Comparison of π-Complexations of Ethylene and Carbon Monoxide

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Ind. Eng. Chem. Res. 1999, 38, 2720-2725

SEPARATIONS Comparison of π-Complexations of Ethylene and Carbon Monoxide with Cu+ and Ag+ Helen Y. Huang, Joel Padin, and Ralph T. Yang* Department of Chemical Engineering, University of Michigan, Ann Arbor, Michigan 48109-2136

A direct comparison of adsorption of CO and ethylene on AgCl and CuCl is made. An ab initio molecular orbital study using the effective core potential is performed to determine the bond energies and the nature of the bonds between the adsorbates and adsorbents. Experimental results show that both CO and C2H4 adsorb more strongly on CuCl than on AgCl. However, CO adsorbs much more strongly on CuCl than on AgCl, while the difference in heats of adsorption is far less for C2H4 on these two sorbents. Ab initio molecular orbital calculations correctly predict the trends. The natural bond orbital (NBO) theory is used to explain the results. The NBO results show that in the two-way π-complexation bonding, the d-π* backdonation plays a major role in determining the bonding for these systems. Introduction Separations and purifications via π-complexation have been studied with great interest because it is a promising alternative for distillation and a possible solution to difficult separations. The π complexation is a subgroup of chemical complexation where the mixture is contacted with a second phase containing a complexing agent.1 The advantage of separation by π complexation is that the π-complexation bonds are stronger than those by van der Waals forces alone, so it is possible to have a high selectivity and a high capacity for the component to be bound. Yet, the bonds are still weak enough to be broken by simple engineering means such as raising the temperature or decreasing the pressure of the system. Separation and recovery of CO from its mixtures by absorption using π complexation has been practiced in industry for a long time. Aqueous solutions containing various Cu+ and Ag+ salts have been used, with Cu+ salts being more successful. Examples are copperammonium salts2 and CuAlCl4, i.e., the COSORB process.3 These are absorption processes involving gasliquid systems. More recently, cyclic adsorption processes involving gas-solid systems have been developed, because of success in the development of solid sorbents, e.g., CuCl-alumina4 and CuCl-carbon.5 Olefin-paraffin separations represent a class of most important and also most costly separations in the chemical industry. Cryogenic distillation has been used for over 60 years for these separations.6 They remain the most energy-intensive distillations because of the close relative volatilites.7 The most important olefinparaffin separations are for the binary mixtures of ethane-ethylene and propane-propylene. A number of alternatives have been investigated.8 The most promis* Corresponding author. Tel.: (734) 936-0771. Fax: (734) 763-0459. E-mail: [email protected].

ing one appears to be separation via π complexation.6,8 Aqueous solutions of various Ag+ and Cu+ salts have been used. The most successful absorption system has been that of Keller et al.6 using a AgNO3 solution. Most recently, efficient solid adsorbents have been developed in our laboratory for cyclic adsorption processes. In the work of Yang and Kikkinides,9 two types of solid π-complexation sorbents have been developed: cationexchanged resins and monolayer salts on alumina. AgNO3-SiO2 appears to be most promising.10,11 These sorbents contain highly dispersed Ag+ or Cu+ cations and have high selectivities, high capacities, fast rates, and reversibility. Despite the success and the tremendous efforts made in this area, fundamental questions that remain to be answered are, what is the difference between the bonding of CO and olefins with the cations and what is the difference between Ag+ and Cu+? Ab initio molecular orbital calculations as well as experiments were performed to provide an answer to this question and also to provide a fundamental understanding of the π-complexation bonds for these systems. Ab Initio Molecular Orbital Calculation Details The Gaussian 94 program package12 was used for all of the calculations. The molecular orbital calculations were performed at the restricted Hartree-Fock (RHF) and density functional theory (DFT) using an effective core potential (ECP)13,14 for the Ag, Cu, and Cl atoms at the double-ζ-type valence basis set level, whereas Dunning’s basis set was used for the C, H, and O atoms. DFT. DFT was applied to analyze the adsorbentadsorbate bond features in terms of a natural bond orbital (NBO) scheme. In recent years, the methods originated from DFT have been recognized as being advantageous because they incorporate electron correlation at an affordable computational cost, so it is an efficient tool for studying molecular properties of large

