J. Phys. Chem. 1995, 99, 16571-16575
Threshold Photoionization of Al- (N2)x and Al- (COZ)~ Complexes: Evidence for Solvation-Induced Reactions L. R. Brock and M. A. Duncan* Department of Chemistry, University of Georgia, Athens, Georgia 30602 Received: July 12, 1995; In Final Form: September 6, [email protected]
Aluminum atom van der Waals complexes of the form Al-(N2), and Al-(C02), are prepared in a molecular beam environment using a laser vaporization pulsed nozzle cluster source. These species are studied with threshold laser photoionization in a time-of-flight mass spectrometer. Ionization thresholds as a function of cluster size reveal cluster bonding energetics and evidence for intracomplex reactions. The ionization potential of A1-N2 (5.805 eV) is combined with the dissociation energy of Al+-N2 to yield the bond energy (466 cm-I) of the AI-N2 van der Waals complex. In the Al-(COz), system, the intracluster metal oxidation reaction proceeds spontaneously only after “solvation” by at least five CO:! molecules. The ionization threshold red shift indicates that Do(Al+-C02) 2 2453 cm-I.
Introduction There has been considerable recent interest in metal atom cluster comp1exes.’-l6 Metal-rare gas diatomics and metalsmall molecule complexes have been investigated with electronic spectroscopy.‘-9 Photoionization mass spectroscopy has been applied to larger complexes seeded with metal atoms.’ Aluminum and the aluminum group metals have been the focus of many of these studies. Aluminum is itself of special interest because it is convenient for study in laser vaporization cluster experiments and it is small enough to be tractable for theoretical investigations. The diatomics A1-Ar, A1-Kr, and Al-Xe have been studied with laser-induced fluorescence and resonant twophoton photoi~nization.~-~ A1-Ar, clusters have been studied with threshold photoionization,” and Al-Ar has been the subject of high-resolution photoelectron spectroscopy.I0 In the present report, we extend the study of aluminum complexes to Al-(N2), and Al-(C02), systems. The nitrogen complexes are expected to have purely van der Waals binding.’’ The interaction of aluminum with CO2 has been studied in rare gas matrices,I7in the gas and by theory.22 The insertion reaction to form A10 CO is believed to be slightly end other mi^,^^ but relatively strong complexation is observed in low-temperature matrices.I7 Thus, this system is interesting for the possibility of intracluster reactions. We use threshold photoionization as a function of cluster size to study both of these cluster systems. These studies determine ionization potentials, which allow derivation of dissociation energies, and they produce evidence for novel “solvation-induced” intracomplex reactions. Threshold photoionization has been applied previously to related systems such as Al-Ar,,” Na-(H20),,I3 Na-(NH3),,I3 Cs-(H20),,I4 and Cs-(NH3),.I4 In all of these systems, threshold ionization is possible because the corresponding metal I atoms have relatively low ionization potentials (P’s). Ionization potentials are even lower in the weakly bound complexes than they are in the isolated metal atom because of the stronger binding in the cation complexes compared to the corresponding neutral van der Waals complexes. In “inert” systems, such as A1-Ar,, the trend in ionization potential with cluster size is smooth, with a gradual decrease reflecting the incrementally 1513-14
* To whom correspondence should be addressed. @
Abstract published in Aduance ACS Ahsrrucrs, November 1, 1995.
increasing solvation.’ In potentially reactive systems, such as the sodium or cesium complexes with water or ammonia, the decrease in ionization potential with cluster size is steeper than predicted by simple models, providing evidence for partial charge separation (M+, e-) within the cluster, analogous to solvation in the bulk media.I3-l5 In the present systems, the nitrogen complexes behave in a manner similar to A1-Arr. The red-shifted ionization potentials fall smoothly with cluster size. In the Al-(C02), complexes, however, the ionization potentials are extremely irregular with cluster size, providing evidence for an intracluster oxidation reaction.
