J. Phys. Chem. 1991, 95,6569-6574
6569
Kinetics of the Reaction of Molecular Fluorine wlth Dimethyl Sulfide Andrew A. Turnipseed and John W.Birks* Department of Chemistry and Biochemistry and Cooperative Institute for Research in Environmental Sciences (CIRESJ,Campus Box 216, University of Colorado, Boulder, Colorado 80309-0216 (Received: January 17, 1991; In Final Form: April I O , 1991)
The reaction of molecular fluorine with dimethyl sulfide (DMS)was investigated by using the flow tube technique with mass spectrometry for detection of reactants and products and a diode array spectrometer for collection of chemiluminescence spectra. The reaction was found to proceed with an extraordinarily large rate coefficient of k4 = (1.6 f 0.5) X lo-'' cm3 molecule-' s-' at 298 K. The observation of such a fast rate for the reaction of two closed-shell species is unexpected. It is proposed that the reaction proceeds through a charge-transfer complex with subsequent elimination of H and F atoms or of molecular HF. Once initiation of the reaction occurs, a radical chain reaction leads to chemiluminescent products. A reaction intermediate, thought to be H2C=SFCH3, occurs at mle = 80 when DMS is in excess over F2. Under these conditions, the principal emitting species is H 2 M . When F2is the excess reactant, the intermediate at m/e = 80 is destroyed and the principal emitting species is HCF. A charge-transfermechanism for initiation of F2reactions qualitatively explains the relative responses of classes of compounds in the fluorine-induced chemiluminescencedetector. The F2 + Hfi reaction cm3 molecule-' s-' was measured. was also studied, and an upper limit of k7 5 6.4 X
Introduction
The kinetics of atomic fluorine have long been of interest due to its significance in molecular dynamics studies" and applications to chemical lasers.' Atomic fluorine reactions often produce visible and infrared chemiluminescent products as a result of the high degree of exothermicity and the extremely fast rates of F atom abstraction reactions. Reactions of molecular fluorine are generally thought to proceed by a self-propagating radical chain reaction initiated by F atoms
F + RH--HF+ R R
+ F2-+
RF
+F
(2)
where R is an organic radical or hydrogen. The second step of this sequence provides a large fraction of the total exothermicity. This can be seen by comparing the strengths of bonds broken and formed in each reaction. H F (1 35 kcal mol-') is formed in the first step at the expense of breaking a C-H bond ( ~ 9 kcal 9 mol-'). A C-F bond ( ~ 1 1 6kcal mol-') is formed in the second step, while the relatively weak F2 bond (37.5 kcal mol-') is broken. The net energy released in the initial abstraction step is about 36 kcal mol-', whereas the net available energy in reaction 2 is much higher (=79 kcal mol-'). Therefore, a large fraction of the energy released in these chain reactions is a result of reactions of molecular fluorine. However, kinetics studies of molecular fluorine are scarce and few measured rate coefficients have been reported. Previous work in our laboratory found that when molecular fluorine was added to the effluent of a gas chromatograph (GC) intense visible and near-infrared chemiluminescencewas observable for most organosulfur compounds.' Emissions from vibrationally excited HF(Av=4,v18), CH#(a3A2+X1AI), HCF(A'B+X'Z), and electronically excited SF2have been observed under various condition^.^.^ Chemiluminescent reactions of F2 form the basis of a detector applied to gas chromatography: high performance liquid chromatography,' and supercritical fluid chromatography8 The detector exhibits no significant response to most hydrocarbons, (1) Anlauf, K. G.; Mayotte. D. H.; Pacey, P. D.; Polanyi, J. C. Phys. Lcrr. A 1%7,24,209. (2) Polanyi, J. C.; Sloan, J. J. J . Chem. Phys. 1972, 57, 4988. (3) Kompa, K. L.; Pimental, G. C. J . Chem. Phys. 1967, 47, 857. (4) Nelson, J. K.;Getty, R. H.; Birks, J. W. Anal. Chem. 1983,55, 1767. (5) Glinski, R. J.; Mishalanie, E. A.; Birks, J. W. J. Phorochem. 1987, 37,
217. (6) Glinski, R. J.; Taylor, C. D. Chem. Phys. h i t . 1989, 155, 51 1. (7) Mishalanie, E. A.; Birks, J . W. Anal. Chcm. 1986, 54,918. (8) Foreman, W. T.; Shellum, C. L.; Birks, J. W.; Sievcrs, R. E. J . Chromatogr. 1989, 23.
