By Marie-Luise Blue
Updated Aug 30 2022
Image credit: Leonid Eremeychuk/iStock/GettyImages
The charge of a transition‑metal ion reflects the electrons it has shed during a chemical reaction. Determining that charge requires knowledge of the element, the charges of the surrounding atoms, and the overall molecular charge. All oxidation numbers are integers, and the sum of the atomic charges equals the net charge of the species.
When an atom loses electrons, chemists refer to the process as oxidation. For transition metals, the oxidation state—and thus the ionic charge—can range from +1 to +7. These elements possess partially filled d‑orbitals that make electron loss more facile than in main‑group elements. Some oxidation states are inherently more stable, so they appear more frequently. For example, iron (Fe) can adopt +2, +3, +4, +5, or +6 states, but +2 and +3 dominate in natural and industrial contexts. In chemical formulas, the oxidation state is indicated by a Roman numeral in parentheses (e.g., iron(II) oxide, FeO, where Fe carries a +2 charge).
In a neutral compound, the total charge is zero. Knowing the oxidation state of the ligand atoms lets you solve for the metal’s charge. For instance, in MnCl₂ the two chloride ions each carry –1. The combined –2 charge forces manganese to be +2 to maintain neutrality.
Transition‑metal ions often form complex ions that are either positively or negatively charged. Take the permanganate ion, MnO₄⁻: each oxygen has an oxidation state of –2, giving a total of –8 from four oxygens. The overall –1 charge means manganese must be +7.
Most neutral, soluble transition‑metal salts in water have oxidation states of +3 or lower. Higher oxidation states usually precipitate or hydrolyze to form oxygen‑containing complexes. For example, vanadium(V) salts hydrolyze to produce the hexaaquavanadate(IV) ion, [V(OH)₆]⁺, or the aquavanadate(V) ion, [VO₄]⁻, depending on the environment.
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