J. Phys. Chem. 1996, 100, 12355-12363
12355
Gas Phase Kinetics of Neutral Transition Metal Atoms: Reactions of Hf, Ta, Ir, Pt, and Au with Alkanes and Alkenes John J. Carroll† and James C. Weisshaar* Department of Chemistry, UniVersity of WisconsinsMadison, 1101 UniVersity AVenue, Madison, Wisconsin 53706-1396 ReceiVed: February 9, 1996; In Final Form: May 6, 1996X
The chemical kinetics of the ground-state, neutral transition metal atoms Hf, Ta, Ir, Pt, and Au with alkanes and alkenes is surveyed. Laser-induced fluorescence measures effective bimolecular rate constants at 300 K in 0.5-1.1 Torr of He buffer gas. Pt(5d96s1) is thus far the only ground-state neutral atom in the entire transition metal block that reacts with methane. It also reacts rapidly with linear alkanes, cyclopropane, and the alkenes. Au(5d106s1) reacts only with the largest alkenes studied, and then very slowly. In spite of their 6s2 ground-state configurations, Hf(5d26s2), Ta(5d36s2), and Ir(5d76s2) react rapidly with alkenes including ethene; these atoms do not react with methane and the linear alkanes. An attempt is made to interpret the pattern of reactivity in terms of simple chemical bonding concepts, including electron configuration and spin, guided by earlier extensive electronic structure calculations for the 4d series. It appears that the lanthanide contraction of the 6s orbital strongly influences the chemical reactivity of the 5d-series atoms.
I. Introduction The study of bare transition metal atom reactivity in the gas phase provides a particularly simple view of fundamental metal-hydrocarbon interactions, unencumbered by complications due to solvent and ligands. The electronic structure of atomic metal reactants is well understood,1 in contrast to ligated species in solution phase organometallic chemistry, metal clusters, and metal surfaces. Well-defined experiments in the gas phase can provide benchmarks for the development of both quantitative electronic structure theory and new qualitative insights as well. Early gas phase work focused on reactions of the cations M+ and M2+, many of which are remarkably reactive with alkanes and alkenes.2,3 This work has provided quantitative gas phase metal-hydrogen and metal-carbon bond energies,4 which vary systematically across rows and down columns of the transition metal block. Many mechanistic insights have been achieved as well. For example, we now have a fairly good understanding of how the pattern of low-lying M+ electronic states, including electron configuration and spin, governs chemical reactivity.5 However, the gas phase M+-hydrocarbon potential is in a sense artificially attractive. The long-range, ion-induced dipole force makes barriers to bond insertion and rearrangement unusual, whereas barriers are commonplace in condensed phase bond-breaking chemistry. Accordingly, we and others have been studying reactions of neutral metal atoms with alkanes, alkenes, and oxidants.6-9 In our experiments, laser-induced fluorescence (LIF) monitors the decay of ground-state, neutral metal atom density vs calibrated hydrocarbon density in a fast flow reactor at 300 K and 0.5-1.1 Torr of He buffer gas. We obtain ground-state effective bimolecular rate constants from a well-defined, Boltzmann distribution of collision energy and internal hydrocarbon energy. The total reaction rate may include both bimolecular contributions (from elimination reactions) and termolecular contributions (from collisional stabilization of longlived complexes). We have recently used stimulated emission † Present address: Department of Chemistry, Augsburg College, 2211 Riverside Avenue, Minneapolis, MN 55454. X Abstract published in AdVance ACS Abstracts, June 15, 1996.
S0022-3654(96)00408-X CCC: $12.00
pumping and LIF to study the chemical kinetics of specific excited states of V and Mo.10 A weakness of this work so far is the absence of experimental information about the identity of the products. In earlier work, we surveyed the kinetics of neutral atoms from the 3d and 4d series in reactions with linear alkanes, cyclopropane, and various alkenes.6 In the 4d series, we recently combined our experimental rate constants with electronic structure calculations of Siegbahn, Wittborn, and Blomberg and with statistical estimates of complex lifetimes to gain considerable insight into the identity of the products.11 In addition, the calculations reveal the underlying nature of certain potential energy barriers along the reaction path from M + hydrocarbon through CH or CC insertion intermediates and on to H2 elimination products. In this paper, we extend the kinetics measurements to reactions of neutral metal atoms from the 5d series (Hf, Ta, Ir, Pt, and Au) with the same hydrocarbons. The earlier work on cations from the 5d series shows them to be uniquely reactive; many cations dehydrogenate even methane in low-energy collisions.12 For the neutrals, we have already reported on the uniquely efficient reactions of ground-state Pt with linear alkanes, including methane.13 Calculations by Blomberg, Siegbahn, and Svensson provide quantitative insight into why the barrier between Pt(5d96s1,3D3) + CH4 to the CH insertion product H-Pt-CH3 is so small.13 Here we find Pt to react remarkably rapidly with cyclopropane and the alkenes as well. None of the other 5d atoms studied reacts with alkanes, but all of the atoms except Au react with alkenes, including ethene. The most remarkable new results are the rather efficient reactions of ground-state Hf and Ta with alkenes. This occurs in spite of their energetically well isolated 5dx-26s2 groundstate configurations. We hope these new kinetics results will inspire further theoretical effort in the difficult 5d series, in spite of complications arising from electron correlation, relativistic effects, and spin-orbit effects. II. Experimental Section A. Overview. The kinetics technique uses a sputtering source to produce gas phase transition metal atoms (M) in a fast flow of He buffer gas. LIF monitors the decay of the metal © 1996 American Chemical Society
12356 J. Phys. Chem., Vol. 100, No. 30, 1996
Carroll and Weisshaar
TABLE 1: Laser-Induced Fluorescence Transitionsa atom
transition
Hf Hf Ta Ta Ir Pt Pt Au Au
z3G°3 r a3F2 y3D°2 r a3F2 z4G°5/2 r a4F3/2 y4G°5/2 r a4F3/2 z4D°7/2 r a4F9/2 6p7°4 r 6s3D3 6p4°3 r 6s3D3 6p2P°3/2 r 6s2S1/2 6p2P°1/2 r 6s2S1/2
a
energy
(cm-1)
28584 28267 36014 36828 37515 37591 35322 41174 37359
wavelength (nm) 349.749 353.662 277.588 271.467 266.479 265.945 283.029 242.795 267.595
ln[M/M0] ) -k1nhctrxn
Assignments from ref 1.