chlorinated hydrocarbons, water, or alcohols, so that organosulfur compounds can be detected selectively over these classes of compounds. Recent work has shown that volatile organoselenium, -tellurium, -phosphorous, and -arsenic compounds also exhibit intense chemiluminescence from reaction with F2.9.10 The selectivity of the detector suggests a gas-phase reaction in which F2 reacts directly with the organosulfur compound. If atomic fluorine were involved in the initial step, nearly all organic compounds would be expected to respond due to lack of selectively of F atom reactions. Arrested relaxation studies have shown that F atom reactions with H# and CH'SH both produce vibrationally excited H F up to u = 4,'' which could explain some of the chemiluminescenceobserved in the detector; however, H2S is found not to give a response in the detector, whereas CH$H is very sensitively detected. Those studies also found that the reaction of dimethyl sulfide (DMS)(CH3SCH3)with F atoms exhibits similar characteristics to the analogous reaction with dimethyl ether (CH30CH,).'2 Dimethyl sulfide is one of the more sensitively detected compounds in the detector whereas dimethyl ether shows no response, Thus, the detector exhibits much greater selectivity than would result if the reaction were initiated by F atoms. The goal of the present study was to use the flow tube technique, with both mass spectrometry and chemiluminescencedetection, to try to elucidate some of the major steps of these reactions so as to better understand the selectivity exhibited in the fluorineinduced chemiluminescence detector invented in our laboratory. The reaction conditions within the flow tube are similar to the conditions used in the GC detector. The detector is operated at low pressure (0.5-5 Torr) with a residence time of about 100 ms. Although most experiments were conducted in a shorter time frame than the detector, higher concentrations were used which tends to force the chemistry to completion. The addition of mass spectrometry as a means of detection allows for the detection of nonchemiluminescent species. The identity of these species could only be inferred in previous studies from their emission spectra. Although a wide range of sulfides and thiols respond in the detection system, the reaction of dimethyl sulfide (DMS)with F2 was chosen as a model reaction. DMS also exhibits high sensitivity within our detector and is a relatively simple molecule. The hydrogen analogue of DMS, hydrogen sulfide, is interesting, since (9) Chasteen, T. G.; Silver, G. M.; Birks, J. W.; Fall, R. Chromaropphfa 1990, 181.
(10) Chasteen, T. G.; Fall, R.; Birks, J. W.; Martin, H. R.; Glinski,R. J. Chromatographia, in press. ( 1 1 ) Dill. B.; Heydtmann, H. Chem. Phys. 1978, 35. 161. (12) Duewer, W. H.; Setser. D. W. J . Chem. Phys. 1973,58, 2310.
0022-365419112095-6569$02.50/0 0 1991 American Chemical Society
6570 The Journal of Physical Chemistry, Vol, 95, No. 17, 1991
1-
vacuum
Multichannel Analyzer
1
Diffusion Pumps
Ousnz Window
vacuum
t Figure 1. Diagram of the flow tube/mass spectrometer system used in the present experiments. Inset shows the placement of the flow tube when chemiluminescence spectra were obtained by using the intensified diode array spectrometer (IDARSS).
it shows virtually no response in the F2-inducedchemiluminescence detector. The reaction of F2 with H2S has also been briefly examined in attempts to better understand the mechanism of detector response.
