2644
JERALD A. DEVORE AND H. EDWARD O'NEAL
Heats of Formation of the Acetyl Halides and of the Acetyl Radical by Jerald A. Devore Department of Chemistry, University of Nevada, Reno, Nevada 89607
and H. Edward O'Neal Department of Chemistry, San Diego State College, Sun Diego, California 02116
(Received December 18, 1968)
Heats of hydrolysis of acetyl chloride, bromide, and iodide have been redetermined calorimetrically, and vapor pressures of the halides have been measured over moderate temperature ranges in order to determine heats of vaporization and heats of formation of the liquid and vapor species. In conjunction with kinetic data, the heat of formation of the acetyl radical has been determined which, together with the results of other experi0.4 kcal/mol and firmly establishes a number of ments, places the value for A\Hro(CH3C0,298°K) at -5.8 controversial bond dissociation energies.
*
Introduction Carson and Skinner' determined heats of formation for the liquid acetyl halides by measuring their heats of hydrolysis calorimetrically. Heats of formation of the vapor species were obtained using Trouton's rule estimates for the heats of vaporization. Their values in kilocalories per mole for acetyl chloride, bromide, and iodide, respectively, are aHf"(298, liq) = -65.78 =t 0.02, -53.92 f 0.07, -39.75 f 0.03; AHr"(298, g) = -59.0, -46.6, -31.8; 4Hvap' = 6.8,7.3, 7.9.2 Walsh and Benson3 recently measured the equilibrium constant for the reaction CH3CH0
+ I,
CHaCOI
+ HI
(1)
spectrophotometrically at 208'. Using known and estimated entropies and heat capacities for the species of reaction 1, they arrived at a value of -30.3 t-13 0.5 kcal/ mol for the heat of formation of acetyl io.dide vapor. We have reinvestigated heats of hydrolysis of acetyl chloride, bromide, and iodide and have determined their heats of vaporization from vapor pressure measurements in order to obtain accurate heats of formation of the gaseous species. Of particular interest was the confirmation, in an independent manner, of the heat of formation of acetyl iodide vapor. This heat of formation is important since in conjunction with the kinetic data of O'lSeal and Benson4 on the thermal reaction of acetyl iodide with hydrogen iodide, the heat of formation of the acetyl radical can be calculated. By combining the forward and back activation energies for the reaction
+
+
CHIC01 I 12 CHSCO (2) O'Neal and Benson determined a reaction enthalpy of 14.6 t-13 0.5 kcal/mol at 217". With known5 thermodynamic functions for I., Iz,and CHSCOI (see Table 111) and estimated3 heat capacities for CH360, the heat of format,ion of the acetyl radical in kilocalories per mole is given by 4Hr0zss(CH360)
=
25.6
+ AHr'zss(CH&OI,
The Journal of Physical Chemistry
g)
(3)
A reliable determination of the heat of formation of acetyl iodide therefore determines the acetyl radical heat of formation which in turn can be used to accurately establish many bond dissociation energies in organic compounds. Experimental Section The Joule-type solution calorimeter consisted of a constant temperature box in which was suspended a 660-ml Dewar flask with a well-fitted cork cover. A glass stirrer, calorimeter heater, thermistor probe, and sample bulb passed through the cork cover and into the flask. The thermistor formed one leg of a Wheatstone bridge, and temperature was measured by the recorded bridge imbalance which has been standardized against an NBS calibrated mercury in glass thermometer. All resistors in the bridge and heater circuits were accurate to within 0.05%. Voltages were measured with a E(-3 Leeds and Northrup potentiometer to within =kO.01%. Temperature calibration consisted of 92 temperaturemillivolt (recorder scale deflection) data points between 24 and 27" curve fitted to a second-degree polynomial by the method of least squares. The uncertainty in calorimeter temperature was t-130.003". The heat capacity of the calorimeter system was determined after each run by noting the temperature change resulting from the input of a known amount of electrical energy. The uncertainty in the energy input was EtO.O4%. The definition 1 cal = 4.1840 abs joules was used to convert energy measurements into calories. The heat capacity was assumed to be constant over the range of temperature change. For dilute aqueous solutions near room temperature, a 1" change in tempera-
s. Carson and H. A. Skinner, J. Chem. SOC.,936 (1949). (2) Walsh and Benson (ref 3) have suggested that for the heat of vaporization of acetyl iodide, an additional 1.3 kcal/mol is required to correct the Trouton's rule estimate to room temperature. (3) R. Walsh and S. W. Benson, J . Phys. Chem., 70, 3761 (1966). (4) H. E. O'Neal and 5. W. Benson, J . Chem. Phys., 37, 540 (1962). (6) "JANAF Interim Thermochemical Tables," D. R. Stull, Ed. Dow Chemical Co., Midland, Mich., 1961-1968. (1) A.
