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Mechanistic details for dimethylcadmium decomposition and CdS production by MOCVD
J. . . . . . . .
ELSEVIER
CRYSTAL G R O W T H
Journal of Crystal Growth 182 (1997) 321 328
Mechanistic details for dimethylcadmium decomposition and CdS production by MOCVD Jinzhi Ni, Joseph J. BelBruno* Department of Chemistry, Burke Chemical Laboratory, Dartmouth College, Hanover, NH 03755, USA
Received 10 October 1996; accepted 23 May 1997
Abstract
The MOCVD production of CdS from dimethylcadmium and hydrogen sulfide is shown to be consistent with a homogeneous reaction mechanism, in contrast to the surface catalyzed decomposition of the dimethylcadmium precursor also described in this report. Molecular orbital calculations, including electron correlation, are used to explore the reaction potential energy surface and the results lend additional support to the proposed homogeneous reaction mechanism.
1. Introduction
Most commercially produced CdS is obtained by high-temperature reaction of the elements. Recently, a room temperature gas-phase production method was reported involving dimethylcadmium (DMCd) and hydrogen sulfide [1]. The published report indicated that this new method could potentially produce uniform-quality powder at r o o m temperature and at a high rate of reaction. Although the procedure was successful, the rapid, spontaneous nature of this reaction led to several product deficiencies, including variable particle size and nonstoichiometric material. Two subsequent studies reported on attempts to overcome the product quality difficulties by employing less reactive
*Corresponding author. Fax: + 1 603 646 3946; e-mail: [email protected].
cadmium adducts at elevated temperatures 1-2, 3]. Thin films rather than powders were produced in these subsequent studies. Our own research indicates that the original reaction system, dimethylcadmium and hydrogen sulfide at 300 K, leads to production of CdS films rather than powders on quartz substrates and has the potential to play an important role in the production of CdS detectors. However, in none of these previous studies is the mechanism of the reaction explored. It is uncertain whether CdS formation from D M C d + H2S is homogeneous, as is the case, at least in part, for GaAs [4] or the result of a surface catalyzed process as observed in the production of ZnSe [5]. This question is not a simple one, since the decomposition of the neat dimethylcadmium precursor, is homogeneous at elevated temperatures [10]. A study of the spectroscopy of D M C d physically and chemically adsorbed onto various substrates provides some evidence that D M C d may
0022-0248/97/$17.00 © 1997 Elsevier Science B.V. All rights reserved PII S0022-0248(97)0036 7-9
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~L Ni, J.J. BelBruno /Journal of Crystal Growth 182 (199D 321 328
decompose catalytically at room temperature [6], but decomposition was not the focus of the experiment and the question of whether or not the room temperature decomposition occurs remains an open one. If D M C d does indeed decompose at room temperature, then one may imagine a competitive kinetic process in which the surface catalyzed decomposition of D M C d is in competition with the reaction of D M C d and H2S, whether the latter process occurs by a homogeneous or heterogeneous mechanism. As part of our attempt to understand the nature of this chemistry, we have turned to product measurements and the use of computational methods to assist in sifting through possible kinetic pathways. Our focus, therefore, is on the mechanism, not kinetic parameters. We have tested several different mechanisms and, in concert with ab initio calculations, determined that the available data are consistent with a homogeneous, gas-phase production mechanism for CdS. Several unique structural determinations, for important intermediates in the formation channel, are also presented.
2. Experimental procedure Reactions were carried out in 316 stainless-steel cross-cells equipped with KBr windows and having a 10 cm optical pathlength. Cells were chemically cleaned and oven dried prior to initial use. An all stainless-steel vacuum line, capable of less than 5 × 1 0 - 3 Torr ultimate pressure, was used to introduce the sample into the cell. Sample pressures were monitored with either conventional or capacitance manometers. Final pressures, subsequent to reaction, were measured by means of quantitative F T I R spectroscopy on a Perkin-Elmer 1605 spectrophotometer at 2 c m - 1 resolution. As the cell is used in D M C d + H2S runs, an orange-red material deposits on the walls and is resistant to simple removal methods. Some of the experiments were conducted in cells having this coating. These are noted as "coated" cell experiments in the reaction mixture results. In order to identify this material, a quartz substrate could be inserted into the reaction chamber and the film produced on the substrate later analyzed by the
usual techniques. XPS studies of the coated surface of the substrate confirmed the presence of Cd and S, although the atomic ratios indicated that in addition to CdS, there was excess Cd present in the solids. Similarly, the dissociation of neat D M C d would also produce a coating on the interior walls of the cell. This coating was shown by XPS, after insertion of a quartz substrate, to contain primarily Cd with small amounts of carbon. Dimethylcadmium was obtained from Johnson Mathey Co. as a liquid. The room temperature vapor pressure is approximately 28 Torr and the vapor over the liquid, after pumping away initial gases, is used for these experiments. Hydrogen sulfide and methane (used as a standard for product formation analysis) were obtained from Matheson as gaseous materials and used without any additional purification.
3. Calculations Among the widely used basis sets, those available for cadmium include STO-3G and 3-21G, as well as the L A N E 1 D Z basis set which combines STO-3G functions for first row atoms with an effective core potential for cadmium. Our experience, as well as that of other researchers, indicates that the 3-21G basis set provides the best computational performance, that is the best agreement with experiment, from among these choices [7, 8]. Electron correlation was treated using a Moeller-Plesset perturbative approach, MP2, so that the reported parameters result from a MP2/3-21G theoretical model. The reported calculations employ the G A U S S I A N 92 [9] suite of programs. No constraints were imposed during geometry optimization to locate the global minima; all of the structural parameters were treated as variables. Therefore, geometries and energies calculated by means of this optimization represent fully relaxed points along the potential surface. Frequency calculations were obtained to identify the optimized geometries as saddle points or as minima on the potential energy surface. Molecular orbitals were examined using the natural bond orbital approach in order to localize the functions to particular atoms.