atom concentration vs hydrocarbon number density. The flow reactor has been used for transition metal ion kinetics studies14 and more recently for neutral metal atom studies.6 The flow tube diameter is 7.3 cm. In the present work, the reaction length from the hydrocarbon gas inlet to the detection region is 74 cm. A thermistor indicates that the temperature in the reaction zone is 298 ( 5 K. The bulk of the buffer gas (mostly He with a little Ar to assist sputtering) flows directly through the sputtering source, which operates at a total pressure of about 2.5 Torr. The source is separated from the main flow tube and maintained at a higher pressure by a 1 cm diameter diaphragm. We measure reaction rate constants at three total pressures of buffer gas in the main flow tube, 0.50, 0.80, and 1.10 Torr, as measured midstream by a Baratron capacitance manometer. In all cases the flow characteristics are dominated by the properties of He, whose mole fraction is 0.93 at 0.5 Torr total pressure, 0.96 at 0.80 Torr, and 0.98 at 1.10 Torr. In addition to He, the flow contains about 3 mTorr of Ar (added through the source) and about 60 mTorr of N2 (added downstream of the source region to help quench excited states of the metal atoms). The hydrocarbon gas is added 70 cm downstream of the source through showerhead inlets oriented upstream relative to the He flow. Flow of hydrocarbon is regulated by a needle valve or a flow controller and monitored by a mass flow meter (Tylan). The hydrocarbon gases, ethene (Matheson, >99.5%), propene (Matheson or Liquid Carbonic, >99.0%), 1-butene (Matheson, >99.0%), isobutene (Matheson or Agfa, >99.0%), ethane (Matheson, >99.95%), propane (Matheson, >99.0%), n-butane (Matheson, >99.0%), and cyclopropane (Matheson, >99.0%), are used directly from the bottle. Because flow meter response is not linear with gas flow, we have calibrated the flow meters with each reactant gas. Metal atom reactant number density is monitored by unsaturated LIF using a commercial dye laser (Lumonics HD-300) pumped by a Nd:YAG laser (Lumonics HY-600, 10 Hz, 10 ns FWHM) and frequency doubled into the near-ultraviolet. To avoid saturation, the dye laser beam is filtered to admit only 10-100 µJ/pulse into the detection region. Metal atoms are detected with resolution of the J quantum number. Table 1 summarizes the atomic transitions used to study each metal atom. A wavelength scan checks the purity of the spectrum; lines are then selected for kinetics studies. We observe clean, sharp lines. All of the intense lines in the spectra are due to transitions from the ground term, indicating that most excited atoms have relaxed upstream of the reaction zone. The simplest kinetics mechanism consistent with the data is a single reaction step: k1
M + alkene 98 product
where.6 By measuring the logarithmic attenuation of metal atom number density (M) vs hydrocarbon number density nhc at fixed mean reaction time trxn ) zrxn/VM, we obtain the effectiVe bimolecular rate constant k1 from the pseudo-first-order expression:
(1)
k1 could include both bimolecular and termolecular components. Other more complicated kinetics models are discussed else-
(2)
The ratio M/M0 is measured as the ratio of time-integrated metal atom fluorescence signal with and without hydrocarbon reactant flow; zrxn is the length of the reaction zone; nhc is calculated from the metered hydrocarbon flow; and VM is the mean axial speed of the metal atoms, which is larger than the bulk flow velocity due to loss of metal atoms at the walls. Examples of the logarithmic plots are given in the earlier papers. Linear least-squares fitting produces the effective bimolecular rate constant k1, which could include both bimolecular and termolecular components. Collisional cascade of electronically excited states in the reaction zone has negligible effect on the ground-state rate constants. Quenching to the ground state on a time scale similar to that of ground-state reaction would lead to nonlinear semilog plots, which are not observed. The dynamic range of our rate constant measurements is about 3 × 10-14 to 1 × 10-9 cm3‚s-1. The typical precision of the rate constants is (10%, except for especially slow rates, which are less precise. The absolute accuracy of the rate constants is estimated as (30% due to small uncertainties in the many factors that enter eq 2. B. Mean Reaction Time and Error Estimates. Estimation of the mean reaction time trxn ) zrxn/VM deserves some discussion, since this is not easily measured directly in steadystate flow reactors. The He flow is essentially laminar and viscous with a few percent slip at the walls. The He flow velocity profile then becomes approximately parabolic, V(r) ) 2V0[1 - (r/a)2]. The mean of this distribution is V0, the bulk flow speed, which we can obtain experimentally from the measured flows of He, Ar, and N2, the measured temperature and pressure, and the diameter of the flow tube. Typical values of V0 are 5.6 × 103 cm‚s-1 at 0.50 Torr, 6.2 × 103 cm‚s-1 at 0.80 Torr, and 6.5 × 103 cm‚s-1 at 1.10 Torr, accurate to about 10%. The bulk flow speed increases slightly with increasing total pressure, primarily due to changes in the effective pumping speed of the Roots blower for He. Metal atoms diffuse to the walls and stick, so that the metal atom density vanishes at the walls and also peaks in the center of the flow tube where the He flow velocity is largest. Thus, the mean metal atom velocity VM is substantially larger than V0 by a factor that depends on the radial distribution of metal atoms. In earlier work using a pulsed, laser vaporization source, and LIF detection downstream,6a,c,f we directly measured arrival time distributions of Ti and Ti+ ground states vs distance downstream from the source for total pressure from 0.4 to 1.2 Torr of He. We found VTi ) 9070 ( 250 cm‚s-1 at 0.8 Torr of He, where VTi is taken as the slope of a plot of distance vs arrival time of the peak of the distribution. The velocity thus measured was the same within experimental error for ground-state Ti(3F), for the excited state Ti*(3P), and for the ground-state cation Ti+(4F). This strongly suggests that Ti neutrals and cations both stick to walls with unit efficiency and have quite similar radial density profiles across the flow tube. The measured ratio VTi/ V0 is 1.45 ( 0.07 at all three pressures, 0.50, 0.80, and 1.10 Torr. More recently, we have directly measured the speed of V* and Mo* excited-state atoms in the central 2 cm of the flow tube over short distances (∼1 cm) under the same steady-state
