61 12
Kinetics of the Gas Phase Reaction of Pentafluoroethyl Iodide with Hydrogen Iodide. Enthalpy of Formation of the Pentafluoroethyl Radical and the 7r Bond Dissociation Energy in Tetrafluoroethylene E-Chung Wu and A. S. Rodgers* Contribution from the Department of Chemistry, Texas A & M University, College Station, Texas 77843. Received February 17, 1976
-
Abstract: The kinetics of the gas phase reaction of pentafluoroethyl iodide with hydrogen iodide have been studied spectrophotometrically from 478 to 560 K. The rate-determining step was found to be C2F5I I C2F5 12 for which log ( k / L mol-' s-') = (1 1.2 f 0.3) - (16.8 f 0.8)/0 (0 = 2.3RT in kcal/mol) was determined. This result was combined with established thermochemical data to yield the C-I bond dissociation energy, DHO(CF3CF2-I) = 52.5 1 kcal/mol, and the enthalpy of formation of pentafluoroethyl, AHfo(CF$F2,g,298) = -213 f 1.3 kcal/mol. This also led to an evaluation of the T bond dissociation energy in tetrafluoroethylene, DTO(CF~=CF~)= 52.3 f 2 kcal/mol. This value is shown to be in good agreement with the thermochemical data on tetrafluoroethylene.
+
In 1949, Lacher e t a1.I showed that the heat of addition of chlorine to tetrafluoroethylene was nearly 14 kcal/mol more exothermic than that t o ethylene, and it has since been generally found that addition reactions to tetrafluoroethylene a r e more exothermic than those to e t h ~ l e n e . This ~ . ~ has been almost exclusively attributed t o the destabilization of the C-C T bond by fluorine s u b ~ t i t u t i o n , though ~-~ recently some evidence has been presented indicating that u bond stabilization may be important too.' In this paper, we wish t o report on a kinetic study of the gas phase reaction of pentafluoroethyl iodide with hydrogen iodide, from which we derive t h e heat of formation of the pentafluoroethyl radical and, thereby, t h e ?r bond dissociation energy in tetrafluoroethylene.8 W e shall further show that this value of the T bond energy is in quantitative agreement with t h e thermochemical data for tetrafluoroethylene. Experimental Section
Pentafluoroethyl iodide was obtained from the Pierce Chemical Co. and was vacuum distilled at dry ice temperatures before use. Gasliquid chromatographic analysis on a 0.6 X 305 cm column packed with 30% dimethylsulfolane on Chromosorb W indicated that the sample, after distillatio., was better than 99% pentafluoroethyl iodide. Hydrogen iodide, obtained from the Matheson Gas Co., was similarly distilled from a dry ice bath to a liquid N2 trap and degassed at liquid N2 temperatures before use. Iodine, obtained from the J. T. Baker Chemical Co., was sublimed in vacuo before use. The kinetics of the reaction of C2FsI with HI were followed spectrophotometrically by observing the appearance of 12 at 500 nm. The stoichiometry of the reaction was established from measurements at 225,270, and 500 nm. Consequently, the absorption constants, a = cjRT in OD Torr-', were determined for C ~ F S IHI, , and 12 at these wavelengths and at each reaction temperature. The experimental apparatus and procedure has been described in detail previously and will not be repeated here.9
Results T h e reaction of CF3CF2I with H I is expected t o follow the well-established mechanism given by reactions 1 t o 4.1°
12
+ M e 21 + M
KI,
CF3CF2
+ HI
4
CF3CF2H
+I
Journal of the American Chemical Society
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(394)
98:20
/
*
for which the overall reaction is: CF3CF2I
+ HI
CF3CF2H
+ I2
(5)