,,( Ni, J.J. BelBruno /Journal of Crystal Growth 182 (1997) 321 328
4. Results and discussion The mechanism for dissociation of D M C d at high temperatures has been reported to be homogeneous; other results hint that at lower temperatures a heterogeneous process may occur. The mechanism for the reaction of D M C d and H2S is unknown; however, since the D M C d decomposition may occur in parallel with the CVD formation of CdS, the mechanism for neat D M C d may play an important role in directing the reactive process. We have examined both neat D M C d and mixtures of D M C d with H2S in order to isolate the major constituents of the reaction mechanism of the latter.
4.1. Neat dimethylcadmium Temperatures near 400°C have been found to be sufficiently high to promote a homogeneous gasphase process for the decomposition of neat D M C d [10]. However, the decomposition also occurs in neat mixtures at room temperature. A homogeneous mechanism is unlikely at room temperature due to the 2.4 eV bond strength of the C H 3 - C d bond in D M C d [11]. This bond energy would represent a minimum energy barrier for the dissociation process and, at room temperature, only a fraction of an eV is available, on average, to initiate the dissociation. One must then look, at low temperatures, to a surface catalyzed process which would provide a dissociation pathway with a lower activation barrier. The dissociation of neat D M C d was observed under a number of different experimental conditions. The time evolution for the decomposition of D M C d and the formation of CH4 at room temperature is presented in Fig. 1 for these conditions. In one experiment shown in the figure, the chemically cleaned cell is used, while in a second set of runs a cell coated with the carbon Cd film is the reaction vessel (compare runs a and d, for example). As shown in the figure, the rate of the decomposition is much faster in the clean cell than in the cell with the coating. We have taken this as initial evidence of a surface catalyzed process for the decomposition of D M C d at room temperature. The data presented in Fig. 1 for runs a and d are well-fit by exponential functions (see inset, Fig. 1)
323
and yield rate constants of 1.1 x 10 2 min 1 for the clean cell and 2.1 x 10 -3 rain -1 for the coated cell. Previously reported experiments [12, 13] describing the decomposition of AI(CH3)3 on alumina resulted in production of both 1 : 1 and 1 : 2 adsorbate-surface site complexes where the active site was presumed to be A1 OH. This mixture of multiple surface site and single surface site complexes occurred only in the early stages of the reaction and leads to nonstoichiometric yields of methane at that point in time. The net result was that for each molecule of the AI(CH3)3, an average of 1.5 molecule of methane were generated. Study of Fig. 1 curve a indicates that, for the clean cell, the ratio of methane produced to D M C d lost is approximately 1.5:1, for any given time up to approximately 100 min. Since the absolute pressures of methane and D M C d are measured simultaneously, this analysis simply requires a comparison of the relative amount of methane produced to D M C d lost. At sufficiently long times, the ratio of methane produced to D M C d consumed approaches 2 : 1 , indicating complete dealkylation. Moreover, the initial relative rate of methane production, determined as the slope of the methane production curves in Fig. lb, is greater than the rate of D M C d loss from Fig. 1a. This indicates that we have a surface coverage limited reaction in the early stages of the chemistry. Curve b in Fig. 1 presents the results for a cell in which one experiment had previously been conducted. In this experiment, it is expected that a significant number of surface sites would be occupied by molecules of the carbon Cd product, reducing the probability of a 1 : 2 complex at the surface. Indeed, the results indicate that for the first 2 h, the ratio of methane produced to D M C d consumed is 1 : 1. There are insufficient sites available for formation of the 1 : 2 complex at the surface. Again, at later times, complete dissociation of the D M C d is observed. D M C d has been shown to be reactive with OH-containing organic molecules such as alcohols [14]. As was postulated for the AI(CH3) 3 complex, D M C d is also reactive in the presence o f H 2 0 . S i - O H groups at the surface have been previously proposed as the adsorption sites in a study of the spectroscopy of D M C d on surfaces [6]. As a test of this assumed surface site, a cell coated with the c a r b o n - C d product was immersed
~L Ni, J.~L BeIBruno /Journal of Crystal Growth 182 (1997) 321 328
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time, min Fig. 1. Plots of the loss of D M C d (a) and production of methane (b) as a function of time: (a) initial pressure 9 Torr in the chemically clean chamber; (b) initial pressure of 6.8 Torr in a chamber with one previous experiment; (c) initial pressure of 4.3 Torr in a chamber with product coating; (d) initial pressure of 14 Torr in a chamber with product coating, (e) initial pressure of 6.2 Tort with 30 Torr of H 2 in a chamber with product coating and (f/initial pressure of 9 Torr in a product coated cell into which water has been introduced (see text). Exponential fits to selected curves are shown in the inset.
