J. Phys. Chem. 1986, 90, 4393-4397 vibrational redistribution in this system. Further extensive pressure-dependent studies of alkyl nitrites which address this question are in progress. In a recent study of Bauer and Lazaar, “apparent activation energies” of 15 and 24 kcal mol-’ were obtained for syn-anti conformational exchange in gaseous methyl and ethyl nitrite.17 To account for these strange results, a regional phase space model was proposed. Our results demonstrate that an alternate explanation of these apparent activation energies is in order. For both molecules the Arrhenius activation energy determinations utilized rate constant data that were obtained at systematically varying pressures as well as temperatures. In the case of methyl nitrite, pressures beteen 7 and 120 Torr were analyzed and for ethyl nitrite, pressures between 4.5 and 60 Torr were used. The infinite pressure rate constants and p l 1 2values obtained in the present study and a previously reported study of methyl nitritelg demonstrate that much of the data were obtained in the falloff regime where the order of the reaction is changing rather than in the low-pressure bimolecular regime rendering the analysis i n ~ a l i d . ~ ’
4393
Using systematically decreasing pressures and temperatures in the data analysis resulted in these unrealistic apparent activation energies. The greater discrepancy obtained for the case of ethyl nitrite is consistent with its lower p1/2value.
Acknowledgment. We are glad to acknowledge the National Science Foundation (Grant CHE-83-51698-PYI and CHE-8503074) and the National Institutes of Health (GM-29985-05) for support of this research. Registry No. H 3 C C H 2 0 N 0 ,109-95-5; H3C(CH2)20N0,543-67-9; (H3C)jCCH20NO, 77212-96-5.
Supplementary Material Available: Tables of phase-dependent rate constants for syn anti conformational exchange in ethyl, n-propyl, and neopentyl nitrite (4 pages). Ordering information is given on any current masthead page. (31) Eyring, H.; Lin, S. H.; Lin, S. M. Basic Chemical Kinetics; Wiley: New York, 1980; pp 196-197.
Use of Hydrogen Chloride as a Probe for the Determination of Primary Processes in Acetaldehyde Photolysis Abraham Horowitz Department of Radiation Chemistry, Soreq Research Center, Yaune, Israel 70600 (Received: January 6, 1986)
-
The photolysis of acetaldehyde at 3000 A and 25 f 2 OC was studied in CH3CHO-HCl and CH3CHO-HCI-C02 mixtures. The formation of methane via the reaction CH3 + HCI CH4 + C1 and its pressure dependence were utilized for the evaluation of the low-pressure quantum yields of the primary photodissociation processes: CH3CHO* CH3 + HCO (I); CH3CHO* CH4 + CO (11); CH3CHO* CH3C0 + H (111). The values obtained for q+, dII,and 4111 of 0.92, 0.015, and 0.065, respectively, are in very good agreement with estimates based on CO and H2 determination in mixtures of oxygen with acetaldehyde (Horowitz, A.; Kershner, C. J.; Calvert, J. G. J . Phys. Chem. 1982,86, 3094). These findings indicate that HC1 can be used as a convenient and sensitive probe in studies aimed at the determination of the radical yields in the photolysis of acetaldehyde.
-
-
Introduction Acetaldehyde photochemistry has been extensively studied during the past 50 years.’ As result of these studies it is now well established that, when excited within its first absorption band (3340-2200 A), acetaldehyde decomposes by three primary photodissociation processes: CH,CHO* CH3 + H C O (1)
--
CH3CHO*
CH3CHO* --*
+ CO CH3CO + H CHI
(11) (111)
In this schematic representation CH3CHO* stands for excited acetaldehyde molecules, all not necessarily in the same state. Acetaldehyde is present in relatively large amounts in the polluted atmosphere, where it serves as a source of radicals and as a precursor of peroxyacetyl nitrate (PAN).2,3 The important role of acetaldehyde in atmospheric chemistry and the resultant need for quantitative data on its photodissociation led to several recent s t ~ d i e s ~in- ~which an attempt was made to assess the ~
(1) Lewis, R. S. Adu. Photochem. 1980, 12, 1. (2) Demeqian, J. A.; Kerr, J. A.; Calvert, J. G. Adu. Emiron. Sci. Technol.
