Coordination Chemistry(Part-I)

1. CY 101: CRYSTAL FIELD THEORY
Course Instructor: Dr. SAMYA BANERJEE

2. Books to read
 Inorganic Chemistry: Principles of Structure and Reactivity,
Pearson, by J. E. Huheey, E. A. Keiter, R. L. Keiter and O. K. Medhi
 Concise Inorganic Chemistry: Fifth Edition by J.D. Lee
 Shriver and Atkins' Inorganic Chemistry by P. Atkins, C. H.
Langford, D. F Shriver
 Introduction to Coordination Chemistry by G. A. Lawrance
 Coordination chemistry, by J. R. Gispert

3. Why Important

4. Why Important
Metalloenzymes

5. Why Important
 Carbon monoxide poisoning
 Cyanide poisoning

6. Why Important
Catalysis
Drugs
cisplatin

7. Double Salt and Coordination Complexes
For example: KCl + MgCl2 + 6H2O  KCI·MgCl2·6H2O [Carnallite;
KMgCl3·6(H2O)]
K2SO4 + Al2(SO4)3 + 24H2O  K2SO4·Al2(SO4)3·24H2O (Potash-alum)
CuSO4 + 4NH3 + H2O  CuSO4·4NH3·H2O (tetrammine copper(II)
sulphate monohydrate)
Fe(CN)2 + 4KCN  Fe(CN)2·4KCN (potassium ferrocyanide)
Addition compounds are of two types:
I. Those which dissociate (lose their identity) in solution (Double
salts)
2. Those which retain their identity in solution (Complexes)

8. Alfred Werner
Switzerland
University of Zurich
Zurich, Switzerland
Birth: 1866
(in Mulhouse, then Germany)
Death: 1919
The Nobel Prize in Chemistry 1913
"in recognition of his work on the
linkage of atoms in molecules by
which he has thrown new light on
earlier investigations and opened up
new fields of research especially in
inorganic chemistry"
Alfred Werner: father of coordination complexes

9. 9
In all of these complexes there is no free NH3
(No reaction with acid)
Werner solved the puzzle
Same metal, same ligands, different number of ions when
dissolved

10. Werner’s explanation of coordination complexes
 Metal ions exhibit two kinds of valence: primary and
secondary valences
 The primary valence is the oxidation number (positive
charge) of the metal
 The secondary valence is the number of atoms that
are directly bonded (coordinated) to the metal
 The secondary valence is also termed
the “coordination number” of the metal in
a coordination complex

11. Ligand
 Neutral molecule or ion having a electron pair that
can be used to form a bond to a metal ion
 Monodentate ligand – one bond to a metal ion
 Bidentate ligand (chelate) – two bonds to a metal
ion
 Polydentate ligand – more than two bonds to a
metal ion
 Donor atom – atom in the ligand directly
bonded to the metal atom

13. Polydentate Ligands
• Some ligands have two or more donor
atoms.
• These are called polydentate ligands or
chelating agents.
• In ethylenediamine, NH2CH2CH2NH2,
represented here as en, each N is a
donor atom.
• Therefore, en is bidentate.
Quiz: Me2S

14. Ethylenediaminetetraacetate (EDTA),
has six donor atoms.
Polydentate Ligands
Chelate effect: Chelating agents form more stable complexes than
monodentate ligands.
Reading Assignment: Bridging ligand, π acid type ligand,

15. Bonding in Metal Complexes
(1)Valence bond theory: Developed by Linus Pauling
(2) Crystal field theory
(3) Molecular orbital theory
Hans Bethe
Nobel Prize in 1967
John Hasbrouck Van Vleck
Nobel Prize in 1977

16. 1. Magnetism
2. Color
3. Distorted geometry of octahedral complexes
Drawbacks of Valence bond theory

17. • Focuses on the effect of ligands on the energies of the d orbitals
of metals.
• Assumptions
1. Ligands are negative point charges/dipoles.
2. No orbital overlap is considered.
3. Metal–ligand bonding is entirely ionic:
• strong-field (low–spin):
large splitting of d orbitals
• weak-field (high–spin):
small splitting of d orbitals
4. Ligands approach towards the metal ion along a definite
direction.
Crystal Field Theory

