348
R. C.KENNEDY AND J. B. LEVY
Bistrifllliolrolriethyl Peroxide. 11. Kinetics of the Decomposition to Carbonyl Fluoride and Trifluoromethyl Hypofiuorite Craig Kennedy* and Joseph B. Levy Department of chemi8try, The George Washington University, Washington, D.C. 60006 (Received April 10, 1978)
Publication costs assisted by the Air Force Ofice of Scientific Research, Ofice of Aerospace Research
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The kinetics of the forward and reverse reactions for the equilibrium CFaOOCF3i=!CF30F COFzhave been measured in the temperature range 480-580°K. The results are interpreted in terms of the mechanism CF,OQCFa if 2CFaO (h,kz) ; CFaO Frt CFzO F (k3, kd) ; CFaO F Ft CFBOF ( k b , ke). Activation energies of 46.2 and 45.0 kcal molw1have been found for steps 1and 6 leading to bond dissociationenergies,evaluated at 298"K, of 46.7 =t0.8 and 44.5 f 0.8 kea1 mol-', respectively,for the central oxygen-oxygen bond in bistrifluoromethyl peroxide and the oxygen-fluorine bond in trifluoromethyl hypofluorite. General rak expressions have been determined,by a considerationof kinetic and thermodynamic data, for all six rate constants.
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Introductioln vessel of high surface area concentration to show that the reaction is homogeneous. We have recently' reported measurements of the The experimental results are collected in Tables equilibrium constant, K , for the reaction CFsOOCF3 -+CFaOF GQF2, over the temperature range 500- 1-111.2 The data in Tables I and I1 refer to initial rates of the forward reaction, i.e., the decomposition of the 600°K. The results gave K = 108.' exp -23,80O/RT peroxide; the data in Table I11 refer to the reverse atm and led to a bond dissociation energy, DCF~O-OCF,, reaction, i.e., the formation of the peroxide from carof 40.6 i 5 kea1 mol-'. The system was shown to be bonyl fluoride and trifluoromethyl hypofluorite. free of side products and hence, from the starting comForward Reaction. BistriJluoromethyl Peroxide Alone position and thc total pressure a t any temperature, and with Inert Gases. It proved convenient to examine the compositicn a t that temperature can be calculated. the reaction over a rather wide range of pressures in It was thus clear that the kinetics of the reactions in the temperature range 211-304" and the results are both directions could be studied by a manometric collected in Table I. The general behavior of the data method and we report such a study here. The results is that, for the highest two temperatures, the initial lead to a value for the bond dissociation energy, velocities are first order in bistrifluoromethyl peroxide DCIF~O-OCF~ of 46.7 f 0.8 kcal mol-l and to a proposed over the entire range of pressure studied, i.e., the reaction mech:miarn. quantity 105V0/P0is constant, For the range 221Results 260" first-order behavior is shown by the data for the The velocities a t which bistrifluoromethyl perioxide lower end of the pressure range. Thus for 221" the or a mixture of trifluoromethyl hypofluorite and carquantity 105V0/P0 is constant within experimental bonyl fluoride are transformed into the equilibrium error for pressures below 41 Torr; a t 238" the upper mixtures are siich thak for pressures in the range 1-900 limit is about 120 Torr, etc. For pressures above these Torr it proved possible to make rate measurements limits, the kinetics approach half-order behavior. This manometrically from 204 to 350". It is clear that can be seen by examining data in the last column, whether one SJ udies the peroxide decomposition (here 106Vo/(Po)1'z.Thus the highest four pressures for referred to as the forward reaction) or the hypofluorite221' give fairly constant values in this column while carbonyl fluoride association the kinetics will reflect the the first-order constants vary from 0.13 effect of t,he opposiny reactions as soon as some degree 0.35 x lo-' sec-5. The behavior is shown in Figure 1 of reaction has occurred. Accordingly a good deal of where the log Vo-log Po plots show clearly a change in attention has heen paid to measurements of initial reslope. The line drawn with a slope of 1.0 passes action rates. Measurements have been made for the forward reaction with the peroxide alone, or with added (1) J. €3. Levy and R. C. Kennedy, J. Amer. Chem. Soc., 94, 3302 (1972). carbonyl fluoride or trifluoromethyl hypofluorite, and (2) The accuracy of the initial rate measurements was 5 ~ 5 % . The for the reverse reaction with varying ratios of carbonyl rate constants within the first-order and half-order regimes, L e . , where fluoride to triflnorornethyl hypofluorite. A few experithe particular rate constant showed no systematic trend with pressure, all have average deviations that fall within this error limit, with ments have also been carried out with added inert gases the exception of the data a t 250'. These were all obtained with the to show the absence of total pressure effects and in a tubular reactor for which greater scatter of the data was observed.