10.1021/ie990035b CCC: $18.00 © 1999 American Chemical Society Published on Web 06/02/1999

Ind. Eng. Chem. Res., Vol. 38, No. 7, 1999 2721

molecules such as compounds of transition metals. It is also found that they can provide an accurate description of the metal-ligand interactions.15 In particular, it has been found that a hybrid method consisting of HF and DFT, leading to the so-called self-consistent hybrid (SCH) approaches,16 is even better than the standard DFT. The SCH approaches are able to provide reliable geometric, thermodynamic, and spectroscopic parameters for a wide class of metal-ligand interactions, ranging from covalent bonds to weak noncovalent interactions.17-19 In this work, we selected one of the most useful SCF approaches, the B3LYP20 approach, which is the combination of HF and Becke exchange21 with the Lee-Yang-Parr (LYP) correlation potential.22 ECP. The applications of the all-electron ab initio molecular orbital calculations have been restricted to small molecular models because of its relatively expensive computational cost. In addition, for the heavier elements, relativistic effects must be considered, which have a significant influence on the physicochemical properties of molecules.13 The use of ECP has been a great success in molecular orbital calculations involving transition metals. ECP is simply a group of potential functions which replace the inner shell electrons and orbitals which are normally assumed to have minor effects on the formation of chemical bonds. Calculations of the valence electrons using ECP can be carried out at a fraction of the computational cost that is required for an all-electron (AE) calculation, while the overall quality of the computation does not differ significantly from the AE calculations.13,14 In addition, the relativistic mass-velocity and Darwin terms, which are derived from AE relativistic HF calculations, are implicitly incorporated into the relativistic effective core potentials (RECP) for heavier elements (Z > 36).23-25 Combined with the use of reliable basis sets, it appears to be a very powerful method for dealing with molecules containing heavy transition metals. More recently, Hay and co-workers have shown that ECP can be used reliably in density functional computations as well.26,27 Following this approach, the LanL2DZ basis set was employed for both geometry optimization and NBO analysis in this work. The LanL2DZ basis set is a double-ζ basis set containing ECP representations of electrons near the nuclei for post-third-row atoms. The reliability of this basis set has been confirmed by the accuracy of calculation results as compared with experimental data. Geometry Optimizations and Calulations of Energy of Adsorption. The restricted Hartree-Fock (RHF) theory at the LanL2DZ level basis set was used to determine the geometries of MCl-C2H4 and MClCO (where M ) Ag or Cu) adsorption systems. The optimized structures were then used for energy calculation and NBO analysis at the B3LYP/LanL2DZ level. The bond energies were calculated according to the following expression:

Eads ) Eadsorbate + Eadsorbent - Eadsorbent-adsorbate (1) where Eadsorbate and Eadsorbent are the total energies of the adsorbate (C2H4 or CO) and the bare adsorbent (CuCl or AgCl), respectively, and Eadsorbent-adsorbate is the total energy of the adsorbate-adsorbent system. NBO. Information concerning atomic charge distributions is important in understanding chemical bonds. Mulliken population analysis28 has been widely used. Unfortunately, Mulliken populations fail to give a reliable characterization of charge distribution in many

Figure 1. Pure-component equilibrium isotherms for CO adsorption on CuCl and AgCl salts at 0 °C. Lines are fittings with eq 2.

