Chemical Reactivity and E⁰ Values

Reactivity varies widely across the series

Transition metals as a class do not behave uniformly toward acids and other oxidising agents. Many of them are sufficiently electropositive to dissolve readily in mineral acids, while a small number are effectively "noble" — largely unaffected by a single acid acting on its own.

Among the first-row metals, every one except copper is relatively reactive and is oxidised by 1 M H⁺, although the rate of that reaction can be slow even when it is thermodynamically favourable. Titanium and vanadium illustrate this gap between thermodynamics and practice: both are found to be passive toward dilute, non-oxidising acids at room temperature, despite what a purely thermodynamic reading of their electrode potentials might suggest.

Reading reactivity off the E° values

The $E^{\ominus}$(M²⁺/M) values track a decreasing tendency, moving across the series, to form the divalent cation — a trend rooted in the rising combined first- and second-ionisation enthalpies from one element to the next. As with the M²⁺/M trend generally, Mn, Ni and Zn sit apart from it with more negative values than expected: the half-filled $3d^5$ configuration of Mn²⁺ and the completely filled $3d^{10}$ configuration of Zn²⁺ account for their departures, while nickel's owes instead to its unusually large, strongly negative hydration enthalpy.

The $E^{\ominus}$(M³⁺/M²⁺) values point to a different, complementary kind of reactivity — not metal-versus-acid, but one ion oxidising or reducing another in solution.

Mn³⁺ and Co³⁺ stand out as the strongest oxidising agents among the aqueous M³⁺ ions of the first-row series.

Ti²⁺, V²⁺ and Cr²⁺ sit at the opposite extreme as strong reducing agents — strong enough to liberate hydrogen gas directly from a dilute acid, for example: $2\text{Cr}^{2+}(aq) + 2\text{H}^+(aq) \rightarrow 2\text{Cr}^{3+}(aq) + \text{H}_2(g)$.

The overall picture …