Experimental Section A diagram of the flow tube/mass spectrometer system is shown in Figure 1. The flow tube technique" and its coupling to mass ~pectrometry'~ have been described in detail elsewhere. The main reactor consists of a 80-cm-long Pyrex tube with an inner diameter of 2.51 cm. A large flow of He is introduced into the Pyrex flow tube to establish a known flow rate, pressure, and linear velocity. One reactant is added directly to a fixed side arm in the flow tube. A second reactant is then added through a movable injector consisting of a 140-cm-long, 0.95-cm-0.d. Pyrex tube, concentric with the main reactor flow tube. The pressure within the flow reactor was varied between 1.3 and 4.5 Torr with the linear velocity kept at 1500 f 15 cm s-I. This resulted in total flow rates of 11-38 cm3 (STP) s-I. F2was added through the movable injector, and DMS was added to the fixed side arm in most experiments. These conditions were reversed and found not to have an effect on the reaction kinetics. The injector was translated over 22 cm,resulting in a reaction time of up to 15 ms. The flow partitioning between the main reactor and the movable dual injector was such that the flow through the injector made up 5 1 0 % of the total flow. All flows were measured by using calibrated mass flow meters (Tylan FC260 and FC200, Teledyne-Hastings NALL- 100). and the pressure was measured by a capacitance manometer (MKS Baratron Model 170M). All experiments were carried out at room temperature. The majority of the flow is pumped by a mechanical vacuum pump (Sargent-Welch No. 1375). A small fraction of the flow mixture is sampled through a pinhole into a differentially pumped quadrupole mass spectrometer (UTIlW)operated at an electron impact energy of 20-60 eV. Ion detection was accomplished by means of a high current Channeltron electron multiplier (Galileo Electrooptics No. 47 17) which was maintained at 2300-2900 V. Selected ion monitoring was used to monitor the concentrations of all important species. The detection limits for CH3SCH3and ~, The F2were 6 X 1O'O and 3.0 X 10" molecules ~ m -respectively. relatively high detection limits obtained for F2were the result of a high hydrocarbon background signal over the range of m / e = 38 to mle = 45. Also, the F2+ion appears to fragment to a large extent in the mass spectrometer. (13) Howard, C. J. J . Phys. Chem. 1979.83, 3. (14) Hills. A. J.; Cicerone, R. J.; Calvert, J. G.; Birks, J. W.J . Phys. Chem. 1988, 92, 1853.
Turnipseed and Birks In experiments where chemiluminescence emission was monitored, the flow tube was moved back 5 cm from the mass spectrometer. Low-resolution chemiluminescence spectra were obtained by the placement of an intensified diode array rapid scanning spectrometer (IDARSS, Tracor 1710). This configuration is shown in the inset of Figure 1. It allowed the simultaneous collection of chemiluminescencespectra and mass spectra. Kinetics data from the mass spectrometer is not possible in this configuration due to the perturbation in both the flow profile and the linear velocity caused by the 5-cm gap between the flow tube and mass spectrometer entrance. Although there is a loss in sensitivity by the mass spectrometer, this configuration was useful in detecting nonchemiluminescent products simultaneously while monitoring chemiluminescent emission. When the mass spectrometer was used to gather kinetics data, the flow tube was inserted to within 1 mm of the mass spectrometer entrance. The diode array spectrometer disperses the emission spectrum across 5 12 individual photodiode detectors. Wavelengths were calibrated against a He line source. Resolution of the He 587.6-nm line was 22 nm, fwhm. In order to increase the signal-to-noise ratio, a long scan time (8 s) was used and 25-40 spectra were signal averaged. The transformation of reaction time to flow tube position in the fast flow technique makes long signal averaging times possible. This is an important advantage of this technique for studying chemiluminescent reactions. Cylinder gases for this study included helium (UHP), nitric oxide (CP), and a 9% fluorine/He mixture (Matheson). The fluorine was then used directly or admitted into an evacuated 22-L glass bulb and diluted with He (loo0 Torr) to produce mixtures of 0.5%. When F2concentrations needed to be known accurately, these mixtures were allowed to sit for 2-3 h to mix thoroughly and then used the day of the experiment. Chlorine was obtained in a mixture of 9.8% C12 (HP)/He (UHP) and used directly. Dimethyl sulfide (Aldrich) was prepared from the pure liquid by admitting the vapor into evacuated, conditioned, 22-L glass bulbs. This was then diluted with He (1656 Torr) to produce mixtures of 0.05-3.0%. H2S (CP, >99.6%) from a lecture bottle was evaporated into 22-L glass bulbs and diluted in a similar fashion to dimethyl sulfide.
