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Chapter 10 Coordination Chemistry II: Bonding

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Chapter 10 Coordination Chemistry II: Bonding Chapter 10 Coordination Chemistry II: Bonding 10-1 Experimental Evidence for Electronic Structures 10-2 Theories of Electronic Structure 10-3 Ligand Field Theory 10-4 Angular Overlap 10-5 The Jahn-Teller Effect 10-6 Four- and Six-Coordinate Preferences 10-7 Other Shapes “Inorganic Chemistry” Third Ed. Gary L. Miessler, Donald A. Tarr, 2004, Pearson Prentice Hall http://en.wikipedia.org/wiki/Expedia Theories of Electronic Structure; Crystal field theory ∆E = strong field – weak field ∆E > 0 weak field ∆E < 0 strong field Theories of Electronic Structure; Crystal field theory What determine ? Depends on the relative energies of the metal ions and ligand orbitals and on the degree of overlap. Theories of Electronic Structure; Crystal field theory Spectrochemical Series for Metal Ions Oxidation # ↑→ ∆↑ Small size & higher charge Only low spin aqua complex Pt4+ > Ir3+ > Pd4+ > Ru3+ > Rh3+ >Mo3+ > Mn4+ > Co3+ > Fe3+ > V2+ > Fe2+ Co2+ > Ni2+ > Mn2+ Ligand field theory; Molecular orbitals for Octahedral complexes CFT & MO were combined The dx2-y2 and dz2 orbitals can form bonding orbitals with the ligand orbitals, but dxy, dxz, and dyz orbitals cannot form bonding orbitals Ligand field theory; Molecular orbitals for Octahedral complexes The combination of the ligand and metal orbitals (4s, 4px, 4py, 4pz, 3dz2, and 3dx2-y2) form six bonding and six antibonding with a1g, eg, t1u symmetries. The metal T2g orbitals do Electronnot have in bonding appropriateorbitals provide symmetry the - nonbondingpotential energy that holds molecules together Ligand field theory; Orbital Splitting and Electron Spin Strong-field ligand – Ligands whose orbitals interact strongly with the metal orbitals → large ∆o Weak-field ligand. d0~d3 and d8 ~d10 – only one electron configuration possible → no difference in the net spin Strong fields lead to low-spin complexes Weak fields lead to high-spin complexes Ligand field theory; Orbital Splitting and Electron Spin What determine ? Depends on the relative energies of the metal ions and ligand orbitals and on the degree of overlap. Ligand field theory; Orbital Splitting and Electron Spin Spectrochemical Series for Metal Ions Oxidation # ↑→ ∆↑ Small size & higher charge Pt4+ > Ir3+ > Pd4+ > Ru3+ > Rh3+ >Mo3+ > Mn4+ > Co3+ > Fe3+ > V2+ > Fe2+ Co2+ > Ni2+ > Mn2+ Ligand field theory; Ligand field Stabilization Energy Ligand field theory; Orbital Splitting and Electron Spin Orbital configuration of the complex is determined by ∆o, πc, and πe + + In general ∆o for 3 ions is larger than ∆o for 2 ions with the same # of e-. ∆o > π low-spin ∆o < π high-spin For low-spin configuration Require a strong field ligand Ligand field theory; Ligand field Stabilization Energy Ligand field theory; Orbital Splitting and Electron Spin The position of the metal in the periodic table Second and third transition series form low- spin more easily than metals form the first transition series -The greater overlap between the larger 4d and 5d orbitals and the ligand orbitals -A decreased pairing energy due to the larger volume available for electrons Ligand field theory; Pi-Bonding The reducible representation is Ligand field theory; Pi-Bonding LUMO orbitals:can be used for π bonding with metal HOMO Ligand field theory; Pi-Bonding metal-to-ligand π bonding or π back-bonding -Increase stability -Low-spin configuration -Result of transfer of negative charge away from the metal ion Ligand-to metal π bonding -decrease stability -high-spin configuration Ligand field theory; Square planar Complexes; Sigma bonding Ligand field theory; Square planar Complexes; Sigma bonding ll ⊥ e- from metal 16 e- 8 e- Ligand field