Gas Phase Kinetics of Transition Metal Atoms flow conditions used here.10 The excited states are prepared by stimulated emission pumping (SEP), and their density is probed vs time by LIF, with the viewing region restricted to the center of the flow tube. The result is (1.2 ( 0.2) × 104 cm/s for both V* and Mo* at 1.2 Torr of He. For parabolic flow with no slip, the velocity in the center of the flow tube is 2V0. Thus, our SEP measurements indicate V0 ) 6000 ( 1000 cm/s. These measurements are quite consistent with the bulk flow speed V0 ) 6500 cm/s at 1.1 Torr and with the earlier pulsed work on Ti, which measured the metal atom velocity averaged across the entire flow tube, VM = 9420 cm/s ) 1.45V0. Finally, measurements of likely pure bimolecular rate constants (such as M + O2 f MO + O)6a,c at 0.5, 0.8, and 1.1 Torr of He confirm that our relatiVe reaction times are correct for a given metal at the three pressures studied. Within the precision of (10%, the measured rate constants do not change with pressure. As in previous work,6 in this study we assume VM is the same as the measured values for Ti for all neutral metals. That is, for all atoms we use VM ) 8160 cm‚s-1 at 0.50 Torr of He, 9070 cm‚s-1 at 0.80 Torr, and 9420 cm‚s-1 at 1.10 Torr. There remains the possibility that VM varies significantly from metal atom to metal atom; we have carried out the pulsed arrival time measurements only for Ti. The different metal atoms likely have significantly different diffusion coefficients DM in He. However, as long as the metal atoms stick to the walls, differences of even a factor of 2 in metal atom diffusion coefficient affect VM and the rate constant obtained from eq 2 remarkably little. Diffusion causes a nearly constant fractional loss in metal atom density. By measuring M/M0 vs nhc at fixed distance z (rather than vs z or trxn at fixed nhc), we are sensitive only to the additional fractional loss due to reaction. On a more subtle level, diffusion and reaction do influence the radial distribution of metal atoms in the flow tube and, thus, VM and trxn. As described by Ferguson et al.,15 VM exceeds V0 because both the radial velocity profile and the metal atom density profile are peaked in the center of the flow tube; i.e., there are more metal atoms where the flow is fastest. For planar flow, V(r) ) constant, the radial distribution of metal atoms can be decomposed into fundamental modes which are zeroorder Bessel functions J0(kir/a), where ki is the ith root of J0. If a combination of modes is present initially, the amplitude of each mode decays exponentially vs distance z as exp(-ki2DMz/ a2V0), where V0 is the bulk flow speed, a the flow tube radius, and DM the diffusion constant of the metal atom in He. Since λ12 ) 5.8, λ22 ) 30.5, λ32 ) 74.9, etc., for our flow reactor, the radial distribution quickly relaxes to the fundamental mode over a distance of about 10 cm. For the realistic parabolic velocity profile, but still neglecting reactive losses, the radial modes are no longer the J0(kir/a), but they can be obtained numerically. The first few characteristic constants analogous to the λi2 become γ12 ) 3.7, γ22 ) 22.3, γ32 ) 57.0, etc. The fundamental mode, which looks quite similar to J0, again quickly dominates. Reagents are added sufficiently far (70 cm) downstream of the source orifice to allow the parabolic velocity distribution to fully develop and the fundamental radial mode to dominate. In that limit, the numerical work finds VM = 1.6V0; i.e., trxn in eq 2 must be decreased by a factor of 1.6 to account for the faster transport of metal atoms down the flow tube. Our actual measurements on Ti and Ti+ find VM/V0 ) 1.45, a correction factor about 10% smaller than the model of Ferguson. Chemical reaction beginning at z ) 0 can further distort the radial distribution from the fundamental mode appropriate for the parabolic velocity profile in the absence of reaction; i.e., it can mix in higher modes at z ) 0. The extent of this mixing
J. Phys. Chem., Vol. 100, No. 30, 1996 12357 depends on the ratio of reactive losses to diffusive losses, which Ferguson quantifies as the dimensionless parameter ξ ) Qk/ πγ12DMV0. Here Q (molecules/s) is the flow of reactant gas and k (cm-3‚s-1) is the bimolecular rate constant. Qk is proportional to the reaction rate, while γ12DM is proportional to the rate of loss to diffusion. For ξ , 1, diffusion dominates reaction and reaction does not distort the density profile or the correction factor. The diffusion coefficients of our metal atoms in He are unknown, but we can estimate them from the well-established formula for the interdiffusion of hard spheres,16 which involves only the temperature, the total number density, the He/M reduced mass (all of which we know), and the atomic radii rHe and rM. The only factor in the diffusion coefficient D that changes appreciably among metal atoms goes as (rHe + rM)-2. We take rHe ) 1.0 Å from transport data.16 We use ns, the mean radius of the valence ns orbital from Hartree-Fock calculations,17 as our estimate of rM for the various metal atoms. The metal radii decrease substantially across each transition series and change rather little down a column. Examples are rTi ) 2.00 Å, rZr ) 2.16 Å, rHf ) 2.16 Å, rNi ) 1.62 Å, rPd ) 1.79 Å, and rPt ) 1.83 Å. The formula gives the sensible estimate DTi ) 450 cm2‚s-1 at 1.0 Torr of He. The other DM estimates all lie in the range 375-600 cm2‚s-1 at the same pressure. On a relative scale, we can infer with some confidence that the various DM values vary in the approximate range 0.83DTi to 1.32DTi at fixed He pressure. Differences in total pressure or in the intrinsic diffusion constant between metal atoms alter Ferguson’s dimensionless parameter ξ. The key issue is how much the ratio VM/V0 changes with ξ due to changes in the radial distribution of metal atoms. Over the range explored numerically by Ferguson (0 e ξ e 4), the correction factor for trxn varied only 1-2%; i.e., VM/V0 is remarkably insensitive to pressure, diffusion coefficient, and extent of reaction. Two effects tend to cancel. Reaction introduces higher order diffusion modes, which enhance diffusion losses and tend to flatten the radial distribution and decrease VM. However, reaction makes the fundamental mode itself more peaked near the center of the flow tube, which tends to increase VM. Using the estimate DTi ) 450 cm2‚s-1 at 1 Torr, we find that in our experiment ξ varies from 0 to 3.5 when the hydrocarbon density is increased sufficiently to decrease the measured metal atom signal by a factor of e3, roughly the dynamic range of our experiment. For other metals in the worst case (1.1 Torr of He, DM ) 375 cm2‚s-1), ξ becomes only as large as 4.6, and the model would still predict a change of only a few percent in the correction to trxn. Thus, we estimate that use of the measured VTi for all other metals should lead to systematic errors no larger than about 5%, which is not highly significant on the scale of the overall estimated uncertainty ((30%). We compare rate constants measured in our flow reactor with results from other laboratories in section IV. III. Results We have studied the reactions of five neutral, ground-state transition metal atoms with nine hydrocarbons at 0.80 ( 0.05 Torr of He and 300 ( 5 K. The atoms and levels probed (Table 1) are as follows: Hf (5d26s2,3F2) and Ta (5d36s2,4F3/2) from the left-hand side of the 5d series; and Ir (5d76s2,4F9/2), Pt (5d96s1,3D3), and Au (5d106s1,2S1/2) from the right-hand side. These atomic states1 are labeled by electron configuration, total electron spin S, orbital angular momentum L, and total angular momentum J, in spite of the substantial jj coupling in the 5d series. The hydrocarbon collision partners include linear and branched alkenes, methane and linear alkanes, and cyclopropane. The
12358 J. Phys. Chem., Vol. 100, No. 30, 1996
Carroll and Weisshaar
TABLE 2: Effective Bimolecular Rate Constants k1 (10-12 cm3‚s-1) for Reactions of Transition Metal Atoms with Hydrocarbons at 300 ( 5 K and 0.80 ( 0.05 Torr of Hea
TABLE 3: Effective Bimolecular Rate Constants k1 (10-12 cm3‚s-1) vs He Pressure for Reactions with Ethene and 1-Butene at 300 ( 5 Ka
reactant
Hf (3F2)
Ta (4F3/2)
Ir (4F9/2)
Pt (3D3)
Au (2S1/2)
reactant (PHe)
ethene propene 1-butene isobutene cyclopropane methane ethane propane n-butane
28.2 ( 2.8 89 ( 9 92 ( 9 146 ( 15 NR NR NR NR
8.4 ( 0.8 47 ( 5 53 ( 5 109 ( 11 NR NR NR NR
62.4 ( 6.2 114 ( 11 122 ( 12 150 ( 15 NR NR NR NR
376 ( 38 426 ( 43 435 ( 44 429 ( 43 136 ( 14 3.08 ( 0.31 39.9 ( 4.0 93 ( 9 162 ( 16
NR NR 0.3 ( 0.2 0.2 ( 0.2 NR NR NR NR NR
ethene (0.5 Torr) (0.8 Torr) (1.1 Torr) 1-butene (0.5 Torr) (0.8 Torr) (1.1 Torr)