Results Initially, to verify that the flow tube system was in working order, the rate coefficient for the reaction Fz + N O FNO + F (3) +
was measured under pseudo-first-order conditions where [NO] = (1&100)[F2] and the decay of F2at m / e = 38 monitored with time. N O concentrations were varied in the range (1-10) X 1OIs molecules ~ m - Thirty ~ . measurements were made and the observed rate coefficient was found to be (1.2 f 0.3) X cm3molecule-' s-I. This is in relatively good agreement with the previously measured15 value of 1.5 X lo-'* cm3 molecule-' s-I. The kinetics of the reaction Fz + CH3SCH3 products (4)
-
was studied under pseudo-first-order conditions such that [F2]= (10-25)[CH3SCH3]. A concentration of (6-9) X 10" molecules cm-3 of dimethyl sulfide was added to the flow reactor through either the fixed side arm or the movable injector. Fluorine was added via the remaining source to produce flow tube concentrations ~ o. effect on in the range of (0.90-2.3) X lOI3 molecules ~ m - N the kinetics was observed upon switching the DMS source from the side arm t o the movable injector. Confirmation of pseudofirst-order conditions was obtained by monitoring the F2 ( m / e = 38) signal. The F2signal was observed not to vary more that 10%over the entire reaction zone. Figure 2 shows examples of typical decay plots at various F2concentrations. Corrections for wall loss of DMS and F2were unnecessary, as wall loss was found to insignificant ( [DMS]. Diode array spectra under these conditions (Figure 6b) show strong emission from HCF(A'Z+X'Z) and possibly the Av = 4 sequence of vibrationally excited H F as identified by Glinski et al. from a similar F2/organosulfur reaction.s However, due to the large variety of possible products it is not possible to isolate any distinct reaction steps under these conditions. The rate coefficient of for the reaction Fz + HIS products (7) +
was measured under conditions where [H2S] >> [F2]. H$ concentrations within the flow reactor were in the range of (6.*15.0) X IO" molecules ~ m - ~ Under . these conditions, no reaction could be detected over a reaction time of 25 ms. An upper limit for the rate coefficient of reaction 7 was placed at k,(298) 5 6.4 X cm3 molecule-' s-I. (16) Glinski. R.J.; Getty, J. 359.
N.;Birks, J. W.Chem. Phys. Lcrr. 1985, 117,
i 200
300
400
500
608
100
800
Wavelength, nm
Figure 6. (a) Mass spectra for [DMSIo = 3.04 X lOI3 and [ F z ] = ~ 2.8 X I O l 4 molecules cm-' and a reaction time of 11 ms. Peaks at m / e = 18 and m/e = 28 are due to a small air leak into the system. Peak marked with * is due to background hydrocarbons. (b) IDARSS spectra obtained for [DMSIo = 1.3 X 10'' and [F2I0= 6.7 X 10" molecules cd.
