Lecture Notes: Redox Reactions

Key Definitions

• Reducing Agents: Electron donors.
• Oxidizing Agents: Electron acceptors.

Processes

• Oxidation: Process of electron loss.

  • Example: ( \text{Zn} \rightarrow \text{Zn}^{2+} + 2\text{e}^- )
  • Results in an increase in oxidation number.

• Reduction: Process of electron gain.

  • Example: ( \text{Cl}_2 + 2\text{e}^- \rightarrow 2\text{Cl}^- )
  • Results in a decrease in oxidation number._

Rules for Assigning Oxidation Numbers

1. Uncombined Elements: Oxidation number = 0 (e.g., Zn, Cl(_2), O(_2), Ar).
2. Compounds: Sum of oxidation numbers = 0 (e.g., NaCl: Na = +1, Cl = -1).
3. Monoatomic Ions: Oxidation number = ionic charge (e.g., ( \text{Zn}^{2+} = +2 ), ( \text{Cl}^- = -1 )).
4. Polyatomic Ions: Sum of oxidation numbers = charge on ion (e.g., ( \text{CO}_3^{2-} ): C = +4, O = -2).
5. Invariable Oxidation Numbers:
   • Group 1 metals: +1
   • Group 2 metals: +2
   • Al: +3
   • H: +1 (except in metal hydrides, e.g., NaH)
   • F: -1
   • Cl, Br, I: -1 (except with O and F)
   • O: -2 (except in peroxides like ( \text{H}_2\text{O}2 ))

Redox Equations and Half Equations

• Example Reaction: ( \text{Br}_2(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow \text{I}_2(\text{aq}) + 2\text{Br}^-(\text{aq}) )
  • ( \text{I}^- ) oxidized (loses electrons)
  • ( \text{Br}_2 ) reduced (gains electrons)
• Naming Agents: Use full name (e.g., "Bromine water" not just Br or ion)._

Half Equations

• Reduction Half Equation: Electrons on the left.
• Oxidation Half Equation: Electrons on the right.

Balancing Redox Equations

1. Identify oxidation changes (e.g., ( \text{Zn} \rightarrow \text{Zn}^{2+} )).
2. Add electrons to balance oxidation number changes.
3. Ensure sum of charges is equal on both sides.

Complex Half Equations

• In Acidic Conditions

  • Balance using ( \text{H}^+ ) and ( \text{H}_2\text{O} ).
  • Example: ( \text{MnO}_4^- \rightarrow \text{Mn}^{2+} )
    1. Add 5 electrons: ( \text{MnO}_4^- + 5\text{e}^- \rightarrow \text{Mn}^{2+} ).
    2. Add water: ( \text{MnO}_4^- + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} ).
    3. Add ( \text{H}^+ ): ( \text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} ).

• Combining Half Equations

  • Reduction: ( \text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} ).
  • Oxidation: ( \text{C}_2\text{O}_4^{2-} \rightarrow 2\text{CO}_2 + 2\text{e}^- ).
  • Multiply to equalize electrons, add, and cancel.
  • Example: ( 2\text{MnO}_4^- + 16\text{H}^+ + 5\text{C}_2\text{O}_4^{2-} \rightarrow 2\text{Mn}^{2+} + 10\text{CO}_2 + 8\text{H}_2\text{O} ).

• Additional Example: ( \text{SO}_4^{2-} \rightarrow \text{SO}_2 )

  • Balance electrons, add ( \text{H}_2\text{O} ), add ( \text{H}^+ ), ensure charge balance.