cases.29-31 Natural population analysis (NPA) seems to be a promising alternative to the conventional Mulliken population analysis. It gives a better description of the electron distribution in compounds of high ionic character, such as those containing metal atoms.32 NPA includes a series of calculations, such as determination of the natural atomic orbitals (NAOs), natural hybrid orbitals (NHOs), natural bond orbitals (NBOs), and natural localized molecular orbitals (NLMOs). It performs population and energetic analyses that pertain to localized wave function properties. It is very sensitive for calculating localized weak interactions, such as charge transfer, hydrogen bonds, and weak chemisorption. Therefore, the NBO program33 included in Gaussian 9412 was used in this work for studying the electron density and charge distribution of the adsorption systems. Experimental Section Equilibrium adsorption isotherms were measured using a Micromeritics ASAP 2010. The ASAP 2010 utilizes a volumetric system to obtain adsorption isotherms. Equilibrium isotherms were measured at 0 and 25 °C to obtain isosteric heats of adsorption data. Surface area measurements were made also using the Micromeritics ASAP 2010 and were performed using nitrogen adsorption at 77 K. The ethylene and carbon monoxide gases used for this experiment were obtained from Matheson (CP grade, minimum purity 99.0%). The gases were used without further purification. The silver and copper halides used in this work were obtained from Strem Chemicals. The samples were degassed in vacuo at 100 °C before each experiment to remove adsorbed impurities. Results and Discussion Experimental Results. To determine the adsorption behaviors of CO and C2H4 on CuCl and AgCl, purecomponent equilibrium isotherms were measured at 0 and 25 °C using the Micromeritics ASAP 2010. The selection of these salts allowed us to maintain the same anion, and this also simplifies the ab initio molecular orbital study. The equilibrium data were fitted with the Langmuir isotherm with two parameters shown as:

q)

qmpbpP 1 + bpP

where qmp and bp are the fitting parameters.

(2)

2722 Ind. Eng. Chem. Res., Vol. 38, No. 7, 1999 Table 1. Fitting Parameters for C2H4 and CO on AgCl and CuCl at 0 and 25 °C and BET Surface Area (SA) 0 °C

25 °C

adsorbent

adsorbate

qmp (mmol/g)

bp (1/atm)

qmp (mmol/g)

bp (1/atm)

SA (m2/g)

C2H4 C2H4 CO CO

AgCl CuCl AgCl CuCl

0.077 0.088 0.054 0.055

0.227 0.349 0.635 2.62

0.042 0.078 0.053 0.054

0.140 0.104 0.210 0.550

11 8

Table 2. Energy of Adsorption for MCl-C2H4 and MCl-CO Systems (M ) Ag or Cu) adsorbent

adsorbate

theoretical ∆H (kcal/mol)

experimental ∆H (kcal/mol)

C2H4 C2H4 CO CO

AgCl CuCl AgCl CuCl

11.20 15.74 9.64 16.56

6.9 8.3 7.5 10.2

Figure 3. Pure-component equilibrium isotherms for C2H4 adsorption on CuCl and AgCl salts at 0 °C. Lines are fittings with eq 2.

Figure 2. Normalized isotherms for CO adsorption on CuCl and AgCl salts at 0 °C. Lines are fittings with eq 2.

The pure-component equilibrium isotherms for CO on CuCl and AgCl at 0 °C are shown in Figure 1. The equilibrium data were fitted to eq 2. Fitting parameters are shown in Table 1. Equilibrium capacities for CO on CuCl and AgCl at 1 atm and 0 °C were measured at 0.040 and 0.021 mmol/g, respectively. On the basis of Figure 1, it can be seen that the affinity of Cu+ cations for CO is greater than that of Ag+ cations. This behavior is also supported by the isosteric heats of adsorption data shown in Table 2. The heats of adsorption for CO on CuCl and AgCl were calculated at 10.2 and 7.5 kcal/ mol, respectively. Another method of reporting the data is to normalize the data with respect to surface area. BET surface area measurements for CuCl and AgCl salts are reported in Table 1. The normalized equilibrium isotherms for CO on CuCl and AgCl at 0 °C are shown in Figure 2. Once the data have been normalized, it seems that the affinity of CuCl on CO is nearly 2.5 times greater than that of AgCl. While CuCl adsorbs CO much more strongly than AgCl does, this behavior does not translate to other adsorbates such as C2H4. The pure-component equilibrium isotherms for C2H4 on CuCl and AgCl at 0 °C are shown in Figure 3. The equilibrium data were fitted to eq 2. Fitting parameters are shown in Table 1. Equilibrium capacities for C2H4 on CuCl and AgCl at 1 atm and 0 °C were measured at 0.024 and 0.015 mmol/g, respectively. Again, the CuCl surface adsorbs C2H4 at a significantly higher rate than AgCl does. This behavior is also supported by isosteric heats of adsorption data. The heats of adsorption for C2H4 on CuCl and AgCl were calculated at 8.3 and 6.9 kcal/mol, respectively. The normalized equilibrium iso-

