J . Phys. Chem. 1992, 96, 7324-1328
7324
Intramolecular Proton Transfer as a Preliminary Step for Profon Dissociation in 2-Na phthob3,b-disulfonate A. Masad and D.Huppert* School of Chemistry, Raymond and Beverly Sackler Faculty of Exact Sciences, Tel Aviv University, Tel Aviv 69978, Israel (Received: November 26, 1991; In Final Form: April 23, 1992)
Time-resolved and steady-state fluorescence techniques were employed to study the intramolecular and intermolecular proton-transfer processes occurring in the first excited electronic state of 2-naphthol-3,6-disulfonate.It was found that the first step after excitation by a short light pulse is an intramolecular proton transfer from the hydroxy group to the adjacent sulfonate group. Subsequently, at a slower rate, the proton is transferred to the solvent. The intramolecular proton transfer is mediated by water molecules. The solvated proton has a f ~ t probability e to recombine geminately with the parent molecule via either of the proton-accepting moieties, the hydroxy group or the sulfonate groups. We used the exact transient numerical solution of the DebyeSmoluchowski equation to quantitatively describe the proton geminate recombination reaction,
Introduction Proton-transfer (PT) reactions are among the most common processes in chemistry.' F6rster2 and Weller3 were the first to explain the difference between the ground-state and the excited-state ionization constants by means of an excited-state proton transfer. Since that time, considerable research effort has focused on understanding the reaction dynamics and mechanism of proton transfer, and short laser pulses are commonly used for this study. The weak acids, hydroxynaphthols and their derivatives, have long been considered attractive probes for investigations in which both time-resolved and steady-state spectroscopy techniques are employed (see reviews in refs 4-6). Most of the data analysis has been based on rate equations, where the proton-transfer process exhibits an exponential time dependence. We have presented a different approach' in which we look at the PT process as a transient, nonequilibrium dissociation of an excited-state molecule. We have suggested a two-step reaction model (Scheme I). The
SCHEME I ROH*
DSE & [RO*--H+] RO*-+ H+ k, Q
first chemical step is described by back-reaction boundary conditions with intrinsic rate constants kd and k, and is followed by a diffusional step in which the hydrated proton is separated from the contact radius, a, to infinity. The second step is described by the exact transient (numerical) solution of the Debye-Smoluchowski equation (DSE)*s9with the above-mentioned boundary conditions. Our previous work dealt mainly with 8-hydroxy1,3,6pyrenetrisulfonate (HPTS).7J09"This molecule was found to undergo an almost ideal case of reversible diffusion-influenced reaction. It is easy to monitor this molecule because in its deprotonation form it has a charge of -4 and exhibits a large Coulomb attraction potential. This potential enhances the geminate recombination which is manifested in a long-time nonexponential tail in the ROH* fluorescence decay profile. However, there are other systems in which the chemical step is not a single, reversible step. Our aim in this work is to elucidate the proton dissociation mechanisms of one of these systems, 2-naphthol-3,6-disulfonate (2-np3,6). This belongs to a family of naphthols whose hydroxyl group is adjacent to a sulfonate group. Naphthols which belong to this group have been investigated in the past using both steady-stateand transient techniq~es.l~-'~ These studies excluded the proton geminate recombination and the possibility of an ultrafast intramolecular proton transfer and concluded that an intramolecular hydrogen bond exists in both the ground state and the excited state. Here, we consider the possibility of the existence of an intermediate, reversible step where the oxygen from the sulfonate group accepts the proton from the hydroxyl group and only then transfers it to the water. This process competes with
the direct, reversible, proton transfer. Our experimental system enables us to monitor the transient behavior on the picosecond time scale and our theoretical model to determine the different rate constants which are being involved in the overall process. For comparison, we also studied the time-resolved fluorescence of a second proton emitter, 2-naphthol-6,8disulfonate(2-np-6,8). Both compounds are 3 times negatively charged upon deprotonation. The time-resolved fluorescence of 2-naphthol-6,8-disulfonate should show a direct proton transfer to the solvent similar to HPTS.