Discussion Initial studies of the reaction between DMS and F2indicate that the reaction proceeds with a rate coefficient of (1.9 f 0.5) X lo-" cm3 molecule-' s-I. However, the significant positive intercept in the plot of kI4 vs [F,] and the lower value of the rate coefficient obtained from the slope (k4= (1.56 f 0.12) X lo-" cm3 molecule-' s-I) suggests that secondary chemistry is influencing the observed removal rate of dimethyl sulfide. It is likely that the reaction initially produces an F atom which then either recombines on the wall of the flow tube or destroys another dimethyl sulfide molecule. In case of no wall loss of F2, the rate of decay of dimethyl sulfide would be greater by a factor of 2, and the rate coeficient would Likewise be overestimated by a factor of 2. The appearance of ClF when C12was added to the reaction mixture indicates that F atoms are being produced within the reaction. By the production of a highly reactive F atom, it is possible that a self-propagating chain reaction is initiated which can destroy the dimethyl sulfide rapidly and cause an overestimation of the rate coefficient. To establish whether such a chain reaction could be occurring, a simple model of the chemistry was undertaken. Table I shows the reactions and the rate coefficients used for this simulation. In many cases, the rate coefficients were unknown for these reactions and were estimated by comparison to similar reactions (Le., F atom abstraction reactions generally accur at cm3 molecule-' d). From the near collisional rates, k a: model, it was found that the addition of secondary chemistry did affect the measured rate coefficient by accelerating the destruction of DMS. By using concentrations of [DMS] = 6 X 10'' and [F2] = 1 x 10') molecules ~ m - it~ was , found that a rate coefficient of 1.1 X lo-'' cm3molecule-' PI would result in an inferred value of 1.9 X lo-'' cm3 molecule-' s-I. A slight negative curvature was noticeable in the simulated decay plots; however, this curvature was small enough that it would be difficult to verify experimentally. It should also be noted that wall loss of F atoms and organosulfur
Reaction of Molecular Fluorine with Dimethyl Sulfide
The Journal of Physical Chemistry, Vol. 95, No. 17, 1991 6513
TABLE I: Rate Coefficients Used for Modeling the Fz + DMS
TABLE 11: Critical Distances for Cbrrge-Tnmfer Complexes a d Reported Detection Limits in F2-Iduced C h e m i l u " e c
Reaction
k, cm3 molecule-' reaction (4) Fz + CHSSCH, F + CH3SFCH3
-- +
S-'
k,
(12) C H S F C H , H CHSFCH, ( 6 ) . F + bH3SCH3 H F + kHzSCH3 (15) H CH3SCH3 4 CH3 CH3SH (14)H+FzdHF+F (20) CH3 Fz --L CH3F + F (21) CHZSCH3 F, CHzSFCH3 F (22) CHpSFCH3 F2 CH,FzCH3 F (23) F CHJSFCH, 4 H F CH2SFCH3 (24) F CH3SH 4 H F CH3S (16) H CHJS Hz + CH2S (17) F CH$ HF CH2S (25) F H2 4 H F H
+
+
+
+ + + +
+
+
+
+
-
+
+
+
+
+
+ +
+
2 X IO3 s-' 2 x 10-10 9.8 X 3.7 x IO-" 2 x 10-1' 2 x lo-" 2 X lo-'' 2X 2 x 10-'0 2 x 10-10 1 x 1o-'O 1.8 X IO-"
note ..... a
b
dimethyl disulfide
C
trimethylphosphine m-xylene
d e
f
g g C
c
c c
e
#Studied reaction. bAssumed lifetime 5 X lo4 s based on similar complexes of F + olefins.zs eAssumed to be nearly collisional. dYoleota,T.; Straw, 0. P. J . Phys. Chem. 1979,83, 3196. 'Foon, R.; Kaufman, M.Prog. Read. Kiner. 1975, 2 , 8 1 . /Jones, W. E.; Skolnik, E. G. Chem. Rev. 1976, 76, 563. 'Assumed to be similar to CH, + F2 reaction. radicals could be significant and affect the kinetics by removing reactive radicals from the reaction mixture, thereby reducing the amount of secondary chemistry. Any loss of radicals to the wall increases the accuracy of the measured rate constant. Thus, we are confident that the true rate constant lies between 1.1 X and 1.9 X lo-" and we recommend (1.6 f 0.5) X lo-" cm3 molecule-' s-' for the rate coefficient for the initial F2 + DMS reaction. The error allows for both the upper and lower limits established by kinetics modeling, and the central value is identical with the result obtained from the slope of the plot of kt4vs [DMS]. In principle, a more accurate measurement of the rate coefficient could be determined by using excess DMS over Fzras was attempted in the present study. Simulations indicated that, under conditions of excess DMS, the input rate coefficient could be extracted accurately from the model. The excess DMS limited the effects of secondary chemistry by scavenging F atoms in these simulations. However, due to the lack of sensitivity for F,, the value obtained for k4 under these conditions, k4 = (1.7 f 0.8) X 10-" cm3molecule-' s-l, is less precise but in good agreement with the value obtained by using Fz in excess. The observation of such a high rate for the reaction of two closed-shell molecules is unexpected. At present, we can still only speculate at the possibilities. It has been found from studies of the G C detector that chemiluminescence is only observed for compounds with low ionization potentials (19.0eV).497-9 Since molecular fluorine has the highest electron affinity (3.0 eV)" of any stable molecule, it is possible that the reaction is initiated by the transfer of an electron from the sulfur compound to F,, forming a charge-transfer complex.