Experimental Section The complexes for this study are produced by laser vaporization using a XeCl excimer laser (Lumonics TE-860) at 308 nm on a rotating aluminum rod sample. The pulzed nozzle source uses a modified Newport BV-100 beam valve. A “cutaway” type sample holder produces an expansion beginning at the sample, which minimizes metal-metal recombination and maximizes cooling and van der Waals c ~ n d e n s a t i o n . ~Pure ~ expansions of nitrogen or carbon dioxide are used for these experiments. Photoionization is accomplished with a Nd:YAG pumped dye laser system (Spectra-Physics PDL-2 or Lumonics HD-500 system) frequency doubled into the ultraviolet with a BBO crystal or with the fixed frequency output of an ArF excimer laser at 193 nm (Lumonics TE-860). The various ionized complexes are detected using a homemade time-of-flight mass spectrometer. The molecular beam machine and its various components have been described previously.2s
Results and Discussion Figure 1 shows an example of photoionization of Al-(N2), complexes at 5.36 eV (231 nm). At this wavelength, complexes extending out to x 2 60 or so are detected, with peaks or discontinuities in the intensity distribution at x = 12, 32, and 54. These “magic numbers” were noted in the previous mass distributions measured by Whetten and co-workers’ for ionization at 6.42 eV (193 nm; ArF excimer laser), and they were interpreted to arise from icosahedral close packing of Nz molecules about the AI atom. These magic numbers have been observed previously in metal-seeded rare gas clusters such as Al+-Ar,‘1.12and Mg+-Ar,.I6 At the 6.42 eV ionization energy used by Whetten for Al-(N2),, which we have reproduced, all 0 1995 American Chemical Society
Brock and Duncan
16572 J. Phys. Chem., Vol. 99, No. 45, 1995
I.P. of AlNi Field = 53 V/cm ,
Al(N2)x Photoionization at 5.36 eV
Energy (cm-') 750 1500 Mass (amu) Figure 1. Mass spectrum of complexes produced by photoionization of AI-(N2), at 5.36 eV (231 nm). Clusters out to x t 60 are detected, with truncations in the intensity distribution at x = 12, 32, and 54 attributed to close-packing shell structures. Complexes for x < 6 are not ionized at this wavelength. A small amount of Nzi is detected, which is believed to result from electron impact ionization by photoelectrons generated when the UV laser hits metal surfaces in the ion source. which are accelerated in the source field.
Al(C02)x Photoionization XI
x=l 2 1 1 3
Mass (amu) Al+(COz)x Ion Cluster Distribution AI XI
Mass (amu) Figure 2. Upper trace shows the distribution of complexes produced by photoionization of AI-(C02), at 5.80 eV (214 nm). Complexes out to x = 12 are detected, but there is no ionization of complexes for x = 6, 7. The lower trace show the distribution of cation complexes produced directly from the source without any photoionization. clusters are detected down to the smallest possible sizes (e.g., x = 1). However, at the lower energy shown in Figure 1, the smaller clusters in the range of x = 1-5 are not detected.
Tunable photoionization studies, as described below, confirm that these smaller complexes have higher ionization potentials and are therefore not ionized at the 5.36 eV energy used for Figure 1. At other ionization energies in this vicinity, various similar spectra are produced with missing peaks corresponding to those species not ionized by the energy of light employed. As discussed below, larger complexes are generally expected to have lower ionization potentials, and so this behavior is understandable. A totally different behavior is observed for ionization of Al(C02)r complexes. As shown in Figure 2 (upper trace), ionization at the 5.80 eV energy (214 nm) produces signal for x = 1-5 and 8-12 complexes but no signal for x = 6, 7 complexes. Indeed, photoionization at the higher energy of 6.42 eV (193 nm) produces essentially the same mass distribution. The complexes with six or seven C02 molecules are not detected. This bimodal distribution can be shifted in relative
45500 45700 45900 46100 46300 46500
Energy (cm-I) Figure 3. Ionization thresholds of AI-N2 and Al-CO2 measured with a tunable UV laser. The ionization onset in the nitrogen complex rises sharply up to full intensity within about I O cm-', while the onset in the COz complex rises gradually over about 500 cm-I.
