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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 # ↑→ ∆↑
Only low spin aqua complex
Small size & higher charge
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 Electron
not havein bonding
orbitals provide
the
appropriate
symmetry
potential energy that holds
- nonbonding
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 lowspin 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
CO > CN- > PPh3 > NO2- > phen > bipy > en σ donor only
NH3 > py > CH3CN > NCS- > H2O > C2O42OH- > RCO2- > F- > N3- > NO3- > Cl- > SCNS2- > Br- > Iπ 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
isoelectronic
The magnitudes of
both the σ and π
parameters ↓ with
↑ 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.
Low-spin square planar
Large # of bonds formed in
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