15602
J. Phys. Chem. 1995,99, 15602-15607
Factors Affecting 0 Bond Activation in Simple Systems: Measurement of Experimental Binding Energies of Fe+(H2)1-6 Clusters John E. Bushnell, Paul R. Kemper, and Michael T. Bowers” Department of Chemistry, University of Calijomia, Santa Barbara, Calijomia 93106 Received: April 20, 1995; In Final Fom: July 25, 1995@
Dissociation energies and entropies for H2 loss from Fe+(H2), clusters ( n = 1-6) have been determined via temperature-dependent equilibrium measurements. DO= -AHO, = 10.8 f 0.6, 15.7 f 0.7, 7.5 f 0.4, 8.6 f 0.4, 2.2 f 0.3, and 2.3 f 0.3 kcaYmol for n = 1-6, respectively. The trends in these binding energies are similar to those observed for Co+(H2), and V+(H2), with some important exceptions. The small value of the first binding energy relative to the second provides strong evidence for curve crossing from ground state Fe+ (6D, 3d64s’) to the Fe+ (4F,3d7) surface as the first H2 ligand approaches. Comparison with recent theoretical calculations has led us to discard steric hindrance as a primary factor in explaining the decrease in binding energy as ligation number increases in favor of metal specific covalent interactions. The most important of these are metal s-d hybridization, CI donation from H2 to M+, and dz back donation from M+ to the CI* orbital on H2.
Introduction Activation of u bonds by transition metal centers is a very active area of research. Of particular interest is the activation of the simplest bonds, those in H2 and CH+ Because of their simplicity these systems can be studied in depth using ab initio electronic structure theory and the synergistic interplay of theory and experiment can hopefully reveal the dominant factors governing this important activation process. Historically, the H-H bond in H2 has readily yielded to activation by most transition metals in solution while the C-H bond in methane has not.’ On one level this is curious since the H-H bond and the CH3-H bond have identical bond dissociation energies (BDE).2 While there are undoubtedly a number of reasons for this difference in ease of activation, the two most important points are probably the fact H-H has a low lying u* state of convenient symmetry (while CH3-H does not) and the fact that the insertion transition state for CH3-H will be substantially higher in energy than that for H-H due to the strong directional character of the sp3 hybridized orbital on ~ a r b o n . ~One consequence of the relative inertness of the C-H bond in alkanes is that even if a reagent is found that can activate the bond, it is extremely difficult to turn off the reaction before unwanted sequential reaction products are generated. Hence, the “Holy Grail” of selective C-H bond activation in alkanes, and especially methane, is still far off.4 One of the most promising areas for exploring the factors governing u bond activation by transition metal centers is reaction of bare metal ions in the gas phase. These systems offer the opportunity to systematically study the reactivity of the metal center itself and see how this reactivity changes across and down the periodic table. Of paramount importance is the fact that the charged nature of the metal ion allows the powerful methods of mass spectrometry to be brought to bear on the problem. Hence, ions can be readily formed, reacted, and detected and consequently both rate constants and equilibrium constants measured. The fact that the metal ion is charged also means long range polarization forces come into play, reducing the energy of initial insertion transition states and allowing gas phase studies of reactivity for metal ions in instances where metal neutral atoms are completely ~nreactive.~ @
Abstract published in Advance ACS Abstracts, October 1, 1995.