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The Journal of Phpical Chemistry, Vol. 76, No. $8,197B
Table 11: Initial Rates of Decomposition of Bistrifluoromethyl Peroxide in the Presence of Added Carbonyl Fluoride and ?'rifluoromethyl r it^ at 260" lnntis? pemx1de 'Csmp,
06
22.I
22 1
gror 9 U I $3, PU,
Tr-r
20oa 103.. 0 39 "51 20.0 10.0
sa 1Q 860"
172 98 5 ~
41 0 20.0
15.5
238
1Q.5 10.0 5.5 83 La
4418 285 200 5'3 5'3k 51oc 2%)
10 250
260
288
3688 20Oa 808 52 205 10a 8270 617 4106 310 100 100" 48.6 10.0 120
1010" 54.0 25.0 17.0 10.0a 5 .:io 304
4.06 59.0 52.0 4 80
106V0,
TQSP. sec-1
17.5 10.1 7.10 4.05 2.60 108 108 43.0 34.4 20.4 10.4 7 45 5.35 4.90 3 00 E30 495 333 262 x 23 120 120 120 51 2 25.5 1340 1Q20 512 400 178 79.0 5550 4920 4000 3330 1670 1640 820 170 19900 16500 7950 400F 2850 1640 840
655
3(Eooo. 27000 2490
10"o/Po, sec -1
0.087 0. 10 0.18 0.20 0.26 0.13 0.12 0.25 0.35 0.49 0.52 0.48 0.51 0.49 0.55 0.76 1.12 1.17 1.31 2.05 2.40 2.40 2.40 2.56 2.55 3.70 5.10 6.40 7.95 8.90 7.90 6.74 7.98 9.75 10.8 16.7 16.4 16.1 17.0 166 165 162 159 163 164 158 163 507 520 520
lo~vn/(PO)"q Torrl'2 sec-1
1.24 1.00 1.10 0.87 0.83 3.66 3.44 3.27 3.48 3.20 2.32 1.91 1.67 1.56 1.25 21.9 23.6 19.1 18.6 16.0 17.0 17.0 17.0 11.4 8.0 70.0 73.0 57.5 56.5 39.7 25.0 193 200 195 189 167 164 146 53.8 1820 1650 1460
800 680 519 363 328 3890 3750 1I40
a Experiments performed in aluminum tubular reactor, A : V = 8.3. All others pcrfxnned in aluminum cylindrical reactor, A : V = 0.51 (see Experimental Section). Nitrogen (450 Torr) added. Carbon tetrafluoride (450 Torr) added.
through the data at fhe lower pressures while the data a t the higher vreswres fall near the line drawn with a
CFrOOCFs, Torr
100 100 100 100 100
GFsOF,
COFz, Torr
Torr
0
0 0 10
50 0 0 0
50 200
Initial velocity, V O ,Torr sec-1
Vo/(CFaOOCBx)o,
0.0165
0.000165 0.000163 0.000113 0.00130900 0.0000210
0.163 0.0113
O.OO90O 0.00210
9ec - I
~
~
_
Table 111: Kinetic Data for the CF30F-COFz Reaction from Initial Velocities
Temp, O C
200
249
260
28 1
--Initial CFaOF
40 40 40 40 20 40 60 10 10 10 10 40 40 40 40 20 40 80
reactant pressures, TorrCOF, Inert gas
10 40 80 160 80 160 240 10 30 60 90 160 160 160 240 80 160 320
[%'o/(CFaOFo)1 X 108 sec-1
3.55 5.82 6.65 6.72 79.5 93.0 95.0 68 111
134 300 Nz 300 0%
158 230 230 230 235 590 1220 1310
slope of 0.5. The pressures a t which the slopes change increase with increasing pressure, so that, as discussed above, for the pressures examined here, the first-order region is clearly defined from 221" up and is, in fmt, the only regime described by our data above 260". Experiments in an aluminum tubular reactor in which the surface area concentration was 16 times that in the aluminum cylindrical reactor are also shown in Table I from 211 t o 260". The rpsults support the conclusion that heterogeneous contributions t o the rate are negligible. Also shown in Table I are experiments at 238" in which substantial amounts a i nitrogen and carbon tetrafluoride were added to the system without affecting the rate. This supports the conclusion that the data shown in Table 1 are independent of total pressure. It may be expected3 that this would be true for the present pressure regime for a rriolecule of the complexity of bistrifluorometkyi peroxide. The first-order rate constants determined in the above manner are plotted as the A r r h e ~ i ~function s in (3) S. W. Benson, "Thermochemical Kinetics,'' Wiley, New Yo&, N. Y . . 1968. The Journal of Phyaieal Chemistry, Vol. 76, N o . 2?2?> 1979
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_
R. C. KENNEDY AND J. €4. LEVY
3482
6Dl
21 .
Figure 2, Arrhenius plot for t h e pyrolysis of bistrifluorometh yl peroxide.