Experimental Section Transient fluorescence was detected using time-correlated singlephoton counting (TCSPC). We used a continuous wave (cw) mode-locked Nd/YAG pumped dye laser (Coherent Nd/ YAG Antares and Coherent 702 dye laser) providing a high repetition rate (100 kHz-3.8 MHz) of short pulses (1-ps fwhm) at a wavelength of 580620 nm as a sample excitation source. The frequency was doubled (to 290-310 nm) using a KDP Type I crystal. We typically employed a repetition rate of 700 kHz. The TCSPC detection system is made up of a monochromator (American Holographic DB-IO), a Hamamatsu 156421-01 photomultiplier (S-20 photocathode), a Tennelec 854 TAC, and a Tennelec 454 discriminator. The TAC output was processed by a multichannel analyzer (Nucleus Inc. PCA 11) and an IBM personal computer which was also used for data storage and processing. The overall instrument response function was determined by reflecting the 600-nm pulse from ground glass. The detected laser profile for a 20-11s full scale of the TCSPC system has a sharp primary peak of approximately 55-ps fwhm, superimposed on a long tail whose amplitude is about 0.6% that of the main peak. Measurements were taken at 10-, 20-, and 50-11s full scale. The reagents used were 2-naphthol-3,6-disulfonate(Reidelde-Haen) and 2-naphthol-6,8-disulfonate(Eastman Kodak) (at concentrations of about 2 X M); methanol, spectroscopic grade (Merck), and DzO, 99.8% isotopically pure (Aldrich), were used with no further purification. The deionized water had a resistance of >10 MS2. The fluorescence spectra of 2-naphthol-3,6-disulfonateand 2-naphthol-6,8-disulfonateconsist of two bands. In water, the maxima of the acidic (ROH*) band are located at 377 nm and 388 nm for 2-naphthol-3,6-disulfonateand 2-naphthol-6,8-disulfonate, respectively. The band maximum of the basic (RO*-) form of both compounds occurs at 458 nm. In this study, we monitored the ROH* fluorescence at 370 nm and the RO*fluorescence at 470 nm. Results and Discussion A. Diffusing Reversible Proton. We examined two naphthol derivatives, 2-naphthol-3,6-disulfonatein which the OH is adjacent
0022-3654192f 2096-7324SO3.00f 0 0 1992 American Chemical Society
The Journul of Physicul Chemistry, Vol. 96, No. 18, 1992 7325
Proton Dissociation in 2-Naphthol-3,6-disulfonate
SCHEME 11 ROH*
kd
DSE
n , [RO*-***H+] t------\ RO*- +
H+
TABLE I: Parametem for Fitting the Numerical Solution of the DSE to the Experimental Time-Resolved Fluorescence of 2-~g3,6Shown in "2 I
0.0
20
4.0
6.0
8.0
.
l0.0
TIME (n3 Figure 1. ROH* fluorcPcene decay curvcs of (a) 2-ng3,6 and (b) 2np6,8 in H 2 0 measured at 370 nm.
to the SO3- and 2-naphthol-6,8-disulfonate(2-np6,8) in which the OH and the SO3- are on two different rings. To reveal the influence of OH-SO3- proximity, we conducted two sets of measurements, monitoring the transient fluorescence decay of the naphthol form (ROH*): (1) in H20/MeOH mixtures of varying proportion and (2) in DzO.The data were fitted (as in our group's earlier ~ o r k ' , ~ * subject ~ ~ ) to the assumption that the proton is diffusive and reacts reversibly with the excited RO*-. This was done by solving the time-dependent Smoluchowski equation for the relative translational diffusion of the proton:
wherep(r,f) is the probability density function for finding a proton at time t at a distance r from the origin, J is the diffusional flux, and D is the diffusion coefficient. V = RD/r,where RD = z 3 / is the Debye radius, z is the RO*- charge number, E is the dielectric constant, kB is the Boltmann factor, and L is the diffusion operator. The solution was subject to the Coulomb attraction force, the back-reaction boundary condition at the contact distance, and an initial bound state. In the general case, the boundary condition at the contact distance is given by
J(u,t)
kd[l
- S(r)] - k,4*U2p(U,t)
where
%McOH,M 103kd,p-I 103kl, ps-1 103k2,ps-I 1o3kl, ps-l A 103k,, ps-l i03k ps-' A D,' ps-l
L2
RD,b A X2
0 0.58 15 4.5 3.0 2.70 3.0 0.93 21.3 3.56
10.0 0.40 7.0 3.5 3.5 1.30 3.0 0.71 21.7 3.66
18.0 0.20 3.0 2.4 2.5 0.80 3.0 0.61 22.7 3.09
30.8 0.10 1.2 1.9 2.2 0.28 2.7 0.42 25.1 2.22
46.3 0.03 0.7 1.5 2.0 0.17 2.0 0.33 29.4 2.30
OCalculated from conductivity data taken from ref 23. bDielcctric constants arc taken from ref 21. CRcducedx2 formula taken from ref 30.
