Standard Electrode Potentials of Transition Metals

Redox reactions always involve the transfer of electrons from one reacting species to another, and the two halves of any redox reaction are inescapably coupled: whenever one substance is oxidised (loses electrons), some other substance in the same reaction must simultaneously be reduced (gain those electrons). The substance that is itself oxidised, by donating electrons to the other species, is termed the reducing agent; the substance that is itself reduced, by accepting those electrons, is termed the oxidising agent. The relative oxidising and reducing power of a given element or ion is quantified in terms of its standard electrode potential.

The standard electrode potential, E⁰, of a half-cell is formally defined as the standard emf (electromotive force) of a complete cell in which molecular hydrogen gas, under standard pressure (1 atm) and at standard temperature (commonly quoted as 273 K, i.e. 0°C, in this reference framework), is oxidised to solvated protons at the standard hydrogen electrode -- the universal zero-point reference against which every other half-cell's potential is measured.

If a metal's standard electrode potential (specifically the reduction potential, as conventionally tabulated) is large and negative, that metal is a powerful reducing agent: a large negative E⁰ means the metal loses electrons (is oxidised to its cation) very readily, i.e. the reverse (oxidation) reaction is strongly thermodynamically favoured. Working through the tabulated M²⁺/M reduction potentials for the first-series transition metals (Figure 4.7(a)): the values start quite negative at titanium (E⁰ = -1.63 V) and, allowing for real element-to-element irregularity, trend toward less negative values moving across to copper, which alone among this set shows a positive reduction potential (E⁰ = +0.34 V). This means elemental copper metal is thermodynamically more stable than the Cu²⁺ ion under standard conditions -- copper genuinely resists being oxidised by dilute non-oxidising acids, which is why copper (unlike, say, zinc or iron) does not dissolve in dilute HCl or dilute H₂SO₄ to liberate hydrogen gas.

The general rising (less-negative) trend from Ti to Cu shows two clear deviations from smoothness: the E⁰(M²⁺/M) values for manganese and for zinc are BOTH more negative than the smooth trend line connecting their neighbours would predict. This is explained, once again, by the extra thermodynamic stability that arises from a half-filled d⁵ configuration in the Mn²⁺ ion, and from a completely filled d¹⁰ configuration in the Zn²⁺ ion -- in both cases, that extra stability of the +2 ion makes it thermodynamically more costly (more negative E⁰, i.e. a stronger drive toward the reduced/metallic side) to reduce that ion all the way back to the metal, compared with the smooth trend.

Transition metals in unusually high oxidation states characteristically tend to act as oxidising agents. A concrete worked example: Fe³⁺ is a moderately strong oxidant, and it is capable of oxidising metallic copper to Cu²⁺ ions. The thermodynamic feasibility of this specific reaction is predicted directly by comparing the two relevant standard electrode potentials: Fe³⁺(aq) + e⁻ ⇌ Fe²⁺, E⁰ = +0.77 V, against Cu²⁺(aq) + 2e⁻ ⇌ Cu(s), E⁰ = +0.34 V. Because the Fe³⁺/Fe²⁺ couple has the more positive (higher) reduction potential, Fe³⁺ is thermodynamically capable of oxidising Cu to Cu²⁺, with Fe³⁺ itself being reduced to Fe²⁺ in the process -- the reaction proceeds spontaneously as written under standard conditions.

The standard electrode potential quoted specifically for the M³⁺/M²⁺ half-cell (as distinct from the M²⁺/M half-cell already discussed) gives the relative stability between a metal's +3 and +2 oxidation states directly. Tabulating these values for six 3d metals (Figure 4.7(b)): the NEGATIVE values found for titanium (E⁰ = -0.37 V), vanadium (E⁰ = -0.26 V) and chromium (E⁰ = -0.41 V) indicate that, for these three elements, the higher (+3) oxidation state is thermodynamically preferred over the lower (+2) state -- a negative E⁰ for the M³⁺/M²⁺ reduction couple means the reduction (gaining an electron to go from +3 down to +2) is not thermodynamically favourable, so the +3 form persists. Consequently, if one wants to reduce a comparatively 'stable' ion such as Cr³⁺ down to Cr²⁺, a strong reducing agent is required -- one with a suitably large negative reduction potential of its own, such as metallic zinc (E⁰ = -0.76 V), which is strong enough to overcome chromium's resistance to reduction. …