HEATSOF FORMATION OF ACETYL HALIDESAND ACETYLRADICAL ture results in about a 0.02% change in heat capacity. The greatest relative error was thus in the measurement of temperature, and the over-all uncertainty in the heat of hydrolysis measurements was estimated to be k0.370. The calorimeter was checked for systematic errors by measuring the heat of neutralization of HzSOain aqueous NaOH, a reaction whose enthalpy is known fairly a c c ~ r a t e l y . ~ -For ~ the reaction
+
H2S04. (0.2284)11~0 2n'aOH. (604.0)HzO +
NazSOa.(1210.2)H20
+ 2Hz0
2645
tized decomposition of the sample. Temperature was measured to within k0.03" using a mercury in glass thermometer graduated in 0.1" intervals. The average error in nulling the Bourdon gauge and reading the manometer was k0.4 mm.
Results Heats of hydrolysis for the reaction CHSCOX
+ sHzO + (CHSCOOH
(4)
an average of four determinations gave AHT= 48.36 =t 0.20 kcal/mol. The heat of reaction calculated from existing thermodynamic data6-* was 48.23 kcal/mol. The difference between the two values is less than the uncertainty in the calorimetric measurements or the uncertainty in the existing thermodynamic data for the reaction.'" Reagent grade acetyl chloride, bp 50-52", and acetyl bromide, bp 75-77", were purified by repeated bulb-tobulb distillations under high vacuum. Calorimeter samples weighing between 0.5 and 2 g taken from the middle fraction of the final distillate were condensed into thin-walled Pyrex bulbs which were then sealed under about 0.5 atm of dry nitrogen. Samples of the halides were analyzed gravimetrically as the silver salts. An average of three determinations each gave 99.98 0.08 and 99.96 It 0.06% of theoretical for the chloride and bromide, respectively. Acetyl iodide was prepared by the method used by O'Seal and Benson.* The middle fraction of the final distillate was stored under vacuum at Dry Ice temperature and in the absence of light. The product was pale yellow, and no further color change was observed during several weeks of storage under these conditions. Calorimeter samples were prepared from a portion of the stored material which had been further purified by several bulb-to-bulb distillations and ranged from colorless to very pale yellow. It was necessary to make calorimeter runs with the iodide samples immediately after preparation since, even in the absence of light, appreciable coloration indicative of decomposition was observed in the sample bulbs after about 1 hr. Hydrolyzed samples of the iodide were titrated for Ia- using standard iodate and thiosulfate solutions and found to be 99.87% of theoretical. The impurity was most likely acetic acid. Vapor pressures of the halides were measured with a specially constructed Bourdon gauge nulled against a mercury manometer. I n operation, about 5 g of sample was further purified by several vacuum bulb-to-bulb distillations and condensed into the sample chamber of the gauge which was then immersed in a salt water-ice bath. The entire apparatus was wrapped with asbestos paper to help control temperature and to prevent light-sensi-
*
+ HX) (X - 1)HzO *
(5)
were measured with the calorimeter, and heats of formation of the liquid acetyl halides a t 298°K were calculated from the equation
