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Thermochemical Properties of Iodoacetone. Intramolecular Electrostatic Interactions in Polar Molecules' Richard K. Solly,2 David M. Golden, and Sidney W. Benson Contribution from the Department of Thermochemistry and Chemical Kinetics, Stanford Research Institute, Menlo Park, California 94025. Received December 22, 1969
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Abstract: Equilibrium constants for the system CH3COCH3 IZe CH~COCHZI HI have been measured spectrophotometrically in the gas phase. From the van't Hoff plot of these equilibrium constants over the temperature range 314-430°, the best straight line obtained by a least-squares fit was log K = (3.84 =t1.1)/2.303R (13.03 f 0.76)/0, where 0 = 2.303RTin kilocalories/mole. Correcting these values to room temperature, using an estimate of E,' = 2.0 gibbs/mol yields = 2.3 gibbs/mol and AHfo2s8 = 12.1 kcal/mol. By combining these with known values of So and AHf"for CH3COCH3,12, and HI, So~g~(CH3COCH~I(g)) = 85.7 gibbs/moland AHfozg8 (CH3COCHzI(g))= - 31.0 kcal/mol are obtained. A(S0298)i,t and A(AHf0298)on substitution of a hydrogen atom by an iodine atom are compared for a series of compounds containing hydrocarbon, ether, and ketone groups. As expected from the atom substitution rule, A(S0298)int is constant at 12.1 gibbs/mol with a small variation of +0.8 gibbs/mol. However, the range of A(AHr'z98) is from 9.6 kcal/mol to 20.7 kcal/mol. An explanation for this large variation is presented as evidence for intramolecular coulombic interaction between partial charges in the polar molecules.
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hermochemical data and kinetic parameters may be obtained from a gas-phase spectroscopic study of the equilibrium3 RH
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RI
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As the number of organic iodides for which the thermochemical properties are known becomes more extensive, it is possible to estimate the corrections required for the atom substitution rule of estimating thermochemical properties. In this paper, we present the thermochemical data for iodoacetone and consider the effect of electrostatic interactions on the enthalpy of organic iodides. Experimental Section Materials. CH3COCH3 was distilled and degassed at liquid nitrogen temperatures prior to use. No impurity could be detected by gas chromatography. Iz was degassed at - 10" and sublimed into a storage vessel attached to the inlet manifold. The CHICOCHZIproduced in the reaction mixture had the same retention time for gas chromatographicanalysis on a silicone oil column as a sample prepared from CICHaCOCHaand HL4 Apparatus and Procedure. This has been described previously.* Equilibrium concentrations were determined from the optical density of 12 at 450-500 nm and from the absorbance change at 230 nm. A pure sample of ICHzCOCH3could not be obtained in the gas phase for optical density calibrations. The absorbance at 230 nm is equal to
where I is the cell path length, [AIo the initial concentration of acetone, [IA] the concentration of iodoacetone, and A[12]the change in 4 concentration, which is assumed to be equal to A[A], [HI], and [IA]. Expressed as rates
where q is a constant at a given temperature.6 Values of q were calculated from eq 3 using AIIz]/Atmeasured in the visible prior to equilibrium. Using q, values of (A[I& were calculated from (AAzro)e after equilibrium had been established. Good agreement was obtained between equilibriumconstants calculated from (AA230)e and those obtained from (A[Iz])e measured in the visible. This agreement supports the assumption that the reaction stoichiometry is unchanged from the initial stages to equilibrium. There was no decomposition of the ICHZCOCH3in the gas phase in the time required to reach equilibrium. At equilibrium, the absorbance was constant for a time of at least five half-lives of the reaction. In this time, there was no pressure change and no noncondensable gases were formed, further supporting the stability of ICHzCOCH3. However, if the reaction time was extended to many multiples of the equilibrium time, some decomposition occurred, as was shown by small changes in the absorbance, the formation of noncondensable gases, and a pressure increase of the reaction mixture.
Results and Discussion Equilibrium data for the system CH3COCH3 I 2 Ft. C H ~ C O C H Z I H I are shown in Table I . Variation of the equilibrium constant, Ke4 = [ICH2COCH3]. [HI]/[CH3COCH3][12], is within 3 0 z for a 100-fold variation in the [CH3COCH3]/[12] ratio. For all runs, [ICH2COCH31 was assumed to be equal to [ H I ] . As the equilibrium lies well on the side of CH3COCH, and 1 2 , no attempt was made to add H I to the initial reaction mixture. This would decrease -A[12] t o the same order of magnitude as the errors in measuring the concentrations. A van't H o f f plot of this data yielded a reasonable straight line (see Figure l), the equation of which was calculated by least squares: log K = (3.84 f 1.1)/2.303R - (13.03 f 0.76)/0, where 0 = 2.303ipT and the errors quoted are standard deviations. No Cpo values are known for ICH2COCH3 and there are insufficient groups for group additivity estimates to be applied. We will use ACPo = (ACp0646 ACp029s)/2 = 2 f 1 gibbslmol for the system ___ RH I2 RI H I , by comparison with ACPo = 2.18 gibbs/mol for R = C H 3 C 0 6 and an estimate of
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(1) This work was supported in part by Grant No. 00353-05 of the Air Pollution Division, Public Health Service, (2) Postdoctoral Research Associate. (3) D. M. Golden and S. W. Benson, Chem. Reo., 69, 125 (1969).
(4) R. K. Solly, D. M. Golden, and S. W. Benson, Int. J . Chem. Kin., 2, 11 (1970).
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may be pressure dependent at the experimental concentrations
