is inversely related to its conventional acid strength. The spontaneous and the acid-catalyzed addition of nucleophiles to dimethylketene is assumed to occur via cyclic transition state^.^^^ Bothe et ala9measured the rate constants of hydration of arylketenes. The substrates were generated in situ by a photochemical Wolff rearrangement, and the rates of hydration were followed in aqueous solution by the change in conductivity after a UV flash (eq 2).
High rates of reaction were reported with first-order rate constants of 4 X lo3 and 5 X lo4 s-l for p-methyl- and p-nitrophenylketenes, respectively. These reactions were 0022-3654/80/2084-3270$0 1.OO/O
shown to be independent of pH in the range 4.0-10.8. A solvent kinetic isotope effect of kH20/kDzo= 1.8-2.0 was measured. Tidwell and co-workers’” investigated the hydration of di-tert-butylketene reporting that it undergoes general and specific acid-catalyzed hydration in a 50 5% acetonitrile/ H20 mixture at 25 “C with kH+= 5.57 M-l s-l and kH+/kD+ = 2.8. They suggested that the reaction occurs by ratelimiting proton transfer with a rate constant lo4 greater than that to isobutene. The protonation of the carbon may give the acylium ion (eq 3) and protonation of the oxygen a carbocation (eq 4). (t-Bu)ZC=C=O (t-Bu)ZC=C=O
H+
H+
(t-Bu)&HC+=O
(3)
(~-Bu)~C=+COH
(4)
In the present work, we wish to present some evidence concerning the first step of reaction 1 in the case of the unsubstituted ketene. Experimental Section All solutions were prepared by using six-times-distilled water. pH values were adjusted with HC104 or NaOH. Cyclobutanone (Aldrich),acetonitrile, and dioxane (Merck p.a.) were redistilled (purity 1 99% GC). DzO (MerckUvasol) was redistilled twice from a quartz vessel and had a conductivity of 0.45-0.8 PUS. The conductivity measurements were carried out by using a combination of a UV-flash photolysis setup (half-width of the flash 5 P S ) coupled with an apparatus for conductivity measurements described earlierll (risetime 100 ns). 0 1980 American Chemical Society
Ths JouMl Of physleal Ch9dSby, Vd. 84, NO. 24, 1980 3271 6
Ink 5
4
3
W
e 1. condw3Mly signal (adinate. 20 mVldiv) vs. time (abscissa. 20 ms/div) as observed in flash photolysis of an aqueous solution of cyclobutanone (10.' M) at 21 "C and pH 6.5. 2 h
[.-'I
f 5
3.1
50 40
30
1.2
3.3
3.4
3.6
3.7
20
10 4
5
6
7
8
9
10
PH
ne*r 2. Rate COMtant 01 conducaVny lncreaseas a funCmn of pH as observed after 5-FSUV flash in argcn-salwatedaqueous solutlon mtainlng 3 X
IO-' M cyWuianone at
20 ' C .
Results During photolysis of cyclobutanone in aqueous solution, ketene is formed among other products (eq 5).12J3 This
reaction was used to nroduce the ketene in situ bv flash photolysis. The concentration of cyclobutanone was typically 3.0 X M (e- at 272 nm = 20 M-'c d ) . The kinetics of reaction 1m t h R, = R, = H was followed after the flash by measuring the bulldup of conductivity (Figure 1)caused by the diasociated acetic acid. The dissociation of acetic acid is complete within a few microseconds. The half-life of the conductivity increase a t natural pH is 16 f 10% nu at 20 'C, from which a rate constant of 44 s-I for reaction 1 is calculated. On varying the light intensity by a factor of 35, we observed no change in the buildup rate of conductivity. This, as well 88 kinetic analysis of the time dependence of the conductivity buildup, indicated that the reactions were of strictly first or pseudo first order. After reaching a maximum the conductivity remained constant for at least 30 8.
