Labile and inert metal ions - Kinetic effects

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

3. 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+ Fe2+ Co2+ Ni2+ Cu2+ Zn 2+ dnd5d6d7d8d9d10 LFSE (o) 0 2/5 4/5 6/5 3/5 0

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 orbitalsfor 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) Fe3+ + e Fe2+ It is related to the free energy change and the redox equilibrium constant for the reduction process G =  nDEo F = - 2.3 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)63++ eFe(H2O)62+Eo = 0.77 V Fe(CN)63+ eFe(CN)64Eo = 0.36 V Heme(Fe3+)+eHeme(Fe2+) Eo = 0.17 V Fe(III)cyt-c + e- Fe(II)cyt-cEo = 0.126 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 O2 + 4 H+ + 4 e- 2 H2O Eo = 1.23 V (1.0 M H+) E = 0.82 V (pH 7)

20. 2 H+ + 2 e- H2Eo = 0.00 V (1.0 M H+) E = -0.413 V (pH 7)