Z Ni, J.Z BelBruno/Journal of Crystal Growth 182 (1997)321-328
in distilled water, dried with nitrogen gas and evacuated for 1 h. The net effect would be to make a small amount of adsorbed water available to a subsequent D M C d dissociation experiment and, if surface sites were important, to increase the reaction rate over that observed for a "dry coated" cell. Water was clearly present in the original, chemically cleaned cell since it had not been vacuum baked at high temperatures. The result, shown as curve f in Fig. 1, yields a first-order rate constant for D M C d decomposition, 1.2 x 10- 2 min-- 1 that is approximately equal to that in the clean cell. This is an additional evidence of the heterogeneous nature of the dissociation process. In the absence of this added water, the rate constant is slower and the production of methane variable depending upon the extent of the coating coverage (see curves c and d). Finally, hydrogen gas was added to the D M C d reaction cell as a scavenger of methyl radicals. As shown in Fig. 1, curve e, the net effect is an increase in both D M C d decomposition and methane formation. The increase in the rate of C H 4 formation is readily attributed to a scavenging process, while the increase in D M C d decomposition rate may be due to reaction of an H' radical with D M C d in the gas phase. The hydrogen radical is known to be much more reactive than the methyl radical and the latter is believed to play an important role in homogeneous decomposition of D M C d . Overall, the experimental evidence is indicative of an efficient surface catalyzed process for D M C d decomposition at low reactant temperatures. A general mechanism consistent with this assertion is the following: CH3CdCH3 + 2 S ~ O H ,~ (CH3)zCd-(OH-Sf):,
(1) (CH3)zCd (OH Sf)2 -+ C d - ( O -
Sf)2 + 2CH4,
(2) where S f - O H is the surface site which involves an - O H at the actual active site on the surface. At later times in the course of film growth, a D M C d molecule will adsorb at a single site, since extensive site coverage will preclude a more elaborate surface complex. In that case, the adsorbed material will be Sf-O CdCH3 and this would lead to the observed carbon-containing films.
325
4.2. Dimethylcadmium/H2S reaction The reaction between D M C d and hydrogen sulfide was also maintained at approximately 300 K. Two reaction mechanisms may be easily conceptualized: a surface catalyzed process in line with the decomposition of D M C d or a homogeneous gasphase reaction leading to formation of CdS and subsequent deposition on the surface. For a surface catalyzed process, the rate of the reaction would be expected to be controlled by the adsorption/decomposition rate of D M C d as described above. That is, we would expect that the rate of the reaction would be determined by the adsorption rate of D M C d with the subsequent reaction with HzS a rapid step. The time and surface availability dependence determined above for D M C d would also be evident in the reactive process. Instead, the reaction is extremely fast and is essentially complete upon addition of hydrogen sulfide to the D M C d previously placed into the cell. It is not possible to observe changes in the pressures of D M C d , HzS or CH4 on the time scale of the F T I R measurements. The reactions are completed before the first scans. The results indicate that the reaction is independent of cell conditions (clean, "coated", etc.); there is no surface site dependence. We conclude that the reaction must be gas phase, however, there are a number of possible mechanisms and there is also the need to break a bond (CH3 Cd) via an input of 2.4 eV, whereas the available average thermal energy is only 0.03 eV. The initiation step cannot, therefore, be the unimolecular decomposition of D M C d . We present two possible reaction schemes. The major difference between the two schemes is the creation of a stable intermediate in Scheme I, while Scheme II proceeds by direct formation of radicals and products. The latter mechanism involves formation of radicals, CH3Cd and SH, while the former does not. Scheme I: CH3CdCH3 + H2S --+ CH4 + [-CH3Cd-SH],
(3)
[CH3Cd SH] --+ Cd S + CH4,
(4)
Scheme II: C H 3 C d C H 3 Jr- H2S ~ CH 4 q- CH3Cd + SH,
(5)
d. Ni, .L.L BelBruno / Journal of Co,stal Growth 182 (1997) 321 328
326 CH3Cd
+ SH --+ Cd-S + C H 4.
(6)
The experimental data for production of methane and loss of D M C d are consistent with either mechanism, since both mechanisms predict two molecules of methane and one of CdS per D M C d molecule. However, an experimental run that included H2 as a scavenger provided additional information. Hydrogen should scavenge radicals. Moreover, the hydrogen atom generated in this scavenging process may react with gaseous D M C d to produce CdCH3 which should be extremely reactive with H2S. The results indicate that the presence of hydrogen has no effect on the yield of the chemistry as a function of H2S pressure. Our conclusion is that the formation of the C H 3 C d - S H intermediate is consistent with the obvious low activation barrier for formation of CdS and with the rapidity of the process. Additional evidence is presented from the results of the ab initio calculations that follow. 4.3. Ab initio results
Optimized geometries for several of the unique Cd-containing molecules and radicals are shown in Fig. 2. Other molecules/radicals were required for
thermodynamic and kinetic analysis, but those molecules (H2S, HS, CH~, etc.) have been welldocumented in the literature and their geometries are not reported here, even though the calculations were completed along with those shown in the figure and their energies are included in the calculations that follow. The ab initio calculations may be used to compute the absolute energies of the reactants and products given as part of Scheme I and Scheme II. These data are contained in Table 1. F r o m these results, it is possible to compute the entha!py change for any reaction shown in the proposed mechanisms and relevant results are contained in Table 2. It is seen in the table that the first step in the mechanism denoted as Scheme II is endothermic by approximately 50 kcal m o l - 1. This is consistent with the reported bond energy of the C H 3 Cd bond in dimethylcadmium (2.4 eV). The overall heat of reaction for reactions 5 and 6 is exothermic. On the other hand, the first step in Scheme I is exothermic by approximately 41 k c a l m o l 1, with the overall energy change equivalent to that of Scheme I| so that some of the exoergicity will be available to the intermediate and promote further reaction to the final CdS product. Additional computational exploration along the potential energy surface connecting the reactants
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(d) Fig. 2. Ab initio structures for unique molecules in Schemes I and ll: (a) DMCd; (b) D M C d H2S transition state to (c) CH3CdSH and (d) CH3Cd. All calculations are fully optimized at the MP2/3-21G level of theory. Energies are shown in Table 1.