Figure 2. The rata constant is designated as IC, and ha8 the value4 k- = 1 0 1 6 . 2 =k 0 . 1 1 0 - 4 6 , 2 0 0 zk 330/4.57Tsee-1
Forward Reackion. Bistrifluoromethyl Peroxide with The data for the effect of the reaction products on the initial rate of decomposition of the peroxide are &own in Table 11. Clearly carbonyl
Products Added.
The Journal
o>fPhysical
Chemistry, Vol, Y6,N o . 83,19Y2
I
I
19
1 8.
I O h
LOG P (TORR)
Figure 1. Log V-log P plots for t h e pyrolysis of bistrifluoromethyl peroxide: e, cylindrical reactor; A, cylindrical reactor (inert gas added); 0 , tubular reactor.
I
2.0
Figure 3. Arrhenius plot for t h e decomposition of trifluoromethyl hypofluorite with excess carbonyl fluoride: 0, present work; 0 , ref 9.
fluoride inhibits the reaction powerfully while trifluoromethyl hypofluorite is without significant effect. These data furnish a very important clue to the reaction mechanism, since they require that carbonyl fluoride plays a different role than trifluoromethyl hypofluorite in the reaction. Reverse Reaction. I n Table I11 are shown data for the reverse reaction. The general pattern of these data is that, for a given pressure of trifluoromethyl hypofluorite, the initial velocity increased with increasing carbonyl fluoride pressure but reached a limiting plateau value at some critical carbonyl fluoride pressure. The limiting velocities have been treated on the assumption of first-order kinetics in trifluoromethyl hypofluorite and support this assignment. The addition of inert gases in rather substantial amounts has, as shown by the data a t 260", no effect on the velocity so that it is clear that the reaction is in the first-order region. Measurements of this sort were made from 204 to 282" and are plotted in Figure 3 as the Arrhenius function. The rate constant is designated, for reasons that will become clear below, a8 ks. = 1 0 1 4 . 9 k 0.310-45,OOO =k 300/4 57T sec-l
Discussion Reaction Mechanism. The most reasonable opening step in the pyrolysis of bistrifluoromethyl peroxide is (4) Treatment of data for Arrhenius plots in the present paper i s by the method of averages! ( 5 ) H. W. Salzberg, J. I. Morrow, 8. R. Cohen, and M. E. Green, "Physical Chemistry, A Modern Laboratory Course," Academic Press, New York, N. Y . ,1909,pp 6-40.
BISTRIFLUOROMETHYL PEROXIDE
3483
certainly cleavage of the central oxygen-oxygen bond. This follows not only by analogy with dialkyl peroxides, but also from the fact that the pyrolysis of bistrifluoromethyl peroxide in the presence of hexafluoropropenc results in low molecular weight polymers with trifluoromethoxyl end The steps that follow are indicated clearly by the effects of carbonyl fluoride on both the pyrolysis of the peroxide and that of trifluoromethyl hypofluorite. The r n e ~ h a n i s m ~that - ~ appears to be required by the present results is kl
CF3OC)CFs *- 2CF30 kz
z. COFz + F -+ F
Kinetics of the Initial Velocity of the Forward Reaction. When the concentrations of trifluoromethyi hypofluorite and carbonyl fluoride are considered to be negligible, (f) reduces to the expression for the initial rate of the forward reaction, Vo.
vo =
kl(CFaOOCF3) k2 (kl(CFs00CF.y) I + - k3 ( k2
2) (h)
kz(CF30) k ~ ( C F a 0 )-~ ____ - vz IC3 ka(CFaO) V3
4
CP’&
(g)
The second term in the denominator, from expression , d, can be seen to be equal to V z / V 8 i.e,
3
CEO
( C F ~ O F ) ~ ( C O F-Z )hksks ~ -K (CF~OOCFZ), k z h h
5
CFsO 6
The inert gas experiments indicate that steps 1 and 6 are in the first-order region (hence that steps 2 and 5 are in the second-order region). Step 3 i8 also written as a first-order reaction and this point will be considered below. Kinetic Analysis of the Reaction. Even thaugh the results presented above refer to initial rzlte studies starting from either end of the reaction, it is convenient to derive a general rake expression for the reaction first and then to consider i t for the limiting cases of initial conditions. The velocity of the forward reaction, Le., -d(CF3OOCR)/dt, at any time, is denoted here by V and can be expressed as
V = kl(C.F300CF3) - kz(CF30)’
where the subscripts refer to the elementary steps of the same number in the mechanism. When V2/V3 > 1 the initial rate expression becomes
(4
a\
’
ks(CFs0) - ka(F)(CFzO)
(b)
- ks(CF3OF)
(c)
It can be seen, however, that if V z>> V s ,then kl(CF3OOCF3) >> V Oand (j) becomes
By inserting (d) into (c) the expression for (F) is found.