1974, 4, 1. ( 3 ) Atkinson, R.; Loyd, C. J . Phys. Chem. Ref. Data 1984, 13, 315. (4) Gill, R. J.; Johnson, W. D.; Atkinson, G. H. Chem. Phys. 1981, 58, 29.
0022-3654/86/2090-4393$01.50/0
-
absolute quantum yields and the relative importance of the three channels (1-111) of C H 3 C H 0 decomposition as well as their variation with pressure and excitation wavelength. Results of studies carried out by Weaver et ai.,’ Horowitz et al.,637 Meyrahn et al.,Band Simonaitis and Heickleng indicate that process I is the predominant decomposition channel. In all four studies oxygen was used as a radical scavenger and except in work by Weaver et al. the quantum yields estimates were based on the assumption of complete conversion of H C O into C06-8 and of C H 3 into CH302.’ The present study of the 3000-8, photolysis of the CH3CHOHCl system was initiated in view of the importance of accurate determination of the yields of radicals formed in the primary photodissociative processes in acetaldehyde and in order to verify earlier findings in systems in which oxygen was used as a radical scavenger. It appeared to us that quantum yield estimates based on the highly sensitive FID gas chromatographic determination (5) Weaver, J.; Meagher, J.; Heicklen, .I.J . Photochem. 1976/1977, 6, 111. (6) Horowitz, A.; Kershner, C. J.; Calvert, J. G . J . Phys. Chem. 1982,86, 3094. (7) Horowitz, A.; Calvert, J. G. J . Phys. Chem. 1982, 86, 3105. (8) Meyrahn, H.; Moortgat, G . K.; Warneck, P., presented at the 15th International Conference on Photochemistry,Stanford, CA, 1982; Paper D-4. (9) Simonaitis, R.; Heicklen, J. J . Photochem. 1983, 23, 299.
0 1986 American Chemical Society
4394
The Journal of Physical Chemistry, Vol. 90, No. 18, 1986
Horowitz
TABLE I: Quantum Yield Data from 3000-A Photolysis of Acetaldehyde-Hydrogen Chloride Mixtures"
PCH,CHO, Torr 5 6 10 20 20 20 20 20' 26 20 20 31.2 35.2 50 60 66.7 75.6
PHCL, Torr 2.0 2.0 2.0 0.5 1.O 2.0 2.0 3.5 5.0 2.0 2.0 2.0 2.0 2.0 2.0
10-'5Q 1.195 1.248 1.254 1.244 1.205 1.196 1.257 1.159 1.198 1.264 1.221 1.267 1.175 1.238 1.116 1.271 1.257
irr time, min 130 62 52 64 30 40 40 35 30 30 30 22 33 30 16 38 16
P C ~ 4lo-' , Torr 87.7 52.5 67.6 56.4 69.2 84.8 88.9 75.0 1.12 65.7 70.5 65.0 118.2 103.8 54.3 138.8 63.0
PNC: 136.9 86.4 123.8 154.0 109.7 150.7 146.5 123.0 112.1 103.2 102.0 195.8 170.5 90.2 223.2 81.1
aCH4
Torr
0.902 0.915 0.881 0.341 0.922 0.854 0.852 0.890 0.015 0.834 0.927 0.823 0.815 0.784 0.762 0.691 0.722
%OC
0.603 0.538 0.683 0.547 0.497 0.621 0.509 0.527 0.88 1 0.546 0.387 0.430 0.499 0.473 0.476 0.393 0.565
"Half-bandwidth 66 A, temperature 25 f 2 O C . b u n i t s of quanta cell-' s-'. CPNC= total pressure of products that do not condense at liquid N, temperature. dCalculated from PNC.Le., PCO = PNC - P C H-~P H ~P ;H computed ~ from ref 1. '0.01 Torr of biacetyl added. f 3 Torr of O2 added; +co average = 0.88 f 0.09(2a), directly determined in five runs. TABLE 11: Effect of Carbon Dioxide on the Formation of CH4 and CO in the 3000-A Photolysis of Acetaldehyde (20 Torr) and Hydrogen Chloride (2 Torr)" Prn.. Torr
150 300 380 450 520 600 600
10-'5Inb 1.281 1.173 1.242 1.213 1.192 1.300 1.265
irr time, min 40 40 58 60 50 61 50
PpH..