J. Ni, J.J. BelBruno / Journal (?f Crvstal Growth 182 (1997) 321-328
Table 1 Absolute energies (hartrees a) for the MP2/3-21G ab initio structures Molecule
Energy
Zero point energy
Cd DMCd CH3Cd CH3 HzS HS CH4 CdS CH3CdSH DMCd H2S
5439.92349 - 5518.87061 - 5479.36439 - 39.41824 - 396.76662 - 396.11889 - 40.07551 - 5835.51796 - 5875.63087 5915.61640
0.07225 0.03517 0.03039 0.01421 0.00561 0.04642 0.00085 0.04376 0.08819
"1 hartree (atomic unit) = 627.5095 kcal tool-1
Table 2 Calculated reaction enthalpies (kcal mol ~) for Scheme 1 and Scheme II Reaction
Enthalpy
(3) (4) Net
- 41.1 25.7 15.4
(5) (6) Net Transition state formation
49.7 - 65.5 15.8 14.1
Scheme I
Scheme II
327
geneous gas-phase process. The n a t u r e of the t r a n s i t i o n state shown in Fig. 2 is of interest. The t r a n s i t i o n state c o n t a i n s a four center cyclic system with C d S H - C b o n d s in the process of f o r m a t i o n a n d breaking. T h e r e a c t i o n c o o r d i n a t e is a vibrational m o d e in which the C d - S a n d C H b o n d s are decreasing in length and the S - H a n d C d - C b o n d s are lengthening. The t r a n s i t i o n state is l o c a t e d tow a r d the p r o d u c t side of the surface. The r e a c t i o n leading to f o r m a t i o n of the m e t h y l c a d m i u m - S H a d d u c t m a y be c h a r a c t e r i z e d [15] as a c o n c e r t e d process initiated by the o x i d a t i v e a d d i t i o n of H2S to the D M C d , followed by reductive e l i m i n a t i o n of CH~. Such a process is known, b u t not c o m m o n in o r g a n o m e t a l l i c chemistry. However, o r g a n o c a d m i u m complexes are k n o w n to u n d e r g o unique mechanistic processes.
5. Conclusions W e have shown by c o m p a r i s o n of e x p e r i m e n t a l results a n d ab initio calculations that the s p o n t a n eous reaction between d i m e t h y l c a d m i u m a n d hyd r o g e n sulfide is consistent with a h o m o g e n e o u s m e c h a n i s m t h r o u g h a CH3 C d SH intermediate. In a d d i t i o n , the d e c o m p o s i t i o n of D M C d was observed to be consistent with a surface c a t a l y z e d process at r o o m t e m p e r a t u r e .
Acknowledgements a n d p r o d u c t s of r e a c t i o n 3 allows for l o c a t i o n a n d o p t i m i z a t i o n of the t r a n s i t i o n state for this reaction. It is s h o w n in Fig. 2 a n d lies 14.1 kcal m o l - t a b o v e the reactants. This b a r r i e r is nearly a factor of four less than t h a t e s t i m a t e d by the b o n d cleavage energy of D M C d . A b a r r i e r of 14 kcal mol 1 is "typical" in chemical kinetics a n d at 15 T o r r each of H2S a n d (CH3)2Cd, the p r e d i c t e d A r r h e n i u s rate constant (ignoring steric factors) is 10t5 c m - 3 s 1, a relatively fast rate constant, consistent with the o b s e r v e d r a p i d reactivity of the reaction mixture. W e conclude, therefore, that Scheme I is consistent with b o t h the e x p e r i m e n t a l o b s e r v a t i o n s a n d the c o m p u t a t i o n a l results a n d that the r e a c t i o n of D M C d a n d H2S is a h o m o -
S u p p o r t was p r o v i d e d by the P i t t s b u r g h Superc o m p u t i n g Center ( G r a n t CHE950013P), which is funded by the N a t i o n a l Science F o u n d a t i o n .
References [1] A.E. Stanley, Statutory Invent, Registration, US H459, 1988. [2] M.P. Halsall, J.J. Davies, J.E. Nicholls, B. Cockayne, P.J. Wright, G.J, Russell, J. Crystal Growth 91 (1988) 135. [31 P.J. Wright, B. Cockayne, A.C. Jones, E.D. Orrell, P. O'Brien, O.F.Z. Kahn, J. Crystal Growth 94 (1989) 97. [4] M. Tirtowidjojo, R. Pollard, J. Crystal Growth 93 (1988) 108.
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,Z Ni, J.J. BelBruno / Journal o f Crystal Growth 182 (1997) 321 328
[5] Sz. Fujita, A. Tanabe, T. Sakamoto, M. Isemura, Sg. Fujita, J. Crystal Growth 93 (1988) 259. [6] E. Sanchez, P.S. Shaw, J.A. O'Neill, R.M. Osgood, Chem. Phys. Lett. 147 (1988) 153. [7] J.J. BelBruno, J. Mol. Struct. (Theochem.) 358 (1995) 125. [8] J.J. BelBruno, Heteroatom Chem. 7 (1996) 39. [-9] M.J. Frisch, G.W. Trucks, M. Head-Gordon, P.M.W. Gill, M.W. Wong, J.B. Foresman, B.G. Johnson, H.B. Schlegel, M.A. Robb, E.S. Replogle, R. Gomperts, J.L. Andres, K. Raghavachari, J.S. Binkley, C. Gonzalez, R.L. Martin, D.J. Fox, D.J. DeFrees, J. Baker, J.J.P. Stewart, J.A. Pople
[10] [11] [12] [13] [14] [15]
(Eds.), GAUSSIAN 92, Revision A, Gaussian, Inc., Pittsburgh, PA, 1992. J.B. Mullin, S.J.C. Irvine, D.J. Ashen, J. Crystal Growth 55 (1981) 92. R.L. Jackson, Chem. Phys. Lett. 117 (1989) 315. A.C. Dillon, A.W. Ott, L.A. Okada, S.M. George, Surf. Sci. 334 (1994) 125. S.M. George, A.W. Ott, J.W. Klaus, J. Phys. Chem. 100 (1996) 13121. P.R. Jones, P.J. Desio, Chem. Rev. 78 (1978) 335. R.F. Heck, Organotransition Metal Chemistry, Academic Press, New York, 1974.