theory; Tetrahedral Complexes; Sigma bonding The reducible representation is A1 and T2 Ligand field theory; Tetrahedral Complexes; Pi bonding The reducible representation is E, T1 and T2 Angular Overlap LFT → No explicit use of the energy change that results Difficult to use other than octahedral, square planar, tetrahedral. Deal with a variety of possible geometries and with a mixture of ligand. → Angular Overlap Model The strength of interaction between individual ligand orbitals and metal d orbitals based on the overlap between them. Angular Overlap: Sigma-Donor Interactions The strongest σ interaction There are no examples of complexes with e- in the antibonding orbitals from s and p orbitals, and these high-energy antibonding orbitals are not significant in describing the spectra of complexes. → we will not consider them further. Angular Overlap: Sigma-Donor Interactions Angular Overlap: Sigma-Donor Interactions Stabilization is 12eσ Angular Overlap: Pi-Acceptor Interactions The strongest π interaction is considered to be between a metal dxy orbitals and a ligand π* orbital. Because of the overlap for these orbitals is smaller than the σ overlap, eπ < eσ. Angular Overlap: Pi-Acceptor Interactions Angular Overlap: Pi-Acceptor Interactions Angular Overlap: Pi-Donor Interactions In general, in situations involving ligands that can behave as both π acceptors and π donors (such as CO, CN-), the π acceptor nature predominates. Angular Overlap: Pi-Donor Interactions Angular Overlap: Pi-Acceptor Interactions Angular Overlap: Types of the ligands and the spectrochemical series Spectrochemical Series for Ligands - - σ donor only CO > CN > PPh3 > NO2 > phen > bipy > en - 2- NH3 > py > CH3CN > NCS > H2O > C2O4 - - - - - - - OH > RCO2 > F > N3 > NO3 > Cl > SCN S2- > Br- > I- π acceptor (strong field ligand) π donor(weak field ligand) Angular Overlap: Magnitudes of eσ eπ and ∆ Metal and ligand Angular Overlap: Magnitudes of eσ eπ and ∆ Angular overlap parameters derived from electronic spectra eσ is always larger than eπ. overlap The magnitudes of both the σ and π parameters ↓ with isoelectronic ↑ size and ↓ electronegativity of the halide ions. overlap Angular Overlap: Magnitudes of eσ eπ and ∆ Can describe the electronic energy of complexes with different shapes or with combinations of different liagnds. The magnitude of ∆o → Magnetic properties and visible spectrum. Angular Overlap: The Jahn-Teller Effect There cannot be unequal occupation of orbitals with identical orbitals. To avoid such unequal occupation, the molecule distorts so that these orbitals no longer degenerate. In other words, if the ground electron configuration of a nonlinear complex is orbitally degenerate, the complex will distort to remove the degeneracy and achieve a lower energy. Angular Overlap: The Jahn-Teller Effect Angular Overlap: Four- and Six-Coordinate Preference Angular overlap calculations Only σ bonding is considered. Large # of bonds formed in Low-spin square planar the octahedral complexes. Angular Overlap: Four- and Six-Coordinate Preference Angular Overlap: Four- and Six-Coordinate Preference How accurate are these predictions? Their success is variable, because of there are other differences between metals and between ligands. In addition, bond lengths for the same ligand-metal pair depend on the geometry of the complex. The interactions of the s and p orbitals. The formation enthalpy for complexes also becomes more negative with increasing atomic number and increasing ionization energy. By careful selection of ligands, many of the transition metal ions can form compounds with geometries other than octahedral. Angular Overlap: Other shapes 1 1 1 Strength of σ–interaction 1 1 2+3/4 9/8 9/8 0 0 Angular Overlap: Other shapes Trigonal-bipyramidal ML5 (D3h) σ-donor only.
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