io-' Torr 73.9 53.3 84.4 76.9 59.8 73.6 58.7
PNr,
Torr 120.1 90.0 131.2 131.9 94.9 121.1 94.3
@PH.
amd
0.695 0.547 0.564 0.509 0.483 0.459 0.447
0.434 0.377 0.3 13 0.364 0.284 0.296 0.27 1
"Half-bandwidth 66 A, temperature 25 f 2 OC. b u n i t s of quanta cell-' s-'. c P Nis~the total pressure of products that do not condense at liquid
N2 temperature. dOCocomputed from PNc.
+
of methane formed in the CH3 HCl reaction will be more accurate than those derived from the CH3CHO-02 system in which potential complications could arise from triplet C H 3 C H 0 quenching, rapid consumption of oxygen, and conversion of a small fraction of formyl radical into products other than CO(HC03). Experimental Procedures
Materials. Acetaldehyde (BDH, Analar), acetone (Biolab, Analar), biacetyl (Kodak), hydrogen chloride (Matheson), and carbon dioxide (Matheson, Coleman) were degassed and purified by trap-to-trap distillation. Oxygen (Hoechst, prepurified) was used as received. Apparatus and Procedures. The 3000-A photolysis of acetaldehyde was carried out in a cylindrical 21.6-cm-long, 4.6-cm-i.d. quartz cell joined to a grease-free vacuum system. Pressure transducers (Celesco, Model P7D) were used for the measurement of reactant pressures and of the liquid N2 noncondensable product gases. The light source was an Owam HBO 200-W high-pressure mercury lamp. The 3000-A region was isolated with a high-intensity quarter-meter grating monochromator (Schoeffel, G M 252) equipped with 1180 g/mm grating blazed a t 2400 A and with a reciprocal linear dispersion of 33 .&/mm. Slits were set at 2 mm, giving an effective half-bandwidth of 66 A. Half of the cell volume (370 cm3) was illuminated, and the light transmitted through it was monitored with an RCA 935 photodiode mounted at the rear of the cell and connected to an Oriel detection system, Model 7072, coupled to a data logger. Acetone (20 Torr, 125 "C) was em= 1.l0 Acetaldehyde abployed for actinometry, taking apc0 sorption followed the Beer-Lambert law with € = log (Zo/l)/cl = 10.01 L mol-' cm-l. After photolysis, the reactant and product gases were Toepler pumped through two traps held in liquid N,. Gases that did not condense at this temperature (CHI, CO, H,, and 0,) were (10) Calvert, J. G.; Pitts, J. N., Jr. Phorochemistry; Wiley: New York, 1966.
transferred to a precalibrated volume in which their total pressure was measured. In most runs only methane was determined by using a 5A molecular sieve (1/8 in. X 8 ft) column held at 80 "C and a GC equipped with a FI detector (HP 5750). In some cases the same instrument with a TC detector was used for C O determination. In these runs the molecular sieve column was held at 40 O C . Results
The photolysis of CH,CHO-HCl mixtures was studied over a wide pressure range (5-75.6 Torr). The results of these experiments are summarized in Table I. In all runs HC1 consumption was around 5%, while C H 3 C H 0 conversion did not exceed 2% and in most cases was lower than 1%. In order to compare present results with earlier findings a few runs were carried out in CH3CHO-02 mixtures. The results, the average of five determinations, are also presented in Table I. It should be noted that h0in all runs except those with added O2 was indirectly estimated from the pressure of products that did not condense at liquid N, temperature. Consequently, these Q0 values are not very accurate (about f l 5 % accuracy). The possibility of the Occurrence of secondary reactions as a result of biacetyl accumulation was examined in a run in which it was added to a CH3CHO-HCl mixture. As can be seen from Table I, biacetyl addition at a concentration at which it might have been formed in the CH3CHO-HCl system did not reduce @CH4.