AHf"(CHSCOX, 1)
AHf"(CH3COOH
+
HX).(x - 1)HzO - AHr"(Hz0)
- AH,
(6)
where X = C1, Br, I. Heats of formation of the reaction products were calculated using data taken from RossinP for solutions of the individual acids. For the concentrations of interest, heats of mixing are negligible. The heat of formation used for water was -68.317 kcal/mol." Results of these experiments are shown in Table I. For liquid acetyl chloride and bromide, the heats of formation given in Table I are in almost exact agreement with the results of Carson and Skinner.' For acetyl iodide, the results differ by 0.5 kcal/mol. Considering the difficulty of avoiding at least some decomposition in the iodide samples prior to calorimetric determinations, it seems likely that the acetyl iodide heat of formation determined in this work represents the best lower bound for the exact value. This cannot be stated with absolute certainty, however, so that we have averaged our resalt!; with those of Carson and Skinner to give AHf02,,(CH&OI, 1) = -39.48
* 0.27 kcal/mol
Vapor pressure measurements for the three halides are given in Table 11. The vapor pressure of acetyl chloride was measured over the temperature range -0.84 to 50.58'. A Ieast-squares fit of the data gives log P (mm) = 5.880 - 7.352 X 102T-'
- 7.605 X
1 0 4 P 2 (7)
(6) F. D. Rossini and D. D. Wagman, et al., "Selected Values of Chemical Thermodynamic Properties," NBS Circular 500, Feb 1, 1952. (7) H. hl. Papee, W. J. Canady, and K. J. Laidler, Can. J. Chem., 34, 1677 (1956). (8) W. F. Giauque, E. W. Hornung, J. E. Kunnler, and T.R. Rubin, J . Amer. Chem. SOC.,8 2 , 62 (1960). (9) S. R. Gunn, J . Phys. Chem., 69, 2902 (1965). (10) Gunn (ref 9) has presented a detailed discussion of the problems inherent in this and other reactions commonly used as standards for solution calorimetry. (11) G. N. Lewis and M. Randall, "Thermodynamics," 2nd ed, revised by K. S. Pitaer and L. Brewer, McGraw-Hill Book Co., Inc., New York, N. Y., 1961, p 672.
Volume 73, Number 8 August 1909
2646
JERALD A. DEVORE AND H. EDWARD O'NEAL
Table I: Calorimetric Results for the Heats of Hydrolysis and Heats of Formation of the Liquid Acetyl Halides Moles of HpO/mole of CHaCOX
wt of CHaCOX,
Moles of
g
CHaCOX
1 2 3 4
1.212.9 1500.4 1115.1 836.7
0.7165 0.4344 0.5845 0.7790
Acetyl Chloride 0.009128 24.740 0.005534 25.351 0,007446 24.991 0.009924 24.888 AffrOm(CHaCOC1,1) = -65.77
5 6 7 8
677.3 614.1 797.0 1004.6
2.0097 1.6975 1.2810 1.0172
9 10 11
896.1 997.9 1913,8
1.5747 1.4141 0.7373
Run
a
Tz.
C, (system),
-AHrr
OC
oal/"C
- AHPzss,
kcal/mol
kcal/mol
220.86 164.44 168.40 167.93
22.68 22.30 22.77 22.56
65.66 66.10 65.57 65.76
Acetyl Bromide 0.016345 24.751 26.475 0.013806 24.756 26.683 0.010418 24 693 26.148 0.008273 24.777 25.927 Afftom(CHaCOBr, 1) = -53.92 f 0.07 kcal/mol
221.12 168.06 165.90 167.38
23.32 23.46 23.16 23.28
53.88 53.75 54.09 54.00
Acetyl Iodide 0.009266 24.744 25.969 0.008321 25.215 26.311 0.004338 25.161 25.730 Affrom(CH&OI, 1) = -39.21 =k 0.05 kcal/mol"
170.29 169.88 171.63
22.50 22.37 22.52
39.20 39.33 39.20
TI, O C
25.677 26.101 25.998 26.222 =I= 0.11 kcal/mol
I
Corrected for 0.13% impurity.