TABLE Ik Parameters for Fitting the 2-np6,8 Fluorescence Decay in HzO/McOH ma P d y Shown h F i g ~ r e3 % McOH, M 103kd,ps-' 103k,, ps-l A x2b 0 10.0 18.0 30.8 46.3 63.0
17.0 9.0 5.5 3.3 1.6 0.75
23.0 19.0 18.0 16.0 16.0 8.7
2.54 2.65 2.52 2.63 2.60 2.52
D (AZps-I) and RD(A) are given in Table I. Reduced x2 formula taken from ref 30.
described by Scheme 11. The appropriate boundary condition of the DSE in this case is J(u,t) = kd[ROH*] + k4[RO*-**.HSO3] - (k, k 3 ) 4 ~ ~ ~ p ( ~ , t )
+
The differential population changes of the three species in each time step are given by d[ROH*] = -[kd kf kl][ROH*] + k2[RO*--HSO3] + k,4ra2p(u,t)
+ +
S(t) = 4rJmp(r,t),,2 dr
is the suNival probability of a dissociated pair, 1 - S(r) = [ROH*], S(t) = [RO-1, and u is the contact distance. In the above scheme, the adjustable parameters of the actual data lit are the dissociation and recombination rate parameters at contact. kd determines the initial decay rate, while (mainly) RDand k, determine the amplitude of the long-time tail. The relative (translational) diffusion coefficient of the proton and the anion was calculated as the sum of the individual diffusion coefficients," and RD(7') was calculated from the experimentalvalue e( Z);21 z was -3 for RO*-, and u was 7 A. u = 7 A roughly corresponds to bare naphtholate plus one or two solvation layers. The above boundary condition satisfies the general case of reversible proton transfer. For HPTS7*'O*"and 2-np6,8, we found excellent overall agreement between the experimental and theoretical data. However, the above boundary conditions are insufficient for the case of 2-np3.6. Figure 1 shows the time-resolved fluorescence decay curves of the naphthol forms 2-np3,6 and 2-np6,8. Scheme I cannot account for the complex fluorescence decay of 2-np3,6 composed of an initial fast component followed by a relatively slow component. We suggest that the reason for this complex behavior is a reversible ultrafast intramolecular proton transfer between the OH group and the adjacent SO3- group. This process can be
+ k2 + k4][ROS-.**HS03]+ kl [ROH*] + k34ru2p(u,r) d[RO*-] = -kft[RO*-] - [k,+ k3]4ra2p(u,t)+ d[RO*-**.HSO3] = -[kit
kd[ROH*]
+ k4[RO*-.*.HS03]
where kfand kf,are the radiative rate constants for ROH* and RO*-,respectively. By inserting these new rate constants, we obtained good agreement between the experimental and theoretical data over the whole time range. Since we are aware of the fact that there are six adjustable parameters, we examined the transient behavior under various conditions: in H20/MeOH mixtures and in D2O. We used our previously established analyses under similar conditions of HFTS,'J@ll a-naphthol,u and 1-np3,612as guidelines for determining the different rate constants. B. Water/Methmd Mixtures: Composition Dependence of the Rate Coastaots. Figures 2 and 3 display the time-resolved data of 2-np3,6 and 2-np6,8, respectively, measured at room temperature for several waterlmethanol mixtures. The fits for 2-np6,8 were made using the model shown in Scheme I and those for 2-np3,6 using that represented in Scheme 11. The physical parameters for 2-np3,6 are collected in Table I and those for 2-np6,8 in Table 11. In both cases, the diffusion coefficients for the proton and the aromatic anion were calculated from HCl conductivity data for H20/MeOH mixtures.23 The static dielectric
Masad and Huppert
7326 The Journal of Physical Chemistry, Vol. 96, No. 18, 1992
s
1; g:
0
d
0.0
4.0
2.0
0 2 .
