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

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+

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)

Mn2+ Fe2+ Co2+ Ni2+ Cu2+ Zn 2+ dn d5 d6 d7 d8 d9 d10 LFSE (o) 0 2/5 4/5 6/5 3/5 0

Ligand field stabilization energy (LFSE)

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

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.

Jahn-Teller effect in crystalline CuCl2 lattices

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

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.

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

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-c Eo = 0.126 V

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.

Latimer Diagrams

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

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)

2 H+ + 2 e- H2 Eo = 0.00 V (1.0 M H+) E = -0.413 V (pH 7)