CH3SCH3
+
F2
-C
[
'F$ \ CH3 CH3
I'
compound dimethyl sulfide dimethyl selenide dimethyl telluride
(8)
The driving force in this mechanism is electrostatic attraction between F2 and DMS. An equation from BensonI8 6E = IP(DMS) - EA(Fz) - ( e 2 / r ) (1) can be used to determine the stability of this metathesis complex, where bE is the energy difference in the reaction, IP(DMS) is the ionization potential of dimethyl sulfide, EA(Fz) is the electron affinity of molecular fluorine, e is the charge on an electron, and r is the distance between the two molecules. For the electron to (17) Franklin, J. L.;Harland, P. W. Annu. Reu. Phys. Chem. 1974, 25, 485. (18) Benson, S. Thermochemical Kinetics, 2nd ed.; Wiley: New York, 1976.
toluene
tetrahydrofuran hexene
iodohexane propionaldehyde hydrogen sulfide acetonitrile water
re, A
detection limit,' pg
2.63 >2.63 >2.63 2.64 2.57 2.58 2.47 2.24 2.22 2.33 2.06 1.94 1.56 1S O
16 30 7 15 0.5 1700 840 2500 8.4 x 104 1.7 X 10' 8.4 x 10s >8 X 1OIo >8 X 10Io
>8
X 1O'O
Detection limits are from refs 4 and 7-9. be transferred, the reaction must be either exothermic or thermoneutral (Le., 6E I0). By setting 6E equal to zero, a critical distance, r,, can be calculated to give the maximum distance a t which the charge-transfer complex can be stable and the electron can be transferred: r, =
14.4
(11) IP(DMS) - EA(F2) The ionization potential of DMS is 8.48 eV,19 and the electron affinity for Fz is 3.0 eV.17 Substitutin these values in the above equation, a critical distance of 2.6 is calculated. The requirement of formation of an intermediate charge-transfer complex would explain the slow reaction of HzS. The ionization potential of H,S is 10.4 eV and the calculated critical distance, r,, is calculated to be 1.9 A, which requires a very close approach between the two species before a reaction can occur. The repulsive interaction between the two molecules would limit such encounters to strong collisions and thus high temperatures. Table I1 is a tabulation of the calculated critical distances, r,, for a range of compounds and the reported detection limits obtained by F2-induced chemiluminescence. As can be seen, a significant chemiluminescence response (indicating a reaction is occurring with a measurable rate) is only exhibited by compounds which have a critical distance of greater than about 2.3 A. Compounds containing large atoms such as sulfur, iodine, and phosphorous are known to expand their octet and form extra bonds to highly electronegative elements such as fluorine (Le., SF,, IFS, and PFJ. Also, it is well known that complexes between reduced sulfur compounds and radicals such as OH form in the gas phase." Formation of an intermediate complex may be facilitated by the use of the empty d orbitals on the sulfur atom. This provides one possible explanation for why oxygenated analogues, which have no available d orbitals, do not exhibit