0022-365419512099-15602$09.0010
In our laboratory we have pursued activation of both R-H bonds in alkanes,6,’ including methane,*and the H-H bond9-” by transition metal ion centers. Because of its simplicity, the H2 ligandl2-l3has received more attention by theorists than CHq or larger alkane^.'^,'^ Our emphasis in the H-H activation process has centered on measuring very accurate enthalpies and entropies for the association reactions represented by eq 1 for n = 1-7 for first row transition metal ions. So far we have M+(H*),-,
+ H2
#e
M+(H2),
(1)
published work on Sc+, V+, and Co+ with each system portraying quite different characteristics. Of these three metals only Sc+ inserted into the H-H bond (for n = 1) and then only via a cluster-assisted mechanism.’’ Theory had predi~ted’~ a large barrier to insertion of Sc+ into H2 and hence the need for “clustering” to assist the process. In V+, the fist four H2 ligands have essentially equal bond dissociation energies, the BDE dropped substantially for the fifth ligand, and then substantially increased for the sixth H2 ligand. We interpreted this unusual behavior to indicate that a major geometry change occurs between addition of n = 4 and n = 5 (induced to minimize repulsion between the fifth H2 ligand and the singly occupied cg u orbital) followed by a spin change on Vf (from quintet to triplet) in going from n = 5 to n = 6. This interpretation has subsequently been confirmed by theory.I6 Finally, the first system studied, Co+(H*),, was seemingly the simplest with observation of rather strong painvise addition of H2 ligands: -AHo, % 18, 10, and 4 kcal/mol for the first, second, and third pairs, respectively, which we suggested was due to increased steric hindrance with increased ligation (we will revisit this point later in this paper.) From these initial studies one thing is clear: the electronic configuration on the metal ion strongly influences the details of the bonding with the H2 ligands. While this may not be surprising to many, it was contrary to the notion that the bonding between metal ions and simple noninserted ligands was primarily electrostatic in nature. To more directly test this lafter notion we measured9 AHo, for Na+ and K+ for n = 1 and 2. These binding energies were very small (-2.3 and -1.4 kcal/mol for Na+ and K+, respectively.) Consequently, even when ion size is considered, the much larger binding energies to transition 0 1995 American Chemical Society
Factors Affecting
0 Bond
J. Phys. Chem., Vol. 99, No. 42, 1995 15603
Activation
metal ions clearly indicate that “pure” electrostatic effects play a minor role in the bonding to transition metal centers. In all systems reported to date the 4s orbital has been either vacant (V+ and Co+) or hybridized (Sc+). Numerous studies have shown that population of the large 4s orbital has a destabilizing effect on the M+-ligand bond strength, mainly due to (Pauli exclusion principle induced) electrostatic repulsion. This effect is difficult to quantify and unambiguously study, however. In this paper a first attempt is made to do so. The Fe+ ion has a 3d64s1,6D ground state in contrast to the 3d8, 3F ground state of Co+. Of particular interest is whether or not the 3d7, 4F first excited state of Fe+, just 6.5 kcal/mol above the ground state, will dominate the bonding as it does with simple a l k a n e ~or ~ .whether ~ Fe+ will remain on the ground state surface as it ligates with H2. We will also compare our results to ab initio calculations that are in progressI7 in an attempt to extract the factors important in the bonding.
Methods The instrument,I8 equilibrium experiment, and data analysis have all been described p r e v i o u ~ l y . ~ - ”Sources ~’~ of error have also been discussed extensively.Iob In this section, a brief overview is given with emphasis on the differences encountered in these experiments. The Instrument. A variety of methods were used to form Fe+ reactant ions. Initially, Fe+ ions were formed via surface ionization (SI) on a hot filament (Re ribbon or W wire, T 2500 K) or by electron impact (EI) of FeC12 or FeC13 vapor. The FeC12 and FeC13 vapor was introduced by heating a sample of FeC12 or FeCl3 powder in a small oven attached to the source. This allowed comparison between “low-energy” SI formed Fe+ and “high-energy” E1 ions. No measurable differences in the various equilibria could be detected. Unfortunately, the chloride sources of Fe+ proved to be unstable. Attempts to make Fe+ via SI of Fe(C0)s failed due to the large amount of heat produced by the SI filament. Even with source cooling, the heat rapidly built up to the point where the Fe(C0)5 thermally decomposed before reaching the SI filament, and the resulting CO pressure quenched the ion signal. We were, however, able to make Fe+ from E1 on Fe(C0)s. The equilibria obtained agreed with E1 and SI results from the chlorides within the uncertainty of the measurements. As an extra caution against impurities in the source, alternate runs were performed using the 56Fe and 54Fe isotopes. These results were in close agreement with each other. The Fe+ ions were accelerated to 5 keV, mass selected with a double focusing, reverse geometry mass spectrometer, decelerated to approximately 3-5 eV, and injected into a reaction cell containing H2 at about 1 x lOI7 molecule/cm3 (3 Torr at 300 K.) The ions were quickly translationally thermalized via collisions with Hz, and they drifted through the 4 cm long reaction cell under the influence of a small uniform electric field. The electric field was kept small enough that the ion thermal energy was not significantly perturbed. The H2 pressure in the reaction cell was monitored directly with a capacitance manometer. Cell temperatures were varied using a flow of heated or cooled NZ and temperatures measured using a thin film platinum resistor suspended in the bath gas. Ions exiting the cell were accelerated slightly (2-5 eV), quadrupole mass analyzed, and collected using standard ion-counting techniques. The quadrupole was computer scanned over the mass range of interest, and the baseline-resolved peaks were then integrated to give the relative ion intensities. The Equilibrium Experiment. After the temperature and H2 pressure were observed to be stable within the reaction cell,
-
product/parent ion ratios were measured as a function of reaction time. This time was varied by changing the drift voltage across the cell. As the drift time was increased, the product/parent ion ratios eventually become constant indicating that equilibrium had been reached. The reactions were probed out to 1 ms (-1 x 105 collisions) to ensure that equilibrium had been established. This measurement also ensured that the drift field did not significantly pertub the ion thermal kinetic energy. For selected experiments, the pressure of H2 was varied by a factor of 2 with no significant change observed in AGO, The ion ratios were converted to equilibrium constants using eq 2,
where P H is ~ the hydrogen pressure in torr and M+(Hdn and M+(H2)n-~are the measured intensities of the cluster ions of interest. The standard free energy change was calculated using eq 3,
AGO = -RT In K;:
(3)
where R is the gas constant and Tis the temperature. A plot of AGO us temperature gives a straight line with intercept = AW,and slope of -ASo, where T is the average experimental temperature. This plot is functionally equivalent to a van’t Hoff plot but is more convenient for our data analysis. In order to obtain the desired cluster bond dissociation energies (BDE), the AW, must be converted to W , (EBDE). This was done by calculating AGO as a function of temperature using statistical mechanics. The bond lengths, frequencies, and dissociation energy used in the calculation were varied until the experimental and calculated functions agreed. In all cases, the vibrational frequencies were varied over a wide range to see the effect on the resulting AW,(-DO) and these uncertainties are included in the error limits. A number of potential sources of error are present in these experiments, e.g., pressure and temperature inaccuracies and mass discrimination and resolution. The effect of these factors on our dissociation energies and entropies is discussed extensively in refs 10b and 19. The net result is that the bond dissociation energies are essentially unaffected by these uncertainties. The entropies are affected to a small extent by any mass discrimination in the quadrupole mass analyzer, but in these experiments such mass discrimination should be very small since the parent/product mass difference is only 2 amu. In several of the M+(H2), systems, high-level molecular structure calculations have been performed. In these cases, dissociation entropies can be very accurately calculated. In all such cases, our experimental entropies are within f 2 cal/Kmol of the theoretical entropies. This is probably the best estimate of our maximum experimental error. Deactivation of excited states was a concem in our experiments on Co+(H2),, with very long reaction times required to reach the true equilibrium between Co+ and CoH2+ due to persistence of excited states of 4s’3d7 electron configuration even after thousands of collisions with H2.Iob This was presumably due to the lack of curve crossings between the weakly bonding 4s13d7excited states and the relatively strongly bound 3d8 ground state of Co+. In the case of Fe+, the strongly bound 3d7 first excited state should provide ample curve crossings to the weakly bonding 4s13d6 ground state. In previous drift-cell experiments in He buffer gas, excited state Fe+ was seen to deactivate slowly, while no deactivation was
15604 J. Phys. Chem., Vol. 99, No. 42, 1995
Bushnell et al.
-
TABLE 2: Derived Binding Energie (-AH:)for M+(H,),-~ + H,
M+(H,),
n
V+
Fe+
co+
1 2 3 4 5 6
10.2 f 0.5 10.7 f 0.5 8.8 f 0.4 9.0 f 0.4 4.2 f 0.5 9.6 f 0.5
10.8 f 0.6 (16.5)b 15.7 f 0.7 7.5 +C 0.4 8.6 f 0.4 2.2 f 0.3 2.3 & 0.3
18.2 rt 1.0 17.0 f 0.7 9.6 f 0.5 9.6 f 0.6 4.3 f 0.7 4.0 f 0.7
a In units of kcaymol. Binding energy with respect to lowest d7 configuration.
0
100
200
500
400
300
600
Temperature (K)
-
Figure 1. Experimental free energy change (AGO, in kcal/mol) plotted H2 Fe+(H2),. Lines through vs temperature (in K) for Fet(H2),-l idata represent linear least squares fits. TABLE 1: Experimental AH; and AS; Values for Fe+(H,),-I n 1 2 3 4 5 6 a
-
w;
12.5 f 0.2 17.0 f 0.2 8.4 f 0.1 9.1 f 0.1 2.6 f 0.1 2.7 f 0.1
+ H, =+
Fe+(H,),
- AS:b 21.5 f 0.4 25.2 f 0.5 19.1 f 0.4 24.9 f 0.4 17.9 f 0.5 18.1 f 1.1
temp rangeC 360-580 375-580 240-455 240-400 85-240 85-125
In units of kcaymol. In units of caVmo1.K. In K.
