oxidizing and reducing agents

Every redox reaction is an electron handover between two partners, and the names just describe who plays which role. An oxidizing agent is the electron thief: it takes electrons from something else, and in doing so it is itself reduced. A reducing agent is the electron donor: it gives electrons away, and is itself oxidized. The slightly dizzying part is that the agent suffers the opposite fate to the name it bears — the oxidizing agent gets reduced, the reducing agent gets oxidized — because the agent is doing the verb to its partner, not to itself.

Picture chlorine gas dropped into a solution of bromide ions. Chlorine is a strong oxidizing agent: each Cl2 molecule grabs two electrons (one onto each chlorine, going from 0 to -1), and in giving up those electrons the bromide is oxidized to bromine. The mnemonic OIL RIG keeps it straight — Oxidation Is Loss, Reduction Is Gain (of electrons). The agent that causes the loss is the oxidizing agent; the agent that causes the gain (by being the source) is the reducing agent.

Knowing which substances are strong oxidizers or reducers is the heart of practical redox chemistry. Fluorine, O2, MnO4-, Cr2O7^2-, and Ce4+ are powerful oxidants; the alkali metals, hydrogen, carbon, and reactive metals like zinc and magnesium are good reductants. The strength is not a fixed personality trait, though — it depends on conditions. MnO4- is a ferocious oxidant in acid but much milder in base, and a metal that reduces water can be inert toward a weaker oxidant. Standard reduction potentials are simply a numerical ranking of this oxidizing power.