intensity with the instrumental focusing (Le., the mass spectrometer deflection plates), but no instrumental settings make it possible to detect the x = 6, 7 complexes by low-power photoionization. However, if we use high-power photoionization (> 20 mJ/cm2 pulse) from the ArF excimer laser, a small amount of signal in the x = 6 and 7 mass channels is observed, presumably from multiphoton ionization of the neutrals or from fragmentation of larger complexes after ionization. (We cannot distinguish between these two processes.) The observation that these complexes are "missing" is surprising, since the effects of simple solvation should be to lower the IP as a function of complex size. There are several possible explanations for this unusual intensity distribution. For example, the missing mass peaks at x = 6 and 7 could arise because these complexes are not present in the molecular beam. However, it is difficult to imagine how these complexes could be missing while larger complexes are formed. Another possibility is that the cations of these complexes are unstable. It is well-known in cluster studies that the intensities of mass peaks from neutral photoionization may depend on the properties (concentration, stability) of the neutral as well as the cation formed after ionization. To investigate this as a possibility, we have produced the corresponding cation complexes directly from the laser vaporization plasma without any photoionization. For this experiment, like the neutral experiment, we use the "cutaway" type nozzle geometry to optimize production of metal atom complexes. The resulting cluster distribution is shown in the lower trace of Figure 2. The distribution is smooth, with major peaks assigned to Al+-(C02), masses and smaller mass peaks assigned to A10+(COZ)~ and Al2+. However, there is no evidence of irregularities in the x = 6, 7 vicinity. Therefore, we conclude that there is no unusual instability in the A1+-(C02)x species which could explain our photoionization results. The final possibility is that the x = 6 and 7 complexes do exist in the molecular beam but that their ionization efficiency is low, perhaps because of some chemistry within these clusters which produces species with unusually high ionization potentials. We investigate this possibility further below via the size dependence of ionization potentials in these complexes. Figure 3 shows the results of tunable ultraviolet laser ionization potentials measured for Al-N2 and Al-CO2. As shown in the figure, these ionization potentials are qualitatively
Threshold Photoionization of Al-(N2), and Al-(C02), different in appearance. The threshold of AI-N:! is sharp, rising up to the full intensity obtained within about 10 cm-I. The AI-C02 threshold, however, is much broader, rising up to its full intensity over about 500 cm-I. Sharp ionization thresholds like that measured for A1-N2 indicate that the ion and neutral have similar structures and that the vertical ionization potential measured is also the adiabatic ionization potential. We expect that the Franck-Condon factors for this ionization threshold will be somewhat affected because the bond length in the ion (charge-quadrupole interaction) should be shorter than that in the neutral (van der Waals interaction). Apparently, this effect is not too great for Al-N2, resulting in the sharp ionization onset. A broad threshold like that observed for Al-C02 implies a more significant change in geometry than the neutral cation bond contraction in the Al-N2 case. This more complex onset indicates that the ionization threshold measured represents the vertical IP, which is actually an upper limit on the adiabatic ionization potential. It also suggests that the neutral and cation have significantly different geometries. The ionization thresholds shown in Figure 3 are measured at specific values of the ion source acceleration field because ionization potentials are expected to vary with the electric field. Field ionization is expected to red-shift these ionization potentials by approximately 6E’l2,where E is the electric field present in the ionization region in VIcm’O and the shift is in cm-I. We have therefore measured these thresholds as a function of the electric field. The Al-N2 threshold remains sharp under all fields studied, with the threshold energy scaling as 6E’I2. We extrapolate a plot‘ of threshold versus E’” to a zero-field ionization potential of 46 821 cm-’ (5.805 eV), which is 1458 cm-’ to the red from the atomic aluminum IP. For the Al-CO2 ionization potential, the gradually rising threshold is very weak near its onset, and no electric field shift