Chapter 19 d-Metal complexes: electronic structure and spectra

Chapter 19 d-Metal complexes: electronic
structure and spectra
Electronic structure
19.1 Crystal-field theory
19.2 Ligand-field theory
Electronic-spectra
19.3 electronic spectra of atoms
19.4 electronic spectra of complexes
19.5 Charge-transfer bands
19.6 Selection rules and intensities
19.7 Luminescence

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

Experimental Evidence for Electronic Structures
Thermodynamic Data
Magnetic Susceptibility
Electronic Spectra
Coordination Numbers and Molecular
Shapes

Experimental Evidence for Electronic Structures;
Thermodynamic Data
One of the primary goal of a bonding theory is to
explain the energy of compound.
The energy is openly not determined directly by
experiment.
Thermodynamic measurements of enthalpies and
free energies of reaction are used to compare.
Bonding strength Stability constants(formation constants)

Experimental Evidence for Electronic Structures;
Thermodynamic Data
What is the stability constants
The equilibrium constants for formation of
coordination complex.

Experimental Evidence for Electronic Structures;
Thermodynamic Data
Stability constants
HSAB concepts
Thermodynamic values
Prediction of properties, structures

Experimental Evidence for Electronic Structures;
Thermodynamic Data
HSAB concepts
The gist of this theory is that soft acids react faster and form
stronger bonds with soft bases, whereas hard acids react
faster and form stronger bonds with hard bases, all other factors
being equal.
The classification in the original work was mostly based on
equilibrium constants for reaction of two Lewis bases competing
for a Lewis acid.
Hard acids and hard bases tend to have:
small size
high oxidation state
low polarizability
high electronegativity
energy low-lying HOMO (bases) or energy high-lying LUMO
(acids).

Experimental Evidence for Electronic Structures;
Thermodynamic Data
HSAB concepts

Experimental Evidence for Electronic Structures;
Thermodynamic Data
Chelating Ligands
Entropy Effect
Chelate Effect
Five or six membered ring
en vs methyl amine
Figure in head….
Stability….

Experimental Evidence for Electronic Structures;
Magnetic Susceptibility
The magnetic properties of a coordination
compound can provide indirect evidence of the
orbital energy level.
Hund’s rule the max. # of unpaired e - .
Diamagnetic: all e - paried repelled by a magnetic field
Paramagnetic: all e - paried attracted into a magnetic field
Magnetic Susceptibility: Measuring Magnetism

Experimental Evidence for Electronic Structures;
Magnetic Susceptibility
Magnetic Susceptibility
Gouy method
A sample that is to be tested is suspended
from a balance between the poles of a
magnet. The balance measures the
apparent change in the mass of the
sample as it is repelled or attracted by the
magnetic field.

Experimental Evidence for Electronic Structures;
Magnetic Susceptibility
In physics and applied disciplines such as electrical
engineering, the magnetic susceptibility is the degree of
magnetization of a material in response to an applied
magnetic field.
Electron spin Spin magnetic moment (m s )
Total spin magnetic moment Spin quantum # S (sum of m s )
Isolated oxygen atom 1s 2 2s 2 p 4
S = +1/2 +1/2 +1/2 -1/2 = 1
Electron spin Orbital magnetic moment (m l )
Total orbital magnetic moment Orbital quantum # L (sum of m l )
Max. L for the p 4
L = +1 +0 -1 +1 = 1

Experimental Evidence for Electronic Structures;
Magnetic Susceptibility
Two sources of magnetic moment – spin (S) and Angular (L) motions of electrons
Spin quantum number
Angular momentum quantum number
The equation for the magnetic moment
Contribution from L is small in first transition series
2.00023 ≈ 2

Experimental Evidence for Electronic Structures;
Electronic Spectra
Give a direct evidence of orbital energy level
Give an information for geometry of complexes

Theories of Electronic Structure
Valence bond theory
Crystal field theory
Ligand field theory
Angular overlap method

Theories of Electronic Structure;
Valence bond theory
Hybridization ideas
Octahedral: d 2 sp 3
d orbitals could be 3d or 4d for the first-row
transition metals. (hyperligated, hypoligated)

Theories of Electronic Structure;
Valence bond theory
Fe(III)
Isolated ion; 5 unpaired e -
In O h compound; 1 or 5 unpaired e -
Co(II)
Low spin
Low spin
High spin
High spin

Theories of Electronic Structure;
Crystal field theory
Crystal field theory (CFT) is a model that describes the
electronic structure of transition metal compounds, all of
which can be considered coordination complexes.
CFT successfully accounts for some magnetic properties,
colours, hydration enthalpies, and spinel structures of
transition metal complexes, but it does not attempt to
describe bonding.
CFT was developed by physicists Hans Bethe and John
Hasbrouck van Vleck in the 1930s.
CFT was subsequently combined with molecular orbital
theory to form the more realistic and complex ligand field
theory (LFT), which delivers insight into the process of
chemical bonding in transition metal complexes.