.
,
8.0
6.0
10.0
8.0
4.0
0.0
TlME (n3 ,
1
20.c
16.0
120
TlME (nr) ,
,
,
1
.:t-p---] 2
8
6
I
IO
12
I4
16
18
W
Figure 2. ROH* experimental fluorescence decay curves of 2-np3,6 measured at 370 nm (dots) with the numerical solution of the DSE (full curve) in H20/MeOH solutions of various compositions. (a) Top to bottom: 33%,20%, 0% vol fraction of McOH. (b) Top to bottom: 66%, SO%, 33% vol fraction of MeOH. The lower panels represent the residuals of 2-np3,6 in water (a) and in a solution of 50% McOH (b). Parameters for the fits are collected in Table I.
'32
1: g;
P
P
0.0
3.0
n.0
9.0
6.0
I50
TlME (n3 I
--I
-
OSO O
I
I
I
I
I
I
I
I D
I
2
4
6
B
IO
12
I4
16
18
W
Figure 3. Same as in Figure 2 for 2-np6,8. Methanol contents are (top to bottom) SO%, 33%, and 20% by volume. The lower panel represents the residual of 2-np6,8 in a solution of 33% (vol) MeOH. Parameters for the fits are given in Table 11.
constant of the solvent, e, was taken from recent time-domain reflectometry data.2' z was held constant (-3) in all mixture compositions. We found in our previous work with HPTSZ2that the association of the RO-anion (z = -4) and the Na+ counterion is negligible in solutions of up to 80% (V)or greater methanol content. The dependence of proton-transfer reaction rates, of both molecules, on the solvent composition is shown in Figure 4. In the case of 2-np3,6, we see that k,,the intramolecular protontransfer rate, is an order of magnitude larger than kd, the intermolecular proton dissociation rate, in all of the mixtures. We also obseroe that the three proton dissociation rates have a similar dependence on the solvent composition. From this observation, we conclude that the intramolecular route is dominant in all of the water/methanol mixtures. We further discern that for the two molecules the proton dissociation rates decrease exponentially as a function of the methanol mole fraction of the mixtures. The dependence of the proton recombination rates on the composition of the solution is much weaker. The same behavior was obtained for HPTS in our group's early workz4and in our recent work with 1-np3,6.'2%22
10-
20
40 60 MOL % MeOH Flgwe 4. (a) Solution composition dependence of the proton dissociation rates kd (n),k, ( 0 ) ,and k, (V) and the proton recombination rates k, (A), k2 (0),and k, ( 0 )for 2-np3,6. (b) Solution composition dependence of kd ( 0 )and k, (a) for 2-np6,8. 0
Most of the work on excited-state intramolecular proton transfer (ESIPT) has dealt with two types of reactions: those in which there exists a hydrogen bond between the hydrogen of the donor group and the atom of the acceptor group and those in which the proton is far away from the acceptor and requires a mediator. In the former case, an intramolecular hydrogen bond provides a
The Journal of Physical Chemistry, Vol. 96, No. 18, 1992 7327
Proton Dissociation in 2-Naphthol-3,6-disulfonate
s
molecular proton transfer from sulfonate to water has the lowest isotope effect, kH+/kD+ 1.8. The isotope effect for the deuteron recombination either to the RO-group or to the SO3- group is small: kH+/kD+ 1.2. The magnitude of this isotope effect is consistent with our previous finding of a single-step proton transfer to the s01vent.’~ There is a large difference in isotope effect between the proton-transfer rate to the solvent from OH (3.6) and from S03H (1.8). As mentioned above, the hydroxy aromatic compounds studied so far exhibit almost the same isotope effect of -3.’J9 This relatively large value probably reflects both the molecular properties of the donor, including the aromatic conjugation, and those of water. The sulfonate group is less affected by the aromatic ring and hence has a smaller isotope effect. As we showed in the water/methanol cosolvent study, the intramolecular proton-transfer rate is mediated by the solvent. The intermediate isotope effect of 2.2 might be due to the small isotope sensitivity of the proton acceptor, SO3-.
-
0
d
2.0
0.0
ao
60
4.0
M
4 4
i 2
3
4
i
6
7
8
9
!
