Ionisation enthalpies

Ionisation enthalpy rises from left to right along each series because the nuclear charge grows as the inner d set fills.

The rise is gentle: successive ionisation enthalpies of transition elements do not climb as steeply as those of main-group elements, and the spread across a series is much smaller than across a main-group period. First values for Sc to Cu lie between 631 and 762 kJ mol⁻¹; Zn is 906.

When d-block atoms form ions, the ns electrons leave before the (n−1)d electrons. So Fe (3d⁶ 4s²) becomes Fe²⁺ 3d⁶, and every M²⁺ and M³⁺ ion of the row is a pure dⁿ ion.

The irregular first ionisation enthalpies come from the way removing one electron shifts the relative energies of 4s and 3d. The 3d electrons shield the 4s electrons fairly well, so radii and first ionisation enthalpies change only a little.

Three terms set an ionisation enthalpy: attraction to the nucleus, electron-electron repulsion and exchange energy. Exchange energy grows with the number of parallel-spin pairs in a degenerate set and stabilises it, so a configuration that would lose exchange energy on ionisation holds its electron more tightly.

The steady rise breaks at d⁵: the second ionisation enthalpy of Mn (making Mn²⁺) and the third of Fe (making Fe³⁺). Mn⁺ is 3d⁵ 4s¹ while Cr⁺ is d⁵, so Mn's second value (1509) is below Cr's (1592); Fe²⁺ is d⁶ while Mn²⁺ is d⁵, so Fe's third value (2962) is below Mn's (3260).

Second ionisation enthalpies are unusually high for Cr (1592) and Cu (1958), where M⁺ is d⁵ and d¹⁰. Zn's is comparatively low (1734), because it removes a 4s electron and leaves a stable d¹⁰ ion.

Third ionisation enthalpies are generally high; the very high values for Ni, Cu and Zn explain why these metals rarely go above +2.

Ionisation enthalpies are a guide to the stability of oxidation states, but the full picture is complex and does not reduce to a simple rule.