Theories of Electronic Structure;
Crystal field theory
Repulsion between d-orbital electrons and ligand electrons
Splitting of energy levels of d-orbitals

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory
Electrostatic approach
In an Octahedral field of ligand e - pairs; any e -
in them are repelled by the field.
Crystal field stabilization energy (CFSE);
the actual distribution vs the uniform field.
Good for the concept of the repulsion of
orbitals by the ligands but no explanation
for bonding in coordination complexes.

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory

Theories of Electronic Structure;
Crystal field theory
Why are complexes formed in crystal field theory
Crystal Field Stabilization Energy (CFSE)
Or Ligand Field Stabilization Energy (LFSE)
the stabilization of the d orbitals because of
metal-ligand environments

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
Pt 4+ > Ir 3+ > Pd 4+ > Ru 3+ > Rh 3+ >Mo 3+ >
Mn 4+ > Co 3+ > Fe 3+ > V 2+ > Fe 2+
Co 2+ > Ni 2+ > Mn 2+
Only low spin aqua complex

Ligand field theory;
Molecular orbitals for Octahedral complexes
CFT & MO were combined
The d x2-y2 and d z2 orbitals can form bonding orbitals
with the ligand orbitals, but d xy , d xz , and d yz orbitals
cannot form bonding orbitals

Ligand field theory;
Molecular orbitals for Octahedral complexes
The combination of
the ligand and metal
orbitals (4s, 4p x , 4p y ,
4p z , 3d z2 , and 3d x2-y2 )
form six bonding and
six antibonding with
a 1g , e g , t 1u symmetries.
The metal T 2g orbitals
do Electron not have in bonding
appropriate orbitals provide symmetry the
- nonbonding
potential 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.
d 0 ~d 3 and d 8 ~d 10 – 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
Pt 4+ > Ir 3+ > Pd 4+ > Ru 3+ > Rh 3+ >Mo 3+ >
Mn 4+ > Co 3+ > Fe 3+ > V 2+ > Fe 2+
Co 2+ > Ni 2+ > Mn 2+

Ligand field theory;
Ligand field Stabilization Energy

Ligand field theory;
Orbital Splitting and Electron Spin
Orbital configuration of the complex is
determined by ∆ o , c , and e
In general ∆ o for 3 + ions is larger than ∆ o for 2 + ions
with the same # of e - .
∆ o > low-spin
∆ o < 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 bonding with metal
HOMO

Ligand field theory;
Pi-Bonding
metal-to-ligand bonding
or back-bonding
-Increase stability
-Low-spin configuration
-Result of transfer of
negative charge away from
the metal ion
Ligand-to metal bonding
-decrease stability
-high-spin configuration

Ligand field theory;
Square planar Complexes; Sigma bonding

Ligand field theory;
Square planar Complexes; Sigma bonding
ll
⊥
16 e - e - from metal
8 e -

Ligand field theory;
Tetrahedral Complexes; Sigma bonding
The reducible representation is
A 1 and T 2

Ligand field theory;
Tetrahedral Complexes; Pi bonding
The reducible representation is
E, T 1 and T 2

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 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 12e

Angular Overlap:
Pi-Acceptor Interactions
The strongest interaction is considered to
be between a metal d xy orbitals and a ligand *
orbital.
Because of the overlap for these orbitals is
smaller than the overlap, e < e .

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 acceptors and donors (such
as CO, CN - ), the 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 - > PPh 3 > NO 2- > phen > bipy > en
NH 3 > py > CH 3 CN > NCS - > H 2 O > C 2 O 4
2-
OH - > RCO 2- > F - > N 3- > NO 3- > Cl - > SCN -
S 2- > Br - > I -
donor only
acceptor (strong field ligand)
donor(weak field ligand)

Angular Overlap:
Magnitudes of e eand ∆
Metal and ligand

Angular Overlap:
Magnitudes of e eand ∆
Angular overlap
parameters derived
from electronic
spectra
e is always larger
than e . overlap
isoelectronic
The magnitudes of
both the and
parameters with
size and
electronegativity of
the halide ions.
overlap

Angular Overlap:
Magnitudes of e eand ∆
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
Only bonding is considered.
Angular overlap calculations
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 –interaction
1
1
2+3/4 9/8 9/8 0 0

Angular Overlap:
Other shapes
Trigonal-bipyramidal ML 5 (D 3h ) -donor only

Homework
Exercise 10-1~10-11
Problem 2, 6a, 6c,11, 13, 16.