1. Labile and inert metal ions - Kinetic effects
Water exchange rate constants (s-1) for selected metal centers

2. Approximate half-lives for exchange of water molecules from the first coordination sphere of metal ions at 25 oC
Metal ion
t1/2 , sec
Li+
2 x 10-9
V2+
9 x 10-3
Sn2+
< 7 x 10-5
Na+
1 x 10-9
Cr2+
7 x 10-10
Hg2+ K+
Mn2+
3 x 10-8
Al3+
0.7
Mg2+
1 x 10-6
Fe2+
2 x 10-7
Fe3+
4 x 10-3
Ca2+
Co2+
Cr3+
3 x 105
Ba2+
3 x 10-10
Ni2+
2 x 10-5
Co3+
7 x 105
Cu2+
Zn2+

3. The Irving-Williams Series.
Relative Stability of 3d Transition Metal Complexes
The Irving-Williams Series.
The stability order of complexes formed by divalent
3d transition metal ions.
Mn2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+
M2+ + L ↔ ML2+ (K1)

4. Mn2+ Fe Co Ni Cu Zn 2+
dn d5 d6 d7 d8 d9 d10
LFSE (o) 0 2/5 4/5 6/5 3/

5. Ligand field stabilization energy (LFSE)

7. M2+(g) + nH2O [M(H2O)6]2+ DHhydration

8. Jahn-Teller Effect
Spontaneous loss of degeneracy of eg and t2g orbitals for certain dn configurations
Octahedral
Tetragonal
Some metal ions (e.g. Cu(II), d9 and Cr(II), high-spin d4) attain enhanced electronic stability when they adopt a tetragonally distorted Oh geometry rather than a regular Oh geometry. They therefore undergo a spontaneous tetragonal distortion (Jahn-Teller effect). The net stabilization of the eg electrons for Cu(II), is shown above.

9. Jahn-Teller effect in crystalline CuCl2 lattices

10. Electronic spectrum of Ti3+ (d1)
Dynamic Jahn-Teller effect in electronic excited state of d1 ion

11. Redox Potentials of Metal Complexes
A redox potential reflects the thermodynamic driving force for reduction.
Ox e Red Eo (Reduction potential)
Fe e Fe2+
It is related to the free energy change and the redox equilibrium constant for
the reduction process
G =  nDEo F = RT logK
The redox potential of a metal ion couple (Mnn+/M(n-1)+) represents the relative stability
of the metal when in its oxidized and reduced states.
The redox potential for a metal ion couple will be dependent on the nature of
the ligands coordinated to the metal.
Comparison of redox potentials for a metal ion in different ligand environments provides
information on factors influencing the stability of metal centers.

12. The effect of ligand structure on the reduction potential (Eored) of a metal couple
Ligands the stabilize the higher oxidized state lower Eo (inhibit reduction)
Ligands that stabilize the lower reduced state increase Eo (promote reduction)
Ligands that destabilize the oxidized state raise Eo (promote reduction)
Ligands that destabilize the reduced form decrease Eo (inhibit reduction)
Hard (electronegative) ligands stabilize the higher oxidation state
Soft ligands stabilize the lower oxidation state
Negatively charged ligands stabilize the higher oxidation state

13. Fe(phen)33+ + e Fe(phen)32+ Eo = 1.14 V
Fe(H2O) e Fe(H2O)62+ Eo = V
Fe(CN)63 + e Fe(CN)64 Eo = V
Heme(Fe3+) + e Heme(Fe2+) Eo = V
Fe(III)cyt-c + e Fe(II)cyt-c Eo = V

14. Soft 1,10-phenanthroline stabilizes Fe in the softer lower Fe(II) state - i.e. it provides greater driving force for reduction of Fe(III) to Fe(II)
Hard oxygen in H2O favors the harder Fe(III) state. - resulting in a lower driving force for reduction of Fe(III) to Fe(II)
Negatively charged CN- favors the higher Fe(III) oxidation state (hard - hard interaction) - i.e. it provides a lower driving force for reduction.

15. Latimer Diagrams

16. Changes in free energy are additive, but Eo values are not.
If ΔGo(3) = ΔGo(1) + ΔGo(2),
since ΔGo = − nEoF,
n3 (Eo)3F = n1(Eo)1F + n2(Eo)2F,
and hence
(Eo)3 = n1(Eo)1 + n2(Eo)2 n3

19. Dependence of Reduction Potential on pH
O H e H2O
Eo = 1.23 V (1.0 M H+)
E = 0.82 V (pH 7)

20. 2 H e H2 Eo = 0.00 V (1.0 M H+)
E = V (pH 7)