This is the rate expression that the data shown for the higher pressures in Table I and Figure 1 are approaching. Although the high-pressure points in Figure 1 correspond fairly well to the lines drawn with a slope of one-half, it is clear that they do not satisfy the
V
=z
’Ir == Icb(CF3O)(F)
It is then possible to 13olve,from eq a, for (CF30). ki(CF300CFa)
kz
-
kZ
It is then possible to insert (d) and (e) into (b) and, after some rearrangement of terms, to arrive at the expression
(6) H. L. Roberts, J . Chem. Soc., 4538 (1964) (7) Independent studies of the forward reaction alone have been reported by Descamps and Forst8 and of the reverse reaction by Czarnowski and Schumacher.8 The experimental results of Descamps and Forsts are similar to ours but their interpretation differs. Czarnowski and SchumacherDalso found similar results to ours for the reverse reaction and proposed the same steps as are proposed here: they investigated the reaction only over a l o Dinterval, from
(f)
223 t o 233O. (8) B. Desoamps and W. Forst, Abstracts of Papers Presented before Division of Fluorine Chemistry, 162nd National Meeting of the
When V = 0, Le., atJequilibrium, (f) reduces to the expression for the equilibrium constant.
American Chemical Society, Washington, D. C., 1971. (9) J. Czarnowski and H. J. Schumacher, 2. Phys. Chem. (Frankfurt am Main), 73, 68 (1970). The Journal of Physical Chemistru, Vol. 76, No. 23, 197%
3484
R. C. KENNEDY AND J. B.LEVY Table IV: Calculated Values of lea from Initial Velocities in the Pyrolysis of Bistrifluoromethyl Peroxide Critical
Temp, OC
211 221 238 250 260 a
peroxide pressure, Po*, TorrQ
94 f 5 140 f 5 210 f 10 340f 10 5 2 0 f 20
kt kl X lo6, sec-1
0.209 0.544 2.61 7.42 17.2
x
10-4, Torr-'
ka = (1/2Po*krkz)i'e,
8ec-*
8ec-i
8.28 8.11 7.84 7.68 7.58
2.85 I 0 . 1 5 5.56 4. 0 . 1 14.6 f 0 . 3 31.1 f 0 . 5 58.0 i 1 . 1
See text for explanation of Po*.
(V5/V4) and the effect of increasing (COF2) is to make the denominator approach a value of one. Thus the Figure 4. Arrhenius plot for ka. kinetics observed with the addition of increasing amounts of carbonyl fluoride, i.e., first-order dependence on trifluoromethyl hypofluorite, are assigned to step 6 condition that k1(CF300CF3)/Vo >> 1. For example, for the highest pressure at 221°, ~ I ( C F ~ O O C F ~ ) / V ~ and it is for this reason that the corresponding Arrhenius expression given in the Results for the reverse re4.5. action is denoted as kg. This interpretation is supTo evaluate k3, recourse was had to the observation ported by the close correspondence of that expression that, as the pressure of bistrifluoromethyl peroxide was for kg with the corresponding rate constant measured increased, a point would be reached a t which the obindependently by Czarnowski and S c h ~ m a c h e r . ~ served velocity would have fallen to */2 kl(CF300CF,). By plotting initid velocity us. initial peroxide pressure = 1014.610-43,500/4.."7Tsec-l this point can be accurately determined. At this point, The data points reported in the latter work are shown from expression 11 in Figure 3. Kinetics for Noninitial Conditions. From expressions f and g the general expression for the rate of disappearance of peroxide, V , can be written Since for the initial velocity of the forward reaction it V=k1X is always true that IO? T
vo = ks(CF30)
(m)
it follows tha6
k8 = (k2V0*)"' = (3//2klk2(CF300CF3)*)'/z (n) where (CF3CIiOCF3);* is the concentration of peroxide at which the imitial velocity is just one-half that calculated on the basis of eq I, the simple first-order expression, and Vo" is the corresponding initial velocity. Values of k3 foumcl in this way are tabulated in Table IV and plotted in Figure 4. They yield the rate expression k3 -. 1014,."f 0.210-31,OOO f 5W)/4.5?T see-1 (0)
Kinetics of the Inntial Velocity of the Reverse Reaction. When (CF300CF~)is considered negligible, expression f reduces to
where -Vo := initial velocity of the reverse reaction. [kak~/k&4(CF@)I The denominator is now 1
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The Journal of .F'hpakal Chem;st:strg, Vol. 76,No. .B,1973
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Since all the constants are known, it is possible to test this expression over the entire course of the reaction. This is done for the two temperatures 269 and 308" in Figures 5 and 6. The lines drawn have been constrained to pass through the origins and have been drawn with slopes that correspond to values of IC1 calculated from the Arrhenius expression for kl given in the Results section. The points fall well on the lines and give good support to the mechanism. Rate Parameters for Individual Reactions. There is sufficient information available in the present results to allow calculation of all six rate constants of the rnechanism. Expressions for kl, k3, and kg have already been given. The assumption can be made that the activation energies Ez and E5 are zerolo since they correspond to recombination reactions of simple free radicals. The calculation of k2 follows from the expression (10) S. W. Benson, J . C h e w Educ., 42,502 (1965).