energies at 0 K (-AHo,, extracted using the statistical mechanical fitting procedure outlined in the last section) are given in Table 2 along with previous results for V+ and Cof for comparison. Since the binding energies of H2 with Fe+ roughly follow the pairwise behavior observed for Cof(H2),, the results for Fef(H2), will be discussed two at a time. Fe+(HZ)l,,. The f i s t HZligand binds with Fe+ about 5 kcaV mol less strongly than the second (see Table 2), an effect not observed for Co+ and V+. One important difference in these three systems is their electronic configurations: Co+ and V+ have 3dn ground states while Fe' is 3dn-'4s'. Consequently, the ground state of Fe+ (6D 3d64s1)is less amenable for bonding due to the large highly repulsive 4s orbital. Calculations indicate a binding energy of only 2-3 kcal/mol for the 3d64s' configuration.17 Hence, the observed binding energy of 10.2 kcaV mol is much larger than expected for 3d64s' Fe+ ions. The first excited state of Fe+ is the 4F 3d7 state which is only 5.7 kcal/mol above the ground state (averaged over J for both states2'). If we assume the Fe+ (6D, 3d64s') ion crosses over to the Fef (4F, 3d7) surface as it approaches the H2 ligand, then the diabatic dissociation energy of Fe+H2 Fe+(4F, 3d74s') H2 would just be the observed BDE plus the 6D 4F promotion energy. This yields a value of 16.3 kcaVmo1, slightly larger than the second binding energy of 15.7 kcaVmol. Thus, after accounting for the 6D 4F promotion energy, the first two H2 ligands added to Fe+ would have nearly the same binding energies. This is what was observed for Cof and V+ and strongly supports the proposed curve-crossing mechanism. Theory indicates that the first H2 ligand binds in a sideways geometry rather than end-on due to the positive quadrupole moment of H2 and in order to enhance back donation of charge from & orbitals on the metal to the H2 (a*)0rbita1s.I~ To minimize the repulsion and maximize the electrostatic interaction between Fe+ and H2, the 4s orbital on Fe+ should be unoccupied. The & (dYJ orbital needs to be doubly occupied for back donation to the u* orbital of H2. Theory has shown that back donation from a singly occupied & orbital is negligible due to the resultant loss in d-d exchange energy.I2 Ab initio calculations have been carried out on Fe+ with up to six H2 ligands by Maitre and Baus~h1icher.l~Their ground state geometries are shown in Figure 2. The calculations on Fe+H2 show the ground state to be 4B2 with the dYzand dx2 orbitals doubly occupied.22 The dZ2-+ orbital mixes with the 4s orbital in order to reduce the charge along the bond axis, allowing a closer approach of the ligand. The first excited state of Fe+H2 (4A2with dyzand dxzdoubly occupied) is calculated to be only about 2 kcaVmol above the ground state. At equilibrium, this results in a first excited state population of 6-15% over the experimental temperature range. In fact, the experimental entropy could not be adequately accounted for without including this excited state in the analysis, which effectively increases the entropy for the formation of this cluster. The second excited state of Fe+H2 (4B1 with dyzand dXydoubly +
observed for excited state CO+.~ODeactivation is expected to be more facile in H2, and in fact a dramatic increase in deactivation was seen for excited states of V+ in He when a small amount of D2 was added.Ioa Ion chromatographyZostudies indicated that no excited states of 3d7 electronic configuration were present. Our assumption that all the excited states of Fe+ were effectively deactivated is strongly supported by the fact that identical results were obtained for Fe+ formed by both SI and EI. One additional source of error is collision-induced dissociation (CID) of the clusters as they are accelerated into the quadrupole. This process most likely occurs to a small extent for weakly bound clusters, and its importance has previously been d i s c u s ~ e d . ~ Such - ' ~ ~ processes can have a minor effect on measured values of ASo, but have essentially no effect on AHo, or AZP,. The reason for this is that the fraction of clusters undergoing CID is largely independent of temperature since the pressure outside the cell is very nearly temperature independent. Consequently, any loss of clusters via CID appears as a constant factor in the equilibrium constant. That this fraction of CID is small is apparent from experiments where the cell pressure was changed by a factor of 2, and the measured value of AGO showed little or no change. This result argues against significant CID which should scale with cell pressure.
Results and Discussion Our experimental data is plotted in Figure 1 as AGO (from eqs 2 and 3) vs temperature. As noted above, AS", and AZP, are taken from the slopes and intercepts, respectively, of these plots. These values are given in Table 1 along with error limits equal to 2 standard deviations. The derived binding
-
-
+
J. Phys. Chem., Vol. 99, No. 42, 1995 15605
Factors Affecting a Bond Activation
n= 1
n=2 h
n=3
n=5
n=4
n-6
Figure 2. Calculated structures for Fe+(HZ), (n = 1-6) from Maitre and Bauschlicher (ref 17). The structure for n = 3 is planar; the n = 4 ligand adds perpendicular to this plane; the n = 5 structure is a slightly distorted square pyramid; n = 6 is close to octahedral.