T h e stoichiometry of this reaction was established a t temperatures from 500 t o 560 K by measuring the I2 formed a t 500 nm, the H I and CF3CF2I consumed a t 225 and 270 nm, and the overall pressure change during t h e reaction a t various reaction times. These data are summarized in Table I and, within experimental error, established reaction 5 a s the only process taking place. A steady-state treatment of t h e mechanism, neglecting reaction 4, gives
In order to obtain the rate constants kl and kdk3, this equation was rearranged to
Thus, a plot of [HI]/[I*] vs. [RI] [HI]/([12]1/2R~2) should give a straight line with slope equal t o k,K1,1/2 and intercept equal to k d k 3 . In each experiment, approximately 15 values of p12 (0.5 to 2.5 Torr) and time t were determined and p12was fit by least squares to a fourth-order polynomial in time. T h e standard deviation of these fits was less than the estimated experimental error in PI,. This curve was then used to determine values of RI, a n d p l , a t various times t and thus [HI]/[I2] and [RI] [HI]/[12J1/2Rlz.T h e d a t a for all experiments a t a given temperature were then used to determine values of klK121/2 and k d k 3 by a least-squares procedure. T h e results for four separate experiments a t 477.7 K and the resulting least-squares line are shown in Figure 1. The values of k l and k d k 3 obtained by such plots a t various temperatures a r e given in the fifth and sixth columns of Table 11. The first three columns of this table give the temperature and the initial pressures (Torr) of H I and CF3CF2I for experiments a t each temperature t o show the pressure range and ratio of reactants used. The Arrhenius plot for k , is shown in Figure 2 and a least-squares analysis yields log ( k l / l . mol-'
3
+
S-I)
= (11.2 f 0.3)
- (16.8 f 0.8)/8
T h e errors a r e one standard deviation, and 8 = 2.303RT in kcal/mol. T h e logarithm of the A factor compares favorably
September 29,1976
6113
10ST
Figure 2. Arrhenius plot for the reaction C2F5I iI
Figure 1. A plot of [HI][I2]-' vs. [C2F51][HI][I~]-'/2 rate-' at 477.7 K. The initial conditions in Torr are: (A)HI = 63.5, C2F5I = 23.0; (0)HI = 39.1, C2F51= 19.3; (X) HI = 21.7, C2F51= 40.9; ( 0 )HI = 20.1, CzFsI
log ( k d k j ) = (0.68 f 0.05) HI at Various Temperatures
K
497 518 518 538 561 561 56 1
750 300 400 290 100 120 130
2.0 3.2 4.1 3.1 3.4 3.5 3.2
-APHI, Torr 1.8 3.0 4.0 3.1 3.4 3.5 3.1
- A P c ~ F ~ I , APtot,~, Torr Torr 2.1 3.2 4.1
0.0 0.p 0.0 0.0 0.0 0.0 0.0
3 .O
3.4 3.4 3.5
AHr0(1,298) = 16.4 f 1 kcal/mol
T , OK ~~~
561.4
538.0
518.0
497.2 477.7
p m , Torr ___
+ HI
P(C2FsI), Torr
-
CF3CF2H
(iii)
Discussion (a) u Bond Dissociation Energies. The C-I bond dissociation energy in C F ~ C F Zand I the enthalpy of formation of pentafluoroethyl (CF3CF2) may be obtained by combining eq iii with AHfb(CF&F&g,298) = -240.0 f 1 kcal/mol, determined previously by Wu, Pickard, and R o d g e r ~ , ~ AHro(12,g,298) = 14.9 and AHfo(I,g,298) = 25.5 kcal/mol.ll This yields:
with values reported for ethanes, namely 11.O for C2HsIioand 11.5 for CF3CH2L9 The values for kdk3 are scattered; however, it has been repeatedly shownI0 that E3 - E2 = 1.0 f 1 kcal/mol and, asTable 11. Kinetic Study of the Reaction CF3kF2I
+ (1.0 f l)/O
Again, log (AdA3) compares favorably with vdlues previously determined. Namely, 0.60 for c2HsIl0 and 0.65 for CF3CH21.9 Taken in conjunction with the assumption that E2 = 0.0 f 1 kcal/mol,Io the value for El yields: El - E2 = AUro(l,518K)= AHr0(1,518K) = 16.8 f 1 kcll/mol. Estimatingg acp,"(l) = -1 f .5 cal K-' mol-', from 518 to 298 K one then obtains:
Table I. Stoichiometric Study of the Reaction of CF3CF2I with
T,O
C2F5 + 12.
suming that is the case here, one obtains
= 65.8.
