Redox Chemistry

Oxidation and reduction happen together: electrons lost by one species are gained by another. In CuO(s) + H₂(g) -> Cu(s) + H₂O(g), copper(II) oxide loses oxygen and is reduced; hydrogen gains oxygen and is oxidised. During addition of hydrogen to an alkene, the organic compound gains hydrogen and is reduced.

Oxidation state is an accounting number. An uncombined element has state 0. A monatomic ion has its ion charge. In the common compounds here, oxygen is usually -2 and hydrogen +1; exceptions must be supplied or recognised from context. The states sum to zero in a neutral compound or to the overall charge in an ion.

An oxidising agent accepts electrons and causes another species to be oxidised. It is reduced. A reducing agent supplies electrons and is oxidised. In the zinc/copper-ion reaction, Cu²⁺ is the oxidising agent and Zn is the reducing agent; copper metal is a product, not the agent added at the start.

For the iodide test, 2I^-(aq) -> I₂(aq) + 2e^- shows iodide losing electrons. For manganate(VII), use the acidified reagent specified: conditions affect its products and colours. A colour change is an observation; the assignment of oxidising or reducing behaviour is the inference.

Electrolysis passes electric current through a molten or dissolved ionic compound, the electrolyte, causing chemical changes at electrodes. In the solid, ions are fixed in a lattice; once molten or dissolved they can move. This supports the ionic model. Electrons move through metal wires; ions carry charge through the electrolyte.

For a molten binary ionic compound, there are only its constituent cations and anions to discharge. The metal forms at the cathode and the non-metal at the anode, with inert electrodes. For molten lead(II) bromide: Pb²⁺ + 2e^- -> Pb and 2Br^- -> Br₂ + 2e^-. Balance atoms and charge in each half-equation.

In aqueous electrolysis, consider the solute ions and the water. With inert electrodes, a metal less reactive than hydrogen, such as copper, is usually deposited at the cathode. For a very reactive metal ion such as Na⁺, hydrogen forms instead. At the anode, sulfate is generally not discharged; oxygen forms from water/hydroxide. Halide concentration matters: concentrated chloride favours chlorine, while dilute chloride solution is treated here as giving oxygen.

To purify copper, use impure copper as the anode, pure copper as the cathode, and aqueous copper(II) sulfate as electrolyte. At the anode, Cu(s) -> Cu²⁺(aq) + 2e^-. At the cathode, Cu²⁺(aq) + 2e^- -> Cu(s). Copper transfers from the impure electrode to the pure one; some impurities collect as anode sludge. Detailed industrial operation is not required.

The anode loses mass and the cathode gains mass. Copper ions removed at the cathode are replenished at the copper anode, so concentration remains approximately constant in the simple model. This contrasts with an inert anode, where oxygen forms and the blue solution fades.

For copper electroplating, make the object to be coated the cathode, use copper as the anode, and choose a solution containing Cu²⁺. Clean the object first so the coating adheres. Electroplating can improve appearance or provide a protective surface; the coating metal and intended use determine the benefit.

A simple cell contains two different electrodes in an electrolyte connected through an external circuit. A more reactive metal tends to lose electrons more readily. In a zinc/copper cell using dilute acid, zinc is oxidised: Zn(s) -> Zn²⁺(aq) + 2e^-. Electrons travel through the wire towards copper, where hydrogen ions can gain them: 2H⁺(aq) + 2e^- -> H₂(g). The electrolyte carries ionic current and completes the circuit.

Hydrogen can be obtained from water by electrolysis or from hydrocarbons. In a hydrogen fuel cell it reacts with oxygen to produce water and electricity directly: 2H₂(g) + O₂(g) -> 2H₂O(l). Water is the point-of-use product, but overall environmental impact depends on how the hydrogen and electricity were produced. Detailed fuel-cell construction is not required.