ELSEVIER
CRYSTAL G R O W T H
Journal of Crystal Growth 182 (1997) 321 328
Mechanistic details for dimethylcadmium decomposition and CdS production by MOCVD Jinzhi Ni, Joseph J. BelBruno* Department of Chemistry, Burke Chemical Laboratory, Dartmouth College, Hanover, NH 03755, USA
Received 10 October 1996; accepted 23 May 1997
Abstract
The MOCVD production of CdS from dimethylcadmium and hydrogen sulfide is shown to be consistent with a homogeneous reaction mechanism, in contrast to the surface catalyzed decomposition of the dimethylcadmium precursor also described in this report. Molecular orbital calculations, including electron correlation, are used to explore the reaction potential energy surface and the results lend additional support to the proposed homogeneous reaction mechanism.
1. Introduction
Most commercially produced CdS is obtained by high-temperature reaction of the elements. Recently, a room temperature gas-phase production method was reported involving dimethylcadmium (DMCd) and hydrogen sulfide [1]. The published report indicated that this new method could potentially produce uniform-quality powder at r o o m temperature and at a high rate of reaction. Although the procedure was successful, the rapid, spontaneous nature of this reaction led to several product deficiencies, including variable particle size and nonstoichiometric material. Two subsequent studies reported on attempts to overcome the product quality difficulties by employing less reactive
*Corresponding author. Fax: + 1 603 646 3946; e-mail: [email protected].
cadmium adducts at elevated temperatures 1-2, 3]. Thin films rather than powders were produced in these subsequent studies. Our own research indicates that the original reaction system, dimethylcadmium and hydrogen sulfide at 300 K, leads to production of CdS films rather than powders on quartz substrates and has the potential to play an important role in the production of CdS detectors. However, in none of these previous studies is the mechanism of the reaction explored. It is uncertain whether CdS formation from D M C d + H2S is homogeneous, as is the case, at least in part, for GaAs [4] or the result of a surface catalyzed process as observed in the production of ZnSe [5]. This question is not a simple one, since the decomposition of the neat dimethylcadmium precursor, is homogeneous at elevated temperatures [10]. A study of the spectroscopy of D M C d physically and chemically adsorbed onto various substrates provides some evidence that D M C d may
0022-0248/97/$17.00 © 1997 Elsevier Science B.V. All rights reserved PII S0022-0248(97)0036 7-9
322
~L Ni, J.J. BelBruno /Journal of Crystal Growth 182 (199D 321 328
decompose catalytically at room temperature [6], but decomposition was not the focus of the experiment and the question of whether or not the room temperature decomposition occurs remains an open one. If D M C d does indeed decompose at room temperature, then one may imagine a competitive kinetic process in which the surface catalyzed decomposition of D M C d is in competition with the reaction of D M C d and H2S, whether the latter process occurs by a homogeneous or heterogeneous mechanism. As part of our attempt to understand the nature of this chemistry, we have turned to product measurements and the use of computational methods to assist in sifting through possible kinetic pathways. Our focus, therefore, is on the mechanism, not kinetic parameters. We have tested several different mechanisms and, in concert with ab initio calculations, determined that the available data are consistent with a homogeneous, gas-phase production mechanism for CdS. Several unique structural determinations, for important intermediates in the formation channel, are also presented.
2. Experimental procedure Reactions were carried out in 316 stainless-steel cross-cells equipped with KBr windows and having a 10 cm optical pathlength. Cells were chemically cleaned and oven dried prior to initial use. An all stainless-steel vacuum line, capable of less than 5 × 1 0 - 3 Torr ultimate pressure, was used to introduce the sample into the cell. Sample pressures were monitored with either conventional or capacitance manometers. Final pressures, subsequent to reaction, were measured by means of quantitative F T I R spectroscopy on a Perkin-Elmer 1605 spectrophotometer at 2 c m - 1 resolution. As the cell is used in D M C d + H2S runs, an orange-red material deposits on the walls and is resistant to simple removal methods. Some of the experiments were conducted in cells having this coating. These are noted as "coated" cell experiments in the reaction mixture results. In order to identify this material, a quartz substrate could be inserted into the reaction chamber and the film produced on the substrate later analyzed by the
usual techniques. XPS studies of the coated surface of the substrate confirmed the presence of Cd and S, although the atomic ratios indicated that in addition to CdS, there was excess Cd present in the solids. Similarly, the dissociation of neat D M C d would also produce a coating on the interior walls of the cell. This coating was shown by XPS, after insertion of a quartz substrate, to contain primarily Cd with small amounts of carbon. Dimethylcadmium was obtained from Johnson Mathey Co. as a liquid. The room temperature vapor pressure is approximately 28 Torr and the vapor over the liquid, after pumping away initial gases, is used for these experiments. Hydrogen sulfide and methane (used as a standard for product formation analysis) were obtained from Matheson as gaseous materials and used without any additional purification.