NR means no reaction observed, i.e., k1 < 3 × 10-14 cm3‚s-1. Dashes beside cyclopropane indicate that a reaction was observed, but the rate was too samll to preclude the possibility of impurities dominating the measurement. Uncertainties refer to the precision of the measurements; the Au source was unusually unstable. Absolute accuracies are typically estimated as (30%. a
effective bimolecular rate constants in 0.8 Torr of He at 300 K are collected in Table 2. Each rate constant at 0.8 Torr is the mean of at least two individual determinations. The error estimates in Table 2 refer to the precision of the 0.8 Torr rate constants. They represent the larger of (10%, the range of two measurements, or (1 standard deviation of the mean of three measurements. The absolute accuracy of the observed rate constants is estimated as (30%, limited by the uncertainties in trxn and gas flow calibrations, incomplete reagent mixing, source instability, and random noise in the data. For these neutral reactions, the natural benchmark rate for comparison is the hard-spheres collision rate, estimates of which lie in the range (2-4) × 10-10 cm3‚s-1.18 The sensitivity limit of the apparatus is k1 g 3 × 10-14 cm3‚s-1, so we have a dynamic range of about four decades. Of the five metal atoms studied, only Pt reacts with methane and the linear alkanes ethane, propane, and n-butane, as reported earlier.13 Cyclopropane reacts very rapidly with Pt and not at all with Au. Using nominal >99% purity cyclopropane, we observed slow apparent reactions with Hf, Ta, and Ir (0.33 × 10-12, 0.13 × 10-12, and 0.4 × 10-12 cm3‚s-1, respectively). We do not report these rates in Table 2 since there is some chance that alkene impurities at the 0.5% level could dominate the slow observed rates. However, the rates are stable across several lecture bottles, and attempts to find alkene in the cyclopropane cylinders by mass spectrometry failed. We suspect these rates are truly due to slow cyclopropane reactions. In contrast to the alkanes, the reactions with alkenes are quite efficient for all atoms except Au. Going from ethene to propene, the reaction rates increase substantially for Hf and Ta, but only modestly for Ir and Pt. Rates with alkenes increase only moderately from propene to the butenes. Platinum is easily the most reactive transition metal atom we have studied in the 3d, 4d, and 5d series. The Pt + C2H4 rate is already some 1.5 times faster than our hard-spheres estimate. Gold is generally quite inert, but we did discern very slow reactions with the butenes. Since butenes are likely more reactive than impurities, we tend to trust these slow rates. The Au data were noisier than data for the other metal atoms and the rate constants are small, so the Au + butene rate constants are less accurate than the others. In an attempt to gain insight into whether the reactions are bimolecular or termolecular, we measured selected rate constants over the limited He pressure range 0.5-1.1 Torr (Table 3), again at 300 K. Each entry at 0.5 and 1.1 Torr in Table 3 represents one or usually two determinations of k1 from 10-15 data points of metal atom LIF signal vs flow of hydrocarbon. If a termolecular reaction were in the linear, low-pressure regime of effective bimolecular rate constant vs He pressure, the
Hf (3F2)
Ta (4F3/2)
Ir (4F9/2)
Pt (3D3)
Au (2S1/2)
9.6 ( 1.9 59 ( 18b 387 ( 77 29 ( 9b 28.2 ( 2.8 8.4 ( 0.8 62 ( 6 376 ( 38 NR 60.3 ( 18b 351 ( 105b 28.1 ( 8.4b 97 ( 29b 92 ( 9 89 ( 27b
57 ( 11 53 ( 5 -
116 ( 35b 122 ( 12 435 ( 44 116 ( 35b -
0.3 ( 0.2 -
NR means no reaction observed, i.e., k1 < 3 × 10-14 cm3‚s-1; dash means not studied. Error limits are estimates of precision. Absolute accuracies estimated as (30%. b Denotes cases for which we have made only one determination of the rate constant. Stated uncertainties are (30%. a
measured rate constant k1 would increase by a factor of 2.2 as the He pressure increased from 0.5 to 1.1 Torr. None of the rates with ethene or 1-butene showed significant pressure dependence within the estimated relative uncertainty of (20%. As we have learned from the 4d series, this does not rule out a saturated termolecular mechanism, as discussed earlier.6 Only the Pt + CH4 rates reported earlier13 showed a modest increase with He pressure, a clear signal of a termolecular component of the rate. IV. Discussion A. Comparison with Previously Measured Rate Constants. A significant number of rate constants measured in our laboratory6 can be directly compared with work from other laboratories. Lian et al.19 studied Mo reactions with alkenes and other gases in He buffer. Parnis et al.20 studied Zr, Nb, and Ta reactions in He. Both of these studies used the same pulsed laser vaporization, LIF detection method on the flow tube at the NRCC in Ottawa. Senba et al.21 used a steady-state discharge source and LIF in a flow tube similar to ours to study Ti and V reactions, including excited-state reactions, also in He. They calibrated the reaction time using a pulsed laser vaporization source. The degree of quantitative agreement is mixed, as summarized in Table 4. For Mo + C2H4, Zr + C3H8, Nb + C3H8, and Ta + C3H8, the Ottawa groups find slow reactions, while we find no reaction. This may be due to slow termolecular association reactions at the higher He pressures employed, 1.8-4.0 Torr in Ottawa vs our 0.8 Torr. More disturbing are the consistently larger rate constants obtained in Ottawa for faster (presumably bimolecular) reactions at He pressures quite similar to ours. The Ottawa rates are uniformly about a factor of 2 faster than the Madison rates for the same or similar He pressure, which strongly suggests a systematic error. The measurements on Ti and V from Himeji, which are calibrated against a pulsed source, agree quite well with ours. A clear advantage of pulsed methods6f is that the reaction time can be measured directly. Because of these factor of 2 discrepancies, in section II-B we discussed in detail how we have calibrated our reaction times for Ti using a pulsed source and assumed the same reaction times for other atoms, based on Ferguson’s detailed modeling of the flow characteristics. It is possible that this analysis fails. To explain the discrepancy, Mo, Zr, Nb, and Ta would have to move down our flow tube about twice as quickly as Ti, so that our measured rate constants would be systematically a factor of two too small. However, we believe our analysis is sound and our rates are accurate to the stated (30%. The five atoms in question are all from the left-hand side of the transition metal block and have very similar
Gas Phase Kinetics of Transition Metal Atoms
J. Phys. Chem., Vol. 100, No. 30, 1996 12359
TABLE 4: Comparison of Rate Constants with Previous Work
TABLE 5: Comparison of Effective Bimolecular Reaction Rates (Units of 10-12 cm3‚s-1) of Congeners from the 3d, 4d, and 5d Series with Selected Hydrocarbons at 300 K in 0.8 Torr of Hea
reaction
PHe (Torr)
methoda
k (10-12 cm3‚s-1)
ref
Mo + C2H4
1.8 0.8 0.8 0.8 4.0 0.8 0.45 0.8 4.0 0.8 0.42 0.8 9.0 0.8 0.45-1.0 0.8 0.7 0.8 0.7 0.8
LV SS LV SS LV SS LV SS LV SS LV SS LV SS LV SS SSb SSc SSb SSc
0.073 NR 0.78 0.4 0.076 NR 140 59 0.17 NR 770 314 0.04 NR 22 8.4 6.3 6.2 8.2 9.6
20 6e 20 6e 21 6e 21 6e 21 6e 21 6e 21 this work 21 this work 22 6b 22 6b
Mo + C3H6 Zr + C3H8 Zr + C2H4 Nb + C3H8 Nb + C2H4 Ta + C3H8 Ta + C2H4 Ti + C3H6 V + C3H6
a LV, laser vaporization source, flow tube at NRCC, Ottawa, Canada. This is a pulsed method in which the reaction time is measured directly. SS, steady-state discharge source. Reaction time must be calibrated indirectly. b Flow tube at Himeji Institute of Technology, Japan. c Flow tube at University of WisconsinsMadison.