Figure 4. Normalized isotherms for C2H4 adsorption on CuCl and AgCl salts at 0 °C. Lines are fittings with eq 2.

therms for C2H4 on CuCl and AgCl at 0 °C are shown in Figure 4. It can be observed that the difference in adsorption on the CuCl and AgCl surfaces is greater for CO than for C2H4. On the basis of the normalized isotherms (Figures 2 and 4), it can be seen that CO adsorbs, per unit area, almost twice as much as C2H4. This difference cannot simply be attributed to their size differences. The kinetic diameters of CO and C2H4 are reported in the literature as 3.8 and 3.9 Å.34 The difference in adsorption per unit area is more likely due to their adsorption geometry. While C2H4 adsorbs with its π bond parallel to the surface, CO adsorbs with its π bond perpendicular to the surface, as will be shown shortly in the results on geometry optimization. This particular alignment of the π bond allows CO to form very close packing and also to minimize the effects of close-range repulsion generated between neighboring adsorbed molecules. Optimized Geometries. After the geometry optimization of each structure, frequency calculations were

Ind. Eng. Chem. Res., Vol. 38, No. 7, 1999 2723

Figure 5. Optimized geometries of CO-MCl adsorption systems (M ) Ag or Cu): (1) CO and AgCl before adsorption; (2) CO and AgCl after adsorption; (3) CO and CuCl before adsorption; (4) CO and CuCl after adsorption. Bond distances are in angstroms.

Figure 6. Optimized geometries of C2H4-MCl adsorption systems (M ) Ag or Cu): (1) C2H4 and AgCl before adsorption; (2) C2H4 and AgCl after adsorption; (3) C2H4 and CuCl before adsorption; (4) C2H4 and CuCl after adsorption. Bond distances are in angstroms and angles in degrees.

done, and no imaginary frequencies were found on each structure, which means that the structures are true minima. The equilibrium structures as well as the optimized bond parameters of the adsorption systems are shown in Figures 5 and 6. Figure 5 shows the optimized geometries of the C2H4-AgCl and C2H4-CuCl adsorption systems. Geometries of each species before and after adsorption are shown. The optimized geometry for free ethylene is very close to the reported experimental data, with the calculated C-C bond length ) 1.33 Å, C-H bond length ) 1.075 Å, and H-C-H angle ) 116.4° (not shown). The experimental values35 are 1.34 Å, 1.09 Å, and 117.8°, respectively. The equilibrium structure found for both AgCl-C2H4 and CuCl-C2H4 complexes has C2v symmetry, with the metal atom