3. Calculations Among the widely used basis sets, those available for cadmium include STO-3G and 3-21G, as well as the L A N E 1 D Z basis set which combines STO-3G functions for first row atoms with an effective core potential for cadmium. Our experience, as well as that of other researchers, indicates that the 3-21G basis set provides the best computational performance, that is the best agreement with experiment, from among these choices [7, 8]. Electron correlation was treated using a Moeller-Plesset perturbative approach, MP2, so that the reported parameters result from a MP2/3-21G theoretical model. The reported calculations employ the G A U S S I A N 92 [9] suite of programs. No constraints were imposed during geometry optimization to locate the global minima; all of the structural parameters were treated as variables. Therefore, geometries and energies calculated by means of this optimization represent fully relaxed points along the potential surface. Frequency calculations were obtained to identify the optimized geometries as saddle points or as minima on the potential energy surface. Molecular orbitals were examined using the natural bond orbital approach in order to localize the functions to particular atoms.
,,( Ni, J.J. BelBruno /Journal of Crystal Growth 182 (1997) 321 328
4. Results and discussion The mechanism for dissociation of D M C d at high temperatures has been reported to be homogeneous; other results hint that at lower temperatures a heterogeneous process may occur. The mechanism for the reaction of D M C d and H2S is unknown; however, since the D M C d decomposition may occur in parallel with the CVD formation of CdS, the mechanism for neat D M C d may play an important role in directing the reactive process. We have examined both neat D M C d and mixtures of D M C d with H2S in order to isolate the major constituents of the reaction mechanism of the latter.
4.1. Neat dimethylcadmium Temperatures near 400°C have been found to be sufficiently high to promote a homogeneous gasphase process for the decomposition of neat D M C d [10]. However, the decomposition also occurs in neat mixtures at room temperature. A homogeneous mechanism is unlikely at room temperature due to the 2.4 eV bond strength of the C H 3 - C d bond in D M C d [11]. This bond energy would represent a minimum energy barrier for the dissociation process and, at room temperature, only a fraction of an eV is available, on average, to initiate the dissociation. One must then look, at low temperatures, to a surface catalyzed process which would provide a dissociation pathway with a lower activation barrier. The dissociation of neat D M C d was observed under a number of different experimental conditions. The time evolution for the decomposition of D M C d and the formation of CH4 at room temperature is presented in Fig. 1 for these conditions. In one experiment shown in the figure, the chemically cleaned cell is used, while in a second set of runs a cell coated with the carbon Cd film is the reaction vessel (compare runs a and d, for example). As shown in the figure, the rate of the decomposition is much faster in the clean cell than in the cell with the coating. We have taken this as initial evidence of a surface catalyzed process for the decomposition of D M C d at room temperature. The data presented in Fig. 1 for runs a and d are well-fit by exponential functions (see inset, Fig. 1)
323
and yield rate constants of 1.1 x 10 2 min 1 for the clean cell and 2.1 x 10 -3 rain -1 for the coated cell. Previously reported experiments [12, 13] describing the decomposition of AI(CH3)3 on alumina resulted in production of both 1 : 1 and 1 : 2 adsorbate-surface site complexes where the active site was presumed to be A1 OH. This mixture of multiple surface site and single surface site complexes occurred only in the early stages of the reaction and leads to nonstoichiometric yields of methane at that point in time. The net result was that for each molecule of the AI(CH3)3, an average of 1.5 molecule of methane were generated. Study of Fig. 1 curve a indicates that, for the clean cell, the ratio of methane produced to D M C d lost is approximately 1.5:1, for any given time up to approximately 100 min. Since the absolute pressures of methane and D M C d are measured simultaneously, this analysis simply requires a comparison of the relative amount of methane produced to D M C d lost. At sufficiently long times, the ratio of methane produced to D M C d consumed approaches 2 : 1 , indicating complete dealkylation. Moreover, the initial relative rate of methane production, determined as the slope of the methane production curves in Fig. lb, is greater than the rate of D M C d loss from Fig. 1a. This indicates that we have a surface coverage limited reaction in the early stages of the chemistry. Curve b in Fig. 1 presents the results for a cell in which one experiment had previously been conducted. In this experiment, it is expected that a significant number of surface sites would be occupied by molecules of the carbon Cd product, reducing the probability of a 1 : 2 complex at the surface. Indeed, the results indicate that for the first 2 h, the ratio of methane produced to D M C d consumed is 1 : 1. There are insufficient sites available for formation of the 1 : 2 complex at the surface. Again, at later times, complete dissociation of the D M C d is observed. D M C d has been shown to be reactive with OH-containing organic molecules such as alcohols [14]. As was postulated for the AI(CH3) 3 complex, D M C d is also reactive in the presence o f H 2 0 . S i - O H groups at the surface have been previously proposed as the adsorption sites in a study of the spectroscopy of D M C d on surfaces [6]. As a test of this assumed surface site, a cell coated with the c a r b o n - C d product was immersed
~L Ni, J.~L BeIBruno /Journal of Crystal Growth 182 (1997) 321 328
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Z Ni, J.Z BelBruno/Journal of Crystal Growth 182 (1997)321-328
in distilled water, dried with nitrogen gas and evacuated for 1 h. The net effect would be to make a small amount of adsorbed water available to a subsequent D M C d dissociation experiment and, if surface sites were important, to increase the reaction rate over that observed for a "dry coated" cell. Water was clearly present in the original, chemically cleaned cell since it had not been vacuum baked at high temperatures. The result, shown as curve f in Fig. 1, yields a first-order rate constant for D M C d decomposition, 1.2 x 10- 2 min-- 1 that is approximately equal to that in the clean cell. This is an additional evidence of the heterogeneous nature of the dissociation process. In the absence of this added water, the rate constant is slower and the production of methane variable depending upon the extent of the coating coverage (see curves c and d). Finally, hydrogen gas was added to the D M C d reaction cell as a scavenger of methyl radicals. As shown in Fig. 1, curve e, the net effect is an increase in both D M C d decomposition and methane formation. The increase in the rate of C H 4 formation is readily attributed to a scavenging process, while the increase in D M C d decomposition rate may be due to reaction of an H' radical with D M C d in the gas phase. The hydrogen radical is known to be much more reactive than the methyl radical and the latter is believed to play an important role in homogeneous decomposition of D M C d . Overall, the experimental evidence is indicative of an efficient surface catalyzed process for D M C d decomposition at low reactant temperatures. A general mechanism consistent with this assertion is the following: CH3CdCH3 + 2 S ~ O H ,~ (CH3)zCd-(OH-Sf):,
(1) (CH3)zCd (OH Sf)2 -+ C d - ( O -
Sf)2 + 2CH4,
(2) where S f - O H is the surface site which involves an - O H at the actual active site on the surface. At later times in the course of film growth, a D M C d molecule will adsorb at a single site, since extensive site coverage will preclude a more elaborate surface complex. In that case, the adsorbed material will be Sf-O CdCH3 and this would lead to the observed carbon-containing films.