0
Figure 5. ROH* experimental fluorescence decay curves of 2-np3,6 measured at 370 nm (dots) with the numerical solution of the DSE (full curves) (A) in H20and (B) in D20.The lower panel represents the residual of 2-np3,6 in D20.Parameters for the fits are displayed in Table 111.
0.58 11 4.5 2.0 2.7 3.0 0.93 3.56
0.16 5.0 2.8 2.5 1.5 2.5 0.62 2.50
3.6 2.2 1.61 1.2 1.8 1.21
Reduced x2 formula taken from ref 30. suitable reaction coordinate for a fast proton transfer. Here, hydrogen-bonding solvents such as alcohols and ethers decrease the proton-transfer rate due to the creation of solute/solvent complexes. In the second case, a suitable solvent which forms hydrogen bonds with the solute can catalyze the ESIFT pr0cas.2~ The first instance involves systems in which the proton donor is an oxygen, as in hydroxyanthraquinones and its derivatives, or a nitrogen, as in aminoanthraquinones and its Konijnenberg et al.29have done extensive work with molecules that belong to the second category such as 2-aminopyridine and 7-azaindole. Ours is the first investigation into the influence of solvent on the ESIPT from a hydroxyl group to a sulfonate group as occurs in 2-np3,6. From our experimental work, it can be seen that the intramolecular proton transfer is water catalyzed and barely takes place in methanol. C. Isotope Effect. Figure 5 shows the time-resolved ROH* and ROD* (marsured in HzO and DzO, respectively) fluorescence data of 2-np3,6 and their computer fits. Table I11 summarizes the computer-fitting parameters. The data show that the dissociation rates are more affected by deuteration than the association rates. These results are in accord with our previous results with HPTS in D20. The rate (kd) of direct proton transfer from the OH m p to the W h l t Changes m a t Signifihllfly where kH’/,kD’ -3.6. This ratio is close to the value of -3, which was obtamed by us for HPTS.’O A similar dependence was found by Krishnan et al.19 for the dissociation rate constant of 1-naphthol-2-sulfonate. They explained this dependence as a change in entropy arising from the lower librational frequencies of liquid D20than HzO. The intramolecular deuteron transfer from OH to SO3- exhibits an intermediate isotope effect of kH+/kD+ 2.2, while the inter-
-
Conclusion In this study, we examined the effect of an adjacent sulfonate group on excited-state proton transfer from hydroxyarene compounds. As a probe molecule, we chose 2 - n a p h t h o l - 3 , 6 t e (2-np3,6). We compared the experimental data of 2-np3,6 with a similar compound, 2-naphthol-6,8-disulfonate(2-np6,8). In electronically excited 2-np6,8, a proton is transferred to water at a rate of 100 ps. The proton geminate recombination process is quantitatively described by the DSE. With 2-np3,6, the fmt step is an intramolecular proton t r a d e r from the hydroxy to the adjacent sulfonate. The back-proton-transfer rate is slower by a factor of 2. The proton is transferred to the solvent mainly via the sulfonate group. In water/methanol mixtures, all of the proton-transfer rates are lower than in water and are strongly dependent on the mixture composition. The smaller the water mole fraction] the smaller the rates. Thus, we conclude that the intramolecular protontransfer process is mediated by water molecules. In the absence of water molecules (pure methanol), all of the proton dissociation rates decrease by at least 3 order of magnitude. Acknowledgment. We are grateful to Professor P.M.Rentzepip and Drs. N. Agmon and E. Pines for helpful discussions, This work is supported in part by Grant 88-00125 from the US.-Israel Binational Science Foundation (BSF) and by a grant from the James Franck Binational German-Israeli Program in LaserMatter Interaction. Registry No. 2-np3,6,46992-11-4; 2-np6,8,68384-00-9; D2, 778239-0.