occupied) was calculated to be over 5 kcal/mol above the ground state.17 This gives a second excited state population of 50.11.3%, and thus this state had very little effect on our data analysis. Two structures for Fe+(H2)2 are calculated to be very close in energy." One, with DU symmetry, is derived from the 4B2 ground state of Fe+H2. The second, with D2h symmetry, is derived from the 4A2 first excited state of Fe+H2. In either structure s-d hybridization is now enhanced since both ligands benefit, yet the cost is spread out between them. Again, both states were included in the analysis. While the existence of low lying cluster excited states affects the calculated entropies, and thus the vibrational frequencies necessary to model the experimental data, they had little or no effect on the derived binding energies reported here. The diabatic binding energies (to the lowest lying d" states) for Fe+(H2)1,2 and Co+(H2)1,2 are very similar, with the Fe+ binding energies slightly smaller. Several factors could affect the relative binding energies including orbital size, amount of s-d hybridization, and orbital occupation. In general, it is expected that electrostatic binding with transition metals should increase in moving to the right along the periodic table as the 3d orbitals contract, allowing for a closer approach. For instance, calculations give larger De's for He and Ar with Co+(3d8) (4.0 and 9.6 kcal/mol, respectively) than with Fe+(3d7) (3.3 and 7.4 k~al/mol)?~The relative amount of s-d hybridization in Fe+(H2)1,2 and Co+(H2)1,2 is unclear. The 4s and 3d orbitals should be slightly closer in size for Fe+ since the 3d orbitals contract faster than the 4s moving to the right along the periodic table. But the energetics of excited 3dfl-'4s states also plays a role. Both Fe+(H2)1,2 and Co+(H2)1,2have singly occupied d a and doubly occupied dn orbitals (for back bonding) making the orbital populations nearly equivalent. Thus, the smaller binding energies for Fe+(H2)1,2 relative to Co+(H2)1,2 could be due to the larger size of Fe+, with s-d hybridization effects difficult to predict. Fe+(H2)1,2 and Co+(H2)1,2have binding energies which are 5 -8 kcaymol larger than the corresponding binding energies for V+(H2)1,2. (See Table 2.) This rather large effect is
primarily due to the presence of back donation from the doubly occupied dn orbitals on Fe+ and Co+ to the u* orbitals of the H2 ligands. As mentioned above, the singly occupied d orbitals on V+ do not back donate significant charge due to the resulting loss of d-d exchange interaction between unpaired d electrons. The larger size of the d orbitals on V+ also plays a role but is almost certainly less important than back donation in explaining the difference in binding energies. Fe+(H2)3,4. The third and fourth Fe+(H2), clusters have binding energies which are significantly lower than the first two. This is similar to the pairwise decrease in binding energies observed for Co+(H2),. The change could be due to a loss of s-d hybridization present in the first two clusters and/or to increased steric interference between ligands. In our previous paper discussing Co+(H2),, we proposed a steric interference model to account for the observed pairwise decrease in binding energy." Subsequent calculations on the larger (n 1. 4) clusters indicate that changes in the s-d hybridization are actually responsible for the decrease between the first and second pairs of ligands, and that steric interference between the H2 ligands is, in fact, very s m a l l 7 Calculations give a C, structure as the ground state for Fe+(H2)3.17This quasi T-shaped structure is derived from the DU structure of Fe+(H2)2 by adding the third H2 to the side (that is, at 90" angles to the first two ligands, see Figure 2). The third H2 is then oriented so that its H-H bond lies in the same plane as one of the first two H2 ligands. This orientation aligns the a * orbital of H2 with a doubly occupied dn orbital of Fe+ maximizing back donation. While the first two H2 ligands do bend slightly away from the 180" angle found for Fe+(H2)2, the structure of Fe+(H2)3 is distinctly "T" shaped (rather than trigonal planar) to allow good overlap with the d orbitals. The fourth ligand binds out of plane at approximately right angles to the first three H2's giving a C2 complex. (See Figure 2 . ) This configuration allows each of the two doubly occupied 3d orbitals on Fe+ to overlap with two of the H2 a * orbitals, thus maximizing the dn2 a* donation in the cluster. This structure, rather than a square planar structure which would be predicted on the basis of steric interference, may indicate some involvement of the 4p, and 4p, orbitals on Fe+, and a resultant polarization of electron density away from the ligands. This possibility is discussed below. The large drop in binding energy between Fe+(H2)2 and Fe+(H2)3 can be understood as a loss in s-d hybridization. While the first two H2 ligands benefit from s-d hybridization, which reduces the electron charge along the bonding axis, the third ligand approaches along the maximum