325
4.2. Dimethylcadmium/H2S reaction The reaction between D M C d and hydrogen sulfide was also maintained at approximately 300 K. Two reaction mechanisms may be easily conceptualized: a surface catalyzed process in line with the decomposition of D M C d or a homogeneous gasphase reaction leading to formation of CdS and subsequent deposition on the surface. For a surface catalyzed process, the rate of the reaction would be expected to be controlled by the adsorption/decomposition rate of D M C d as described above. That is, we would expect that the rate of the reaction would be determined by the adsorption rate of D M C d with the subsequent reaction with HzS a rapid step. The time and surface availability dependence determined above for D M C d would also be evident in the reactive process. Instead, the reaction is extremely fast and is essentially complete upon addition of hydrogen sulfide to the D M C d previously placed into the cell. It is not possible to observe changes in the pressures of D M C d , HzS or CH4 on the time scale of the F T I R measurements. The reactions are completed before the first scans. The results indicate that the reaction is independent of cell conditions (clean, "coated", etc.); there is no surface site dependence. We conclude that the reaction must be gas phase, however, there are a number of possible mechanisms and there is also the need to break a bond (CH3 Cd) via an input of 2.4 eV, whereas the available average thermal energy is only 0.03 eV. The initiation step cannot, therefore, be the unimolecular decomposition of D M C d . We present two possible reaction schemes. The major difference between the two schemes is the creation of a stable intermediate in Scheme I, while Scheme II proceeds by direct formation of radicals and products. The latter mechanism involves formation of radicals, CH3Cd and SH, while the former does not. Scheme I: CH3CdCH3 + H2S --+ CH4 + [-CH3Cd-SH],
(3)
[CH3Cd SH] --+ Cd S + CH4,
(4)
Scheme II: C H 3 C d C H 3 Jr- H2S ~ CH 4 q- CH3Cd + SH,
(5)
d. Ni, .L.L BelBruno / Journal of Co,stal Growth 182 (1997) 321 328
326 CH3Cd
+ SH --+ Cd-S + C H 4.
(6)
The experimental data for production of methane and loss of D M C d are consistent with either mechanism, since both mechanisms predict two molecules of methane and one of CdS per D M C d molecule. However, an experimental run that included H2 as a scavenger provided additional information. Hydrogen should scavenge radicals. Moreover, the hydrogen atom generated in this scavenging process may react with gaseous D M C d to produce CdCH3 which should be extremely reactive with H2S. The results indicate that the presence of hydrogen has no effect on the yield of the chemistry as a function of H2S pressure. Our conclusion is that the formation of the C H 3 C d - S H intermediate is consistent with the obvious low activation barrier for formation of CdS and with the rapidity of the process. Additional evidence is presented from the results of the ab initio calculations that follow. 4.3. Ab initio results
Optimized geometries for several of the unique Cd-containing molecules and radicals are shown in Fig. 2. Other molecules/radicals were required for
thermodynamic and kinetic analysis, but those molecules (H2S, HS, CH~, etc.) have been welldocumented in the literature and their geometries are not reported here, even though the calculations were completed along with those shown in the figure and their energies are included in the calculations that follow. The ab initio calculations may be used to compute the absolute energies of the reactants and products given as part of Scheme I and Scheme II. These data are contained in Table 1. F r o m these results, it is possible to compute the entha!py change for any reaction shown in the proposed mechanisms and relevant results are contained in Table 2. It is seen in the table that the first step in the mechanism denoted as Scheme II is endothermic by approximately 50 kcal m o l - 1. This is consistent with the reported bond energy of the C H 3 Cd bond in dimethylcadmium (2.4 eV). The overall heat of reaction for reactions 5 and 6 is exothermic. On the other hand, the first step in Scheme I is exothermic by approximately 41 k c a l m o l 1, with the overall energy change equivalent to that of Scheme I| so that some of the exoergicity will be available to the intermediate and promote further reaction to the final CdS product. Additional computational exploration along the potential energy surface connecting the reactants
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J. Ni, J.J. BelBruno / Journal (?f Crvstal Growth 182 (1997) 321-328
Table 1 Absolute energies (hartrees a) for the MP2/3-21G ab initio structures Molecule
Energy
Zero point energy
Cd DMCd CH3Cd CH3 HzS HS CH4 CdS CH3CdSH DMCd H2S
5439.92349 - 5518.87061 - 5479.36439 - 39.41824 - 396.76662 - 396.11889 - 40.07551 - 5835.51796 - 5875.63087 5915.61640
0.07225 0.03517 0.03039 0.01421 0.00561 0.04642 0.00085 0.04376 0.08819
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Enthalpy
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Scheme I
Scheme II
327
geneous gas-phase process. The n a t u r e of the t r a n s i t i o n state shown in Fig. 2 is of interest. The t r a n s i t i o n state c o n t a i n s a four center cyclic system with C d S H - C b o n d s in the process of f o r m a t i o n a n d breaking. T h e r e a c t i o n c o o r d i n a t e is a vibrational m o d e in which the C d - S a n d C H b o n d s are decreasing in length and the S - H a n d C d - C b o n d s are lengthening. The t r a n s i t i o n state is l o c a t e d tow a r d the p r o d u c t side of the surface. The r e a c t i o n leading to f o r m a t i o n of the m e t h y l c a d m i u m - S H a d d u c t m a y be c h a r a c t e r i z e d [15] as a c o n c e r t e d process initiated by the o x i d a t i v e a d d i t i o n of H2S to the D M C d , followed by reductive e l i m i n a t i o n of CH~. Such a process is known, b u t not c o m m o n in o r g a n o m e t a l l i c chemistry. However, o r g a n o c a d m i u m complexes are k n o w n to u n d e r g o unique mechanistic processes.
