Labile and inert metal ions - Kinetic effects

Water exchange rate constants (s-1) for selected metal centers

Cr3+ Co3+

Li+ Na+K+

Ca2+ Sr2+Mg2+

Al3+ Ti3+Fe3+ V3+

Cd2+ Hg2+Zn2+Pd2+Pt2+

Ru2+ Cu2+Cr2+Ni2+ Mn2+Co2+ Fe2+

V2+

10-6 10-4 10-2 100 102104 106 108 1010

Approximate half-lives for exchange of water molecules from the first coordination sphere of metal ions at 25 oC

Metal ion t1/2 , sec Metal ion t1/2 , sec 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+ 2 x 10-9

K+ 7 x 10-10 Mn2+ 3 x 10-8 Al3+ 0.7

Mg2+ 1 x 10-6 Fe2+ 2 x 10-7 Fe3+ 4 x 10-3

Ca2+ 2 x 10-9 Co2+ 2 x 10-7 Cr3+ 3 x 105

Ba2+ 3 x 10-10 Ni2+ 2 x 10-5 Co3+ 7 x 105

Cu2+ 7 x 10-10

Zn2+ 3 x 10-8

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)

2

4

6

8

10

Mn Fe Co Ni Cu Zn

log

K1

en

gly

ox

mal

OO

OO

_

_

mal

NH2

NH2

en

NH2

OO_

gly

O

O

O

O _

_ox

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+ Hhydration

Spontaneous loss of degeneracy of eg and t2g orbitals

for certain dn configurations

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

Octahedral Tetragonal

Cu

Cl

ClCl

Cl

Cl

Cl

2.3 Ao

2.9 Ao

2.9 Ao

2.3 Ao

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 forthe reduction process

G = nEo F = - 2.3 RT logK

The redox potential of a metal ion couple (Mnn+/M(n-1)+) represents the relative stabilityof the metal when in its oxidized and reduced states.

The redox potential for a metal ion couple will be dependent on the nature ofthe ligands coordinated to the metal.

Comparison of redox potentials for a metal ion in different ligand environments providesinformation on factors influencing the stability of metal centers.

The effect of ligand structure on the reduction potential (Eo

red) 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)3

2+ Eo = 1.14 V

Fe(H2O)63+ + e Fe(H2O)6

2+ Eo = 0.77 V

Fe(CN)63 + e Fe(CN)6

4 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.

Fe3+ + e- Fe2+

Eo (V)

0.771

Fe3+ + 3e- Fe -0.040

Fe2+ + 2e- Fe -0.44

Fe3+ Fe2+Fe

0.771

-0.040

-0.44

Cu2+ + e- Cu+ 0.15

Cu2+ + 2e- Cu 0.34

Cu+ + e- Cu 0.52

Cu2+ Cu+Cu

0.15

0.34

0.52

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

Fe3+ + e- Fe2+

Eo (V)

0.771

Fe3+ + 3e- Fe -0.040

Fe2+ + 2e- Fe -0.44

Fe3+ Fe2+Fe

0.771

-0.040

-0.44

Cu2+ + e- Cu+ 0.15

Cu2+ + 2e- Cu 0.34

Cu+ + e- Cu 0.52

Cu2+ Cu+Cu

0.15

0.34

0.52

MnO4- + H+ + e- 0.90HMnO4

-

HMnO4 + 3H+ + 3e- MnO2 2.10

MnO2 + 2H+ + e- Mn3+ + H2O 0.95

Mn3+ + e- Mn2+1.54

Mn2++ 2e- Mn -1.19

MnO4- HMnO4

- MnO2 Mn3+ Mn2+ Mn0.90 2.10 0.95 1.54 -1.19

1.69 1.23

1.51

MnO4- + 8H+ + 5e- 1.51Mn2++ 4 H2O

O2 + 4 H+ + 4 e- 2 H2O

)Qlog(0591.0o

nEΕ

)]H[

1log(

4

0591.04

O

o

2

p

EΕ

)]H[

1log(

4

0591.023.1

4Ε

)]H[

1log(4)

4

0591.0(23.1 Ε

HΕ p0591.023.1

Dependence of Reduction Potential on pH

E = 0.82 V (pH 7)

Eo = 1.23 V (1.0 M H+)

2 H+ + 2 e- H2 Eo = 0.00 V (1.0

M H+)

)Qlog(0591.0o

nEΕ

)]H[

1log(

2

0591.000.0

2Ε

)]H[

1log()2(

2

0591.0Ε

HΕ p0591.0

E = -0.413 V (pH 7)