6. Electrochemistry

Oxidation number is the charge an atom would have if bonding electrons were assigned to the more electronegative atom. It is a bookkeeping tool for identifying oxidation and reduction.

Use rules such as elemental substances = 0, monoatomic ions = ionic charge, oxygen usually = −2 and hydrogen usually = +1, then make the total match the species charge.

In H₂SO₄, 2(+1)+S+4(−2)=0, so sulfur has oxidation number +6. In sulfate, the same bookkeeping must total −2.

Oxidation number is not always a real ionic charge. It is an assigned value that helps compare electron loss and gain in covalent compounds.

Oxidation is an increase in oxidation number; reduction is a decrease. The two changes occur together because electrons lost by one species are gained by another.

Assign oxidation numbers before and after the reaction, then identify which element changes. This works even when no free electrons appear in the equation.

In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc changes 0→+2 and is oxidised; copper changes +2→0 and is reduced.

Oxidation does not always mean adding oxygen, and reduction does not always mean removing oxygen. Use the oxidation-number change as the general test.

A redox reaction contains both oxidation and reduction. The oxidised species loses electrons and the reduced species gains electrons.

Use “oxidation” and “reduction” for the changes, then identify the oxidising agent and reducing agent from what each species does.

In Fe²⁺ + Ce⁴⁺ → Fe³⁺ + Ce³⁺, Fe²⁺ is oxidised and acts as the reducing agent; Ce⁴⁺ is reduced and acts as the oxidising agent.

The oxidising agent is itself reduced, while the reducing agent is itself oxidised. Do not assign the labels from the words alone.

An oxidising agent causes another species to lose electrons and is reduced itself. A reducing agent causes another species to gain electrons and is oxidised itself.

Write the electron change for the agent, not just the reaction partner. Stronger agents are better electron acceptors or donors under the stated conditions.

In Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, chlorine accepts electrons and is the oxidising agent; bromide donates electrons and is the reducing agent.

The agent’s name describes what it does to the other species, not the change it undergoes. The oxidising agent does not undergo oxidation.

A Roman numeral in a compound name gives the oxidation number of the named metal ion, especially when the element can form more than one charge.

Use the anion charge and overall neutrality to check the numeral. The numeral is not the number of atoms of that metal in the formula.

Iron(III) chloride contains Fe³⁺ and three Cl⁻ ions, so its formula is FeCl₃. Iron(II) chloride is FeCl₂.

Do not read “iron(III)” as three iron atoms. It specifies the oxidation state of each iron ion in the compound.