Additional experiments, summarized in Table 11, were carried out in the CH3CH0-HCI-CO2 system. In these experiments the effect of COz on methane formation was determined. Discussion
Kinetics and Mechanism of CH,CHO-HCl Photolysis. In the photolysis of acetaldehyde a t 3000 A and in the absence of hydrogen chloride, the reactions shown i n Scheme I take place.
The Journal of Physical Chemistry, Vol. 90, No. 18, 1986 4395
Primary Processes in Acetaldehyde Photolysis
SCHEME I CH3CH0
+ hv
CH3CHO*
+
CH4
H
+ HCO
CH3
+ HCO
CH3
- + -+ +
CH3CHO* CH3CHO*
CH3CHO*
+ HCO CH3 + CH3
CH3
+
(111)
+ CO
(1)
CH3CHO
(2)
C2H6
(3)
+
+ CH3CHO CH4 + CH3CO HCO + HCO CO + CHiO CH3COCH3 CH3CO + CH3 CH3CO + CH3CO (CH,C0)2 CH3COCHO CH3CO + HCO H + CH3CHO H2 + CH3CO
CH3
(11)
CH3CO
CHI
---*
CO
(1)
+
-
+
-
+ HCl CH3 + HCl C1+ CH3CHO
-.+
-
+ CI CH4 + C1 HCl + CH3CO H2
+
(5)
+ HC1-
CH3CHO
+ HC1
(7) (8)
(9)
(10) (11) (12)
(13)
analogous reaction of methyl radicals (reaction 11) is exothermic by 0.8 kcal mol-'. Hence, reaction 3, unlike reaction 4, can be expected to be very slow and incapable of competing with the fast biradical reactions of the acetyl radicals. On the other hand, the reactivity of acetaldehyde in its reaction with chlorine atoms can be expected to be similar to that of tertiary hydrogens; Le., k12 l o k oL mol-' s-I,l2 and therefore C1 atoms should be solely removed by this route. Bearing these considerations in mind, we can turn now to the evaluation of HCl efficiency as a scavenger of methyl radicals formed in the photolysis of acetaldehyde. The purpose of HCI addition to CH3CH0 was to convert completely into methane the methyl radicals formed in the primary photodissociation process (eq I), so that @cH4= 41 + 42. The "efficiency" of conversion of methyl radicals into methane is given by eq 14 and can be computed with the aid of few simplifying
-
@CH4
- 42 -
4'
R11
+
R" + Rl + R4 R1 R2 R3 R4
+ + + + Rs
(14)
assumptions. For our conditions it can be assumed that the two main radical species, present in equal concentrations, are C H 3 C 0 and HCO. We will also assume that all radical termination reactions in the system have the same rate constant of the selfreaction of methyl radicals,I3 Le., ( k , + k,) = k3 = k5 = k6 = k7 = k8 = 2.2 X l o i o L mol-' s-l. Taking k l / k 2= 4,14 k4 = 2.7 X 103,15and k l l = 4.8 X lo6 l 6 L mol-' s-' and remembering that (1 1) Kerr, J. A,; Trotman-Dickenson, A. F. CRC Handbook ofchemistry and Physics, 62nd. ed.; CRC: Boca Raton, FL, 1982. (12) Kerr, J. A. In Free Radicals; Kochi, J. K., Ed., Wiley: New York,
1975; Chapter 1. (13) Shepp, A. J . Chem. Phys. 1956, 24,939. (14) Toby, S.; Kutschke, K. 0. Can. J. Chem. 1959, 37, 672.