Reaction time, API,, min Torr
+
DH0298(CF3CF2-I) = 52.5 f 1 kcal/mol
+
12
k l x 10-3, 1. mol-' s-I
Pi2, Torr
k2/k3
~
20.69 18.51 12.09 8.14 5.45 22.28 22.26 11.44 10.27 5.27 22.29 22.12 21.70 11.35 10.52 5.15 21.53 10.44 10.09 63.48 39.13 21.65 20.47 20.13
5.75 8.34 11.64 18.00 20.47 11.43 5.62 11.55 22.96 22.17 7.06 12.04 3.44 12.64 22.23 22.00 10.08 10.72 21.13 23.00 19.31 40.85 18.65 65.82
0.1 1
0.06 0.07 0.12 0.04 0.07 0.02 0.02 0.02 0.03 0.02 0.03 0.02
67.5 f 2
10.4 f 0.7
42.4 f 1
15.0 f 0.7
0.02 0.08
19.0 f 0.3
1 1 .O f 0.7
0.00 0.00 0.00
10.4 f 0.2
13.3 f 0.5
0.01
0.02 0.03 0.00
0.02 0.05
Wu, Rodgers
5.76 f 0.03
/
14.8 f 0.3
Gas Phase Reaction of Pentafluoroethyl Iodide with HI
61 14 Table 111. Bond Dissociation Energies for the C-X Bond in CzFsX and CzHsX X H F CI Br I
DHO (CF3CF2-X), kcal/mol 103.0 f 126.8 f 84.5 f 68.3 f 52.5 f
The enthalpy of addition of a symmetric molecule X2 to tetrafluoroethylene (reaction 8) may be expressed as eq iv.
DHo(CH3CH2-X),
kcal/mol
lo 1.3 (this work)
98.0 f 1 107.5 (est.)c 81.5 f 1-f 67.7 & 1 53.4 f 1'
Idae le,g 1.3 (this work)
+
X2 CF2=CF2 CF2XCF2X AHr'(8) = DHo(X-X) - DH'(CF2XCF2-X) - DHO(CF2CF2-X) +
D K ' ( C F ~ = C F ~= ) DH'(CF3CF2-X) - DHO(CF2CF2-X) (v) SO that substitution in eq iv yields:
+
AHr'(8) = DH'(X-X) Dr0(CF2=CF2) - DH'(CF2XCF2-X) - DHo(CF,CF2-X)
and AHf0(CF3CF2,g,298) = -213.0 f 1.3 kcal/mol This value and AHfD(CF3CF3,g,298) = -320.9,12 AHf0(CFz=CF2,g,298) = -157.4, and AHf'(F,g,298) = 18.9 kcal/mollI lead to values for the C-F bond dissociation energy in hexafluoroethane and pentafluoroethyl, Le., enthalpies of reactions 6 and 7: CF2CF3
CF3CF2
+F
(6)
CF2 = CF2 t F
(7) thus, DH'(CF3CF2-F) = AHro(6) = 126.8 f 1.3 and DH'(CF2CF2-F) = AHro(7) = 74.5 f 1.3 kcal/mol. With these results, one can compare the values of the C(sp3)-X bond dissociation energies in C2FsX with those in C2HsX for X = H, F, C1, Br, and I. These data are summarized in Table 111. The most outstanding feature of this table is that F for H substitution leaves the C(sp3)-X bond nominally unchanged (2 f 3 kcal/mol) with the striking exception of the C-F bond, which is strengthened by 20 kcal/mol! The value of the C(sp3)-F bond energy in ethyl fluoride is based upon an estimated enthalpy of formation,12 but it is not reasonable to expect this estimate to be in error by 10 to 15 kcal/mol. I n addition, a similar effect has been observed in CFq and CH3F for which DH'(CF3-F) = 130.8 f lI3si4and DH'(CH3-F) = 109.8 f 2 I 0 3 I 3 kcal/mol. Thus, the 20 kcal/ mol difference in the C(sp3)-F bond dissociation energies on complete F for H substitution must be regarded as real. Consequently, an unexpected dichotomy exists in the effect of F for H substitution on the C(sp3)-X bond energies. The explanation for this is at present unknown, but it has certainly stimulated our interest in further work on the