The half-life (t1,3of the conductivity buildup did not depend on the concentration of cyclobutanone (range: 3 x 10-3-1 x 10-1 M-'. No difference was observed whether the aqueous solutions of cyclobutanone were saturated with argon or oxygen. In most experiments the solutions were saturated with argon. No change in the half-life ( t y 2 )of the conductivity buildup was observed by variation of pH in the range of 4.4-9.85 (Figure 2). Because of the intrinsic conductivity
*I. HtO lrolurnc)
npUn 4. Rate cxmtanl of cardu*)vlty hwxease as a hnctlon of Um -Of the solvent ((0) H,OlacetonlbLt; (0)H&/dbXane) a8 observed in argcn-satuated soltdbns containing 3 X IO-* M c y c b b u t a m at 21 O C .
of solutions with 4.4 > pH > 9.85 and the limited intensity of our W flash, it was not possible to expand further the pH range of our investigation. The half-life of the conductivity buildup was shown to be dependent on temperature. An activation enthalpy of 10.3 kcal/mol and an activation entropy of -16 eu were calculated from the Arrhenius law (temperature range: 2.5-52 'C) (Figure 3). By using mixtures of dioxane/water and acetonitrile/ water instead of pure water (range: 045% organic solvent per volume), we found the rate oone.tant of the conductivity buildup ( k = In 2/tIl2) to be approximately proportional to the water concentration in the mixture (Figure 4). Replacing H 2 0 by D20 in the solutions, we found a = 1.9. solvent kinetic isotope effect of k,,/k,
Discussion There are two plausible mechanisms for the hydration of ketenes. The fmt is a proton catalysis with protonation of either carbon (eq 6) or oxygen (eq 7)as a primary step, followed by reaction with water and elimination of a proton.
3272
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J. Phys. Chem. 1980, 84, 3272-3273
H+
H&=C=O
H,C=C=O
-
CH3Cf=0
specific solvent effect is contributing to hydration of the unsubstituted ketene, since the first-order rate constant of the reaction was approximately proportional to water concentration in the mixture. A strongly polar transition state has been proposed for the spontaneous hydrolysis of acetyl fluoride16and acetic anhydride,l’ whereas less polar cyclic transition states have been suggested for the spontaneous hydrolysis of dimethylketene.7i8 The size of the observed solvent kinetic isotope effect for ketene is in the range reported for neutral hydrolysis of various compounds.18
(6)
H+
CH,=C+OH (7) A mechanism according to eq 6 or 7 was suggested to occur in the hydration of di-tert-butylketene in aqueous solution.” Since we found no influence of pH on the rate of acetic acid formation from unsubstituted ketene, reactions 6 and 7 are not considered to be the rate-determining steps for this case in the pH range studied. Because of the expected high rates for the irreversible reaction of water with the cations formed in eq 6 and 7, the latter reactions are not assumed to be the first steps in the hydrolysis. The second possibility for the first step of the hydrolysis is the uncatalyzed nucleophilic addition of water to ketene, whereby a ketene hydrate may be formed as an intermediate. In the case of the unsubstituted ketene, this spontaneous hydrolysis takes place with a rate constant of 0.79 M-l s-l. Comparatively, Tidwell and co-workers found a rate constant of less than lo4 M-l s-l for the uncatalized hydrolysis of di-tert-butylketene. This at least 106-fold decrease in the rate constant of the neutral hydrolysis is surprisingly large and is explained to be due to steric and to electronic factors introduced by the two tert-butyl groups. This low rate constant for the neutral hydrolysis renders the proton-catalyzed hydrolysis of di-tert-butylketene observable.1° It is expected that, above and below the pH range studied in the present work, proton- or base-catalyzed hydrolysis appears also with the unsubstituted ketene. Because of the limitation of the conductivity method, these pH values were not accessible. For the 4-substituted phenylketenes the reaction rate constants were correlated with the cpnvalues of the substituents leading to p = +1.2. This result as well as the independence of the hydrolysis rate constants from pH suggest a nucleophilic addition of water to be the primary step in the hydrolysis also with these ketene^.^ The results of the experiments with dioxane/water and acetonitrile/water mixtures as solvents show that no
Acknowledgment. We thank Professor Dr. C. v. Sonntag, Dr. 5. Steenken, and Dr. G. Koltzenburg for valuable discussions.