can be detected within our signal-to-noise level. Therefore, we make no electric field correction to this threshold position. We chose the threshold as the energy position where signal is first detectable. We do not use any more sophisticated treatment of the threshold shape because it is clear that the threshold measured is vertical and that the adiabatic value lies much lower in energy. The threshold energy thus determined is 45 830 k 50 cm-I. This is 2450 cm-’ to the red of the aluminum atom IP. The red-shifted ionization thresholds measured for AI-N:! and Al-CO2 indicate, as expected, that the cations (Al+-N2, Alf-C02) are more strongly bound than the corresponding neutral complexes. If there is previous thermochemical data for a metal complex cation such as these, this cation data can be used in a simple cycle to derive neutral energetics. For example, the red-shifted complex IP can be combined with the atomic IP and the cation bond energy to obtain the neutral complex bond energy:
J. Phys. Chem., Vol. 99, No. 45, 1995 16573 Photoionization Thresholds of AI(N2)x Clusters
of AI(C0z)x Clusters
9 5.9 0
1 0 1 2
X (No. of C02 Molecules) Figure 4. Photoionization thresholds versus cluster size for the nitrogen and C02 complexes with aluminum. The nitrogen system IP‘s fall gradually and smoothly over the size range indicated. However, the COz complex IP’s fall sharply at x = 1, are constant for x = 2 , 3, rise gradually for x = 4, 5 , rise sharply for x = 6, 7, and are then about constant for x = 8-12. For x = 6, 7 , the I P S are 26.42 eV.
from gas phase reaction systems that the AI-C02 complex is bound by about 9 kcaVmo1 (3100 cm-I), which is roughly consistent with this estimate. Even if we do not use the corresponding Mg+-C02 binding energy to estimate the neutral Al-CO2 binding, the equation above and the strong red-shifted IP observed for Al-CO2 can be used to set energetic limits on the Al+-CO:! ion. The red-shifted IP results from the difference in the neutral and cation binding energies. Therefore, the Al+C02 binding energy must be at least as great as the IP red shift, which results in Do(Al+-C02) 1 2450 cm-I. Thus, the electrostatic bond energy in the Al+-C02 ion is greater than that in the Al+-N2 ion. This is expected because the quadrupole moment of C02 (-14.9 x C m2) is larger than that of N2 (-4.90 x C m2).28 The ionization thresholds observed here can also produce some qualitative insight into the structures of these complexes. A sharp ionization onset, such as that seen for Al+-N2, suggests via Franck-Condon considerations that the ion and the neutral have similar structures. The structure of neither Al-N2 nor Al+-N2 has been measured. However, the quadrupole moment of nitrogen is negative, which results in most favorable chargequadrupole electrostatic interactions for a linear M+-N2 configuration. For example, both theory and experiment agree that Mg+-N2 is linear,29and therefore Al+-N:! is most likely to be linear as well. If this is true, then neutral Al-N2 is also linear. On the other hand, the broad, gradually rising threshold observed for AI-CO:! implies that the ion and the neutral do not have IP(A1-L) DO(Al’-L) - IP(A1) = Do(Al-L) the same structure. Again, because of the negative quadrupole moment of C02, metal ion-CO2 complexes have most favorable The cation dissociation energy for Al+-N:! has been measured charge-quadrupole electrostatic interaction in the linear conby Bouchard and McMahon to be 5.5 & 0.5 kcaVmol(l924 & figuration. Consistent with this, Mgf-C02 is linear.27 There175 cm-’).26 Therefore, the binding energy of the AI-N2 van fore, Alf-C02 is also most likely to be linear, and neutral Alder Waals complex is DO = 466 & 175 cm-I. This is much CO2 is then most likely to be nonlinear. In fact, a nonlinear larger than the bond energy of A1-Ar (122 ~ m - ’ ) .The ~ only Al-02C structure with C2u symmetry and bent O=C=O other metal-nitrogen bond energy which has been measured (oxygens closer to metal) has been suggested previously by is that for In-Nz, which is much larger at 1519 ~ m - ’ . ~ matrix isolation infrared spectroscopy on this species.I7 The Al+-C02 binding energy has not been measured. However, the binding energy for Mg+-C02 is 5150 ~ m - ’If. ~ ~ Figure 4 shows the results of threshold ionization measurements for larger Al-(N2), and Al-(COr), complexes. All the we assume that the Al+ and Mg+ electrostatic interactions are similar, we can estimate that the Al-CO2 neutral binding energy Al-(N2), thresholds up to x = 14 are reasonably sharp, while is about 2700 cm-’. Hackett and co-workersZ1have estimated all the Al-(C02), thresholds up to x = 12 are broad like the
