Redox

Oxidation & Reduction (Oxygen Definition)

• **Oxidation** is the gain of oxygen by an element or compound.
• **Reduction** is the loss of oxygen from an element or compound.
• A **redox reaction** is one where oxidation and reduction occur simultaneously.
• Example: ZnO + C → Zn + CO – zinc oxide is reduced (loses O), carbon is oxidised (gains O).

Oxidation & Reduction (Electron Definition – Extended)

• **Oxidation** is the loss of electrons; **reduction** is the gain of electrons (OIL RIG).
• Oxidation increases the oxidation number; reduction decreases it.
• Half-equations show electron transfer: e.g., Ag → Ag⁺ + e⁻ (oxidation); O₂ + 4e⁻ → 2O²⁻ (reduction).
• Example: Fe + Cu²⁺ → Fe²⁺ + Cu – Fe loses electrons (oxidised), Cu²⁺ gains electrons (reduced).

Oxidation Numbers & Roman Numerals

• **Oxidation number** (state) shows the charge an atom would have if electrons were completely transferred.
• Roman numerals in names indicate oxidation number: e.g., iron(II) = Fe²⁺, iron(III) = Fe³⁺.
• Rules: uncombined elements = 0; monatomic ion = its charge; sum in compound = 0; sum in ion = charge.
• Group 1 always +1, Group 2 always +2, F always –1, H usually +1, O usually –2.

Identifying Redox Reactions by Colour Changes

• **Potassium manganate(VII)** (KMnO₄) is a purple oxidising agent; with a reducing agent it turns colourless.
• **Potassium iodide** (KI) is a colourless reducing agent; with an oxidising agent it turns red-brown (I₂ formed).
• These colour tests are used to detect oxidising or reducing agents.

Oxidising & Reducing Agents (Extended)

• An **oxidising agent** oxidises another substance and is itself reduced (gains electrons).
• A **reducing agent** reduces another substance and is itself oxidised (loses electrons).
• Common oxidising agents: O₂, Cl₂, KMnO₄, H₂O₂. Common reducing agents: C, H₂, KI.
• Example: CuO + H₂ → Cu + H₂O – CuO is oxidising agent (reduced), H₂ is reducing agent (oxidised).

Displacement Reactions & Redox

• In displacement reactions, the more reactive metal loses electrons (oxidised) and the less reactive metal ion gains electrons (reduced).
• Example: Fe + CuSO₄ → FeSO₄ + Cu – Fe is oxidised, Cu²⁺ is reduced.
• Ionic equation: Fe + Cu²⁺ → Fe²⁺ + Cu.

Half-Equations & Balancing (Extended)

• Half-equations show either oxidation or reduction separately with electrons.
• Oxidation half: atoms/ions lose electrons (e.g., Ni → Ni²⁺ + 2e⁻).
• Reduction half: atoms/ions gain electrons (e.g., Ni³⁺ + e⁻ → Ni²⁺).
• Balance atoms and charge: e.g., O₂ + 4e⁻ → 2O²⁻.

Extraction of Metals & Redox

• Carbon is used to reduce metal oxides (e.g., ZnO + C → Zn + CO).
• The metal oxide is reduced (loses oxygen), carbon is oxidised (gains oxygen).
• Metals below carbon in the reactivity series can be extracted by reduction with carbon.

Bohr model of iron atom showing electron shells. Iron can lose 2 electrons (Fe²⁺) or 3 electrons (Fe³⁺), corresponding to oxidation numbers +2 and +3.

Colour change of potassium manganate(VII) from purple to colourless when added to a reducing agent.

Colour change of potassium iodide from colourless to red-brown when added to an oxidising agent.