charge density of this hybridization, cancelling out the benefit by the first two ligands. Thus, not only is the third ligand unable to benefit from s-d hybridization but it reduces this benefit for the first two ligands as well. Once this price is paid by the third H2, the fourth might be expected to bind more strongly since the s-d hybridization has already been disrupted, and we do observe a slightly higher binding energy for the fourth H2 in Fe+(H2),. The third and fourth binding energies in Co+(H2), were observed to be identical and lower than the first two. Steric interference could easily explain this, but it could also be the result of a combination of a hybridization change along with a reduction in back donation between n = 3 and n = 4. Since Co+ has only three doubly occupied d orbitals available for back donation, H2 ligands beyond the third must share back bonding metal orbitals with other ligands. We would expect this effect to begin at n = 3 for Fe+ (since it only has two filled d orbitals available) and the drop in binding energy between n = 2 and n = 3 is in fact slightly larger for Fef(H2), than Co+(H2),. If
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15606 J. Phys. Chem., Vol. 99, No. 42, 1995 the drop in binding energies between n = 1 and 2 and n = 3 and 4 were due to steric interference between ligands, we would expect a larger drop in binding energy for Co+(Hz), which has shorter M+-H2 bond lengths than Fe+(H2),. This is not observed. Fe+(H&,,j. The binding energies for the fifth and sixth Fe+(H2), clusters are again observed to be equal and much lower than the third and fourth. This effect is analogous with the trend observed for Co+(H2),. Calculations show that a slightly distorted square pyramid is the lowest energy structure for Fef(H&.I7 The H2's are oriented so that three ligands can accept charge from one filled d orbital while the other two ligands can accept charge from the second filled d orbital. Fe+(H2)6 is calculated to be approximately octahedral, with H2's oriented so that each of the doubly occupied d orbitals can back donate to three ligands. The larger drop in binding energy between n = 3 and 4 and n = 5 and 6 for Fe+(H2), compared with Co+(H2), is almost certainly because Fe+ only has two filled d orbitals from which to back-donate, while Co+(H2), has three. The larger size of Fe+ could also lead to smaller binding energies, though this effect is expected to be small since the purely electrostatic contributions to binding energy are actually quite small (see below). Factors Affecting Bond Energies in M+(H*),,Complexes. In the Co+(H2), system previously studied, the pairwise decrease in binding energies was explained using a simple model of increasing steric repulsion with increasing coordination.Iob While this model was consistent with the trends in binding energy observed experimentally, subsequent theoretical and experimental evidence now suggests that orbital interactions are more important than simple electrostatic interactions. Electrostatic Contributions (charge-induced dipole and charge-quadrupole attractions) appear to be significant but less important than originally thought. Some steric repulsion must exist but is probably small. However, major contributions from "covalent" effects including HZ M+donation, back donation to H2 (T*orbitals, and s-d hybridization are all critical in explaining the magnitudes and trends observed in these M+(H2), systems. Binding energies for Na+-H2 and K+-H2 were determined to be very small (2.5 and 1.5 kcal/mol, respe~tively).~ These systems were chosen for study because Na+ and K+ are closed shell ions, making them excellent model systems for purely electrostatic interactions. Even after taking into account the smaller size of V+ and Co+, purely electrostatic interactions account for less than half of the observed binding energies in V+-H2 and CO+-HZ.~ The trend in binding energies for V+(H2), is also difficult to explain using a simple model of electrostatic attraction and steric interference. While we did see a pairwise drop in binding energies from V+(H2)1,2to V+(H2)3,4,the change is small, about 2 kcal/mol.'oa However, the fifth binding energy is about 5 kcal/mol less than the fourth, very similar to that observed for Co+(H2),. Presumably, the binding energy of the sixth H2 to Vf would be similar to the fifth except that V+(H2), undergoes a spin change at n = 6, driven in part by the ability to back donate charge from the doubly occupied d orbital in the lowspin configuration.loa This decrease in binding energy between n = 4 and n = 5 may seem surprising for V+(Hz),, in which longer Mf-H2 bond lengths should give significantly less ligand-ligand repulsion, but it can be understood in terms of 3 d a - H ~ repulsion. ~ The V+ 3d4 ground state configuration allows four H2 ligands to approach an empty d a orbital (the dXz-p). The fifth and sixth H2 ligands must approach the half filled dZ2orbital however, with a resulting increase in repulsion. This explanation is supported by preliminary data on the
Bushnell et al.