effects of F for H substitution on bond energies. (b) x Bond Dissociation Energies. The K bond dissociation energy of a monounsaturated compound has been defined as the difference in the bond dissociation energy of a given bond in the saturated compound and in the relevant free radical;8 thus, for tetrafluoroethylene, Dr0(CF2=CF2) = AHro(6) AHro(7)= 52.3 f 2 kcal/mol. This value is significantly less than the K bond dissociation energy in ethylene, Dao(CH2=CH2) = 59.1 f 2 kcal/moL8 but it is not sufficiently less (nominally 7 kcal/mol) so as to wholly account for the differences in the thermochemistry of addition reactions to ethylene and tetrafluoroethylene.' -6 Journal of the American Chemical Society
/ 98:20 /
(vi)
This may be simplified further if one makes the reasonable approximation that: DHO(CF2XCF2-X) = DH'(CF3CF2-X) so that eq vi becomes: AHro(8) = DH'(X-X) t D K ' ( C F ~ = C F ~ ) - 2DH0(CF3CF2-X)
--
(iv)
However, the x bond dissociation energy for tetrafluoroethylene is given by:8
a J. E. Bassett and E. Whittle, J . Chem. Sor., Faraday Trans. 1 , 68,492 (1972). Reference 10; this leads to AHfo(CH3CH2,g,298) = 25.7 f 1 kcal/MOL/ Reference 12 and footnote b. J. W. Coomber and E. Whittle, Trans.Faraday Soc., 63,2656 (1967). e W. G . F. Ford, Ph.D. Dissertation, Texas A&M University, May 1975. J. Chao, A. S. Rodgers, R. C. Wilhoit, and B. J. Zwolinski, J . Phys. Chem. Ref: Data, 3, 141 (1974), and footnote b. g K. c. Ferguson and E . Whittle, J . Chem. Sor., Faraday Trans. I , 68, 306 (1972). D. R. Stull, E. F. Westrum, and G. C. Sinke, "The Chemical Thermodynamics of Organic Compounds", Wiley, New York, N.Y., 1969, and footnote b. D. B. Hartley and S. W. Benson, J. Chem. Phys., 39, 132 (1963).
CF3CF3
(8)
(vii)
There are data on three different reactions of type 8, chlorination, bromination, and polymerization, for which quantitative comparison can be made. (1) Chlorination. Lacher and co-workers]measured the heat of chlorination of tetrafluoroethylene calorimetrically and obtained AHr0(8,C12) = -57.3 f 0.2 kcal/mol at 363 K, which may be corrected to AHr0(8,C12)= -57.4 f 0.2 kcal/ mol at 298 K. When one takes the value of DHo(CF3CF2-Cl) from Table I11 and DHO(C1-C1) = 58.2 kcal/moll' then eq vii yields AH,' = -58.5 kcal/mol in good agreement with experiment. (2) Bromination. The heat of bromination of tetrafluoroethylene was also determined calorimetrically by Lacher and co-workersls who obtained AHr0(8,Br2) = -38.7 f 1 kcal/ mol at 376 K which may be corrected to AHr0(8,Br2)= -38.8 f 1 kcal/mol at 298 K. Again, eq vii, Table 111, and DH'(Br-Br) = 46.01' yield AHr0(8,Br2j= -38.3 kcal/mol in excellent agreement with experiment. (3) Polymerization. For the polymerization of tetrafluoroethylene, X2 becomes the gas phase polymer, -(CF2CF2),-, and DH'(X-X) and DH'(CF2XCF2-X) become identical; thus eq vi becomes: AHpoly' = Dr0(CF2=CF2) - DH' [CF~CFZ-(CF~CFZ),-](viii) No value for DH' [CF3CF2-(CF2CF2),-] is known, but the gas phase heat of polymerization of tetrafluoroethylene has been estimated recently7at AHp,lyo = -37.2 kcal/mol, so that eq viii results in:
DH0[CF3CF2-(CF2CF2),-] = 89.5 kcal/mol This is some 8 kcal/mol stronger than the estimate of the corresponding bond in polyethylene8 and indicates considerable stabilization for the C-C u bond.' This is not, however, an unexpected result in view of the fact that DHo(CF3-CF3) = 98.7 f 1 k c a l / m 0 1 ~while ~ , ~ ~DHo(CH3-CH3) = 88.4 f 1 kcal/mol. I Finally, it would be tempting at this time to conclude that F for H substitution destabilizes the K bond in ethylene, and such a view is further supported qualitatively by the rate parameters found by JeffersI6 for the cis to trans isomerization of 1,2-difluoroethylene;however, we have just completed a determination of the K bond dissociation energy in 1,l-difluoroethylene17and obtained a value of Dr0(CH2==CF2) = 62.1
September 29, 1976
61 15
kcal/mol which indicates that the interpretation of the fluorine substituent effect will not be a simple matter in P bonds either.
Acknowledgment. T h e authors gratefully acknowledge t h e support of this research by a G r a n t from the Robert A. Welch Foundation and one of us (E.C.W.) expresses his appreciation for a Robert A. Welch Postdoctoral Fellowship.
(5) J. P. Chesick, J. Am. Chem. SOC.,88, 4800 (1986). (6) H. E. O'Neal and S. W. Benson, J. Phys. Chem., 72, 1886 (1968). (7) M. Ben-Yehuda, M. G. Katz, and L. A. Rajbenbach, J. Chem. Soc., Fara&y Trans. 1, 70, 908 (1974). (8) S. W. Benson, J. Chem. Educ., 42,502 (1965). (9) E. C. Wu and A. S. Rcdgers, Int. J. Chem. Kinet., 5, 1001 (1973). (10) D. M. Golden and S. W. Benson, Chem. Rev., 89, 125 (1969). (1 1) D. Stull, Ed., "JANAF Thermochemical Tables", The Dow Thermal Laboratory, Dow Chemical Co., Midland, Mich.. US. Department of Commerce, PB No. 168370 and supplements. (12) S. S. Chen. A. S. Rcdgers. J. Chao, R . C. Wilhoit, and B. J. Zwolinski, J. Phys. Chem. Ref. Oat;, 4, 441 (1975). (13) A. S. Rodgers, J. Chao, R. C. Wilhoit, and B. J. Zwolinski, J. Phys. Chem. Ref. Data, 3, 117 (1974). (14) K. C. Ferguson and E. Whittle, J. Chem. SOC.,Faraday Trans. 1, 88,295 (1972). (15) J. R. Lacher, L. Casali. and J. D. Park, J. Phys. Chem., 80, 608 (1956). (16) P. M. Jeffers, J. Phys. Chem., 7 8 , 1469 (1974). (17) J. M. Pickardand A. S. Rcdgers, J. Am. Chem. Soc., following paper in this I
References and Notes J. R. Lacher, J. J. McKinley, C. M. Snow, L. Michel, G. Nelson, and J. D. Park, J. Am. Chem. SOC.,71, 1330 (1949). (2) C. R. Patrick, Adv. Fluorine Chem., 2, 1 (1961). (3) E. C. Wu, J. M. Pickard, and A. S. Rodgers. J. Phys. Chem., 79, 1078 (1975). (4) E. W. Schlag and E. W. Kaiser, Jr., J. Am. Chem. SOC., 87, 1171 (1965). (1)
.
issue.