sizes as judged by ns for the valence ns orbital, which varies only in the range 2.00-2.09 Å. While we cannot suggest why the Ottawa rates might be uniformly twice too fast, the fact that Ottawa’s fastest rate constant of 7.7 × 10-10 cm3‚s-1 is about twice the estimated hard-spheres rate suggests there may be a systematic error. There are no strong, long-range forces between neutral reactants; a harpooning mechanism seems unlikely due to the negligible electron affinity of alkenes. In contrast, our largest rate constants are consistently about 4 × 10-10 cm3‚s-1, reassuringly comparable to hard-spheres estimates. Clearly it would be useful for the Ottawa groups to measure Ti and V rates and for us to measure trxn directly for metal atoms other than Ti. In the meantime, the factor of 2 discrepancies do not affect our qualitative discussion of the factors controlling reactivity. B. Simple Bonding Model and Pattern of Low-Lying Electronic States. Before we discuss electronic effects on reactivity, it is important to appreciate how sensitive 300 K rate constants are to small differences in activation energy. We can use estimated reaction efficiencies to set limits on the activation energies of each observed reaction. For those reactions that occur in the flow tube at 300 K, we first estimate the hardspheres collision rate constant khs from approximate atomic and molecular sizes.18 These estimates of khs lie in the range (24) × 10-10 cm3‚s-1. Since kBT is only 200 cm-1 ) 0.6 kcal mol-1 at 300 K, only those reactions with rather small activation energies will be observed. Assuming the Arrhenius dependence
k1(T) ) A exp(-E/kBT)
(3)
and using the estimated hard-spheres collision rate constant khs as an upper bound on the pre-exponential factor A, we can convert each k1/khs into an upper bound Emax on the activation energy. The observed range k1 ) (0.2-435) × 10-12 cm3‚s-1 corresponds to a range in Emax of about 4.4 to 0.0 kcal/mol. This assumes that only translational energy is effective in overcoming the barrier, which is not unreasonable for the kinds of reaction steps in question. We are neglecting possible contributions from internal energy, which can amount to several kilocalories per mole at 300 K for the larger alkanes.11
atom
ethene
Sc Y
NR 8.2 ( 0.8
9.5 ( 1 141 ( 14
propene
NR NR
propane
cyclopropane -
Ti Zr Hf
NR 59 ( 6 28.2 ( 2.8
6.2 ( 0.6 153 ( 15 89 ( 9
NR NR NR
NR -
V Nb Ta
NR 314 ( 34 8.4 ( 0.8
96 ( 10 360 ( 36 47 ( 5
NR NR NR
-
Cr Mo
NR NR
NR 0.38 ( 0.19
NR NR
NR NR
Co Rh Ir
NR 11.3 ( 1.1 62 ( 61
NR 115 ( 12 14 ( 11
NR 9.4 ( 0.9 NR
NR 14.2 ( 1.4 -
Ni Pd Pt
0.5 ( 0.05 15.0 ( 1.5 376 ( 38
11 ( 4 189 ( 19 426 ( 43
NR 1.22 ( 0.12 93 ( 9
10 ( 1 59 ( 8 136 ( 14
Cu Ag Au
NR NR NR
NR NR NR
NR NR NR
NR NR NR
aNR means no reaction, i.e., k e 3 × 10-14 cm3‚s-1. Uncertainties refer to estimated precision; absolute accuracies (30%. Data from present work plus refs 6, 11, and 13. Dashes under cyclopropane indicate that a slow reaction was observed, but the rate was too small to preclude the possibility of impurities dominating the measurement.
Figure 1. Qualitative potential energy surfaces. The atomic ground state is ill-suited to bond formation, due to its high spin or its dx-2s2 configuration. Attractive diabatic surfaces then correlate with excited state reactants M* + hydrocarbon. The barrier height on the adiabatic ground state surface (heavy line) is related to the promotion energy ∆E and to the slopes of the attractive and repulsive potentials.
In many cases we are now able to compare the reactivity of 3d-, 4d-, and 5d-series congeners with alkanes and alkenes. Table 5 collects the reaction rates with ethene, propene, propane, and cyclopropane for comparison. We now attempt to interpret the observed pattern of reaction rates in terms of simple chemical bonding models and the pattern of low-lying electronic states of the bare metal atoms. A simple picture that has enjoyed good qualitative success in both neutral and cation chemistry is shown in Figure 1. We imagine a reaction path leading from M + hydrocarbon asymptotes to formation of M-hydrocarbon chemical bonds. In M + alkane, the initial chemical step would be CH or CC bond insertion. In M + alkene, it would be either addition to the double bond (forming a π complex or a metallacyclopropane) or CH insertion. Following the initial step, in a termolecular mechanism collisional stabilization at 0.8 Torr of He would occur to form the stable, chemically bound adduct. This becomes more likely the deeper the well and the more vibrational degrees of freedom in the complex, as
12360 J. Phys. Chem., Vol. 100, No. 30, 1996 discussed in detail earlier.11 In especially favorable cases, subsequent rearrangement and bimolecular elimination of H2 might occur. Particularly for neutral transition metal atoms, barriers often arise along the lowest energy, adiabatic path from reactants to chemically bound products. The reason is that the ground state is usually ill-suited for formation of chemical bonds, as indicated by the repulsive diabatic potential arising from ground-state asymptotes in Figure 1. Two cases arise often. First, many ground-state atoms have an unfavorable dx-2s2 configuration. Such states typically have the proper electron spin for bond formation, but double occupancy of the large, valence s orbital renders them chemically inert. This should lead to highly repulsive M + hydrocarbon potentials. Second, some atoms have a more favorable dx-1s1 ground-state configuration, but invariably it is the high-spin dx-1s1 term. Calculations indicate11,22 that at best such terms give rise to weakly bound, longrange complexes with alkanes and alkenes. The high-spin potentials become repulsive at shorter range. As metal and hydrocarbon approach, the repulsive potential from either a dx-2s2 or a high-spin dx-1s ground state may cross more attractive, low-spin potentials arising from appropriate excitedstate asymptotes (Figure 1). The usual candidates are low-spin dx-1s terms; on the right-hand side of the transition series, dx terms may also come into play. In this simple picture, a coupling matrix element, either electrostatic or spin-orbit, mixes the two diabatic potentials to avoid surface intersection. If the coupling is not too large, the result is an adiabatic potential with a barrier between groundstate reactants and chemically bound products. We would expect electrostatic interactions between two potentials of the same spin typically to be larger than spin-orbit interactions between potentials of the same configuration but different spin. Spin-orbit interaction should become stronger in moving from the 3d to the 4d to the 5d series. All else being equal, we would expect the height of the adiabatic barrier to chemical bond formation to be smaller when the metal atom ground state is dx-1s rather than dx-2s2, when the promotion energy to the chemically active excited state is relatively small, and when the strength of the hydrocarbon bond being broken is relatively small. Our discussion will press these simple, qualitative arguments as far as possible, combining atomic data with simple bonding models and more detailed insights gleaned from earlier electronic structure calculations. For representative atoms from the left- and right-hand sides of the transition series, Figures 2 and 3 compare the patterns of low-lying electronic states of particular configuration and spin types, either high-spin or low-spin. The congener trios TiZr-Hf, V-Nb-Ta, Co-Rh-Ir, and Ni-Pd-Pt include all of the 5d-series atoms that have been found to be reactive. For each atom, we always show the lowest dx-2s2 state and both the high- and low-spin states of the lowest dx-1s configuration. We include the lowest dx state when it lies below 40 kcal/mol excitation energy. For simplicity, we omit dx-2sp states, which always lie 40 kcal/mol or more above the ground state, even though they participate significantly in covalent bonding on the left-hand side.22 Much has been written about the underlying causes of the patterns of electronic states and how they vary across a row and down a column of the transition metal block.22-24 The interplay of s and d orbital energies and exchange interactions favoring high-spin states determine the basic pattern. In the 3d series, the 4s orbital energy lies well below the 3d energy, so 3dx-24s2 ground states dominate. In addition, 4s is a much larger orbital than 3d, so 4s character dominates at least the
Carroll and Weisshaar
Figure 2. Comparison of atomic states for selected congeners from the left-hand side of the transition metal block. Shown are the lowest atomic levels of each configuration and spin (h ) high-spin; l ) lowspin), without averaging over J. Many other states exist.1 Energies in kcal/mol.
