Redox Reaction in Terms of Electron Transfer

The classical, substance-based definitions of Section 6.1.1 can all be re-derived from a single, more fundamental idea: oxidation is loss of electrons, and reduction is gain of electrons. Consider magnesium burning, written with charges shown: Mg(s) + 1/2 O2(g) → Mg²⁺ + O²⁻. The neutral magnesium atom loses two electrons to become Mg²⁺, while the elemental oxygen gains those two electrons to become O²⁻; the same pattern holds for Mg(s) + F2(g) → Mg²⁺ + 2F⁻. Each of these single-direction electron changes -- a loss, or a gain -- is called a half reaction, and adding an oxidation half reaction to a reduction half reaction gives the overall redox reaction. A cleaner pair of half reactions is Fe(s) → Fe²⁺(aq) + 2e⁻ (an oxidation half reaction: Fe acts as the reducing agent here) and Cu²⁺(aq) + 2e⁻ → Cu(s) (a reduction half reaction: Cu²⁺ acts as the oxidising agent, accepting the electrons Fe gives up); summing them gives the redox reaction Fe(s) + Cu²⁺(aq) → Fe²⁺(aq) + Cu(s). In this electron language, an oxidant/oxidising agent is redefined simply as an electron acceptor, and a reductant/reducing agent as an electron donor -- a tidier, more general restatement of the classical roles. Displacement reactions -- where an ion or atom in a compound is replaced by an ion or atom of a different element, general form X + YZ → XZ + Y -- are then identified as one whole class of redox reaction, since replacing one metal ion with …