Redox Reactions — Notes

Classical view: oxidation is addition of oxygen or an electronegative element, or removal of hydrogen or an electropositive element; reduction is the reverse. Electron view: oxidation is loss of electrons, reduction is gain of electrons (remember: OIL RIG).

The oxidising agent accepts electrons and is itself reduced; the reducing agent donates electrons and is itself oxidised. Every redox reaction is a pair of half-reactions; electrons lost equal electrons gained. For $\ce{Zn + Cu^2+ -> Zn^2+ + Cu}$: oxidation $\ce{Zn -> Zn^2+ + 2e-}$, reduction $\ce{Cu^2+ + 2e- -> Cu}$.

The oxidation number is the hypothetical charge on an atom if all bonding electrons went to the more electronegative atom. Rules: (1) atoms in an element: 0; (2) simple ion: its charge; (3) the sum for a species equals its charge; (4) group 1 is +1, group 2 is +2, fluorine is −1; (5) hydrogen is +1 (−1 in metal hydrides); (6) oxygen is −2, but −1 in peroxides, −½ in superoxides, +2 in $\ce{OF2}$.

A fractional ON is an average (Fe in $\ce{Fe3O4}$ is +8/3; S in $\ce{S4O6^2-}$ is +5/2). ON is not the same as valence: ON has a sign, and can be zero or fractional. Stock notation gives the metal's ON in Roman numerals: iron(III) oxide, $\ce{Fe2O3}$.

A reaction is redox if any oxidation number changes. The species containing the atom whose ON rises is oxidised (the reducing agent); the species containing the atom whose ON falls is reduced (the oxidising agent). Neutralisation and precipitation reactions are not redox.

An element at its highest ON can only act as an oxidant ($\ce{MnO4-}$, $\ce{Cr2O7^2-}$, $\ce{ClO4-}$); at its lowest only as a reductant ($\ce{S^2-}$, $\ce{I-}$, $\ce{H2S}$); at an intermediate ON as either ($\ce{H2O2}$, $\ce{SO2}$, $\ce{HNO2}$).

• Combination: $\ce{N2 + 3H2 -> 2NH3}$ (redox only if a free element is involved).
• Decomposition: $\ce{2KClO3 -> 2KCl + 3O2}$ (redox if an element is formed; $\ce{CaCO3 -> CaO + CO2}$ is not).
• Displacement: a more reactive metal displaces a less reactive one ($\ce{Zn + CuSO4 -> ZnSO4 + Cu}$) or hydrogen from an acid; a more reactive halogen displaces a less reactive one ($\ce{Cl2 + 2KBr -> 2KCl + Br2}$). Activity series: K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au.
• Disproportionation: the same element is both oxidised and reduced: $\ce{2H2O2 -> 2H2O + O2}$; $\ce{Cl2 + 2OH- -> Cl- + ClO- + H2O}$ (cold, dilute alkali); $\ce{P4 + 3OH- + 3H2O -> PH3 + 3H2PO2-}$.

In a Daniell cell the zinc anode is oxidised, the copper cathode is reduced, and electrons flow from zinc to copper through the wire; a salt bridge completes the circuit.

Oxidation-number method: write the skeleton, assign ON, find the total rise and fall, multiply to equalise, balance the other atoms, O and H last. Example: $\ce{Fe2O3 + 3CO -> 2Fe + 3CO2}$.

Half-reaction (ion-electron) method: write two half-reactions; balance (i) atoms other than O and H, (ii) O with $\ce{H2O}$, (iii) H with $\ce{H+}$, (iv) charge with $\ce{e-}$; multiply to equalise electrons and add. In basic medium, add $\ce{OH-}$ equal to the $\ce{H+}$ on both sides and cancel water.

Acidic: $\ce{MnO4- + 8H+ + 5Fe^2+ -> Mn^2+ + 5Fe^3+ + 4H2O}$. Basic: $\ce{2MnO4- + 6I- + 4H2O -> 2MnO2 + 3I2 + 8OH-}$. Always check atoms and charge.

In acid, $\ce{MnO4- + 8H+ + 5e- -> Mn^2+ + 4H2O}$, so 1 mol $\ce{MnO4-}$ reacts with 5 mol $\ce{Fe^2+}$, with 5/2 mol $\ce{H2O2}$, and 5/2 mol $\ce{C2O4^2-}$ ($\ce{2MnO4- + 5C2O4^2- + 16H+ -> 2Mn^2+ + 10CO2 + 8H2O}$). Permanganate is self-indicating (end point: first permanent pale pink). Acidify with dilute sulfuric acid: hydrochloric acid is oxidised by permanganate and nitric acid is itself an oxidant.

Method: moles of titrant = molarity × volume (L); moles of analyte = moles of titrant × ratio from the balanced equation; molarity = moles ÷ volume of sample (L).