Figure 3. Comparison of atomic states for selected congeners from the right-hand side of the transition metal block. Shown are the lowest atomic levels of each configuration and spin (h ) high-spin; l ) lowspin), without averaging over J. Many other states exist.1 Energies in kcal/mol.
first σ bond to a metal atom. Because of the large 3d/4s size disparity, double bonds involving both 4sσ and 3dπ orbitals must choose a bond length that is optimal for neither σ nor π.24 In the 4d series, the 5s and 4d orbitals are more comparable in orbital energy and also in size. This favors 4dx-15s configurations, which are either the ground state or a low-lying excited state. It also improves the efficacy of sd hybridization, leading to stronger M-C and M-H bonds involving greater d orbital character. In the 5d series for Hf and beyond, the lanthanide contraction has lowered the 6s orbital energy relative to 5d, so that 5dx-26s2 ground states are again most common. Bond energies to carbon and hydrogen are unusually large in the 5d series,4,12,22 in part because the valence 6s and 5d orbitals are quite similar in size. This helps drive chemical reactions of the bare 5d-series atoms.
Gas Phase Kinetics of Transition Metal Atoms One simple reason for the wide variability of metal atom reactivity at 300 K is immediately obvious. The 0-5 kcal/mol range of activation energies probed by the 300 K kinetics experiment is small compared with the range of variation of important electronic excitation energies among the neutral atomic metal atoms. Unless excited-state potentials can plummet sufficiently rapidly, the adiabatic barrier on the groundstate surface will be too large to permit reaction at 300 K. C. Reactions with Alkanes. In the simplest view, each M + alkane reaction rate compares the adiabatic barrier to bond insertion with the roughly 5 kcal/mol cutoff imposed by the 300 K measurements. Roughly speaking, methane, the linear alkanes, and cyclopropane form a hierarchy of reactivity. Only Pt reacts with methane. Those few metals that react slowly with linear alkanes (Pt, Rh, and Pd) react much more rapidly with cyclopropane. Most of the 4d and 5d atoms react with cyclopropane, albeit quite slowly in most cases. The exceptions are Mo and Au. In the 3d series, only Sc, V, and Ni react with cyclopropane. For all 4d series atoms and for Ir and Pt in the 5d series, electronic structure calculations from Siegbahn and co-workers have provided a detailed, highly instructive mechanistic picture.13 In the 4d series, for a given metal atom the calculations consistently show smaller barriers to CH insertion in methane than to CC insertion in ethane, in spite of the fact that the C-C bond is weaker. This is due to the more directional nature of the partially occupied orbitals for methyl as compared with hydrogen. We expect still smaller barriers to CH insertion in ethane, whose C-H bonds are weaker than those of methane. In cyclopropane, the calculations find the smallest barriers of all, with the barrier to CH insertion comparable to the barrier to CC insertion. This occurs even though the C-C bond in cyclopropane is very weak and the C-H bond is actually stronger than in methane. Not surprisingly, barriers to bond insertion are quite sensitive to specific orbital effects and correlate only roughly with the strength of the bond being broken. According to the calculations, the Pd + alkane reactions form long-range, η2 complexes that are collisionally stabilized. This mechanism is apparently unique to Pd because of its unusual 4d10 ground state. The calculations find a modest barrier to insertion of Pd into a C-H bond (4 kcal/mol for Pd + CH4). However, the H-Pd-alkyl intermediate is only weakly bound (2 kcal/mol for H-Pd-CH3). In addition, a large barrier to subsequent hydrogen migration arises [e.g., in H-Pd-C2H5 f Pd(H)2C2H4]. The reason is apparently the large promotion energy to Pd(d9s) configurations (Figure 3). In contrast, the Rh and Pt reactions with linear alkanes involve insertion of the metal atom into a C-H bond over at most a small barrier from ground-state reactants. In these cases, the calculated H-M-CH3 potential well is deep, 19 kcal/mol for Rh and 32 kcal/mol for Pt. The Rh reactions with ethane and larger alkanes likely proceed all the way to bimolecular H2 elimination products. The Pt + CH4 reaction almost surely stops at the H-Pt-CH3 intermediate, which is collisionally stabilized. The Pt reactions with larger alkanes may or may not proceed to H2 elimination. The obvious question is why Rh and Pt, both atoms from the lower-right-hand corner of the transition metal block, are uniquely able to activate C-H bonds of linear alkanes at 300 K. Examination of Figure 3 suggests that the existence of both dx-1s and dx states at low energy is important. In Rh and Pt, the dx-1s ground states avoid the strong, long-range repulsion suffered by more typical dx-2s2 ground states. A low-lying dx-1s state helps lower the barrier to C-H bond insertion, since the
J. Phys. Chem., Vol. 100, No. 30, 1996 12361 resulting pair of σ bonds must ultimately be formed from sd hybrid orbitals. Pt is also aided by the generally greater strength of bonds in the 5d series and perhaps by the degradation of spin as a good quantum number by spin-orbit coupling. As metal and alkane approach, admixture of dx character can further reduce electron-electron repulsion. In comparison with Rh, Pd suffers from its large promotion energy to d9s configurations. In contrast, Co and Ir apparently suffer from their s2 ground states, which render them unable to react with linear alkanes at 300 K. Cobalt has the more favorable excitation energy to the important low-spin dx-1s states, but it is Ir and not Co that reacts slowly with cyclopropane. This is again consistent with the more efficient sd hybridization and much stronger bonding in the 5d vs the 3d series. Nickel has a favorable triplet dx-1s ground state and a very low-lying singlet dx-1s excited state, but it also fails to react with linear alkanes. For Ni, we suspect the barrier to insertion in C-H bonds of linear alkanes is just large enough (g5 kcal/mol) to preclude observation of reaction at 300 K, since Ni reacts fairly efficiently with cyclopropane. Turning to the left-hand side of the transition metal block, we find that no neutral ground state reacts with linear alkanes at 300 K. Most of these atoms have dx-2s2 ground-state configurations; the important low-spin dx-1s excited states lie above 24 kcal/mol in all cases. Evidently the repulsive s2 potential rises too rapidly to cross the attractive dx-1s excited state potential at sufficiently low energy to allow access to the attractive surface at 300 K. We do observe slow reaction of Sc, Y, Zr, Hf, V, and Ta with cyclopropane, suggesting that for these metal atoms the barriers to CH insertion in linear alkanes may fall in the 5-10 kcal/mol range. According to these arguments, Nb presents the most favorable case for C-H bond insertion among ground-state neutral atoms on the left-hand side. It has a high-spin d4s ground state that diminishes long-range repulsion. The low-spin d4s state lies at a relatively benign excitation energy of 24 kcal/mol, yet we observe no reaction with linear alkenes at 300 K. The moderately efficient rate with cyclopropane again suggests that the barrier to CH insertion in linear alkanes may lie just above 5 kcal/mol. Measurements at higher temperature could provide an experimental estimate of the barrier height for comparison with theory. In the center of the transition metal block, Cr(3d54s), Mn(3d54s2), and Mo(4d55s) have extremely stable (energetically well isolated) ground states due to the d-d exchange interaction. The d5s Cr and Mo configurations avoid some long-range repulsion compared with s2 ground states. However, the promotion energy to the lowest state (d5s,5G) that breaks the high-spin coupling of the five d electrons is prohibitively large, 59 kcal/mol in Cr and 47 kcal/mol in Mo. Apparently, this creates large barriers to bond insertion on the adiabatic groundstate surface. Manganese suffers from the 49 kcal/mol promotion energy to the 3d64s(6D) state, which is again related to d-d exchange. Accordingly, Cr, Mn, and Mo do not react with alkanes, even with cyclopropane. Tungsten and Re have not yet been studied. D. Reactions with Alkenes. Almost all of the 4d- and 5dseries ground-state atoms react with alkenes, including ethene (Table 5). The rates typically increase sharply from ethene to propene and then level off for larger alkenes. In the 3d series, only Ni reacts with ethene, but most of the other atoms react with larger alkenes. Co reacts unusually slowly with alkenes, and Cr fails to react with even the larger alkenes. The products of the M + alkene reactions are not known from experiment. In the 4d series, the combination of reaction rate vs He pressure, electronic structure calculations, and RRKM