approaching the π bond along the perpendicular bisector of the C-C bond. Two equal-length metal-carbon bonds were formed with ethylene to give a three-membered ring. Our results of the optimized CuCl-C2H4 are consistent with that of Merchan and co-workers,36 who also found the most stable structure of Cu+-C2H4 to involve the perpendicular approach of Cu+ to the midpoint of the C-C bond, and a value of 2.322 Å was obtained for the optimal distance between the Cu+ and the ethylene plane. A comparison of the bond length of free ethylene with ethylene in the MCl-C2H4 (M ) Ag or Cu) complexes shows that, upon adsorption, the C-C bond length increased from 1.33 to ∼1.35 Å. Also, upon adsorption, the four C-H bonds bent away slightly from the metal, with H-bending angles of 2.6° and 2.74° for AgCl-C2H4 and CuCl-C2H4, respectively. This is not surprising because other bent transition metal-ethylene complexes are also known.37,38 Comparing the geometries of the AgCl-C2H4 and CuCl-C2H4 systems, we can see from Figure 6 that after adsorption CuClC2H4 shows a slightly longer C-C bond, shorter M-C bond length, and more H bending than the corresponding AgCl-C2H4. These are consequences of stronger interactions of CuCl with C2H4 than those of AgCl with C2H4. We have also explained this trend in a previous work.39 According to the NBO analysis results obtained in our previous work, the amount of positive charges on CuCl and AgCl (0.644 and 0.625, respectively) follows the order Cu > Ag. Metals with a more positive charge will be a better electron acceptor to form π complexation with olefins. With a more positive charge than Ag+, Cu+ attracts more electrons from the olefin. This means a stronger σ donation from olefin to the vacant s orbital of Cu+; thus, the bond distance between M and C is shorter for Cu. A shorter M-C distance at the same time leads to more overlap of the π* orbitals of olefin with the d orbitals of the metal, which means a stronger d-π* backdonation between M and C. These two factors result in the above geometric differences between CuCl and AgCl with C2H4. Geometry optimizations of CO on AgCl and CuCl were done on both linear and perpendicular structures. It was found that the linear structure is more stable than the corresponding perpendicular one, so our further considerations are based on linear CO-MCl structures only. Optimized geometries of each species in AgClCO and CuCl-CO before and after CO adsorption are shown in Figure 5. The geometries found for both AgClCO and CuCl-CO complexes have linear structures, with the C atom interacting with Ag+ or Cu+ directly. The bond distance of C-Cu is 2.05 Å, being 0.34 Å shorter than that of C-Ag. When the bond distances of C-Cu and C-Ag in CuCl-C2H4 and AgCl-C2H4, shown in Figure 6, are compared, it can be seen that the distance is only 0.27 Å shorter for the Cu-C bond than for the Ag-C bond. This means that the interactions are much greater for the CO-CuCl/AgCl systems than for the C2H4-CuCl/AgCl systems. This is consistent with the fact that CO-CuCl interaction is much stronger than CO-AgCl interaction, as will be discussed in more detail in the next two sections. Bond Energies. The energies of adsorption calculated using eq 1 are shown in Table 2 along with the experimental data. A good comparison between the theoretical values and experimental values is seen. Also, a perfect consistency with the geometry results can be obtained by comparing the differences in the energies

2724 Ind. Eng. Chem. Res., Vol. 38, No. 7, 1999 Table 3. Summary of the NBO Analysis of the π-Complexation between MCl-C2H4 and MCl-CO (M ) Ag or Cu) Cu Electron Population Changes after C2H4 and CO Adsorption CuCl-C2H4 CuCl-CO

4s (σ donation)

3dxy

3dxz

3dyz

3dx2-y2

3dz2

total d orbital electron change (d-π* backdonation)

0.052 0.117

0 0

0 -0.051

-0.060 -0.051

0 0

-0.019 -0.027

-0.079 -0.129

Ag Electron Population Changes after C2H4 and CO Adsorption AgCl-C2H4 AgCl-CO

5s (σ donation)

4dxy

4dxz

4dyz

4dx2-y2

4dz2

total d orbital electron change (d-π* backdonation)