34885
wbere &he st^^^^^^^ enkropy change and enthalpy change are evttluateei POF the given temperature. A Bimilatr expxewiiim czm be written for the (5) (6) equilibriirni The rliitiiiLiitpi/ A H O ~ ,can ~ be equated to activation energy E1 E J ~ C P ~ &n"Cc'lva
E-
E1
- RT
and
The stmdarci entropy lor the peroxide can be taken from our earlier paper while that for CF30 is here estimated by subtracting from the tabulated value for CF30Fthe difference between the tabulated values for OF, and OF.11 Exitmination of this procedure shows that AH" and ilS0 ar2 virtually unchanged from 500 to 600°K. The value3 used here are 46.2 kcal mol-' an3 34.0 eu xbiol-', respectively. (The values for A.lrPof9s60and Sojd0 obtained in this way for CFZO, and ~7hichwill figurp in thi: further discussion, are -156.8 2 kcal mol - l and :78.4 f 1.0 eu mol-', respectively.) These yield aut ii50"K, AGO1 = 26.5 kcai mol-l and kl,/kz == l O - l O 6 atrn Insertion of the experimental value for kl and conversion to concentration units I n this same way it yield ks = 2.5 X 80" 44-l is found that kg = 3.2 Y IO9M-I sec-I. These results are compared in Tabllc V- with values reported for other sirnrlar radicals 1 t 1s clear that trifluoromethoxy radicals combine at ahoiil; the same rate as alkoxy or alkyl radicaln.
Figure 5 . Plot of t h e rate of decomposition of bistrifluoromethyl peroxide a t 269' us. the rate predicted by the general expression q.
o_ x
--
___o**_l__-_*l
Table V :: Radical Recomlination Rate Constants Log A , M-1
Radicals
GHaO, (GN1)%c:WO, (CH*)$CO CF, CH,, C*H6 CFBO CF,O f F
sec-1
Ref
9.3
12
10.4
3 3 Present work Present work
1.0.5
9,4 9.5
KCCF300CF~~-CCF,0)~CF30F)
The value given earlier for refers to the decomposition of CFIO arid it ie necessary to consider whether the reaction is in the first-order region or in the fall-off region. The evidence is that the reaction is in the firut-order region. Thus the addition of inert gases in rather large amounts, e.g., the data at 260°, Table I, causes no significant effect on the rate in peroxide pressure regions where ks is involved. I n the same way no total pressure eFfccts axe indicated by the trend of the data for 221" where peroxide pressures up to 921 Torr were used. It may also be noted that the preexponen-
Figure 6. Plot of t h e r a t e of decomposition of bistrifluoromethyl peroxide a t 308' os. t h e rate predicted by t h e general expression q.
tial factor shown for ka is reasonable for a molecule of this ~omplexity.~We therefore treat the above ex(11) "JANAF Thermochemical Tables," Thermal Research Laboratory, Dow Chemical Go., Midland, Mich., 1965.
The Journal of Physical Chemastry, Vol. 76, N o . $3, 197'9
R. C. KENNEDY AND J. B. LEVY
3486 pression a9 that for the first-order regime for this reaction rather than the second order. (It may be noted that if the reaction is not quite up into the first-order range, the observed activation energy would still not be expected GO differ from that for the unimolecular reaction by more tEian about 1kcal.) l2 ItJis now possible to calculate k* since the equilibrium constant kajk4can be calculated from the standard entropy and heat content calculated above for CF30 at 55U"K, and the known values for F and CF20. The value for k ~ j k 8found is J c ~ , / )==~ ~106 l i 0 . 2 1 ~ - 2 2 , 9 0 0 j 6 0 0 / 4 . 5 7 1 ' n/il This yiel& jC4 =; 1 0 9 . 4 ~ 0 . 2 1 0 - 8 1 0 0 ~ 5 0 0 / 4 . 6 7 T M-1 see-1. These are the only kinetic parameters we know of for the reaction of F ar,oms with the carbonyl bond. The activation energy for the reaction
N
-+ CH2O
--ic
CH3O
is estimated as about 8 kcal mol-' by Gray, Shaw, and Thynne,13 and the activation energy for the reaction
F
+ CF&3'4F2
+CF3-CF-CF3
is estimated by odgers14 as near zero. It may be noted that tae exothermicity of the present reaction is very close t o that of the H-CH20 reaction, 22 kea1 r n ~ l - ~and , ~mudi ~ less than that of the F-C3F6 reaction, -50 kea1 1 ~ 0 l - l . l ~ ~ i 8 p Y O p O T ~ i O n a t ~ QReaction. n The disproportionation reaction for trifluoromethoxy radicals is written, as step 7 ~