12362 J. Phys. Chem., Vol. 100, No. 30, 1996 modeling of complex lifetimes gave considerable insight. We inferred that the Rh and Pd reactions with ethene form strongly bound, low-spin π complexes that are collisionally stabilized. The Y, Zr, and Nb reactions with alkenes very likely involve initial formation of a π complex followed by CH insertion over a substantial barrier and eventual rearrangement and H2 elimination. Mo forms weakly bound, high-spin π complexes that are inefficiently stabilized. For the 5d series, we can only speculate on the product identity. The Pt + alkene reactions are so efficient that we suspect they involve bimolecular H2 elimination. Simple RRKM lifetime arguments indicate that inefficient reactions with the smallest alkene, ethene, have the best chance to exhibit pressure-dependent rates, yet none of the M + C2H4 reaction rates in the 5d series in fact show a pressure dependence, suggesting a bimolecular H2 elimination mechanism throughout the series. If so, the reactions with larger alkenes are probably also bimolecular. Further theoretical and experimental work is required for more definitive product identification. These results pose intriguing questions about electronic structure. One simple question is why neutral metal atoms react so much more efficiently with alkenes than with alkanes. Part of the answer is surely thermochemical. To insert in a C-H bond of the alkane RH, the metal atom must break a strong C-H bond while forming two much weaker bonds in H-MR. In forming a M-alkene π complex, the π bond of the alkene is only weakened, not broken. Potentially two new dative bonds are formed, one by donation of π2 into an empty σ orbital on the metal and the other by “back donation” from a metal dπ orbital into the empty π* orbital of the alkene. In the strong limit of metallacyclopropane formation, the π bond is actually broken but two new σ bonds are formed. The atomic configuration best suited for such bonding to alkenes is again dx-1s. In the low-spin dx-1s state, the atom can place two electrons into a hybrid orbital sd-, which is formed from the valence sσ and dσ orbitals.22,23 This hybrid maximizes probability in the plane perpendicular to the axis of M-alkene approach, thus minimizing repulsion between the (sd-)2 metal configuration and the π2 configuration of the alkene. The symmetric combination sd+, which can be empty for atoms from the left-hand side, serves as the acceptor orbital on the metal atom. The back donation from metal to π* of the alkene requires one or, optimally, two electrons in a dπ orbital. On the left-hand side, only one such electron is typically available. On the right-hand side, a dx state can engage in the same mechanism without hybridization, using the empty valence s orbital as acceptor. This occurs in ground-state Pd. On the left-hand side, a high-spin dx-1s state can also bind weakly to an alkene, primarily using an empty dσ orbital as acceptor and a half-filled dπ orbital as donor.22a This occurs for ground-state Nb(d4s,6D), which reacts very efficiently even with ethene. In contrast, the Mo(d5s,7S) ground state lacks an empty d orbital and is inert. Rather surprisingly, and in sharp contrast with the M + alkane reactions, most metal atoms with s2 ground states react with alkenes at 300 K. Quite efficient examples include Y, Zr, and Hf from the left-hand side and Ir from the right-hand side. Apparently, the attractive potentials arising from dx-1s excited states fall much more rapidly for M + alkene than for M + alkane, or the potentials from dx-2s2 ground states rise much more slowly, or both (Figure 1). The kinetics data combined with the energy level diagrams of Figures 2 and 3 further suggest that the M + alkene barriers arising from an s2 ground state are much larger in the 3d series than in the 4d or 5d series. This is consistent with the greater
Carroll and Weisshaar s/d size disparity in the 3d series.22-24 For example, the promotion energy to both low-spin and high-spin d4s states of V in the 3d series is substantially smaller than that to the d3s states of Zr in the 4d series, yet V does not react with ethene and reacts only slowly with propene, while Zr reacts quite rapidly with all alkenes. Cobalt provides another example of a 4s2 ground state and very low promotion energies to 3d84s states, but very inefficient reactions with alkenes. The reactivity with alkenes of the 5d-series atoms having 6s2 ground states is quite remarkable in comparison with the 3d and 4d series (Table 5). Tantalum reacts moderately rapidly with ethene, even though its promotion energies to high- and low-spin d4s states are 28 and 59 kcal/mol, respectively. The analogous promotion energies in V are only 6 and 24 kcal/mol, but V does not react with ethene. On the other hand, Ta reacts less rapidly than Nb, indicating the repulsive effect of the 6s2 ground state. The Hf + alkene reactions are particularly efficient and thus rather puzzling. They occur in spite of very large promotion energies (40 and 67 kcal/mol, respectively) to high-spin and low-spin d3s states. The Ir promotion energies, in turn, are similar to those of V, yet Ir reacts rapidly with all alkenes. Indeed, Ir, which has a high-spin d7s2 ground state, reacts as rapidly with alkenes as Rh, which has a presumably more favorable high-spin d8s ground state. In the 5d series, the lanthanide contraction of the 6s orbital stabilizes 6s2 configurations and makes the 5d and 6s orbital sizes quite similar. This helps explain the unusual strength of M-C and M-H bonds in the 5d series, which provides part of the qualitative explanation of the enhanced chemical reactivity. Another factor related to orbital sizes and electron spin may help s2 ground-state configurations to react more efficiently than we might have expected, particularly on the left-hand side of the 4d and 5s series. When the valence d subshell is less than half occupied (e.g., for Sc, Ti, V, Y, Zr, Nb, Hf, and Ta), an s2 ground state has the proper spin (but the wrong configuration) for bond insertion (Figure 2). This means that the repulsive diabatic surface emanating from the low-spin dx-2s2 asymptote will strongly repel an attractive surface from the low-spin dx-1s excited state (Figure 1). The interaction matrix element is electrostatic and thus may be very strong. This repulsion lowers the adiabatic barrier, as suggested in Figure 1. On the righthand side, s2 ground states (e.g., Co, Ni, and Ir) have high spin, i.e., the wrong spin for bond insertion. The spin-orbit coupling is then the relevant matrix element, and we expect a more weakly avoided crossing, particularly in the 3d and 4d series. Furthermore, on the left-hand side, we would expect the magnitude of the electrostatic matrix elements between lowspin dx-2s2 surfaces and low-spin dx-1s surfaces to increase as the sizes of the valence s and d orbitals become more similar (from 3d to 4d to 5d). This helps explain why the 4d-series atoms Y and Zr react more rapidly with alkenes than the 3dseries atoms Sc, Ti, and V. Indeed, the strongly avoided crossing between low-spin potentials helps explain why such s2 ground states can react at all with alkenes at 300 K. It further explains why the 5d-series atoms Hf and Ta, which have undergone the lanthanide contraction, evidently have very low s2 barriers. The data show that the 6s2 ground states of Hf and Ta are more reactive than the 5s2 ground states of Y and Zr, despite the much larger promotion energies to the chemically binding low-spin dx-1s excited state. The low-spin vs highspin concept also begins to explain why Hf and Ir, both from the 5d series, are comparably reactive in spite of the much larger promotion energy in Hf. The 5d76s2 ground state of Ir is highspin. Additional quantitative theoretical work is needed here.