0.061 0.101

0 0

0 -0.021

-0.029 -0.021

0 0

-0.026 -0.037

-0.055 -0.079

of adsorption between the four adsorption systems. It is obvious from the table that CO adsorbed much more strongly on CuCl than on AgCl, with the heat of adsorption being ∼7 kcal/mol larger according to our calculation results, while for the adsorption of C2H4 on these two adsorbates, the difference in the heat of adsorption is only about 4.5 kcal/mol. So, it again shows that the difference in the adsorption of CO on CuCl and AgCl is much larger than that of C2H4 on these two adsorbates. Experimental heats of adsorption of ethylene on monolayer AgNO3-silica gel,10 Ag-exchanged resins,9 and monolayer CuCl-alumina are available. These values are closer to the calculated values in Table 2 and also follow the same trend. NBO Results. The nature of the adsorbent-adsorbate bonding can be better understood by analyzing the NBO results. The main feature of the bonding can be seen from the population changes in the vacant outershell s orbital of the metal (Ag or Cu) and those in the d shells of the metals upon C2H4 or CO adsorption. The NBO analysis of the four adsorption systems, summarized in Table 3, is generally in line with the traditional picture of Dewar40 and Chatt and Duncanson41 for π complexation; i.e., it is dominated by the donation and backdonation contributions. An examination of the table shows that in all cases the M-C interaction is a dative bond, i.e., donation of electron charges from the π orbital of olefin or carbon monoxide to the vacant s orbital of metal and, simultaneously, backdonation of electron charges from the d orbitals of M to the π* orbital of olefin or carbon monoxide. This can be interpreted in more detail as follows. When the adsorbate molecule approaches M+, some electronic charge is transferred from the CdC or CtO π orbital to the valence s orbital of M+; at the same time, electrons in the filled d orbitals of metal are transferred to the symmetry-matched π* orbital of olefin or carbon monoxide. It can be seen from Table 3 that upon adsorption the electron occupancies of the valence s orbitals of Cu and Ag always increase, whereas the total occupancy of their 4d or 5d orbitals always decreases. This is caused by the donation and backdonation of electrons between metal and olefin or carbon monoxide. A comparison of the electron population changes in the s and d orbitals of M before and after adsorption shows that for the CuCl-CO and CuCl-C2H4 complexes the overall charge transfer is backdonation; i.e., the amount of electron backdonation is more than that of the donation, while for AgCl-CO and AgCl-C2H4 systems, the amount of backdonation is less than that of the forward donation. The experimental data given in Table 2 show that the heats of adsorption for CuClCO and CuCl-C2H4 are greater than those for the AgCl-CO and AgCl-C2H4 systems. This indicates that in the covalent bonds between the adsorbents and

adsorbates depend mainly on the overlap of the metal d orbitals with the C hybrid orbitals for the systems studied here, similar to the case with acetylene.42 This is particularly the case for bondings of CO and ethylene with Cu+. Acknowledgment This work was supported by National Science Foundation Grant CTS-9520328. Nomenclature AE ) all-electron bp ) Langmuir constant DFT ) density functional theory E ) energy ECP ) effective core potential HF ) Hartree-Fock NAO ) natural atomic orbital NBO ) natural bond orbital NHO ) natural hybrid orbital NLMO ) natural localized molecular orbital NPA ) natural population analysis P ) pressure q ) equilibrium amount adsorbed qm ) monolayer or saturated amount adsorbed RECP ) relativistic effective core potential RHF ) restricted Hartree-Fock SCF ) self-consistent field

Literature Cited (1) King, C. J. Separation Processes Based on Reversible Chemical Complexation. In Handbook of Separation Process Technology; Rousseau, R. W., Ed.; Wiley: New York, 1987; Chapter 15. (2) Kohl, A. L.; Reisenfeld, F. C. Gas Purification, 3rd ed.; Gulf Publishing: Houston, TX, 1979. (3) Haase, D. J.; Walker, D. G. The COSORB Process. Chem Eng. Prog. 1974, 70 (5), 74. (4) Golden, T. C.; Guro, D. E.; Kratz, W. C.; Sabram, T. E. APCI’s CO VPSA: Design, Scale-up and Plant Performance. In Fundamentals of Adsorption; Meunier, F., Ed.; Elsevier: Amsterdam, The Netherlands, 1998. (5) Tamon, H.; Kitamura, K.; Okazaki, M. Adsorption of Carbon Monoxide on Activated Carbon Impregnated with Metal Halide. AIChE J. 1996, 42, 422. (6) Keller, G. E.; Marcinkowsky, A. E.; Verma, S. K.; Williamson, K. D. Olefin Recovery and Purification via Silver Complexation. In Separation and Purification Technology; Li, N. N., Calo, J. M., Eds.; Marcel Dekker: New York, 1992. (7) Humphrey, J. L.; Seibert, A. F.; Koort, R. A. Separation TechnologiessAdvances and Priorities. DOE-ID-12920-1, 1991. (8) Safarik, D. J.; Eldridge, R. B. Olefin/Paraffin Separations by reactive Absorption: A Review. Ind. Eng. Chem. Res. 1998, 37, 2571. (9) Yang, R. T.; Kikkinides, E. S. New Sorbents for OlefinParaffin Separations by Adsorption via π-Complexation. AIChE J. 1995, 41, 509.