18. Shapes of the d-orbitals

20. Octahedral Ligand Field

21. Ligand negative charge
Is repelled by d electrons, d orbital
energy goes up
Octahedral Ligand Field

22. Electron configurations of some octahedral complexes

23. Octahedral vs Spherically symmetrical

24. Rather than referring to the
energy level of an isolated
metal atom, the weighted
mean of these two sets of
perturbed orbitals is taken
as the zero: this is
sometimes called the Bari
centre.
The difference in energy
between the two d levels
is given either of the
symbols o or 10 Dq.
The eg orbitals are +0.6o
above the average level,
and the t2g orbitals are
-0.4 o below the average
(Figure 7.9).
Octahedral Ligand Field

25. Magnitude of o
The size of the energy gap o between the t2g and eg levels
can be measured easily by recording the UV-visible spectrum
of the complex.
Consider a complex like [Ti(H2O)6]3+. The Ti3+ ion has one d
electron. In the complex this will occupy the orbital with the
lowest energy, that is one of the t2g orbitals.
The complex absorbs light
of the correct wavelength
(energy) to promote the
electron from the t2g level to
the eg level.

26. Magnitude of o
The electronic spectrum for [Ti(H2O)]3+ is given in Figure 7 .11. The steep
part of the curve from 27,000 to 30,000cm-1 (in the UV region) is due to
charge transfer. The d-d transition is the single broad peak with a
maximum at 20,300cm-1 . Since 1 kJmol-1 = 83.7cm-1, the value of o for
[Ti(H2O)]3+ is 20,300/83.7 =
243kJ mol-1 . This is much the same
as the energy of many normal
single bonds.
Solutions containing the hydrated Ti3+
ion are reddish violet coloured.
This is because yellow and green light
are absorbed to excite the electron.
Thus the transmitted light is the
complementary colour red-violet.

27. Factors affecting the magnitude of o
The magnitude of o depends on three factors:
(1) The nature of the ligands
(2) The charge on the metal ion
(3) Whether the metal is in the first, second or third row f transition elements
(4) Geometry of the metal coordination.
_______________________________________________________________
(1) The nature of ligands:
Spectrochemical series

28. (2) The charge on the metal ion
(3) Whether the metal is in the first, second or third row f
transition elements
o for [Cr(H2O)6]3+ and [Mo(H2O)6]3+ are 17,400 cm-1 and 26,110 cm-1, respectively.
Factors affecting the magnitude of o

29. (4) Geometry of the metal coordination
Factors affecting the magnitude of o

30. Crystal Field Stabilization Energy (CFSE)
Because of the crystal field splitting of d orbitals, the singled electron in
[Ti(H2O)]3+ occupies an energy level 2/5 o below the average energy of
the d orbitals. As a result the complex is more stable. The crystal field
stabilization energy (CFSE) is in this case 2/5 x 243 = 97 kJ mol-1

31. Concept of High- & Low-Spin
In accordance with Hund’s rule of maximum multiplicity

32. Concept of High- & Low-Spin
32
Which
is
preferred?

33. Concept of High- & Low-Spin

34. Concept of High- & Low-Spin

35. Concept of High- & Low-Spin

36. Concept of High- & Low-Spin

37. Concept of High- & Low-Spin

38. 20_459
–
–
–
–
–
–
– –
–
–
dz2 dx2 – y2
dxy dyz
dxz
(a) (b)
Tetrahedral Complexes

39. Tetrahedral Complexes

40. 20_461
E
Free metal ion Complex
dz2
dxy
dxz dyz
dx2 - y2
M z
(b)
Free metal ion Complex
dx2 - y2
dxy
dz2
dxz dyz
M
(a)
x
y
E
Square Planar & Linear Complexes
Approach along x-and y-axes Approach along z-axis

41. Different ligands interact more or less, change E spacing
Of D orbitals.

42. Spectrochemical series (strength of ligand interaction)
Cl- < F- < H2O < NH3 < en < NO2
- < CN-
Increasing 
Increasing 

43. As Energy difference increases, electron configuration
changes
“High spin”
“Low spin”
Co(III) is d6