Gas Phase Kinetics of Transition Metal Atoms A detailed view of how metal atom s and d populations change along adiabatic paths to bond insertion would be particularly instructive. V. Conclusion Like their cationic counterparts, the neutral transition metal atoms from the 5d series are unusually reactive, particularly with alkenes. The ability of ground-state Pt to insert into the C-H bond of methane at 300 K is thus far unique among all of the neutral transition metal atoms studied. The ability of Ta and Hf to react with alkenes at 300 K, in spite of well-isolated 5dx-26s2 ground-state configurations, is also remarkable. The neutral atoms from the 5d series tend to contradict our semiquantitative intuition built from previous experience in the 3d and 4d series. In particular, a 6s2 configuration in the third row seems to cause substantially less electron-electron repulsion with alkenes than s2 configurations in the first and second rows do. The lanthanide contraction of the 6s orbital helps to strengthen chemical bonds, to reduce long-range repulsion, and perhaps also to increase the electrostatic matrix elements between low-spin 5dx-16s and 5dx-26s2 configurations to further lower adiabatic barriers. A more complete understanding of these effects awaits further electronic structure calculations. Acknowledgment. We thank the National Science Foundation (CHE-9303918) and the Donors of the Petroleum Research Fund (Type AC) for generous support of this work. References and Notes (1) Moore, C. E. NBS Circular 467; U.S. Department of Commerce: Washington, DC, 1949; Vol. I-III. For estimates of atomic sizes, see: Fischer, C. F. The Hartree Fock Method for Atoms; Wiley: New York, 1977. (2) For a comprehensive review of M+ chemistry, see: Eller, K.; Schwarz, H. Chem. ReV. 1991, 91, 1121. (3) Roth, L. M.; Freiser, B. S. Mass Spectrom. ReV. 1991, 10, 303. (4) (a) Martinho Simoes, J. A.; Beauchamp, J. L. Chem. ReV. 1990, 90, 629. (b) Armentrout, P. B.; Kickel, B. L. In Organometallic Ion Chemistry, Freiser, B. S., Ed.; Plenum: New York, 1995; pp 1-45. (5) (a) Weisshaar, J. C. In State-Selected and State-to-State IonMolecule Reaction Dynamics; C.-Y. Ng, C.Y., Ed.; Wiley: New York, 1992; Part I. (b) Weisshaar, J. C. Acc. Chem. Res. 1993, 26, 213. (6) (a) Ritter, D.; Weisshaar, J. C. J. Am. Chem. Soc. 1990, 112, 6425. (b) Ritter, D.; Carroll, J. J.; Weisshaar, J. C. J. Phys. Chem. 1992, 96, 10636. (c) Ritter, D. Ph.D. Thesis, Department of Chemistry, University of WisconsinsMadison, 1990. (d) Carroll, J. J.; Weisshaar, J. C. J. Am. Chem.
J. Phys. Chem., Vol. 100, No. 30, 1996 12363 Soc. 1993, 115, 800. (e) Carroll, J. J.; Haug, K. L.; Weisshaar, J. C. J. Am. Chem. Soc. 1993, 115, 6962. (f) Tonkyn, R. Ph.D. Thesis, Department of Chemistry, University of WisconsinsMadison, 1988. (7) (a) Mitchell, S. A. In Gas Phase Metal Reactions; Fontijn, A., Ed.; Elsevier: Amsterdam, 1992. (b) Parnis, J. M.; Mitchell, S. A.; Hackett, P. J. Phys. Chem. 1990, 94, 8152. (c) Mitchell, S. A.; Hackett, P. J. Chem. Phys. 1990, 93, 7822. (d) Brown, C. E.; Mitchell, S. A.; Hackett, P. Chem. Phys. Lett. 1992, 191, 175. (e) Blitz, M. A.; Mitchell, S. A.; Hackett, P. J. Phys. Chem. 1991, 95, 8719. (f) Lian, L.; Mitchell, S. A.; Rayner, D. M. J. Phys. Chem. 1994, 98, 11637. (8) Clemmer, D. E.; Honma, K.; Koyano, I. J. Phys. Chem. 1993, 97, 11480. (9) Campbell, M. L.; McClean, R. E.; Harter, J. S. S. Chem. Phys. Lett. 1995, 235, 497. (10) Wen, Y.; Weisshaar, J. C. Manuscripts in preparation. (11) Carroll, J. J.; Haug, K. L.; Weisshaar, J. C.; Blomberg, M. R. A.; Siegbahn, P. E. M.; Svensson, M. J. Phys. Chem. 1995, 99, 13955. (12) (a) Irikura, K. K.; Beauchamp, J. L. J. Phys. Chem. 1991, 95, 8344. (b) Gord, J. R.; Freiser, B. S.; Buckner, S. W. J. Chem. Phys. 1989, 91, 7530. (13) Carroll, J. J.; Weisshaar, J. C.; Siegbahn, P. E. M.; Wittborn, C. A. M.; Blomberg, M. R. A. J. Phys. Chem. 1995, 99, 14388. (14) Tonkyn, R.; Weisshaar, J. C. J. Phys. Chem., 1986, 90, 2305. Tonkyn, R.; Ronan, M.; Weisshaar, J. C. J. Phys. Chem. 1988, 92, 92. Tonkyn, R.; Weisshaar, J. C. J. Am. Chem. Soc. 1986, 108, 7128. (15) Ferguson, E. E.; Fehsenfeld, F. C.; Schmeltekopf, A. L. AdV. At. Mol. Phys. 1969, 5, 1. (16) Present, R. D. Kinetic Theory of Gases; McGraw-Hill: New York, 1958. (17) Fischer, C. F. The Hartree-Fock Method for Atoms; Wiley: New York, 1977. (18) Carroll, J. J. Ph.D. Thesis, University of WisconsinsMadison, 1995. (19) Lian, L.; Mitchell, S. A.; Rayner, D. M. J. Phys. Chem. 1994, 98, 11637. (20) Parnis, J. M.; Lafleur, R. D.; Rayner, D. M. J. Phys. Chem. 1995, 99, 673. (21) Senba, K.; Matsui, R.; Honma, K. J. Phys. Chem. 1995, 99, 13992. (22) (a) Blomberg, M. R. A.; Siegbahn, P. E. M.; Svensson, M. J. Phys. Chem. 1992, 96, 9794. (b) PCI-80 method: Siegbahn, P. E. M.; Blomberg, M. R. A.; Svensson, M. Chem. Phys. Lett. 1994, 223, 35. (c) Widmark, P. O.; Roos, B. O.; Siegbahn, P. E. M. J. Phys. Chem. 1985, 89, 2180. (d) Blomberg, M. R. A.; Siegbahn, P. E. M.; Svensson, M. J. Phys. Chem. 1991, 95, 4313. (23) (a) Sodupe, M.; Bauschlicher, C. W., Jr.; Langhoff, S. R.; Partridge, H. J. Phys. Chem. 1992, 96, 2118. (b) Langhoff, S. R.; Bauschlicher, C. W., Jr. Annu. ReV. Phys. Chem. 1988, 39, 181. (24) (a) Low, J. J.; Goddard, W. A., III. J. Am. Chem. Soc. 1984, 106, 8321. (b) Low, J. J.; Goddard, W. A., III. Organometallics 1986, 5, 609. (c) Carter, E. A.; Goddard, W. A., III. J. Phys. Chem. 1988, 92, 5679. (d) Ohanessian, G.; Goddard, W. A., III. Acc. Chem. Res. 1990, 23, 386. (e) Ohanessian, G.; Brusich, M. J.; Goddard, W. A., III. J. Am. Chem. Soc. 1990, 112, 7179.
JP960408B