-
4Px
-0 7,-
4s
1
-
-
e'
I'
'
Zl
4Pz
4Px
-i/ W
Figure 3. Possible electron donations from
4Py
-
H 2 ligands Fe+ in Fe+shown as linear combinations of H 2 u orbitals and the corresponding metal (acceptor) orbitals, for (a) square planar geometry and (b) C2 (distorted tetrahedral) geometry. The arrows indicate which orbital on Fe+ has the correct symmetry to accept charge from the linear combination shown. (&)4
corresponding Ti+(H2), clusters24 which do not show any decrease in binding energy between Ti+(H&,4 and Ti+(H&. This is expected because Ti+ has only three valence electrons and can thus accommodate six H2 ligands with its empty d orbitals. As discussed earlier, the calculated structure for Fe+(H2)4 (Figure 2) was not square planar as would be expected for a system which balances electrostatic attractions against steric repulsion. However, examination of acceptor orbitals on Fe+ provides a possible explanation. Figure 3a and b shows linear combinations of H2 (T orbitals and the corresponding charge-
J. Phys. Chem., Vol. 99, No. 42, 1995 15607
Factors Affecting o Bond Activation accepting orbitals on Fe+ for square planar and CZgeometries, respectively. In the square planar configuration, the HZligands can donate charge to the empty 4s orbital on Fe+ and possibly to two of the empty 4p orbitals. Donation to the singly occupied dg-9 orbital on Fe+, while allowed, is unfavorable due to Pauli repulsion and loss of d-d exchange energy. In the ab initio CZgeometry, though, the ligands are now able to donate into all three of the 4p orbitals on Fe+. The lowest quartet 3d64p state of Fe+ is 125 kcdmol above the lowest quartet 3d7 state, so involvement of the 4p orbitals is very limited and perhaps better viewed as serving to polarize electron density away from the H2 ligands. For whichever reason, however, the CZground state geometry of Fe+(Hz)4 suggests that the 4p orbitals are involved. Acceptor orbitals can also be used to explain the small drop in binding energy between V+(Hz)z and V+(HZ)Jand the larger drop between V+(H2)4 and V+(H&. V+ has only four valence d electrons and so has one unoccupied d orbital in its high-spin ground state configuration. Whereas all of the d orbitals available for accepting charge in Fe+ and Co+ are singly occupied and s-d hybridization is only effective in the binding of the first two H2 ligands, V+ can accept charge from up to four H2 ligands into its empty d orbital. This results in a favorable interaction with the 3dg+ orbital shown in Figure 3a and a square planar geometry for V+(H2)4.I6 Thus, there is little change in binding energy between the first and second pairs of H2 ligands. However, the fifth H2 can no longer donate into this empty d orbital, and we see a substantial drop in binding energy for V+(H2)5. Ti+ only has three valence d electrons, and we did not observe a pairwise decrease in binding energies for T~+(Hz),.*~ Thus, consideration of available acceptor orbitals appears to be able to account for many of the trends which we have observed for first row transition metals in M+(H2)n complexes and indicates covalent charge transfer interactions are important in these systems.
Conclusions We have determined sequential binding energies and entropies for the clustering of Fe+ with molecular hydrogen, Fe+(Hz), (n = 1-6) via temperature-dependent equilibrium measurements. Binding energies show a pairwise behavior much like that observed for Co+(Hz),, except (1) the first H2 ligand is about 5 kcal/mol less strongly bound than the second, presumably due to the promotion energy needed to reach the lowest 3d74s0(4F) asymptote of Fe+. (2) Ligands beyond the first in Fe+(Hz), are bound 1-2 kcdmol less strongly than the corresponding Co+(H2), species, probably due to a combination of iron's larger size and increased M'(3dn) H~(u*)back donation for Co+ relative to Fe+. While previous data on Co+(Hz), could be adequately explained using a simple model of steric interference between ligands, subsequent theoretical and experimental evidence has prompted a reassessment of factors that determine trends in M+(Hz), binding energies. Important factors include s-d hybridization, availability of filled dn metal orbitals for back-donation of charge to the HZo* orbitals, and availability of empty metal
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orbitals to reduce M+do-Hza repulsion and for accepting charge from the H2 u orbitals.