1.4-
(6)
From the bond dissociation energies," D(CH3CO-H) = 86, D(CH,-H) = 104, and D(H-Cl) = 103.2 kcal mol-', it can be seen that hydrogen abstraction from HCl by acetyl radicals (reaction 3) is endothermic by 17.2 kcal mol-' while the CH3CO
Figure 1. Variation of @.CH4 with HCI pressure in the photolysis of CH3CHO-HC1 mixtures ( C H 3 C H 0 = 20 Torr).
(4)
In HCl containing acetaldehyde mixtures it can be safely assumed that, a t the acetaldehyde pressures and light intensities employed, the main additional reactions that need to be considered are
H
P H C ~(Torr 1
1
-
Figure 2. Stern-Volmer plot of = 2 Torr).
aCH4 in CH3CHO-HCI
mixtures (HC1
in our experiments the light beam occupied only about half of the cell volume (185 cm3), one can show that under the most unfavorable conditions ( C H 3 C H 0 = 75.6 and HC1 = 2 Torr) about 99% of the methyl radical will end up as methane molecules. Of these molecules only about 2% are formed in reactions other than hydrogen abstraction from HCI. Further evidence for the adequacy of HC1 to serve as a probe for the estimation of methyl radical yields in the photolysis of acetaldehyde is provided by the experimental results at 20 Torr of C H 3 C H 0 in which the pressure of HC1 was varied (see Figure 1). It can be seen that upon addition of 0.5 Torr of HC1, @cH4 increases drastically from 0.341 to 0.922 and then levels off. As a final test of the reliability of the present method of estimation of primary quantum yields, experiments were conducted in which, as in earlier studies,6-8 O2 replaced HCl as radical scavenger. In the presence of oxygen formyl radicals are converted into C O via reaction 15 and therefore ac0 = 41 + 42. ConseHCO + 0 2 H02 + C O (15) -+
quently, a t the same pressure of acetaldehyde, we would expect @cH4(HCI) = @co(02). Inspection of the results in Table I shows that within the error limits this expectation is indeed borne out by the experimental findings. Primary Quantum Yields and Pressure Effects. The observed in experiments at constant HCl pressure and in decrease in aCH4 which the pressure of acetaldehyde was increased (see Table I) as well as in experiments with added COz (see Table 11) indicates that the population of the excited state responsible for process I is pressure-dependent. Similar pressure effects were observed in the CH3CHO-02 system, both in steady-state studies in which formation of C O was determined at the end of irradiation and in the flash photolytic experiments in which the formation of the methyl peroxy radical via reaction 16 was m ~ n i t o r e d . ~ CH3 + 0 2 CH3O2 (16)
-
On the basis of their own results as well as evidence from various
other studies of acetaldehyde photolysis, Horowitz and Calvert suggested that at 3000 A acetaldehyde photodecomposition into (15) Kerr, J. A.; Calvert, J. G. J . Phys. Chem. 1965, 69, 1022.
(16) Kerr, J. A.; Parsonage, M. J. Evaluated Kinetic Data on Gas Phase Hydrogen Transfer Reactions of Methyl Radicals; Butterworths: London, 1976.
Horowitz
4396 The Journal of Physical Chemistry, Vol. 90, No. 18, 1986
0
a
200 400 Pco2 (Torr)
600
in CH,CHO-HC1-C02 mixtures Figure 3. Stern-Volmer plot of (HCI = 2 Torr, CH,CHO = 20 Torr).