References and Notes (1) R. N. Lacey In “Chemistry of the Alkenes”, S. Patai, Ed., Wiley, New York, 1964,p 1161. (2) E. Schaumann, Chem. Ber., 109, 906 (1976). (3) N. J. Turro, M.-F. Chow, and Y. Ito, J . Am. Cbem. SOC.,100,5580 (1978). (4) W. E. Truce and P. S. Bailey, Jr., J. Org. Cbem., 34, 1341 (1969). (5) A. Kivinen in “The Chemlstry of Acyl Halides”, S. Patai, Ed., Wiley,
New York, 1972,p 177. (6) N. F. Yaggi and K. T. Douglas, J. Am. Chem. Soc., 99,4844 (1977). (7) P. J. Lillford and D. P. N. Satchell, J. Chem. SOC.B , 889 (1968). (8) D. P. N. Satchell and R. S. Satchell, Cbem. Soc. Rev., 4, 231 (1975). (9)E. Bothe, H. Meiir, D. Schulte-Frohlinde, and C. von Sonntag, Angew. Cbem., Int. Ed. Engl., 15, 380 (1976). (10) Shaik Habibul Kabir, Hani R. Saikaly, and Thomas T. Tidwell, J . Am. Cbem. SOC.,101, 1059 (1979). (11) E. Bothe, Dissertation, Ruhr-Unlversitat Bochum, 1976. (12) R. F. Klemm, D. N. Morrison, P. Gilderson, and A. T. Blades, Can. J. Chem., 43, 1934 (1965). (13)T. Howard McGee, J. Phys. Cbem., 72, 1621 (1968). (14) G. P. Laroff and H. Fischer, Helv. Cblm. Acta, 58, 2011 (1973). (15) E. Bothe and D. Schulte-Frohllnde, submitted for publication. The elimination rate constant of HO,. from CH3CH(OH,)02 aqueous solutlon is Increased by a factor of Z lo4 If the CH3group Is replaced by an OH group. (16) C. A. Burton and J. H. Fendler, J. Org. Cbem., 31, 2307 (1966). (17) A. R. Butler and V. Gold, J. Cbem. Soc., 2305 (1961). (18) P. M. Laughton and R. E. Robertson in “Solute-Solvent Interactions”, J. F. Coetzee and C. D. Ritchie, Eds., Marcel Dekker, New York, 1969, p 399.
Hydrogen Production from the Reaction of Solvated Electrons with Benzene In Water-Ammonia Mixtures R. R. Dewald,” S. R. Jones, and 6. S. Schwartr Department of Cbemistv, Tufts University, Medford, Massachusetts 02 155 (Recelved: September 4, 1979; In Final Form: August 4, 1980)
Product analysis data for the reaction of the ammoniated electron with benzenewater mixtures in liquid ammonia show that the dominant product is evolved hydrogen and not 1,4-cyclohexadiene.
The hydrogenation reactions of unsaturated systems with metal-ammonia reagents (Birch reductions) are of considerable synthetic ~tility.l-~However, surprisingly little is quantitatively known about the mechanism of the Birch reduction^.^ The first reported kinetic studies of the hydrogenation reactions were for the reduction of benzene to 1,4-cyclohexadiene by alkali metals and alcohols in liquid a m m ~ n i a .The ~ overall reaction is C6H6+ 2eam-+ 2ROH
-
C6H8+ 2RO-
(1)
0022-3654/80/2084-3272$0 1 .OO/O
where R = H or C2Hb The mechanism originally suggested by Birch6 appeared to be in agreement with the kinetic data reported by Krapcho and Bothner-ByP However, in a later report, Eastham and eo-workers7 give strong experimental evidence that the reaction of the metal (or solvated electron) with the weak acid (alcohol in their experiments) competes with the benzene reduction. The competing reaction is 2eam-+ 2ROH 2RO- + H2 (2)
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0 1980 American Chemical Society