Acknowledgment. The support of the National Science Foundation under Grant CHE-9421176 and partial support of the Air Force Office of Scientific Research under Grants FA49620-93-1-0134 and 93-1-0305 are gratefully acknowledged. We also wish to thank Drs. Philippe Maitr6 and Charles Bauschlicher for communicating results prior to publication. References and Notes (1) See (a) Homogeneous Transition Metal Catalysed Rections; Moser, W. R., Slocum, D. W., Eds.; ACS Symposium Series 230; American Chemical Society: Washington, DC, 1992. (b) Crabtree, R. H.; Hamilton, D. G. Adv. Organomet. Chem. 1988, 28, 299 and references therein. (2) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, R. H.; Halow. I.: Bailev. S. M.: Chumev. K. L.: Nuttall. R. L. J. Phvs. Chem. Re8 Data 1982,-11, Suppl. 2. (3) Low, J. J.: Goddard, W. A. III. J. Am. Chem. SOC.1984,106,8321; Orgunometallics 1986, 5, 609. Blomberg, M. R. A,; Siegbahn, P. E. M.; Nagashima, U.; Wennerberg, J. J. Am. Chem. SOC.1991, 113, 424. (4) Amdtsen, B. A.; Bergman, R. G . ;Mobley, T. A,; Peterson, T. H. Acc. Chem. Res. 1995, 28, 154. (5) Weisshaar, J. C. Acc. Chem. Res. 1993, 26, 213 and references therein. (6) van Koppen, P. A. M.; Brodbelt-Lustig, J.; Bowers, M. T.; Dearden, D. V.; Beauchamp, J. L.; Fisher, E. R.; Armentrout, P. B. J. Am. Chem. SOC. 1991,113,2359; 1990,112,5663. van Koppen, P. A. M.; Bowers, M. T.; Fisher, E. R.; Armentrout, P. B. J. Am. Chem. SOC.1994, 116, 3780. (7) van Koppen, P. A. M.; Kemper, P. R.; Bowers, M. T. In Organometallic Ion Chemistry; Freiser, B. S., Ed.; Kluwer Academic Publishers: the Netherlands, in press, and references therein. (8) van Koppen, P. A. M.; Bushnell, J. E.; Kemper, P. R.; Bowers, M. T. J. Am. Chem. SOC.1995, 117, 2098. (9) Bushnell, J. E.; Kemper, P. R.; Bowers, M. T. J. Phys. Chem. 1994, 98, 2044. (10) (a) Bushnell, J. E.; Kemper, P. R.; Bowers, M. T. J. Phys. Chem. 1993, 97, 11628. (b) Kemper, P. R.; Bushnell, J.; von Helden, G.; Bowers, M. T. J. Phys. Chem. 1993, 97, 52. (1 1) Bushnell, J. E.; Kemper, P. R.; Maitre, P.; Bowers, M. T. J. Am. Chem. SOC.1994, 116, 9710. (12) Bauschlicher, C. W., Jr.; Partridge, H.; Langhoff, S. R. J. Phys. Chem.1992,96,2475. Maitre, P.; Bauschlicher, C. W. J. Phys. Chem. 1993, 97, 11912; 1995, 99, 3444. (13) Perry, J. K.; Ohanessian, G.; Goddard, W. A,, 111. J. Phys. Chem. 1993,97,5238. Perry, J. K. PhD Thesis, California Institute of Technology, 1994. (14) Blomberg, M. R. A.; Siegbahn, P. E. M.; Svensson, M. J. Phys. Chem. 1994, 98, 2062. (15) Alvarado-Swaisgood, A. E.; Harrison, J. F. J. Phys. Chem. 1985, 89, 5198. Rappe, A. K.; Upton, T. H. J. Chem. Phys. 1986, 85, 4400. (16) Maitre, P.; Bauschlicher, C. W., Jr. J. Phys. Chem. 1995,99,6836. (17) Maitre, P.; Bauschlicher, C. W., Jr. Manuscript in preparation. (18) Kemper, P. R.; Bowers, M. T. J. Am. SOC.Mass Spectrom. 1990, 1, 197. (19) Kemper, P. R.; Hsu, M. T.; Bowers, M. T. J. Phys. Chem. 1991, 95, 10600. (20) Kemper, P. R.; Bowers, M. T. J. Phys. Chem. 1991, 95, 5134. (21) (a) Moore, C. E. Atomic Energy Levels; U. S. National Bureau of Standards: Washington, DC, 1952; Vol. Circ. 467. (b) Sugar, J.; Corliss, C. J. Phys. Chem. Re$ Data 1978, 7, 1194. (22) The d,i orbital arises from a rotation of the principal angular momentum axis from the t (Fe+-H* bond) axis to the x axis. This changes the spatially defined set of d orbitals from {d;?, dx2,dyE,d+,., dxy}to {d,., dxy.dxz,dz>-p,dyr}.This change of axes is made in order to achieve a pure electronic component of the lowest iron 3d7 (a4F) state. (23) Partridge, H.; Bauschlicher, C. W., Jr. J. Phys. Chem. 1994, 98, 230 1. (24) Bushnell, J. E.; Kemper, P. K.; Bowers, M. T. Manuscript in preparation. JP95 11262
