Redox

1. Oxidation Numbers and Their Representation

Oxidation number (oxidation state): A numerical value assigned to an atom in a compound or ion that represents its degree of oxidation (loss of electrons) or reduction (gain of electrons).
Roman numeral notation:
- When an element can exist in multiple oxidation states, the oxidation number is shown using a Roman numeral in parentheses after the element’s name.
- Example: Iron(II) sulfate → Fe²⁺ oxidation state is +2. Iron(III) chloride → Fe³⁺ oxidation state is +3.
Key points for writing oxidation numbers:
- Positive oxidation numbers are written without a sign (e.g., +2 written as II).
- Negative oxidation numbers are rare in naming but are considered in calculations.

A redox reaction involves simultaneous oxidation and reduction.
- Oxidation and reduction occur together because electrons lost by one species must be gained by another.

In terms of oxygen: Gain of oxygen atoms in a substance.
- Example: 2Mg + O₂ → 2MgO (Magnesium is oxidised).
In terms of electrons: Loss of electrons.
- Example: Na → Na⁺ + e⁻ (Sodium loses an electron).
In terms of oxidation number: An increase in oxidation number.
- Example: Fe²⁺ → Fe³⁺ (+2 → +3 is oxidation).

In terms of oxygen: Loss of oxygen atoms from a substance.
- Example: CuO + H₂ → Cu + H₂O (Copper(II) oxide is reduced to copper).
In terms of electrons: Gain of electrons.
- Example: Cl₂ + 2e⁻ → 2Cl⁻ (Chlorine gains electrons).
In terms of oxidation number: A decrease in oxidation number.
- Example: Fe³⁺ → Fe²⁺ (+3 → +2 is reduction).

If a substance gains oxygen or loses electrons, it is oxidised.
If a substance loses oxygen or gains electrons, it is reduced.
In a redox reaction, one species is oxidised and another is reduced in the same reaction.

Uncombined elements: Oxidation number is 0.
- Example: Na, Cl₂, O₂.
Monatomic ions: Oxidation number equals the ion’s charge.
- Example: Na⁺ → +1, Cl⁻ → -1.
Compounds: The sum of all oxidation numbers equals 0.
- Example: H₂O → H is +1, O is -2, sum is 0.
Polyatomic ions: The sum of oxidation numbers equals the ion’s charge.
- Example: SO₄²⁻ → S is +6, O is -2 (4 × -2 = -8, total -8 + 6 = -2).

Acidified aqueous potassium manganate(VII), KMnO₄:
- In acidic solution, MnO₄⁻ (purple) is reduced to Mn²⁺ (colourless/pale pink).
- Example: Fe²⁺ + MnO₄⁻ → Fe³⁺ + Mn²⁺.
Aqueous potassium iodide, KI:
- Iodide ions, I⁻, are oxidised to iodine, I₂ (yellow/brown colour).
- Example: 2I⁻ → I₂ + 2e⁻.

Definition: A substance that oxidises another substance by accepting electrons and is itself reduced.
Examples: KMnO₄, K₂Cr₂O₇, Cl₂, O₂.

Definition: A substance that reduces another substance by donating electrons and is itself oxidised.
Examples: Zn, Fe²⁺, H₂, CO.

Example: Zn + CuSO₄ → ZnSO₄ + Cu
- Zn: Oxidised (Zn → Zn²⁺ + 2e⁻) → Reducing agent.
- Cu²⁺: Reduced (Cu²⁺ + 2e⁻ → Cu) → Oxidising agent.