CH3 and HCO proceeds from a vibrationally rich triplet CH3CH0(TlW).' This species is formed by fast intersystem crossing from an initially formed vibrationally rich singlet CH3CHO(SI'). According to this hypothesis and neglecting the insignificant radiative decay of CH,CHO(TIW),the reduction in acH4 can be accounted for by the following reactions: CH3CHO(TlW)
+
CH3
+ HCO
(kd)
CH3CHO(TlW)+ CH,CHO(So) 2CH3CHO(So) -+
(~Q(CH,CHO))
CH3CHO(TlW)+ HCI
.-*
CH3CHO(So)
+ HCl
(kQ(Hc1))
CH3CHO(TIW)+ COz
-+
CH,CHO(S,)
+ C02
(kQ(co,))
In terms of these reactions the following relation should apply in CH3CHO-HCI mixtures:
@CHI
=I1+
kQ(HCl)[HC1l
$1
- 411
kd
1'
kQ(CH,CHO)[CH,CHO]
(17)
kd
Here 4, stands for the zero-pressure quantum yield of methane formed in process I, and r#qi is the quantum yield of methane formed in process 11, which is not quenched by oxygen. For our conditions this residual value of aCH, is equal to 0.015. The Stern-Volmer fit of the CH3CHO-HCI data (HCI = 2 Torr), shown in Figure 2, gives l/(@c-,-0.015) = (1.097 f 0.037) (4.55 f 0.85) X C H 3 C H 0 (Torr)
+
In a similar fashion, the effect of C 0 2 addition on @cH4 should be described by expression 18. The Stern-Volmer plot of the
k ~ ( H c i[HCl] )
+ ~ Q ( C H , C H O[CH3CHO] )
kQ(C02)[C021 } +
kd
kd
(18)
C 0 2 results is shown in Figure 3. The respective least-squares fit of the data gives 1/ (aCH4
- dT1)
=
(1.25 f 0.15)
+ (1.72f
0.33)
X
C 0 2 Torr
For both fits the stated error limits are two standard deviations. The results of the Stern-Volmer fits can be utilized to evaluate the relative efficiency k Q ( c H , c H o ) / k Q ( c q , ) of CH3CHO(TlW) quenching by acetaldehyde and carbon dioxide, which is given by kQ(cH,cHo) - slope of Figure 2 (19) kp,C02) slope of Figure 3 From the present work we arrive at a kQ(CH3CHO)/kQ(C02) value of 2.67 f 0.71, which is in fair agreement with the result for this ratio of 1.79 obtained by Horowitz and Calvert in the CH3CH0-O2 ~ y s t e r n .The ~ origin of the difference between the two sets
I
,
,
,
,
, ,
, ,
,
, ,
,
,
._ , ,.-: ,~
4397
J. Phys. Chem. 1986, 90, 4397-4402 energy origin may be expected to have longer lifetimes than at 3000 8, and be more prone to quenching by 02.If indeed this is the case and TI quenching does not result in C O formation, as has been suggested by Weaver et al.? then it should be reflected in a differenceS between $, estimates based on CH3 scavenging
by HCl and those derived from HCO scavenging by 02.Studies aimed a t the testing of this hypothesis are now in progress. Registry NO. CH,CHO, 75-07-0; HCI, 7647-01-0; C02, 124-38-9; 02, 7782-44-7; CH.,, 74-82-8; CO, 630-08-0; H2, 1333-74-0.
Reactions of Chlorine wlth Liquid Metals. 4. Tin D. R. Olander,*t M. Balooch: and W. J. Siekhaust Materials and Molecular Research Division of the Lawrence Berkeley Laboratory and Department of Nuclear Engineering, University of California, Berkeley, California 94720, and Chemistry Division of the Lawrence Livermore National Laboratory, Livermore, California 94550 (Received: January 13, 1986)
The reaction of molecular chlorine with solid and liquid tin was studied by modulated-molecular-beam mass spectrometry and Auger electron spectrometry. The temperature range was 350-700 K and equivalent chlorine pressures between 2 X and 5 X lo4 Torr were used. At constant pressure, the production rate of the sole reaction product (SnC12)increased with temperature with no sign of a discontinuity at the melting point. The reaction was nonlinear with respect to chlorine equivalent pressure. The model developed for the lead-chlorine system was applicable to the tin-chlorine system. However, solution diffusion of chlorine in the liquid metal did not appear to be significant for the tin reaction. A comparison of all four metal-chlorine reactions studied in this series is presented.