16574 J. Phys. Chem., Vol. 99, No. 45, 1995
Al-CO2 threshold. Therefore, for larger Al-(N2), clusters, we measure the ionization potentials at one value of the electric field and then shift the quoted value by 6E'12to correct for the effects of field ionization. For larger Al-CO2 complexes, we report the measured values with no electric field correction. As shown, the Al-(N2), thresholds decrease smoothly as a function of cluster size. There is no noticeable discontinuity in the IP value for Al-(N2)12, which is indicated to be a shell closing in Figure 1 . We are limited by signal-to-noise considerations in the present setup to measuring thresholds only up to x = 14. However, at x = 12-14, the threshold values have converged to a nearly constant value of about 5.2 eV. By comparison with this trend, the ionization potentials for AlAr, also exhibited a smooth decrease with size, with a temporary leveling near x = 12, but there was additional structure in the IP trend at larger cluster sizes." The IP's for A1-Ar, in the vicinity of x = 12 were about 5.65 eV. The increased red shifts observed here reflect the stronger electrostatic interactions between Al+ and N2, which is charge-quadrupole in nature, as opposed to the charge-induced dipole interaction most prominent for Al+-Ar,. The smooth decrease in IP with cluster size is expected from the classical model of ionization from a spherical d r ~ p l e t . ' ~ , In ' ~ ,this ~ ~ picture, the IP reduction is proportional to the radius of the droplet, which scales approximately as (n 1)-'13, where n is the number of solvent molecules. Therefore, a plot of IP versus (n l ) - I I 3 should be linear if this simple picture applies. Previous studies of the IP's of solvated metal atom clusters have noted this ~ c a l i n g , ' ~ . ' ~ and a plot of the Al-(N2), IP's is also roughly linear. Therefore, the bonding in AI-(NZ),~ complexes appears to be purely electrostatic in nature. The most interesting result from this threshold study is the size-dependent trend measured for Al-(C02), complexes. As shown in Figure 4, the IP drops sharply at x = 1, is about constant for x = 1-3, and then increases slightly for x = 4 and 5. Complexes for x = 6 and 7 cannot be ionized, as noted above. Complexes for x = 8-12 are about constant at a value nearly the same as that seen for x = 1-3. All of these thresholds are broad, and none are believed to reflect the adiabatic ionization energies. A plot of IP versus (n -t- l)-]I3 as described above is grossly nonlinear. This trend in thresholds suggests that some unusual behavior is present in these complexes. If the interactions in these complexes are purely electrostatic, the ionization potentials observed should be less than or equal to the IP of the aluminum atom. Apparently, the IP's for complexes where x = 6 , 7 are greater than 6.42 eV, which is considerably greater than the aluminum atom value (5.986 eV). For these complexes to have such high ionization potentials, the chemical composition within the complexes must be significantly different from solvated aluminum atoms; Le., there must have been some chemistry inside these complexes. The most likely reaction which could take place within these complexes is the oxidation process,
As noted above, this reaction is slightly endothermic for isolated A1 CO2. The main uncertainty in calculating the thermochemistry is the dissociation energy of N O . The recommended value of Do(Al0) = 121.2 f 2.2 kcal/m01,~~ when combined with known thermochemistry for Al, C02, and CO, yields the result that the reaction is endothermic by 6.0 & 2.2 kcal/mol. Fontijn observed the gas phase reaction as a function of temperature in the range 300- 1880 K and derived an activation energy of about 2.5 kcaVm01.'~Thus, while there may be slight uncertainty in the exact energetics, there is agreement that the
Brock and Duncan
reaction is slightly endothermic for the gas phase atom molecule system, and therefore the same is true for the monomer complex. While the reaction is endothermic for the monomer, it may be exothermic in larger complexes, depending on the details of the solvation energetics of the reactants versus the products. If solvation affects reactants and products equally, then the total energy of the system will be lowered sequentially by the addition of more C02 molecules, but the difference between the reactant and product species will be constant and the reaction will remain endothermic for all complex sizes. However, if the product species are stabilized more by solvation than the reactant species, the reaction will