Higher oxidation states and reactivity

The highest oxidation states show up in fluorides and oxides. Among the halides they are TiX₄, VF₅ and CrF₆. Mn(VII) has no simple halide, but MnO₃F exists. After Mn, the only trihalides are FeX₃ and CoF₃.

Fluorine holds high states through high lattice energy (as in CoF₃) or strong covalent bonds (as in VF₅ and CrF₆). V(V) exists only as VF₅; the other V(V) halides hydrolyse to oxohalides VOX₃. Fluorides are unstable in low states, for example VX₂ (X = Cl, Br, I), and the same holds for CuX.

Every Cu(II) halide is known except the iodide, because Cu²⁺ oxidises iodide: 2Cu²⁺ + 4I⁻ → Cu₂I₂(s) + I₂.

Many Cu(I) compounds disproportionate in water: 2Cu⁺ → Cu²⁺ + Cu. Cu²⁺(aq) wins because its hydration enthalpy is so much more negative that it more than pays for copper's second ionisation enthalpy.

In oxides the top oxidation number equals the group number from Sc₂O₃ to Mn₂O₇. Iron goes no higher than Fe₂O₃, though ferrate(VI), FeO₄²⁻, forms in alkali and soon decomposes to Fe₂O₃ and O₂. Oxocations hold high states too: VO₂⁺ for V(V), VO²⁺ for V(IV), TiO²⁺ for Ti(IV).

Oxygen holds high oxidation states even more firmly than fluorine does, because it can form multiple bonds to the metal: the top fluoride of Mn is MnF₄ but the top oxide is Mn₂O₇, in which each Mn is tetrahedrally surrounded by O with one Mn-O-Mn bridge. Tetrahedral [MO₄]ⁿ⁻ ions are known for V(V), Cr(VI), Mn(V), Mn(VI) and Mn(VII).

Oxidising power rises in the order VO₂⁺ < Cr₂O₇²⁻ < MnO₄⁻, because the lower species they are reduced to become more stable in that order.

Copper aside, the 3d metals are fairly reactive, and 1 M H⁺ oxidises them, sometimes slowly. At room temperature Ti and V resist dilute acids that are not oxidising. The E°(M²⁺/M) values show that the drive to form M²⁺ weakens across the row.

Mn³⁺ and Co³⁺ are the strongest oxidising agents among the aqueous M³⁺ ions. Ti²⁺, V²⁺ and Cr²⁺ are strong reducing agents that release hydrogen from dilute acid: 2Cr²⁺ + 2H⁺ → 2Cr³⁺ + H₂.