1, a
2CFaC)
CF3OF b9
+ COFz
For alkyl 01' alkoxy radicals this reaction appears to proceed with zero activation energy and about as fast as the recombi~iation.~5For perfluoroalkyl radicals l7 evidence is that the disproportionation the available16$ reaction is very slow compared to recombination reactions. It is clear that the above reaction plays no significant role in the present system. Thus, if step 7 is significant, lben step 8 is too; however, if step 8 occurred, limiting first-order behavior would not be found in the pyroiysis of CF30F-COF2 mixtures. On the other hand, if step 7 occurred, the forward reaction would not show a change in kinetic order as the peroxide pressure was increased. If the assumption is made, as is generally done,la that the preexponential factor for k7 is about that for kz, then it is clear that E, > 0. Our data do not permit a quantitative evaluation of E7 but a reasonable lower limit for E7 is probably about 2RT or 2-3 kcal mol-1, Conceivably E7 could be much higher tham this. Bond Dissociation Energies. The evaluation of bond dissociation energies from the kinetics of pyrolysis reactions i s a well-established procedure.l 8 I n the present work the bond dissociation energies for the oxygenThe Journal of Physiczl, Chemistry, Val. 76, N o . 23, 1972
fluorine bond in trifluoromethyl hypofluorite and for the oxygen-oxygen bond in bistrifluoromethyl peroxide are defined by the quantities denoted earlier as A H o l , , and A H O , , ~ , respectively. When these are converted to 298°K,3 the values found are for the peroxide 46.7 f 0.8 kcal mol-' and for the hypofluorite 44.5 f 0.8 kcal m01-l.'~ The latter value can be compounded to the results reported by Cearnowski and Schumachers of 43.1 kcal mol-1 and the value, 47 kcal mol-' estimated by Porter and Cady.20 The kinetics values are felt to be more accurate. The oxygen-oxygen bond dissociation energy in bistrifluoromethyl peroxide is compared to those of other peroxides in Table VI.21 Table V I : Bond Dissociation Energies RQ-OR
w.
Do-o,zss, koal mol-]
Ref
CHS CzHs
36.1 f 1 34.1 3t 4 37.4 46.7 f 0.8 51 70 d= 10
18 18 18 Present work 18 21
t-ci" CF3 H F
The central bond in bistrifluoromethyl peroxide is thus significantly stronger than those in alkyl peroxides. The oxygen-oxygen bonding in peroxides has received considerable attention in recent years, largely because of the gbservation that the central bond in dioxygen I n valence bond terms,22 difluoride is so one would assign the increased bond strength in CFIOOCF3, gs., for example, CH300CH3,to contributions from the resonance structure CF3- o=Q+GF3. The (12) D. M. Golden, R. K. Sally, and 8. W. Benson, J . Phys. Chem. 75, 1333 (1971). (13) P. Gray, R. Shaw, and J. C. J. Thynne, Progr. React. Kinet., 4, 63 (1967). (14) A . S. Rodgers, J . Phys. Chem., 67,2799 (1963). (15) J. A. Kerr and A. F. Trotman-Dickenson, Progr. React. Kinet., f , 111 (1961). (16) G. 0. Pritchard, G. H. Miller, and J. R. Dacey, Can. J . Chem., 39, 1968 (1961). (17) E. J. Lehmann and J. B. Levy, J . Amar. Chem. Soc., 93, 5790 (1971). (18) J. A . Kerr, Chem. Rev.,66,465 (1966). (19) The HTO - "%OS data for bistrifluoromethyl peroxide are given
in the earlier paper.' The limits of error have been increased from 0.3 to 0.8 koal mol-1 to encompass the chance that the correction to 298O might add an error of k0.5 kcai mol-'. (20) R. S. Porter and G. H. Cady, J . Amer. Chem. SOL,79, 5628
(1957). (21) Calculated from the standard heats of formation of OZFPand OF.^ (22) R. H. Jackson, J . Chem. Soc., 4585 (1962). (28) P. H. Kasai and A . D. Kirshenbaum, J . Amer. Chem. Soc., 87, 3069 (1965). (24) S , 3. Turner and R. D. Harcourt, Chem. Commun., 4 (1967). (25) R. D. Spratley and G. C. Pimentel, d . Amer. Chem. SOC.,88, 2394 (1966).