Ind. Eng. Chem. Res., Vol. 38, No. 7, 1999 2725 (10) Rege, S. U.; Padin, J.; Yang, R. T. Olefin-Paraffin Separations by Adsorption: Equilibrium Separation by π-Complexation vs Kinetic Separation. AIChE J. 1998, 44, 799. (11) Padin, J.; Yang, R. T. New Sorbents for Olefin-Paraffin Separations by π-Complexation: 1. Synthesis and Effects of Anions and Substrates. Chem. Eng. Sci. 1999. (12) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Gill, P. M. W.; Johnson, B. G.; Robs, M. A.; Cheeseman, J. R.; Keith, T.; Petersson, G. A.; Montgomery, J. A.; Raghavachari, K.; Al-Laham, M. A.; Zakrzewski, V. G.; Ortiz, J. V.; Foresman, J. B.; Peng, C. Y.; Ayala, P. Y.; Chen, W.; Wong, M. W.; Andres, J. L.; Replogle, E. S.; Gomperts, R.; Martin, R. L.; Fox, D. J.; Binkley, J. S.; DeFrees, D. J.; Baker, J.; Stewart, J. P.; Head-Gordon, M.; Gonzalez, C.; Pople, J. A. Gaussian 94, Revision B.3; Gaussian, Inc.: Pittsburgh, PA, 1995. (13) Hay, P. J.; Wadt, W. R. Ab Initio Effective Core Potential for Molecular Calculations. Potentials for the Transition Metal Atoms Sc to Hg. J. Chem. Phys. 1985, 82, 270. (14) Wadt, W. R.; Hay, P. J. Ab Initio Effective Core Potential for Molecular Calculations: Potentials for Main Group Elements Na to Bi. J. Chem. Phys. 1985, 82, 284. (15) Ziegler, T. Approximate Density Functional Theory as a Practical Tool in Molecular Energetics and Dynamics. Chem. Rev. 1991, 91, 651. (16) Becke, A. D. A New Mixing of Hartree-Fock and Local Density Functional Theories. J. Chem. Phys. 1993, 98, 1372. (17) Barone, V.; Adamo, C. J. Phys. Chem. 1996, 100, 335. (18) Halthausen, M. C.; Heinemann, C.; Cornhl, H. H.; Koch, W.; Schwarz, H. The Performance of Density-Functional/HartreeFock Hybrid Methods: Catalytic Transition-Metal Methyl Complexes MCH3+ (M ) Sc-Cu, La, Hf-Au). J. Chem. Phys. 1995, 102, 4931. (19) Ricca, A.; Bauschlicher, C. W. Successive Binding Energies of Fe(CO)5+. J. Phys. Chem. 1994, 98, 12899. (20) Becke, A. D. Density Functional Thermochemistry. III. The Role of Exact Exchange. J. Chem. Phys. 1993, 98, 5648. (21) Becke, A. D. Phys. Rev. 1988, B38, 3098. (22) Lee, C.; Yang, W.; Parr, R. G. Development of the ColleSalvetti Correlation-Energy Formula into a Functional of the Electron Density. Phys. Rev. 1988, B37, 785. (23) Basch, H.; Topiol, S. Relativistic Effects in Ab Initio Effective Core Potential Studies of Heavy Metal Compounds. Applications to HgCl2, AuCl2 and PtH. J. Chem. Phys. 1979, 71, 802. (24) Kahn, L. R.; Hay, P. J.; Cowan, R. D. Relativistic Effects in Ab Initio Effective Core Potentials for Molecular Calculations. Applications to the Uranium Atom. J. Chem. Phys. 1978, 68, 2386. (25) Lee, Y. S.; Ermler, W. C.; Pitzer, K. S. Ab Initio Effective Core Potentials Including Relativistic Effects. I. Formalism and Applications to the Xe and Au Atoms. J. Chem. Phys. 1977, 67, 5861. (26) Russo, T. V.; Martin, R. L.; Hay, P. J. Effective Core Potentials for DFT Calculations. J. Phys. Chem. 1995, 99, 17085.

