lect21_reel5_2pg

lect21_reel5_2pg - Molecular Orbital Theory and Charge...

Info iconThis preview shows page 1. Sign up to view the full content.

View Full Document Right Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

Unformatted text preview: Molecular Orbital Theory and Charge Transfer Excitations Chemistry 123 Spring 2008 Dr. Woodward Molecular Orbital Diagram H2 H 1s Orbital H 1s Orbital Energy Antibonding Molecular Orbital (Orbitals interfere destructively) Bonding Molecular Orbital (Orbitals interfere constructively) 1 Orbital overlap – Constructive Interference Bonding Overlap 1s Atom 1 Constructive Interference Wavefunction (∠ ) 1s Atom 2 Between the nuclei the wavefunctions add together. Electron density maximized in the internuclear region. Atom (2) Atom (1) This type of overlap leads to formation of a covalent bond. -4 -3 -2 -1 0 1 2 3 4 z (Angstroms) + = Orbital overlap – Destructive Interference Antibonding Overlap Destructive Interference 1s Atom 1 1s Atom 2 Wavefunction (∪ ) Between the nuclei the wavefunctions cancel each other out. Electron density pushed away from internuclear region. Atom (2) Atom (1) This type of overlap works against formation of a covalent bond. -4 -3 -2 -1 0 1 2 3 4 z (Angstroms) + = 2 Pi (π) Bonding side on overlap + Atom 1 pz orbital head on overlap = Atom 2 pz orbital + Bonding pi molecular orbital Constructive Interference + Atom 1 pz orbital Sigma (σ) Bonding Atom 1 py orbital Atom 2 py orbital + Antibonding pi molecular orbital Destructive Interference Bonding sigma molecular orbital Constructive Interference = Atom 2 pz orbital = Atom 1 py orbital = Atom 2 py orbital Bonding sigma molecular orbital Destructive Interference Principles of MO Theory 1. Conservation of Orbitals: The number of Molecular Orbitals is equal to Conservation the number of Atomic Orbitals. 2. Conservation of Electrons: The number of electrons occupying the molecular orbitals is equal to the sum of the valence electrons on the constituent atoms. 3. Pauli Exclusion Principle: Each MO can hold two electrons of opposite spin. 4. Hunds Rule: When orbitals are degenerate (at the same energy) all electron spins are the same direction (up) until we have to start putting two electrons in the same orbital. 5. Principle of Orbital Mixing: The splitting between bonding and antibonding MO’s decreases as: a. The spatial overlap decreases (due to orientation of the orbitals, interatomic distance, or size of orbitals) b. The orbital electronegativities become different 3 [Cr(NH3)6]3+ Octahedron Cr3+ :NH3 5 d-orbitals on Cr d(Cr3+ = d3 ion) 3 electrons in the d-orbitals :N H HH 6 Ligand Orbitals Nitrogen lone pairs (all containing 2 e-) Only sigma interactions are allowed [Cr(NH3)6]3+ Antibonding (σ*) Metal-Ligand MO’s eg orbitals (dz2, dx2-y2) Δ = Crystal Field Splitting Energy t2g orbitals (dxz, dyz, dxy) Energy Energy Metal (Cr) d-orbitals Nonbonding Metal d MO’s Nonbonding Ligand MO’s Ligand (N) lone-pair orbitals Bonding (σ) Metal-Ligand MO’s 4 Absorption Spectra Cr3+ Solutions Antibonding (σ*) Metal-Ligand MO’s 1.4 [Cr(H2O)6]3+ 1.2 [Cr(OH)4(H2O)2]1- absorbance 1.0 [CrO4]2- Δoct 0.8 0.6 Nonbonding Metal d MO’s 0.4 0.2 0.0 250 350 450 550 650 750 850 wavelength (nm) [CrO4]2- t2 orbitals (more antibonding) e orbitals (antibonding) CT Energy Metal (Cr) d-orbitals Nonbonding Oxygen 2p MO’s e orbitals (bonding) 12 Oxygen 2p orbitals (4 oxygens x 3 p orbitals) t2 orbitals (bonding) 5 Absorption Spectra CrO42- Solutions Antibonding Cr 3d orbitals 1.4 [Cr(H2O)6]3+ 1.2 [Cr(OH)4(H2O)2]1- absorbance 1.0 [CrO4]2- 0.8 CT1 0.6 0.4 0.2 0.0 250 350 450 550 650 750 850 wavelength (nm) Nonbonding Oxygen 2p MO’s Absorption Spectra CrO42- Solutions Antibonding Cr 3d orbitals 1.4 [Cr(H2O)6]3+ 1.2 [Cr(OH)4(H2O)2]1- absorbance 1.0 [CrO4]2- CT2 0.8 0.6 0.4 0.2 0.0 250 350 450 550 wavelength (nm) 650 750 850 Nonbonding Oxygen 2p MO’s 6 Charge Transfer Salts, ACrO4 The absorbance of SrCrO4 is similar to a concentrated solution of CrO42- ions. Charge Transfer Excitations and Charge Periodic Trends We can expect charge transfer transitions when we have a d0 cation in a high oxidation state. How does the charge transfer change as we move around the periodic table? 7 CrO42- vs. MnO42- Antibonding (e) Cr dx2-y2, dz2 CT Antibonding (e) Mn dx2-y2, dz2 CT Nonbonding O 2p Nonbonding O 2p [CrO4]2- [MnO4]- As the cation oxidation state increases [i.e. Cr(VI) → Mn(VII)] d-orbitals become more electronegative (lower in energy) CT Energy Gap decreases Absorption shifts to longer wavelengths 8 SrMoO4 – SrCrO4 Series 100 80 60 2- Reflectance 70 [CrO4] absorbance 90 50 SrCrO4 SrCr0.9Mo0.1O4 SrCr0.8Mo0.2O4 40 SrCr0.5Mo0.5O4 SrCr0.2Mo0.8O4 30 SrCr0.1Mo0.9O4 SrMoO4 20 CrO4(2-) 10 SrMoO4 0 250 350 450 550 650 Wavelength (nm) 750 SrCrO4 Orbital Radii – Group 6 Cr 4s r = 1.63 Å Mo 5s Cr 3d r = 0.46 Å Mo 4d r = 1.75 Å r = 0.73 Å W 6s The d orbitals are always much smaller than the s and p, but the 3d orbitals are particularly small W 5d r = 1.65 Å r = 0.78 Å 9 Antibonding (e) Mo dx2-y2, dz2 Antibonding (e) Cr dx2-y2, dz2 CT CT Nonbonding O 2p Nonbonding O 2p [CrO4]2- [MoO4]2- Mo 4d orbitals are larger than the Cr 3d orbitals d-orbitals interact more with O 2p orbitals – more antibonding CT Energy Gap increases Absorption shifts to shorter wavelengths 2nd & 3rd Row Transition Metals eg (σ*) 2nd and 3rd row transition metals •d-orbitals are larger •Metal-ligand antibonding interactions are stronger •eg (σ*) orbitals are more antibonding •Low spin configurations are always observed [Co(H2O)6]3+ Δ = 2.25 eV [Rh(H2O)6]3+ Δ = 4.23 eV 10 ...
View Full Document

This note was uploaded on 06/03/2011 for the course CHM 123 taught by Professor Woodward during the Spring '08 term at Ohio State.

Ask a homework question - tutors are online