Kinetics of the Gas Phase Addition of Bromine to 1,l-Difluoroethylene. 7 Bond Dissociation Energy of 1,l -Difluoroethylene J. M. Pickard and A. S. Rodgers" Contributionfrom the Department of Chemistry, Texas A & M University, College Station, Texas 77843. Received February 17, 1976
Abstract: The reaction of 1,I-difluoroethylene with bromine has been studied spectrophotometrically over the temperature range 550 to 620 K. The experimental results, based on initial rate measurements, are in good agreement with the following mechanism: Br2 e 2Br; Br CH2=CF2 e CH2BrCFr; CH2BrCF2. Br2 s CHzBrCFzBr + Br. A least-squares treatment of the data yields log k(l.3/2 m0l-31~s-I) = (7.8 f 0.1) - (17.8 f 0 . 3 ) / 0 ,where 0 equals 2.303RT kcal/mol. From the observed kinetic data, DH0298 (CF2CH2-Br) was found to be 6.8 f 1.0 kcal/mol. This value was combined with the known C-Br bond dissociation energy in CF3CH2Br to yield the 7r dissociation energy of I,l-difluoroethylene as 62.1 f 1.0 kcal/ mol.
+
+
In the last several years, the influence of substituent effects on the rate of addition of free radicals to unsymmetrical fluoroethylenes has been the subject of several investigations. In particular, the rate of addition of H , CH3, CF3, and CF2Br to 1,l-difluoroethylene has received considerable attention and the Arrhenius parameters a r e well characterized. In the course of our investigations, we have found that radical addition reactions to olefins offer a pragmatic approach t o measurement of a bond dissociation energies. In this study the analysis of the kinetics of the addition of bromine to 1,l-difluoroethylene is given and the implications regarding a bond dissociation energies a r e discussed.
Experimental Section Anhydrous bromine from Mallinckrodt Chemicals was degassed and used without further purification. The 1,l-difluoroethylene from Matheson Gas Products was purified by bulb to bulb distillation from an n-pentane slush. Gas chromatographic analysis indicated that the purity was greater than 99% and the ir spectrum was identical with that previously reported.' All kinetic runs were followed spectrophotometrically using an apparatus previously described.6 The absorption coefficients of bromine, shown in Table I, were determined over the range 550 to 620 K at 440 nm. In a typical run, a known pressure of bromine was expanded into the reaction cell followed by that of 1,l-difluoroethylene. The bromine pressures were varied from 4 to 15 Torr while that of the olefin ranged from 12 to 250 Torr. The initial rate was determined as -dP/dt = A A / ( a A t )
(i)
where (Y is the absorption coefficient of bromine and AA is the change in bromine absorbance during a specific time interval, Af. From this, the apparent rate constant was determined as kap
where P
= ( a - ' A A / A t ) / ( P B d 3 / *(PCH*=CF2)
(ii)
Band ~ P~C H ~ = Care F ~the average pressures within the interval
At.
Results The expected product of the reaction of bromine with 1 , l difluoroethylene would be 1,2-dibromo- 1,1-difluoroethane. In kinetic runs with low conversion, nominally less than 50% of bromine, the total pressure change was equal to the total decrease in bromine. G a s chromatographic analysis of the quenched reaction mixture revealed only unreacted 1,l -difluoroethylene and one peak which had a retention time equivalent to a n authentic sample of 1,2-dibromo- 1,l -difluoroethane. Representative runs a r e tabulated in Table 11. At large extents of reaction, measurable deviations occurred between the total pressure change and that of bromine; consequently, all data for the kinetic determinations were restricted to periods of time corresponding to the disappearance of less than 50% of the bromine. The order of the reaction was determined from a plot of the logarithm of the initial rate vs. the logarithm of bromine and 1,l -difluoroethylene pressures. Order plots a r e shown in Pigures 1 and 2 for bromine and 1,l-difluoroethylene, respectively. From the slopes of these plots, the reaction was found to be first order in olefin and 3/1 order in bromine. Additional evidence
Pickard, Rodgers
/
Gas Phase Addition of Bromine to 1.1-Difuoroethylene