oride is one of the less active metal fluoride catalysts for the addition of fluorine to carbonyl fluoride. Consequently, a 3-in. diameter cylindrical thmmurn vessel of 1000-cc volume was constructed with ail joints a h minum welded and was used far ail of the experiments reported here except for those performed in the aluminum tubular reactor. The reactor busface was pnssivated only with CFgOF at the temperatures used for the rate determinations. The vessel was connected by 0.25-in. o.d. aluminum tubing to B, Pace pressure transducer. In order to test for ~ e t ~ ~ o ~ e n e o ~ s surface effects some experiments were carried out in a tubular reactor constructed rrorn a E-ft Icngth of 0.25-in, 0.d. aluminum tubing. The tubular reactor was considerably more difficult to passivate than the cylindrical reactor had been. Firat the sr;arlttce was treated wit,h 200-mm pressure of CFaOF at %60ufor 24 br several times. Finally, the s u r f ~ c cwas ~ condilioned with 1.5-atm pressure of CROE' at 3610" for 24 hr, and the process was repeated with elemcnbal Fz. reactor had been sufficiently c o n ~ i 1 , ~ othi ~e~ CJF&KCF.?pyrolysia were rssentidly Ihc eairie ~ E iii I the cylindrical reactor (see Table I) All experinienr,~ reported here were performed in these t m aiiimtnum reactors. The preswre-measuring t ~ c ~ and~ the i ~ ~ ~ s T ~ anoiecubr E orbital picture also serves to give constant-ten?perature bath used have been dcscribecl some insight into the bonding in polyoxides. The inpreviously. A. €folie 482 A4 &Ionel v a h e connected vestigatrons of di-k~rkbutyltri- and tetroxides have the system to a small volume Monel. mnasnrf~.dd used to shown that these are both considerably less stable than evacuate and indrod-clcesamples to the vessel. the peroxides, with the3 tetroxides being least ~ t a b l e . ~ B - ~ ~Experiments were performed by filling the thermoBince the borrd xi the brioxide could be pictured as instated vessel with a measinfed prosnurc oi the peroxide volving, rliiteract ion of one nonbonding orbital (from or a previously prepared mixture of known si oxchiom) and one p or sp3 orbital (from RO) while that in etry including i,wo of the three g a m tread' bktrithe tetroxide would involve two of the former types it fluorornet hyl peroxide, t uorornethj 1 ~ ~ p o ~ i ~ o r ~ t e , is quite reasonable t h d the bond strengths would show ssure - h e curves were and carbonyl fluoride. thP ok~servedtrend. Here too the effect of terminal then recorded and the results iaterpri?ted. h s t m t a GF3groups would be predicted to be bond-strengthening n e w s rates were calculated from tangenCi dra-rm to These reiative t o alkyl groups. The available data are shown curves. Experiments in a9hcfn the peroxide deconiin Table VII and are consistent with all of these qualiposition reaction was allowed to proceed to equilibrium t e i5ve produetJtcms in the conditioned reactors gave firlad pressures corresponding Icr those calculated from i , h ~ropor%rd equilibrium constants. - ~~~ h e r ~ i ~Bistrifluoromethyl ~ ~ s . peroxide, trifluoroTable VkI : 0-0 Bond Strengths in Polyoxides (kcal mol-') methyl hypofluorite, and carbonyl fluoride were obR PLOOR R.0OOR ROOOOR tained from PcilinsuXar ChemPLescwch alnd were used as received. Brstrifluoromethyl peroxide (%l.5yo) and t-C& 36 -23a 4P CF, 45 1-30 ? carbonyl fluoride (SS'%) were analyzed by ~ 8 ehros matography aund mass s p e c t r ~ ~ ~ res e~~r~, References 29 and 31. References 28, 29, and 31.
discussion of peroxides in molecular orbital terms22*25 laas described t h R ~ - 0 bond in ROOR in terms of electron donation rroxn R to a T * antibonding orbital of 01ygen which Cwyitsinb an electron from oxygen. Clearly highly ciectrunegative groups, such as fluorine atom, would donate loss electron density to the antib ~ ~ orbital ~ ~ than i nwould ~ a group like methyl and hcnee t h e oxygc i l --o ~ cygen bonding would be weakened le58 in the farrier cnse. For either explanation one strength of the oxygen-oxygen xide would be between eroxides. Ib is worth er support to the impc~rtancti of the elct.i,roiiega~~vvi.lres of the altached ~r~~~~~ in p e ~ ~ x in~ de ~ e s ining the central oxygcno ~ y bond ~ ~ strengths. e ~ ~ group electronegativities cd ihe trifluc y l and methyl groups have bcen 11d 2 .30,26 respectively, while that calcuPatcd as The electronegativity of the for the fluorinc ibi roup i s 37% of the way along from 1 gmap to that of fluorine. The bond y for ~ ~ s ~ r ~ ~ L ~ o r o rperoxide n e t h y 1 is 32yJ>o f the way aloiig Yrorn the value for the corresponclirig bond i n dmethyI peroxide to that in dioxygen
~
h
a
0
(26) J. Hinxe, M , Whitohea,d, and H. !? daffe, I.J . Anzer. &'hem. Soc., 8 5 , 1.47 (4968).