References and Notes (1) Bell, R. P. The Proton in Chemistry, 2nd ed.; Chapman and Hall: London, 1973. (2) Fhter, Th. Z . Elektrochem. 1950, 54, 531. (3) (al A. Weller. Proa. React. Kiner. 1961, I . 189. (bl . , Z . Phvs. Chem. N.F.‘19j8, 17, 224.. (4) Ireland, J. F.; Wyatt, P. A. H. Ado. Phys. Org. Chrm. 1976,12, 139. (5) Kosower, E. M.; Huppert, D. Ann. Rev. Phys. Chem. 1986,37, 127. (6) Lee,J. J . Phys. Chrm. 1990, 94,258. (7) (a) Agmon, N.; Pines, E.; Huppert, D. J. Chem. Phys. 1988,88,5631. (b) Pines, E.; Huppert, D.; Agmon, N. J. Chem. Phys. 1988, 88, 5620. (8) M. Von Smoluchowski, Z . Phys. Chrm. 1917,92, 129. (9) Debye, P. J. Electrochem. Soc. 1942, 82, 265. (10) Pines, E.; Huppert, D. J . Chem. Phys. 1986,84, 3576. (1 1) Pines, E.; Huppert, D. Chem. Phys. Lert. 1986,126, 88. (12) Masad, A.; Huppert, D. Chem. Phys. Len. 1991, 180,409. (13) Henson, R. M. S.; Wyatt, D. H. J. Chem. Soc., Faraday Trans. 2 1975, 669. (14) Schulman, S. G. Anal. Chim. Acta 1976,84,423. (15) Van Gemert, J. T. Austral. J . Chem. 1969, 22, 1883. (16) Zaitsev, N. K.; Demyashkevich,, A. B.; Kuzmin, M. G. Khim. Vysokikh Energ. 1978, 12, 436, [English translation, p 3651. (17) Shapiro. S.L.; Winn, K. R.; Clark, J. H.SpringerSer. Chem. Phys. 1980, 14, 221. (1 8) Krishnan, R.; Fillingim, T. G.; Lee. J.; Robinson, G. W. J. Am. Chrm. Soc. 1990. 112, 1353. (19) Krishnan, R.; Lee,J.; Robinson, G. W. J . Phys. Chem. 1990, 94, 6365. (20) Cornish, B. D.; Speedy, R. J. J. Phys. Chem. 1984, 88, 1888.
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J. Phys. Chem. 1992, 96, 7328-7331
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(21) Mashimo, S.;Kuwabara,S.;Yagihara, S.;Higasi, K. J. Chem. Phys.
1989, 90, 3295.
(26) Brucker, G. A.; Swinney, T.C.; Kelley, D. F. J . Phys. Chem. 1991, 3190. (27) Barbara, P. F.; Walsh, P. K.; Brus, L. E. J . Phys. Chem. 1989, 93, 29. (28) Flom, S.R.; Barbara, P. F. J . Phys. Chem. 1985, 89, 4489. (29) Konijnenberg, J.; Huizer, A. H.; C.A.G.O. Varma, J . Chem. Soc., Faraday Trans. 2 1989, 85. 1539 and references therein. (30) Cramer, H. Mathematical Meihods of Siatistics; Uppsala: Almquist and Wiksells, 1945. 95,
(22) Agmon, N.; Huppert, D.; Masad, A.; Pines, E. J. Chem. Phys. 1991,
95, 10407.
(23) Conway, B. E.; Bockris, J. OM.; Linton, H. J . Chem. Phys. 1956, 24, 834. (24) Kolodney, E.; Huppert, D. Chem. Phys. 1981.63, 401. (25) See: "Spectrmcopyand Dynamics of the Elementary Proton Transfer in Polyatomic Systems". Chem. Phys. 1989, 136.
Reaction of Atomic Bromine with Difluorochloromethane. The Heat of Formation of the CClF2 Radical and the D (CCIF2-H) Bond Dissociation Energy K. Miyokawat and E. Tschuikow-Row* Department of Chemistry, University of Calgary, Calgary, Alberta, Canada T2N 1N4 (Received: December 16, 1991; In Final Form: April 27, 1992)
The gas-phase photobromination of CHClF, (1) in the presence of CHICl (2) as competitor has been studied in the temperature range 80-150 OC at halomethane pressures of -35 Torr and a Br2 pressure of -2.3 Torr. The temperature dependence of the rate constant ratio is found to obey the Arrhenius expression In ( k l / k z )= (4.0885 f 0.0580) - (1 144 f 20)/T. This result is combined with an earlier relative study of CH3Cl(2) v8 C!& (3) and a recent direct determination by kinetic specmmpy of the rate constant for the bromination of ethane (kJ to obtain absolute rate parameters for the reaction CHClF, Br CClF, + HBr. Using a justifiable approximation concerning the magnitude of the activation energy difference for the reverse reactions between any two competitors of similar complexity, and other thermochemical data from the literature, the following quantities have been derived: AH~0298(CC1F2) = -66.7 f 2 kcal mol-I and Do(CClF2-H) = 100.7 f 2 kcal mol-' where the uncertainties are conservative estimates. On the basis of the new value of k3, activation energies for the bromination of CH4 and other halomethanes are reported.
