ACIDS AND BASES: OXIDATION AND REDUCTION 101
(Note that the equation for metal-metal ion systems is a special case
of this general equation since the reduced state is the metal itself
and the concentration of a solid is a constant and omitted from the
equation.)THE EFFECT OF CHANGE OF LIGAND AND pH ON
REDOX POTENTIALSThe data in Tables 4.2 and 4.3 refer to ions in aqueous acid solution;
for cations, this means effectively [M(H 2 O)X]" + species. However,
we have already seen that the hydrated cations of elements such as
aluminium or iron undergo 'hydrolysis' when the pH is increased
(p. 46). We may then assume (correctly), that the redox potential
of the system
Fe^3 + (aq) + e~ -> Fe^2 1aq)
will change with change of pH. In fact, in this example, change of pH
here means a change of ligand since, as the solution becomes more
alkaline, the iron(III) species in solution changes from [Fe(H 2 O) 6 ]^3 +
to [Fe(OH)3(H 2 O) 3 ] (i.e. iron(III) hydroxide). The iron(II) species
changes similarly. The redox half-reaction then becomes
[Fe(OH) 3 (H 2 O) 3 ] + e~ -> [Fe(OH) 2 (H 2 O) 4 ] + OH~
for which E^ is — 0,56 V. compared with E^ = 4- 0,76 V in acid
solution; thus in alkaline conditions, iron(II) becomes a good
reducing agent, i.e. is easily oxidised.
When the water ligands around a cation are replaced by other
ligands which are more strongly attached, the redox potential can
change dramatically, for example for the cobalt(II)-cobalt(III)
system we have(i) [Com(H 2 O) 6 ]^3 + +*--> [Co"(H 2 O) 6 ]^2 + :E^ = + 1.81 V
(ii) [Com(NH 3 ) 6 ]3+(aq) + e~ -> [Con(NH 3 ) 6 ]2+(aq):£^ - +0.1 V
(iii) [Corn(CN) 6 ]^3 -(aq) + e~ -> [Co"(CN) 5 (H 2 O)]^3 -(aq) + CN~:E*= - 0.83 V
Half-reaction (i) means that Co(II) in aqueous solution cannot be
oxidised to Co(III); by adding ammonia to obtain the complexes
in (ii), oxidation is readily achieved by, for example, air. Similarly, by
adding cyanide, the hexacyanocobaltate(II) complex becomes a
sufficiently strong reducing agent to produce hydrogen from water!