Introduction The present work is a continuation of the series on the reactions between low-melting metals and molecular chlorine over the temperature range bracketing the solid-liquid phase transition. Like the previous investigations of In,' Pb,2 and Bi,3 the present study of tin utilizes the modulated-molecular-beam technique with phase-sensitive detection of reaction products by an in situ mass spectrometer. In a separate chamber, the surface of the metal is monitored by Auger electron spectroscopy (AES) under conditions simulating those in the molecular-beam tests. Details of the experimental method and the data analysis procedures are given in ref 1. Results The tin-containing ions observed in the mass spectrometer are SnC12+,SnCl+ and Sn+. All have the same temperature dependence and the same phase angles, indicating that they all originate from the neutral species SnCl2. This molecule is thus the sole product of the reaction. All measurements utilize the SnCl' peak, which has the highest intensity at a nominal electron energy of 75 eV, and are then corrected for the observed fragmentation pattern. The reaction probabilities reported below are based on the SnC12 product signals divided by the sum of the C12+signal and half of the C1+ signal for a surface at room temperature. Signal ratios are converted to flux ratios (i.e., apparent reaction probabilities) by using ionization cross section ratios obtained from ref 4 and other instrumental efficiency factors and transit time effects discussed in ref 5. Figure 1 shows the apparent reaction probability (c) and phase lag ($) of the SnClz product as functions of ternperdure at the maximum incident chlorine beam intensity and for a modulation frequency of 20 Hz. Unlike the previous reactions studied, no discontinuity is observed in either e or $ as the melting point is *Address correspondence to this author at the Department of Nuclear En ineering, University of California, Berkeley, CA 94720. ?Materials and Molecular Research Division of the Lawrence Berkeley Laboratory and the Department of Nuclear Engineering, University of California, Berkeley, CA 94720. *Chemistry Division of the Lawrence Livermore National Laboratory, Livermore, CA 94550.
0022-3654/86/2090-4397$01.50/0
TABLE I: Parameters of the Sn-CI?Reaction Model preexponential activation parameter factor enerav. kcal/mol ll
KkLH KqER/Ns
0.06 1.2 x io-' cm2/s 3.5 X lo-" cm2
11 10
crossed. The reactivity increases with temperature up to 600 K and then levels off a t higher temperatures. The phase monotonically decreases with temperature. The effect of beam intensity on c and $ is shown in Figures 2 and 3 for constant modulation frequency and three surface temperatures, two above and one below the melting point. These results clearly indicate that the overall reaction is nonlinear and that there is no distinction that can be attributed to the phase change of the tin. The frequency dependences of the reaction product vector are shown in Figures 4 and 5 for two temperatures above the melting point and two below. The phase lag decreases with increasing frequency for all temperatures but 600 K, where it remains approximately constant. This is a clear indication of a mechanism containing parallel reaction paths for the production of SnCl2, one of which provides a rapid channel for converting incident Clz to product. The slow step is demodulated a t high frequencies, leaving only the rapid step that exhibits the low phase lag. The surface coverage of chlorine as a function of temperature for an incident flux of 5.5 X 10l6molecules/(cmz s), as measured in the AES apparatus, is shown in Figure 6. The surface appears to be saturated up to about 350 K, after which a rapid decrease with increasing temperature ensues. (1) Balooch, M.;Siekhaus, W. J.; Olander, D. R. J . Phys. Chem. 1984, 88, 3522. (2) Balooch, M.; Siekhaus, W. J.; Olander, D. R. J . Phys. Chem. 1984, 88, 3530. (3) Balooch, M.; Siekhaus, W. J.: Olander, D. R. J . Phys. Chem. 1986, 90, 1671. (4) Mann, J. B. In Recent Developments in Mass Spectroscopy; Ogata, K., Hayakawa, T., a s . ; University Park Press: Baltimore, MD, 1970; p 814. (5) Jones, R. H.; Olander, D. R.;Siekhaus, W. J.; Schwarz, J. A. J. Vac. Sei. Technol. 1972, 9, 1429.
0 1986 American Chemical Society
