Study Notes on Oxidation and Reduction

Oxidation Reduction

Basic Definitions

  • Oxidation: Originally defined as combining with oxygen, e.g.,

    • A burning match: $C + O2 → CO2 + ext{heat}$
    • Rusting iron: $Fe + O_2 → ext{Rust (iron oxide)}$

  • Reduction: Therefore meant to remove oxygen.

Modern Definitions

  • Oxidation and Reduction: Have much broader definitions today. To understand these definitions, we need to look at oxidation numbers.

Oxidation Numbers

  • Oxidation Number: The "apparent charge" on an atom. True electric charge might not always be apparent.
  • For binary ionic compounds, such as $CaO$, elements can be separated and assigned definite charges. Example: $Ca^{2+}$, $O^{2-}$.
  • In polar covalent compounds like HCl and SO₂, atoms do not always have an integer charge. Instead, we assign an oxidation number to each atom.
Redox Reactions
  • If oxidation states of elements change in a reaction, it is called a reduction-oxidation reaction (redox reaction).
  • In a redox reaction, electrons are transferred:
    • Oxidation: Loss of electrons (LEO)
    • Reduction: Gain of electrons (GER)
    • Mnemonic: OIL RIG (Oxidation is Loss, Reduction is Gain)
Agents in Redox Reactions
  • Reducing agent: The molecule that caused the reduction.
  • Oxidizing agent: The molecule that caused the oxidation.

Rules for Assigning Oxidation Numbers

  1. A pure element has an oxidation state (OS) of zero (e.g., Na, Br₂, S₂).
  2. The OS of a monatomic ion is its charge (e.g., $Al^{3+}$ has OS = +3, $S^{2-}$ has OS = -2).
  3. Fluorine in compounds has OS = -1.
  4. Oxygen in compounds has OS = -2 (exceptions: $O^{2-}$ has OS = -1 and in $OF_2$ it has OS = +2).
  5. Hydrogen has OS = +1 (except in hydrides, where OS = -1).
  6. The sum of all OS in a polyatomic ion must equal the charge of the ion. The sum of all OS in a neutral compound is zero.

Understanding Oxidation Numbers

  • Oxidation Number Explained: The charge an atom would have if each shared electron were totally transferred to the more electronegative atom.

Examples of Oxidation States

  • Example reaction: $H2 (g) + C{12} (g) → 2HCl (g)$ with oxidation states:

    • H: +1 → Cl: -1 (where added oxidation states are shown)

  • Other example: $S (s) + O_2 (g) → SO₂ (g)$

    • Oxidized State: +4, Reduced State: -2.
Learning Objectives
  • Students must learn the rules for assigning oxidation numbers and become familiar with the most common oxidation states of elements up to argon.
  • Ability to use changes in oxidation numbers to identify the atoms oxidized and reduced, as well as their respective agents.
  • Notice that the oxidizing and reducing agents may be the same compound or element.

Classifications of Redox Reactions

  • Many simple redox reactions can be classified into subclasses:
    • Combination reactions

Steps for Balancing by Half-Reaction Method

Step-by-Step Process
  1. Write an unbalanced net ionic equation, if it is not already given.
  2. Divide the unbalanced net ionic equation into an oxidation half-reaction and a reduction half-reaction. Assign oxidation numbers to all elements to determine which are oxidized and which are reduced.
  3. Balance the oxidation half-reaction and the reduction half-reaction independently.
  4. Determine the least common multiple (LCM) of the numbers of electrons in both half-reactions.
  5. Use coefficients to write each half-reaction with the LCM of electrons.
  6. Add the balanced half-reactions that include equal numbers of electrons.
  7. Remove electrons from both sides of the equation.
  8. Eliminate any identical molecules or ions that are present on both sides of the equation.
  9. Include spectator ions in the chemical formulas if a balanced chemical equation is required.
  10. If necessary, include states of the compounds.
Balancing under Different Conditions
  • Acidic Conditions:
    • Balance atoms other than O or H.
  • Basic Conditions:
    • Write two half-reactions and balance them independently assuming acidic conditions.
  • Neutral Conditions:
    • Balance O atoms by adding $H_2O$.
    • Balance H atoms by adding $H^+$ ions.
    • Balance charges by adding electrons.
    • If conditions are basic, neutralize added H⁺ with OH⁻.

Example Problem: Balancing in Acid

Given Reaction: Clo̅ + NO₂ → Cl̅ + NO₂⁻
  1. Unbalanced equation is already written.
  2. Write two unbalanced half-reactions:
    • Oxidation: NO₂ → NO₂⁻
    • Reduction: Clo̅ → Cl⁻
  3. Balance for acidic conditions:
    • Oxidation: NO₂ → NO₂⁻ becomes $NO2 + H2O → NO_3^{-} + 2H^+ + e^-$
Reduction Steps:
  • Cleansed half-reaction prepares for balancing:
    • $ClO₄ + 8H^+ + 8e^- → Cl^{-} + 4H_2O $
  1. Multiply the oxidation half-reaction by 8 to align electron counts.
  2. Add the half-reactions:
    • $8NO2 + 8H2O + ClO₄ + 8H^+ + 8e^- → 8NO3^{-} + 16H^+ + Cl^- + 4H2O$
    • Remove 8 e⁻ from both sides.
  3. Confirming the atoms and charges are balanced, simplifications lead to the final balanced equation.