Oxidation-reduction reactions
Redox reactions transfer electrons. They power batteries, rust iron, run our cells, and drive much of industry.
Oxidation and reduction
Oxidation is loss of electrons; reduction is gain of electrons (remember 'OIL RIG': Oxidation Is Loss, Reduction Is Gain). They always occur together: electrons lost by one species are gained by another. In \(\text{Zn} + \text{Cu}^{2+} \to \text{Zn}^{2+} + \text{Cu}\), zinc is oxidised and copper ions are reduced.
Oxidation numbers
Oxidation numbers track electrons. Rules: elements alone are 0; monatomic ions equal their charge; oxygen is usually \(-2\); hydrogen is usually \(+1\); the sum equals the overall charge. In \(\text{SO}_4^{2-}\), sulfur is \(+6\), since \(x + 4(-2) = -2\). An increase in oxidation number means oxidation.
Oxidising and reducing agents
The oxidising agent gets reduced, causing another species to be oxidised; the reducing agent gets oxidised. Oxygen, chlorine, and permanganate are strong oxidising agents; metals like zinc and sodium are good reducing agents.
Balancing redox equations
In acidic solution, use half-reactions: write oxidation and reduction separately, balance atoms other than O and H, add \(\text{H}_2\text{O}\) for O and \(\text{H}^+\) for H, add electrons to balance charge, multiply so electrons cancel, and add. This method handles complicated redox reactions systematically.
Galvanic cells
A galvanic (voltaic) cell separates the half-reactions so electrons flow through a wire, producing electricity. Oxidation occurs at the anode and reduction at the cathode. A salt bridge completes the circuit. The Daniell cell, with zinc and copper electrodes, produces about 1.10 V.
Standard reduction potentials
Each half-reaction has a standard reduction potential \(E^\circ\). The cell voltage is \(E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}\); a positive value means the reaction is spontaneous. Copper (\(+0.34\) V) and zinc (\(-0.76\) V) give \(0.34 - (-0.76) = 1.10\) V.
E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}Electrolysis and corrosion
Electrolysis uses electricity to drive non-spontaneous redox reactions, such as splitting water or electroplating. Corrosion is unwanted redox: iron rusts when oxidised by oxygen in the presence of water. Galvanising coats iron with zinc, which corrodes first and protects the iron.
Worked example
Assign oxidation numbers in \(\text{KMnO}_4\). K is \(+1\), each O is \(-2\), so \(1 + x - 8 = 0\) gives Mn \(= +7\). Permanganate's high oxidation state makes it a powerful oxidising agent, used in water treatment.
Key takeaways
Oxidation is loss of electrons, reduction is gain, and they occur together. Oxidation numbers track electron transfer. Balance redox reactions with half-reactions. Galvanic cells turn spontaneous redox into electricity, with \(E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}\). Electrolysis and corrosion are redox too.
Flashcard
Oxidation Is Loss, Reduction Is Gain
Check your understanding
In the reaction Zn + Cu^(2+) -> Zn^(2+) + Cu, which species is oxidised?