(27) Hay, P. J.; Wadt, W. R. Ab Initio Effective Core Potential for Molecular Calculations: Potentials for K to Au Including the Outermost Core Orbitals. J. Chem. Phys. 1985, 82, 299. (28) Mulliken, R. S.; Ermler, W. C. Diatomic Molecules: Results of Ab Initio Calculations; Academic Press: New York, 1977; p 33. (29) Collins, B.; Streitwieser, A. J. Comput. Chem. 1980, 1, 81. (30) De Profit, F.; Marin, J. M. L.; Geerling, P. On the Performance of Density Functional Methods for Describing Atomic Populations, Dipole Moments and Infrared Intensities. Chem. Phys. Lett. 1996, 250, 393. (31) Luthi, H. P.; Ammeter, J. H.; Almlof, J.; Faegri, K. How Well Does the Hartree-Fock Model Predict Equilibrium Geometries of Transition Metal Complexes? Large-Scale LCAO-SCF Studies on Ferrocene and Decamethylferrocene. J. Chem. Phys. 1982, 77, 2002. (32) Reed, A. E.; Weinstock, R. B.; Weinhold, F. Natural Population Analysis. J. Chem. Phys. 1985, 83 (2), 735. (33) Glendening, E. D.; Reed, A. E.; Carpenter, J. E.; Weinhold, F. NBO Version 3.1; Chemistry Department, University of Wisconsin: Madison, WI, 1995. (34) Breck, D. W. Zeolite Molecular Sieves: Structure, Chemistry, and Use; John Wiley & Sons: New York, 1974. (35) Hehre, W. J.; Radom, L.; Schleyer, P. V. R.; Pople, J. A. Ab initio Molecular Theory; Wiley: New York, 1986. (36) Merchan, M.; Gonzalez-Luque, R.; Nebot-Gil, I.; Tomas, F. A. Pseudopotential Ab Initio Molecular Orbital Study of the Copper (I)-Ethylene System. Chem. Phys. Lett. 1984, 112, 412. (37) Basch, H.; Newton, M. D.; Moskowitz, J. W. The Electronic Structure of Ni- and Ni2-Ethylene Cluster Complexes. J. Chem. Phys. 1978, 69, 584. (38) Ozin, G. A.; Power, W. J.; Upton, T. H.; Goddard, W. A., III. Experimental and Theoretical Studies of Nin(C2H4)m, Synthesis, Vibrational and Electronic Spectra, and Generalized Valence Bond-Configuration Interaction Studies. The Metal Atom Chemistry and a Localized Bonding Model for Ethylene Chemisorbed on Bulk Nickel. J. Am. Chem. Soc. 1978, 100, 4750. (39) Huang, H. Y.; Padin, J.; Yang, R. T. Ab Initio Effective Core Potential Study of Olefin/Paraffin Separations by Adsorption via π-Complexation: Anion and Cation Effects on Selective Olefin Adsorption. J. Phys. Chem. 1999, 103, 3206. (40) Dewar, M. J. S. A Review of the π-Complex Theory. Bull. Soc. Chim. Fr. 1951, 18, C71. (41) Chatt, J.; Duncanson, L. A. Olefin Coordination Compounds. Part III. Infrared Spectra and Structure: Attempted Preparation of Acetylene Complexes. J. Chem. Soc. 1953, 2939. (42) Sodupe, M.; Bauschlicher, C. W., Jr. Theoretical Study of the Bond of the First and Second-Row Transition-Metal Positive Ions to Acetylene. J. Phys. Chem. 1991, 95, 8640.

Received for review January 14, 1999 Revised manuscript received April 28, 1999 Accepted April 30, 1999 IE990035B