Oxidation state (also called oxidation number) is chemistry's accounting system for electrons. It assigns a hypothetical charge to each atom in a compound by assuming all bonds are ionic — that is, all shared electrons belong to the more electronegative atom. This might sound artificial, but it is incredibly useful for balancing redox reactions, predicting compound formulas, and understanding how electrons flow during chemical changes. Think of it as a way to keep track of who "owns" the electrons in any given situation.
Here is how it works in practice. In water (H₂O), oxygen is more electronegative, so it is assigned both shared electrons from each O-H bond, giving it an oxidation state of -2. Each hydrogen loses its electron in this accounting, giving it +1. The numbers add up to zero (2 times +1 plus -2 = 0), which must be the case for any neutral compound. In ions, the oxidation states sum to the ion's charge. For the sulfate ion (SO₄²⁻), sulfur has an oxidation state of +6 while each oxygen is -2, totaling -2 overall.
Oxidation states become especially interesting with transition metals, which can display multiple oxidation states. Iron, for example, commonly appears as Fe²⁺ (ferrous, as in FeO) or Fe³⁺ (ferric, as in Fe₂O₃). This is why iron forms two different types of rust with different colors — black rust contains Fe²⁺ and red-brown rust contains Fe³⁺. Manganese is even more versatile, showing oxidation states from +2 to +7, each producing compounds with dramatically different colors. When an atom's oxidation state increases, we say it is oxidized (it lost electrons); when it decreases, it is reduced (it gained electrons). The mnemonic "OIL RIG" — Oxidation Is Loss, Reduction Is Gain — has helped generations of students remember this.