Table I1 : Vapor Pressure Data for the Acetyl Halides Acetyl chloride T, OK P, mm
272.31 272.91 273.43 279.56 283.96 286.85 288.97 295.00 297.15 299.29 305.08 308.66 312.23 317.38 319.98 323.73
144.8 146.6 150.0 186.0 219.4 244.3 264.3 327.1 352.8 379.8 455.7 502.7 558.3 639.0 693.0 756.3
Scetyl bromide P , mm
T ,OK
275.55 282.84 289.21 294.12 297.79 301.97 306.10 310.53 313.63 318.53 324.95 328.84 333.65
44.5 62.5 81.3 102.9 120.1 141.2 159.5 198.5 218.2 286.7 314.9 357.3 404.4
Acetyl iodide P , mm
T ,OK
276.63 280.04 286.03 289.31 295.32 301.35
9.9 12.7 17.3 20.9 28.5 37.6
where T = OK. The heat of vaporization at 298°K computed from the Clapeyron equation together with eq 7 is AH,,, = 5.70 f 0.04 kcal/mol. The uncertainty attached to the value is an estimate of the probable error and was calculated with a met,hod given by Margenau and Murphy12by assuming a linear relationship between log P and l/T. Equation 7 gives a normal boiling point of 50.2" for acetyl chloride and a boiling point of 50.8" at 769 mm, in good agreement with the value of 51.2' reported by Carson and Skinner at that pressure. Kireev and Popov13 previously measured the vapor pressure of acetyl chloride from -20 t o 50". The heat The Journal of Physical Chemistry
of vaporization computed from their data is 7.80 kcal/ mol at 298°K. The vapor pressure of CH3COC1has also been measured by Greenwood and Wade14over the temperature range -37 to 18". A least-squares fit of their data to an equation similiar to eq 7 gives 6.96 kcal/mol for the heat of vaporization at 298°K. In addition, Mathews arid Fehland15 measured the heats of vaporization of CH3COCl and CH3COBr directly at the normal boiling points. Their values for the chloride and bromide, respectively, are 6.84 and 7.25 kcal/mol. I n the absence of reliable heat capacities of the liquid species, however, it is difficult to accurately adjust these values t o room temperature. Since there are fairly large differences between the reported values, it was decided t o use the heat of vaporization determined in this work for computing the heat of formation of acetyl chloride vapor. However, further vapor pressure measurements may be necessary to confidently establish the heat of vaporization. The vapor pressure of acetyl bromide was determined over the temperature range 2.40-60.48' and the resulting least-squares fit of the data gives log P (mm) = 5.350 - 4.018 X 102T-l - 1.707 X 105T-2 (8) (12) H.Margenau and G. Murphy, "The Mathematics of Physics and Chemistry," 2nd ed, D. Van Nostrand Co. Inc., New York, N. Y., 1956,p 616. (13) V. A. Kireev and A. A. Popov. Zh. Obshch. Khim., 5, 1399 (1935). (14) N. N. Greenwood and K. Wade, J . Chem. SOC.,1527 (1966). (16) J. H.Mathews and P. R. Fehland, J. Amw. Chem. SOC.,53, 3212; (1931).
2647
HEATSOF FORMATION OF ACETYL HALIDESAND ACETYLRADICAL The reported vapor pressures of acetyl bromide were terminated at 60.48" because, above this temperature, a fall-off in vapor pressure increase with increasing temperature was observed. The most likely explanation is that the bromide undergoes surface-catalyzed reactions above 60" to form less volatile products which lower the vapor pressure by a Raoult's law colligative property effect. The heat of vaporization for acetyl bromide calculated from eq 8 is 7.08 f 0.07 kcal/mol a t 298°K. The vapor pressure of acetyl iodide was measured from 3.48 to 29.88' and the data fitted to the polynomial log P (mm) = 1.383
+ 1.888 X
Table I11 : Thermodynamic Functions for the Acetyl Halides"
1 0 3 ~ 4- 5.515
x
1 0 5 ~ 4 (9)
As with acetyl bromide, the measurements were terminated a t 29.88' because of vapor pressure fall-off. This behavior, nevertheless, is in accord with expected stabilities. Acetyl chloride is known to be quite stable, even a t its boiling point, whereas acetyl iodide, as has often been reported, becomes appreciably colored after standing at room temperature for a short time. The heat of vaporization computed for acetyl iodide from eq 9 is 8.29 f 0.12 kcal/mol at 298°K. Heats of formation for the gaseous acetyl halides resulting from these measurements are AHr2s,(CH3COC1,g)
=
-60.07
the functions for the entire rigid molecule, Le., with internal rotation frozen. Contributions to the thermodynamic functions arising from hindered internal rotation were calculated using the procedure of Pitzer and Gwinn.lg Functions for the other acetyl halides were also computed in a similiar manner in order to provide a fairly complete set of thermodynamic data for the gaseous species. The results of the calculations are shown in Table 111.
