MA

Topic 10 - Redox Reactions (1)

Redox Reactions Overview

  • Redox reactions involve both oxidation and reduction processes where electrons are transferred between substances.

Definitions

  • Oxidation:

    • Gain of oxygen or loss of electrons.

    • For example, a substance that transitions from a lower to higher oxidation state is oxidized.

  • Reduction:

    • Loss of oxygen or gain of electrons.

    • A substance that transitions from a higher to lower oxidation state is reduced.

  • Redox Reaction: A chemical reaction in which both oxidation and reduction occur simultaneously.

  • Oxidizing Agent: A substance that oxidizes another substance by accepting electrons and undergoing reduction.

  • Reducing Agent: A substance that reduces another by donating electrons and undergoing oxidation.

Oxidation Numbers

  • Oxidation Number: A number assigned to an element in a compound representing the number of electrons lost or gained.

    • Rules:

      1. An element alone has an oxidation number of 0.

      2. In compounds, the sum of oxidation states equals 0; in polyatomic ions, it equals the charge of the ion.

  • Analytical methods can be used to determine oxidation states, as in:

    • Cl2 = 0

    • O in -2 state (like in H2O).

Oxidation and Reduction Examples

  • Zinc and Copper Reaction (

    Zn(s) + CuO(s) → ZnO(s) + Cu(s)

    • Zn is oxidized (losing electrons), CuO is reduced (gaining electrons).

  • Iron Reduction Reaction

    • Fe2O3 + 3C → 2Fe + 3CO

    • Fe is reduced, carbon is oxidized.

Key Equations

  • Ionic Equation: An equation that shows only the ions involved in a reaction, omitting spectator ions (ions that remain unchanged).

Identifying Agents in Reactions

  • Carbon reduces oxygen in its reaction with O2 to create CO2 (reducing agent).

  • Sodium reduces chlorine in its reaction with Cl2 to create NaCl (reducing agent).

  • Iron(III) oxide acts as an oxidizing agent, oxidizing aluminum to Al³ while being reduced itself.

  • Sulfate ions ([SO₄]²⁻) do not participate in the redox process and act as spectator ions.

Role of Electrons in Redox Reactions

  • Oxidation States and Their Changes:

    • Oxidation increases oxidation state (e.g., Zn in ZnO).

    • Reduction decreases oxidation state (e.g., O in H2O).

Reactive Groups and Their States

  • Group 1 Metals: In reactions, they increase their oxidation state from 0 to +1.

  • Group 2 Metals: Increase from 0 to +2.

  • Group 3 Metals: Increase from 0 to +3.

  • Group 5 Elements: Have variable oxidation states.

  • Fluorine: Gains an electron (oxidation state changes from 0 to -1).

Naming Compounds Using Oxidation States

  • Typically expressed in Roman numerals following the element name:

    • Example: Iron(III) oxide represents iron at +3 oxidation state.