chemiluminescence. However, oxygen analogues also have much higher ionization potentials than S,Se, Te, and P compunds, and according to our charge-transfer complex theory do not respond in the chemiluminescence detector because of a very slow initiating reaction. It has been observed that F2does add to compounds containing large atoms such as iodine compounds. Crossed molecular beam studies of Iz, HI, IC1,2' and CH31z2with F2 have shown that the complexes, IIF, HIF, FICI, and CHJF are formed. The reaction 1, + Fz IIF F (9)
1
-
+
has been shown to occur with a rate constant of 1.9 X cm3 molecule-' s-I at room temperature.,, It was found that there is a significant barrier (4.2 kcal mol-')24 to the formation of IIF (19) Weast, R. C. Handbook of Chemistry and Physics, 63rd ed.; CRC Press: Bcca Raton, FL, 1982-83. (20) Tyndall. G.; Ravishankara, A. R. Inr. J . Chem. Kinel., submitted for publica tion. (21) Valentini, J. J.; Coggiola. M.J.; Lee, Y. T. Faraday Discuss. Chem. SOC.1977. 62, 232. (22) Farrar, J. M.;Lee, Y. T. J. Am. Chem. Soc. 1974, 96, 7570. (23) Whitefield. P. D.; Davis, S . J. Chem. Phys. Lett. 1981, 83, 44.
6514 The Journal of Physical Chemistry, Vol. 95, No. 17, 1991 in this reaction. The calculation of the critical distance, r,, for an F2-12 charge-transfer complex leads to a value of 2.3 A, which is near the cutoff for compounds exhibiting chemiluminescence in our detector. The activation barrier observed in this reaction may be due to the fact that these molecules must come close together for an electron to be transferred in order to initiate the reaction. It also should be noted that the species IIF has been implicated in the formation of excited state IF(B3n) observed in previous studies of the reaction of I2 F225by means of the reaction2' F
+ IIF
-
+
+
IF (B3n) IF
(10)
Once the DMS.F2 complex is formed, atomic fluorine may be eliminated, analogous to the proposed mechanism for F2 + olefins.26
Turnipseed and Birks H
+ CHpSCH3
+
CH3 + CHSSH
(15)
may be favored, so that HF(Au-4) emission is reduced considerably, as is observed. At high DMS concentrations (ZlOI4 molecules ~ m - ~ emission ), from the triplet state of thioformaldehyde is clearly observable at about 694 nm. Secondary chemistry must be involved to produce this emitter. The reactions H + CH3S H2 + CH2S (16) F
+ CH3S
-
+
HF
+ CH2S
(17)
are exothermic by 62 and 94 kcal mol-I, respectively, and either reaction could be responsible for the formation of CH2S(A3Z). At high concentrations when F2 is in excess, the primary emitter appears to be HCF(A3Z-+X3Z), in agreement with the past work by Glinski et aL5 This emitter has also been identified in F2/CH4 flames,30as well as in reactions of F atoms with organoiodine compo~nds.~lPrior work in the reactions of organic iodides with F atoms observed HCF(A3Z-X3Z) emission and attributed it to the termolecular combination" CH + F + M HCF* + M (18)
-.