become exothermic at some critical complex size. It is then important to notice that the monomer endothermicity of about 6 kcal/mol (2100 cm-l) is energetically smaller than the estimated binding energy of the Al-CO2 complex (2700-3100 cm-' ; see above discussion). Therefore, only a small differential in solvation energies for the product species (A10 CO) versus the reactants (A1 C02) would be enough to make the overall process exothermic. Since entropy effects are expected to be small at the low temperatures of these complexes, an exothermic reaction is therefore likely to be spontaneous, and the result is a "solvation-induced" intracomplex reaction. Preferential solvation of the product species in this reaction is completely reasonable because both diatomic products (A10 and CO) have significant dipole moments and are likely to have more favorable electrostatic interactions with the C02 solvent species than the reactants. Apparently, after the addition of six or more C02 molecules, the solvation of the products exceeds that of the reactants and the intracomplex reaction proceeds spontaneously. Solvation-induced reactions have been noted recently in the study of metal ion cluster complexes, such as Mg+-(H20),,31*32Mg+-(CH30H),,33 and Mg+-(NH3),.33 In all of these systems, the monomer reaction is significantly endothermic (20-40 kcal/mol), but the reactions are observed to occur spontaneously in ground state complexes when x 2 5. The same kind of differential solvation has been invoked to explain these reactions. However, the present system would be the first example, to our knowledge, of a solvationinduced reaction for a neutral system. The solvation-induced reaction described here is completely consistent with the ionization potential data for this system. According to the size-dependent IP's, the species Al-(C02), are missing because of high ionization potentials at the sizes x = 6,7. The ionization potential of AlO is 9.53 eV.34 Therefore, small complexes which contain A10 rather than A1 as the chromophore to be ionized are expected to have significantly higher ionization potentials, and this is the best explanation for the missing cluster peaks at x = 6 and 7 . Cluster peaks for x > 7 are detected, but the ionization potential data and masses are not sufficient to identify the structure of these complexes. (Al-(C02), and A~O-(CO~),-ICOhave the same masses.) It could be possible that the solvation-induced reaction is size selective and only occurs for x = 6, 7 . The larger complexes would then be ionized at low energy because they contain solvated aluminum atoms, like the smaller complexes. On the other hand, the larger complexes could also have reacted and still be ionized at low energy. The solvation of AlO-containing complexes will also lower their ionization potentials, and eventually this effect will lead to ionization of these species at low energies. However, if larger complexes do indeed represent reactedsolvated AlO-(CO&l CO species, with ionization potentials in the 5 eV range, the solvation has lowered the IP from about 9.5 eV to about 5.8 eV. A solvent shift of about 3.7 eV from the addition of only one or two C02 molecules
Threshold Photoionization of Al-(NZ), and Al-(COZ),
J. Phys. Chem., Vol. 99, No. 45, 1995 16575
seems unlikely. It may well be then that the reactivity is size specific for x = 6,7. A final possibility, since the larger clusters are always weak in intensity, is that we only ionize the unreacted fraction of the beam. Larger reacted complexes, which would presumably have high IP’s, would not be detected. To explore these various possibilities more fully, it would be interesting to probe the structure of the larger clusters. However, probes such as collision-induced dissociation or photodissociation may themself cause rearrangements in the clusters which confuse the issue. It may be possible in the future to conduct infrared spectroscopy or high-resolution photoelectron spectroscopy on these clusters in a suitable nonintrusive way. It is interesting to mention is that the chemistry suggested here is only observed for the neutral van der Waals complexes. It is not found for the corresponding cations. This is presumably because the thermochemistry is different, and the electrostatic interactions are vastly different, in the cation clusters. The neutral aluminum-C02 system may be fortuitous in its combination of thermochemistry and solvation interactions. The reaction endothermicity is so marginal that even weak van der Waals interactions can cause a significant difference in the solvation of reactants (A1 C02) versus products (A10 CO).
