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Sub-Topics
1
2
3
4
5
6
• Introduction to Transition Metals
• Crystal Field Theory, Spectrochemical Series, Ligand Field Stabilisation Energies
• Classical Complexes
• Molecular Orbital Theory of Bonding
• Jahn – Teller Distortions, Square Planar Complexes
• Magnetism: Spin only formula and Electronic Spectroscopy
Lecture 7
Molecular Orbital Theory
Molecular Orbital Theory
Orbital Overlap
A bond can only be formed when two atomic orbitals
of two atoms overlap
Orbitals must be of similar energy
To simplify the theory in our context, the 4s and 4p
orbitals on the metal atom are ignored.
Recall the different types of donor/acceptor behaviour
Sigma Donor
Ligands point along
axis
Example, Cl-
Sigma Donor
For the Cl- ion:
Electrons are in a p-orbital
Can be donated to form a sigma bond by
overlapping with eg orbitals
Metal based eg raised in energy
No net overlap with t2g
Hence energy of t2g unchanged
Sigma Donor
Molecular Orbital Diagram
Effect of Sigma Donation
What happens as strength of sigma donation increases?
Energy of interaction inter-action
Splitting between the bonding and anti-bonding orbitals gets
larger
eg moves higher in energy
Ligand field splitting (Δo) increases
Pi donor
Certain p-orbitals can
interact with t2g to give
pi bonds
Overlap is possible
Example Cl-
No net overlap with eg
sets
Pi donor
Molecular Orbital Diagram
Effect of Pi Donation
What happens when Pi Donor interactions increases?
Energy of inter-actions increases
Splitting between bonding and anti-bonding orbitals increases
t2g moves higher in energy
Ligand field splitting (Δo) decreases
Complete Molecular Orbital Diagram
Combination of sigma and pi donor interactions
Pi acceptor
Metal orbitals involved are the same as pi donor
Symmetry of atomic orbital overlap determines
sigma or pi interaction
Only interacts with t2g
Pi acceptor
Requirements
Ligand must have an empty or partially filled
orbital of pi symmetry
Accepting electron density
Typically anti-bonding orbitals
Results in Back Donation
Examples, CO and CN-
Pi acceptor
Molecular
Orbital
Diagram
Pi acceptors
increases Δo
Lecture 8
Effects of Ligand Field Splitting
The following properties of transition metal
complexes are effected under Ligan d Field Splitting
(ie introduction of ligands)
Thermochemical Data
Hydration Energies of the M(H2O)62+ ions
Redox potentials for M3+/M2+
Lattice Enthalpies
Ionic Radii
Coordination Geometries
Effects of Ligand Field Splitting
Hydration Energy
What factors effect the hydration energy?
Zeff, Ionic Radius, Ligand Field Effects?
In each case we are going to deal with the hexaaquo ion M(H2O)62+
Ligand Field Symmetry: Octahedral
Calculate Ligand Field Stabilisation Energies – reflects the double hump!
Electrode Potentials for M3+/M2+
Ligands alter the redox potential
Example Fe3+/Fe2+couple
Ligand
(o-phentanthroline)3
(H2O)6
E (V)
+1.10
+0.71
Ligand field effects alter the redox potential
Differences in Δo for water and other ligand
Changes in spin state and hence LFSE
Our Favourite Friend Again…
Energy Cycle
Determining change in energy through
introduction of a ligand
Electrode Potentials
Ligand changes alter redox potential
Stabilises one oxidation state over another!
LFSE is a small component in Δgcomplex
Not to do with spin state as remains same
Ligand
Spin State
(o-phenanthroline)3 Both Low Spin
(H2O)6
Both High Spin
Electrode Potentials
LFSE plays a part
Ligand
Δo / cm-1
o-phen
20000
H2O
11000
Complex more stable with o-phen through LFSE
Metal-Ligand bond energies favour M2+ o-phen
Ionic Radii
What do we already know
Decrease as you go across row () as Zeff
increases
Electrons are placed in t2g, non-bonding orbitals
Why increase to Mn2+?
Filling of eg orbitals are of anti-bonding character
Lattice Enthalpies
Sub-Topics
1
2
3
4
5
6
• Introduction to Transition Metals
• Crystal Field Theory, Spectrochemical Series, Ligand Field Stabilisation Energies
• Classical Complexes
• Molecular Orbital Theory of Bonding
• Jahn – Teller Distortions, Square Planar Complexes
• Magnetism: Spin only formula and Electronic Spectroscopy
Lecture 8
General Principles
High Co-ordination seen at beginning of
transition series
Metal atoms have larger radii
Fewer electrons making it easier to accept electrons
in sigma donation
SIX CO-ORDINATION
Low Co-ordination seen at end of transition
series
Metal ions are smaller
Rich in d-electrons
NOT SENSIBLE TO ASSIGN COORDINATION DUE
TO LIGAND FIELD EFFECTS
Ligand Field Effects
Jahn-Teller Distortions
Cr and Cu did not give regular octahedral
environments in metal oxides MO (M2+)
Co-ordination Symmetries
Ni, Pd, Pt either triad, square planar, or
tetrahedral
Jahn-Teller Distortions
If ground electronic state of a non-linear complex is
orbitally degenerate the complex will distort to
remove the degeneracy
When do you get Jahn-Teller Distortions
Odd number of electrons in the eg or t2g other than 3
in the t2g
Notes
Nature of distortion is not defined
Axial lengthening through to square planar
Structural distortions more evident with odd number
of eg
Antibonding Orbitals
Tetragonal Distortion
How is orbital degeneracy lifted?
Sub-Topics
1
2
3
4
5
6
• Introduction to Transition Metals
• Crystal Field Theory, Spectrochemical Series, Ligand Field Stabilisation Energies
• Classical Complexes
• Molecular Orbital Theory of Bonding
• Jahn – Teller Distortions, Square Planar Complexes
• Magnetism: Spin only formula and Electronic Spectroscopy
Magnetism
Magnetic properties of a complex can tell us:
• Electronic State of the transition metal
• Co-ordination geometry
Free atom or ion
Consists of orbital angular momentum of the electron
L, and the spin angular momentum of the electron S.
Complex
Orbital angular momentum is quenched!
Spin-only formula
μ = 2 𝑆(𝑆 + 1)
μ = 2 𝑁(𝑁 + 2)
Unit: Bohr Magneton B.M or μB
Magnetic Moment of a complex can be interpreted to number of unpaired delectrons it has
Deviations from Spin-Only
Presence of orbital angular momenum
Electron circulation not fully quenched
High spin/Low spin crossover
Two ways this can happen
Change in the ground state
Low spin at low temperature and high spin at high
temperature
Electronic States for high and low spin
close in energy
Therefore thermal population of the two spin
states with temperature
Lecture 9
Electronic Spectra
Selection Rules
Spin
Lowest Intensity
∆𝑆 = 0
Orbital Momentum
∆𝐿 = ∓1
Laporte
𝑢 ↔ 𝑔 𝐴𝑙𝑙𝑜𝑤𝑒𝑑
𝑔 ↔ 𝑔 𝑎𝑛𝑑 𝑢 ↔ 𝑢 𝐹𝑜𝑟𝑏𝑖𝑑𝑑𝑒𝑛
𝑇ℎ𝑒𝑟𝑒𝑓𝑜𝑟𝑒 𝑡2𝑔 ↔ 𝑒𝑔 𝐹𝑜𝑟𝑏𝑖𝑑𝑑𝑒𝑛
Charge Transfer
Highest Intensity
Both spin and orbitally allowed