Apparatus and Procedure. The initial experiments of the investigation were carried out in a 500-cc spherical i\/ronel vessel. However, there was evidence that a heterogeneous wall reaction influenced the observed rates. Previous work bad shown that aluminum flu-
(27) 11. Pading, ""ature of the Chemiml hr:d," 3d ed, Cornell University Press, Xthacs., N. Y . , 1960, p 93. (28) P. D. Bartlet,t and I?. Giinther, 1.Arner. Ch.em. Sac., 88, 3288 (1966). (29) P. D. Brtrtlett and G. Guaraldi, ibid., 89, 4999 ('1967). (30) P. G. Thompson, ibid., 89,4316 (1967) (31) K. Adrbmic, ,J. A. Howard, and K. 1J. .Ingold, Can. 3.Chcm., 47, 3803 (1969) ~
I
The Journal of Physical Chemistrg, Vol. 76,N o . 23, 1972
HENRYE. WIRTHAND FREDERICK K. BANGERT
3488 Analysis of the trifiuoromethyl hypofluorite by idiometric titration1' showed 89% purity. The nature of the experiments was ruch that only bistrifluoromethyl peroxide and c a b o n fluoride presented any difficulties as possible impurities. Since the measurements made for experiments wiLh trifluoromethyl hypofluorite were initial rate nwmirements, the presence of small (-1%) amounts of bistrifluoromethyl peroxide presented no problem. Since, in addition, these experiments were carried out with substantial amounts of added carbonyl fluoride, the presence of small amounts of carbonyl fluoride as :in impurity in the trifluoromethyl hypo-
fluorite was likewise not a source of difficulty. The procedure adopted was to use the trifluoromethyl hypofluorite, as received, and to correct for the fact that it was only 89% pure.
Acknowledgment. This work was supported by Grant No. AFOSR 70-1939 of the Air Force Office of Scientific Research, Office of Aerospace Research, USAF. The United States Government is authorized to reproduce and distribute reprints for Governmental purposes notwithstanding any copyright notation hereon.
Apparent Molal Volumes of Sodium Chloride and Magnesium bhridie in Aqueous Solution
by Henry E. Wirth* and Frederick K. Bangert Department of Chemistry, Syracuse University, Syracuse, N e w York 13810 (Received .November 19, 1971) Publication costs assisted by Syracuse University
The apparent molal voluines of sodium chloride and magnesium chloride in aqueous solution at 25" were determined for the concentration range 0-4 m using a modified form of the Geffcken dilatometer. The data obtained at low concentrations (ionic strength less than 0.4) were treated by the extended forms of the Debge-IIuckel and Friedman equations. The calculated apparent molal volumes at infinite dilution of magnesium chloride were 14.47 (Debye-Huckel) and 14.52 cm3/mol (Friedman).
Introduction Preliminary t o the determination of the volume changes on mixing solutions of magnesium chloride and sodium chloride, the apparent molal volumes of magnesium chloride in aqueous solution were determined at concentrations ulp t o 4 m. As a check on the precision of the dilatometer method used, the apparent molal ium chloride were determined. Experimental Section A. stock scilution approximately 4 m in sodium chloride (Fisher Certified) was made up and analyzed by evaporation and drying to constant weight a t 220". A similar stoclr solution of magnesium chloride (Baker and Adamsnn) was analyzed by precipitation of the chloride as silver chloride, All weights were corrected t o vacuum. The densities of the stock solutions were determined by the sinker method, using a "supplementary The sinkers used were made of Vycor and had approximate volumes of 300 cm3. The apparent molal volumes of dilute solutions were T h e Journal of Physical Chemistry, Vol. 76, N o . 23, 1972
determined using a modified Geff cken dilatometer.2-4 The apparent molal volume ( 4 ) is obtained from the observed change in volume (Av) when stock solution is added to water by means of the relation (1)
where n2 is the number of moles of electrolyte added, and $ref is the apparent molal volume of the electrolyte in the stock solution. The molality of the solution is 1000 nz/(m'l n z l ) , where ?n'lis the weight of the water to which the stock solution (containing n2 moles of salt and m1 grams of water) is added. With the Geffcken apparatus a series of additions can be made, so that a
+
(1) H. E. Wirth and F. N. Collier, Jr., 1.A m e r . Chem. Soc., 72, 5292 (1950). (2) W. Geffcken, A. Kruis, and L. Solana, Z . P h y s . Chem., Abt. R (Leipzio), 35, 317 (1937). (3) €1. E. Wirth, R . E. Lindstrom, and J , N. Johnson, J. P h y s . Chem., 67,2339 (1963). (4) H. E. Wirth, ibid., 71,2922 (1967).