Conclusions Weakly bound complexes of the aluminum atom with nitrogen and carbon dioxide molecules are produced in a molecular beam environment and detected with threshold laser photoionization. The mass distributions for Al-(N2), are structured, with magic numbers corresponding to icosahedral close packing structures. The mass distributions for Al-(C02), complexes are bimodal, with missing peaks at x = 6, 7. Sizedependent ionization potentials derive thermochemical data for the small complexes, which suggest that both neutral and ionic aluminum complexes are more strongly bound with COZ than they are with N2. Ionization potentials also suggest that there is unusual solvation-induced chemistry in the Al-(COz), complexes. The simplest explanation consistent with all the data is that a solvation-induced metal oxidation reaction occurs for Al-(C02), complexes larger than x = 5 . This is the first example of such a solvation effect in neutral metal seeded complexes. This system is light enough that ab initio theory could address the potential energy surfaces and many-bodied effects causing this novel chemistry.
Acknowledgment. This research was supported by the U.S. Air Force Office of Scientific Research in the High Energy Density Materials (HEDM) program (Grant F49620-94-1-0063). Additional support was received from the National Science Foundation (Grant CHE-9307907). References and Notes (1) Breckenridge, W. H.; Jouvet, C.; Soep, B. In Advances in Metal and Semiconductor Clusters; Duncan, M. A., Ed.; JAI Press: Greenwich, CT, 1995; Vol. 111.
(2) Cardiner, J. M.; Lester, M. I. Chem. Phys. Lett. 1987, 137, 301. (3) Callender, C. L.; Mitchell, S. A.; Hackett, P. A. J . Chem. Phys. 1989, 90, 5252. (4) Heidecke, S. A.; Fu, Z.; Colt, J. R.; Morse, M. D. J . Chem. Phys. 1992, 97, 1692. ( 5 ) Fu, Z.; Massick, S.; Kaup, J. G.; Benoist d’Azy, 0.;Breckenridge, W. H. J . Chem. Phys. 1992, 97, 1683. (6) Hwang, E.; Huang, Y.-L.; Dagdigian, P. J.; Alexander, M. H. J . Chem. Phys. 1993, 98, 8484. (7) (a) Stangassinger, A,; Scheuchenpflug, J.; Prinz, T.; Bondybey, V. E. Chem. Phys. Lerf. 1993,209,372. (b) Stangassinger, A,; Scheuchenpflug, J.; Prinz, T.; Bondybey, V. E. Chem. Phys. 1993, 178, 533. (8) Callender, C. L.; Mitchell, S. A,; Hackett, P. A. J . Chem. Phys. 1989, 90, 2535. (b) Hackett, P. A.; Balfour, W. J.; James, A. M.; Fawzy, W. M.; Shetty, B. J.; Simard, B. J . Chem. Phys. 1993, 99, 4300. (9) Brock, L. R.; Duncan, M. A. J . Chem. Phys. 1995, 102, 