f
0.12 kcal/mol
AHr,,,(CH3COBr, g) = -46.84 i 0.10 AHr2,,(CH3COI, g) = -31.19
f
0.39
To determine the heat of vaporization to the ideal gas, the ratio of vapor fugacity to liquid activity, f,,Jal, rather than vapor pressure must be used in the Clapeyron equation (cf. ref 11,p 533). For the acetyl halides at 1atm pressure and near room temperature, f g / Pis estimated to differ from unity by about 1%. As a result, the heats of vaporization to the ideal gas may differ from those given above by amounts roughly equal to the assigned uncertainties of the values. The heats of formation given for the vapor species should then differ from the standard-state values by less than the over-all uncertainties of the measurements.
Discussion The heat of formation determined by Walsh and Benson3for acetyl iodide was obtained by measuring the Gibbs free energy of reaction 1 at 208" and is dependent on bond additivity rule estimates of the entropy and heat capacity of CHsCOI(g). Ramsey and Ladd16 recently obtained the infrared spectrum of acetyl iodide and have made assignments of the fundamental vibration frequencies. I n addition, Moloney and Krisher" have determined rotational constants and the barrier to internal rotation for CH3COI from microwave data. Using the results of these authors, we have computed thermodynamic functions for CH3COI(g). The rigidrotor-harmonic oscillator model's was used to calculate
OK
gibbs/ mol
Ha - Ho', kcal/ mol
298.15 300 400 500 600 700 800
70.56 70.66 75.71 80.18 84.23 87.94 91.37
298.15 300 400 500 600 700 800
298.15 300 400 500 600 700 800
so T,
-(a0
-
Eoo)/ T, gibbe/ mol
Cp0, gibba/ mol
CHaCOCl(g)b 3.529 3.559 5.321 7.330 9.554 11.960 14,525
58.70 58.80 62.40 65.52 68.31 70.85 73.21
16.25 16.30 18.91 21.22 23.21 24.81 26.27
73.63 73.73 78.91 83.46 87.55 91. .28 94.73
CH&OBr(g) 3.663 3.694 5.449 7.539 9.788 12.216 14.793
61.34 61.42 65.16 68.39 71.24 73.83 76.24
16.76 16.81 19.29 21.50 23.42 25.05 26.46
75.69 75.79 81.07 85.67 89.81 93.48 97.05
CHsCOI(g) 3.753 3.785 5.626 7.695 9.970 12.418 15.019
63.10 63.18 67.01 70.28 73.20 75.74 78.28
17.13 17.18 19.60 21.77 23.64 25.24 26.62
9
References for vibrational frequencies and structural constants are as follows: CH&OCl, ref 16, 20; CHsCOBr, ref 16, 21, 22; CHsCOI, ref 16, 17. Thermodynamic functions for CHsCOCl (g) have been reported previously by J. Overend, R. A. Nyquist, J. C. Evans, and J. Potts, Xpectrochem. Acta, 17, 1205 (1961).