-
+
Introduction
There is a growing interest in the physical and chemical properties of hydrochlorofluorocarbons (HCFCs) as alternatives to the widely used chlorofluorocarbons (CFCs). These HCFCs are expccted to have a relatively shorter lifetime in the troposphere as they react with hydroxyl to yield chlorofluorocarbon radicals which are then converted to compounds less harmful with respect to ozone depression in the stratosphere.' Therefore, it seems relevant to investigate the thermochemical properties of chlorofluorocarbon radicals. Since difluorochloromethane (HCFC-22) is considered to be a prominent candidate among the alternatives to CFC-11 and -12, we report here the kinetics of the photobromination of CHClF2 and thermochemical data for the CClF2 radical estimated from this and other kinetic data.
Experimental Section With one exception, all chemicals were obtained commercially: CHClF2 and CH3C1from Matheson; CH2BrCl from Columbia Organic Chemicals Co.; and ACS grade Br2from Fisher Scientific Co. Prior to their use, all samples were subjected to the usual trap-to-trap distillation and degassing under vacuum at liquid nitrogen tem'pemture until the impurity levels were below the GC detection limit. CBrClF,, needed for calibration (seebelow), was prepared by the bromination of CHClF2. For this purpose a mixture of CHClF2 (47 Torr) and Br2 (1 Torr) was irradiated until GC analysis of the mixture showed the absence of Br2 (under the experimental conditions the detection limit for Br2 was found to be 0.075 Torr, indicating that at least 92.5% of Br2 had reacted with CHCIF2). Kinetic experiments were carried out in a greaseless static system, and the details of the experimental apparatus and procedure have beem described elsewhere.2 Reaction temperatures ranged from 80 to 150 OC and were maintained within 0.1 OC 'Present address: Hitotsubashi University, Kunitachi, Tokyo, Japan.
0022-3654/92/2096-7328S03 .OO/O
by circulating an ethylene glycol/water solution (approximately 4:l) through the outer jacket of the cylindrical Pyrex reactor. The irradiation time was varied from 2.5 to 30 min depending on the reaction temperature to keep the formation of the undesirable secondary bromination products as low as possible. Product analysis was carried out by isothermal gas chromatography (1 75 "C) with an electron capture detector (ecd)and a Durapak R column of 0.85-m length. Calibration curves of relative peak area vs pressure for the products CBrClF2and CH2BrCl were determined by GC analyses of known amounts of these gases diluted with CHClF2, yielding the linear relationships S(CH2BrCI) = (3177 f 50)P(CH2BrCl) - (9.34 0.95) X 104; s = 1 x 105 to 5 x 105 S(CBrClF2) = (2354 f 28)P(CBrClF2) - (1.31 f 0.08) X 104; s = 1 x 104 to 15 x 104
*
where S denotes peak area (counts, range as indicated) and P the partial pressure of the compounds (mTorr). In the bromination experiments the observed peak area varied from 1 X lo5 to 5 X lo5 counts for CH2BrCland from 2.5 X 10" to 6.5 X 10" counts for CBrClF2. Thus,the uncertainty of the product ratios arising from the calibration factors is evaluated to be 4.2-5.9%. Preliminary experiments confirmed the absence of dark reactions: when a CHClF2/CH3Cl/Br2(1oO:2O:6) mixture was kept in a shielded reactor for 30 min at 150 OC, no products were found. For the irradiated mixtures the observed carbon-containing products were CBrCIF2and CH2BrCIexcept for the experiments at 80.0 OC where trace amounts of CHBrzCl (secondary bromination product of CH,Cl) were also detected, owing to the long photolysis times required at this low temperature. Results and Discussion Kinetic P a m " . The kinetcs of gas-phase photobromination of CHClF2 between 80 and 150 OC has been studied by the competitive method using CH3Cl as an external reference and 0 1992 American Chemical Society