This radical could then eliminate an H atom to form a C=S bond, also analogous to olefinic system^:^'.^^ CH3SFCH3
4
H
+
"\
/
F
C=S H'
+
(12) 'CH,
However, atomic hydrogen has not been directly identified in the present study. The sulfurane shown above could generate the ion observed at m / e = 80. In systems where dimethyl sulfide is in excess, atomic fluorine and hydrogen can react with the excess DMS so that the sulfurane product is observable. If F2 is in excess, the sulfurane product most likely reacts further with F2 or with atomic species produced in the earlier steps, especially F atoms, to be consumed. It should also be noted that the complex could undergo rearrangement and eliminate H F in one step as initially proposed by Nelson et aL4
This reaction is sufficiently energetic to populate high vibrational levels of H F and result in the H F vibrational overtone emission observed in our detector. Since the lifetime of the hypothesized complex is most likely short (I1ps), the large degree of rearrangement necessary would tend to argue against this type of mechanism. However, this mechanism cannot be ruled out and may constitute a small fraction of the reaction pathway. Perhaps a more likely source of the observed H F ( u l 8 ) vibrational overtone e m i s s i ~ nis~ .the ~ reaction of hydrogen atoms with molecular fluorine. The reaction H
+ F2 -.HFt + F
(14)
is known to populate H F vibrational levels up to u = 9 and to produce intense chemiluminescenceemission.29 When DMS is in excess. the reaction (24) Kahler, C. C.;Lee, Y . T.J . Chem. Phys. 1980, 73, 5122. (25) Birks, J. W.; Gabelnick, S.D.; Johnston, H. S.J . Mol. Specrrosc.
1975, 57, 23.
Under the conditions of the low pressure and excess F2 in the present study, the most likely reaction to produce HCF* is CH + F2 HCF* + F (19)
(26) Miller, W. T.;Koch, S.D. J. Am. Chem. Soc. 1957. 79, 3084. (27) Parson, J. M.; Lee, Y . T. J . Chem. Phys. 1972, 56,4658. (28) Shobatake. K.: Parson. J. M.: Lee, Y. T.: Rice. S.A. J . Chem. fhvs. 19j3.59, 1416. (29) Mann, D. E.; Thrush, B. A.; Lide, Jr., D. R.; Ball, J. J.; Aquista, N. J . Chem. Phys. 1961, 34, 420.
which is exothermic by 78 kcal mol-I and can easily supply the energy necessary to produce the chemiluminescence observed. In conclusion, we have found that the reactions of F2 with dimethyl sulfide occur at an extraordinarily fast rate for the reaction of two stable molecules. We hypothesize that this is due to the ability of fluorine and DMS to form a charge-transfer complex, which can then eliminate products. The charge-transfer complex theory allows us to predict what classes of compounds are likely to be detectable using the fluorine-induced chemiluminescence detector. Further studies of the kinetics of reactions of F2 with compounds having a range of ionization potentials (Le., varying critical distances, rc) is needed to test the proposed mechanism. It would also be of interest to study the kinetics of reactions of compounds in which the sulfur atom exists in different oxidation states or has different bonding to carbon (Le., CS2 or COS). The selectivity of the F2-induced chemiluminescence detector appears to be explained in terms of the rate of the initiating reaction. This reaction is fast for low ionization potential compounds such as DMS. Once the reaction is initiated, production of atomic fluorine (and possibly H atoms) can create an exothermic chain which results in excited-state species. These reactions are nonspecific to any class of compounds, and the various emitters observed depend upon the reaction conditions and the particular analyte. In conditions where excess F2 is used, many reactions occur, and it is difficult to assign the formation of a specific emitter to a specific reaction. However, it may be possible to determine the reaction responsible for the production of excited-state thioformaldehyde. This emitter is only observed under conditions of excess sulfur compound, which helps to suppress the radical chain reaction. Studies with organoselenium compounds in their reaction with F2 has led to the identification of excited-state selenoformaldehyde in systems analogous to ones in which thioformaldehyde emission is observed.32 Further studies of organotellurium, organoarsenic, and organophosphorous compounds may reveal the presence of other, as yet unknown, emitters, whose spectroscopy is of interest. (30) Patel, R. I.; Stewart, G. W.; Castleton, K.; Gole, J. L.; Lombardi, J. R. Chem. Phys. 1980,57,461. (31) Braynis, H. S.;Whitehead, J. C. J . Chrm. Soc., Faraday Trans. 2 1983,79, 1 1 13. (32) Glinski, R. J.: Mishalanie. E. A.: Birks. J. W. J . Am. Chem. Soc. 1986, 108, 531.