'
The entropy calculated for CH3COI(g) at 298°K is 75.69 gibbs/mol. This value agrees closely with the (16) J. A. Ramsey and J. A. Ladd, J . Chem. SOC.,B,711 (1968). (17) M. J. Moloney and L. C. Krisher, J . Chem. Phys., 45, 3277 (1966). (18) Ref 9, Chapter 27. (19) K. 8. Pitzer and W. D. Gwinn, J . Chem. Phys., 10, 428 (1942). (20) K. M. Sinnott, ibid., 34, 851 (1961). (21) L. C. Hall and J. Overend, Spectrochirn. Acta, 32A, 2535 (1967). (22) L. C . Krisher, J . Chem. Phys., 33, 1237 (1960). Volume 73, Number 8 August. 1969
2648
JERALD A. DEVOREAND H. EDWARD O'NEAL average of these two independent determinations gives a mean value of
Table IV : Thermodynamic Data Ho481
Species
CHaCHO
I* CHaCOI HI
-
AHtoZg,(CH&O,g) = -5.8
AXfo,g8,
S029e,
8'481,
Ho9w,
kcul/ mol
gibbel mol
gibbs/
ked/
mol
mol
Ref
-39.67 14.92 -30.98 6.30
63.05 62.28 75.69 49.35
70.40 66.57 84.12 52.70
2.841 1.618 3.535 1.282
24, u 5 b 5
Calculated using structural constants of ref 23 and fundamental vibration frequencies assigned by Pitaer and Weltner, ref 24. See text. a
With the heat of formation of the acetyl radical firmly established, a number of bond dissociation energies can now be computed (see Table V). Table V: Bond Dissociation Energies a t 298°K D(CHaC0-B),
Compound
CHaCOCHa CHaCHO CHaCO-COCHa CHaCOOH CHsCOCl C H aC 0Br CHiCOI
bond additivity rules estimate3 of 76.0 gibbs/mol. The experimentally determined Gibbs free energy of reaction 1 together with the thermodynamic data of Table IV yields a heat of formation for acetyl iodide of AHf0,,,(CH3COI,g) = -30.78
f
0.42 kcal/mol
a
where the uncertainty is an estimate of the over-all error. Since this value was obtained in a manner independent of the heat of formation of CHaCOI determined in this work and both agree within error limits, we recommend a mean value of AH~",,,(CH3COI,g) = -30.98
f
0.6 kcal/mol
Walsh and Benson3 also determined an enthalpy of 14.3 f 1.2 kcal/mol a t 208" for the reaction
1.
+ CHSCHO I_ CH36O + H I
(10)
From this result they computed a value for the acetyl radical heat of formation of -6.2 f 1.2 kcal/mol. An
The Journal of Phvsical Chemistry
a Ref 25. Ref 28.
Products
CH3CO CHaCO CHaCO CHaCO CHaCO CH&O CHaCO
This work.
+ CH3 +H + CHaCO + OH + C1 + Br +I
Ref 26, 27.
kcal/mol
Ref
80.0 86.0 67.0 108.0 83.2 67.8 50.6
a, b, c a, b, d
b, e a, b, d
b, d b, d b, d
Ref 11, pp 672-686.
Acknowledgments. The authors wish to express their thanks to the National Science Foundation and to the U. S. Air Force Office of Scientific Research (AFOSR68-1354) for partial support of this research.
f 0.20 kcal/mol
Substituting this value into eq 3 yields a heat of formation for the acetyl radical of AHfo,,(CH360, g) = -5.4
f 0.4 kcal/mol
(23) R. W. Kilb, C . C . Lin, and E. B. Wilson, Jr., J. Chem. Phys., 26, 1695 (1957). (24) K. S. Pitzer and W. Weltner, J . Amer. Chem. SOC.,71, 2842 (1949). (25) 8. W. Benson, "The Foundations of Chemical Kinetics," McGraw-Hill Book Co., Inc., New York, N. Y., 1960, pp 662-664. (26) G. C. Fettis and A. F. Trotman-Dickenson, J. Chem. SOC.,3037 (1961). (27) D.M. Golden, R. Walsh, and 8. W. Benson, J. Amer. Chem. Soc., 87, 4053 (1965). (28) H. E. O'Neal and 8. W. Benson, J. Chem. Phys., 36, 2196 (1962).