9498. (IO) Willey, K. F.; Yeh, C. S.; Duncan, M. A. Chem. Phys. Lett. 1993, 211, 156. (11) Schriver, K. E.; Hahn, M. Y.; Persson, J. L.; LaVilla, M. E.; Whetten, R. L. J . Phys. Chem. 1989, 93, 2869. (12) Gantefor, G.; Siekmann, H. R.; Lutz, H. 0.;Meiwes-Broer, K. H. Chem. Phys. Lett. 1990, 165, 293. (13) (a) Hertel, I. V.; Huglin, C.; Nitsch, C.; Schulz, C. P. Phys. Rev. Len. 1991.67, 1767. (b) Nitsch, C.; Schulz, C. P.; Gerber, A,; ZimmermanEdling, W.; Hertel, I. V. Z. Phys. D: At. Mol. Clusters 1992, 22, 651. (14) Misaizu, F.; Tsukamoto, K.; Sanekata, M.; Fuke, K. Chem. Phys. Lett. 1992, 188, 241. (15) Hashimoto, K.; He, S.; Morokuma, K. Chem. Phys. Lett. 1993,206, 297. (16) Velegrakis, M.; Luder, C. Chem. Phys. Lett. 1994, 223, 139. (17) Le Quere, A. M.; Xu, C.; Manceron, L. J . Phys. Chem. 1991, 95, 303 1. (18) Klabunde, K. J. Chemistry of Free Atoms and Particles; Academic Press: New York, 1980. (19) Fontijn, A.; Felder, W. J . Phys. Chem. 1977, 67, 1561. (20) Costes, M.; Naulin, C.; Dorthe, G.; Vaucamps, C.; Nouchi, G. Faraday Discuss. Chem. SOC.1987, 84, 75. (21) Parnis, J. M.; Mitchell, S. A.; Hackett, P. A. Chem. Phys. Lett. 1988, 151, 485. (22) Selmani, A.; Ouhal, A. Chem. Phys. Lett. 1992, 191, 213. (23) J . Chem. Phys. Ref. Data 1983, 12, 967. (24) Pilgrim, J. S.; Yeh, C. S.; Berry, K. R.; Duncan, M. A. J. Chem. Phys. 1994, 100, 7945. (25) LaiHing, K.; Wheeler, R. G.; Wilson, W. L.; Duncan, M. A. J . Chem. Phys. 1987, 87, 3401. (26) Bouchard, F.; McMahon, T. Private communication. (27) Yeh, C. S.; Willey, K. F.; Pilgrim, J. S.; Duncan, M. A. J. Chem. Phys. 1993, 98, 1867. (28) Rigby, M.; Smith, E. B.; Wakeham, W. A,; Maitland, G. C. The Forces Bemeen Molecules; Clarendon Press: Oxford, 1986. (29) Robbins, D. L.; Brock, L. R.; Pilgrim, J. S.; Duncan, M. A. J . Chem. Phys. 1995, 102, 1481. (30) Wood, D. M. Phys. Rev. Lett. 1981, 46, 749. (31) (a) Misaizu, F.; Sanekata, M.; Fuke, K.; Iwata, S. J . Chem. Phys. 1994, 100, 1161. (b) Sanakata, M.; Misaizu, F.; Fuke, K.; Iwata, S.; Hashimoto, K. J . Am. Chem. Soc. 1995, 117, 747. (c) Watanabe, H.; Iwata, S.; Hashimoto, K.; Misaizu, F.; Fuke, K. J . Am. Chem. SOC.1995, 117, 755. (32) Harms, A. C.; Khanna, S. N.; Chen, B.; Castleman, A. W. J . Chem. Phys. 1994, 100, 3540. (33) Robbins, D. L.; Scurlock, C. T.; Duncan, M. A. J. Phys. Chem., to be submitted. (34) Huber, K. P.; Herzberg, G. Molecular Spectra and Molecular Structure IV. Constants of Diatomic Molecules; Van Nostrand Reinhold: New York, 1979.