5. Conclusions W e have shown by c o m p a r i s o n of e x p e r i m e n t a l results a n d ab initio calculations that the s p o n t a n eous reaction between d i m e t h y l c a d m i u m a n d hyd r o g e n sulfide is consistent with a h o m o g e n e o u s m e c h a n i s m t h r o u g h a CH3 C d SH intermediate. In a d d i t i o n , the d e c o m p o s i t i o n of D M C d was observed to be consistent with a surface c a t a l y z e d process at r o o m t e m p e r a t u r e .
Acknowledgements a n d p r o d u c t s of r e a c t i o n 3 allows for l o c a t i o n a n d o p t i m i z a t i o n of the t r a n s i t i o n state for this reaction. It is s h o w n in Fig. 2 a n d lies 14.1 kcal m o l - t a b o v e the reactants. This b a r r i e r is nearly a factor of four less than t h a t e s t i m a t e d by the b o n d cleavage energy of D M C d . A b a r r i e r of 14 kcal mol 1 is "typical" in chemical kinetics a n d at 15 T o r r each of H2S a n d (CH3)2Cd, the p r e d i c t e d A r r h e n i u s rate constant (ignoring steric factors) is 10t5 c m - 3 s 1, a relatively fast rate constant, consistent with the o b s e r v e d r a p i d reactivity of the reaction mixture. W e conclude, therefore, that Scheme I is consistent with b o t h the e x p e r i m e n t a l o b s e r v a t i o n s a n d the c o m p u t a t i o n a l results a n d that the r e a c t i o n of D M C d a n d H2S is a h o m o -
S u p p o r t was p r o v i d e d by the P i t t s b u r g h Superc o m p u t i n g Center ( G r a n t CHE950013P), which is funded by the N a t i o n a l Science F o u n d a t i o n .
References [1] A.E. Stanley, Statutory Invent, Registration, US H459, 1988. [2] M.P. Halsall, J.J. Davies, J.E. Nicholls, B. Cockayne, P.J. Wright, G.J, Russell, J. Crystal Growth 91 (1988) 135. [31 P.J. Wright, B. Cockayne, A.C. Jones, E.D. Orrell, P. O'Brien, O.F.Z. Kahn, J. Crystal Growth 94 (1989) 97. [4] M. Tirtowidjojo, R. Pollard, J. Crystal Growth 93 (1988) 108.
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[5] Sz. Fujita, A. Tanabe, T. Sakamoto, M. Isemura, Sg. Fujita, J. Crystal Growth 93 (1988) 259. [6] E. Sanchez, P.S. Shaw, J.A. O'Neill, R.M. Osgood, Chem. Phys. Lett. 147 (1988) 153. [7] J.J. BelBruno, J. Mol. Struct. (Theochem.) 358 (1995) 125. [8] J.J. BelBruno, Heteroatom Chem. 7 (1996) 39. [-9] M.J. Frisch, G.W. Trucks, M. Head-Gordon, P.M.W. Gill, M.W. Wong, J.B. Foresman, B.G. Johnson, H.B. Schlegel, M.A. Robb, E.S. Replogle, R. Gomperts, J.L. Andres, K. Raghavachari, J.S. Binkley, C. Gonzalez, R.L. Martin, D.J. Fox, D.J. DeFrees, J. Baker, J.J.P. Stewart, J.A. Pople
[10] [11] [12] [13] [14] [15]
(Eds.), GAUSSIAN 92, Revision A, Gaussian, Inc., Pittsburgh, PA, 1992. J.B. Mullin, S.J.C. Irvine, D.J. Ashen, J. Crystal Growth 55 (1981) 92. R.L. Jackson, Chem. Phys. Lett. 117 (1989) 315. A.C. Dillon, A.W. Ott, L.A. Okada, S.M. George, Surf. Sci. 334 (1994) 125. S.M. George, A.W. Ott, J.W. Klaus, J. Phys. Chem. 100 (1996) 13121. P.R. Jones, P.J. Desio, Chem. Rev. 78 (1978) 335. R.F. Heck, Organotransition Metal Chemistry, Academic Press, New York, 1974.