Redox & Oxidation-Reduction Notes

Oxidation Numbers: Representation & Meaning

  • An element’s oxidation number = the charge it effectively carries in a compound ➔ relates directly to electrons lost (positive number) or gained (negative number).
    • Expressed with a Roman numeral in brackets after the element name.
    • Iron(II) oxide → Fe has oxidation number +2+2.
    • Copper(I) chloride → Cu has oxidation number +1+1.

Redox Reactions: Core Definitions

  • Redox = Reduction + Oxidation occurring simultaneously.
  • Classical (oxygen-based) definitions
    • Oxidation = gain of oxygen.
    • Reduction = loss of oxygen.
    • Example: H<em>2+CuO  →  Cu+H</em>2O\mathrm{H<em>2 + CuO \;\rightarrow\; Cu + H</em>2O}
    • H2\mathrm{H_2} gains O ➔ oxidised.
    • CuO\mathrm{CuO} loses O ➔ reduced.

Identifying Redox via Oxygen Transfer

  • Look for species that gain O (oxidised) and those that lose O (reduced).
    • Example: H<em>2+ZnO  →  Zn+H</em>2O\mathrm{H<em>2 + ZnO \;\rightarrow\; Zn + H</em>2O}
    • H<em>2\mathrm{H<em>2} oxidised (now H</em>2O\mathrm{H</em>2O}).
    • ZnO\mathrm{ZnO} reduced (now Zn\mathrm{Zn}).
  • Another example: Fe<em>2O</em>3+3CO  →  2Fe+3CO2\mathrm{Fe<em>2O</em>3 + 3CO \;\rightarrow\; 2Fe + 3CO_2}
    • Fe<em>2O</em>3  →  Fe\mathrm{Fe<em>2O</em>3 \;\rightarrow\; Fe} → Fe loses O ➔ reduction.
    • CO  →  CO2\mathrm{CO \;\rightarrow\; CO_2} → C gains O ➔ oxidation.

Electron-Transfer Definitions (Extended syllabus)

  • Oxidation = loss of electrons ➔ oxidation number increases.
  • Reduction = gain of electrons ➔ oxidation number decreases.
  • Diagnostic method: split into ionic half-equations.
    • Overall: Cl<em>2+2KI  →  2KCl+I</em>2\mathrm{Cl<em>2 + 2KI \;\rightarrow\; 2KCl + I</em>2}
    • Cl2+2e−  →  2Cl−\mathrm{Cl_2 + 2e^- \;\rightarrow\; 2Cl^-} (reduction; Cl gains e−e^-).
    • 2I−  →  I2+2e−\mathrm{2I^- \;\rightarrow\; I_2 + 2e^-} (oxidation; I^- loses e−e^-).

Rules for Assigning Oxidation Numbers (Extended)

  • (a) Element in free/uncombined state: 00.
    • Zn\mathrm{Zn}, Cl<em>2\mathrm{Cl<em>2}, H</em>2\mathrm{H</em>2} all 00.
  • (b) Monatomic ion: oxidation number = ionic charge.
    • Zn2+\mathrm{Zn^{2+}} → +2+2; Cl−\mathrm{Cl^-} → −1-1; H+\mathrm{H^+} → +1+1.
  • (c) Sum of oxidation numbers in a neutral compound = 00.
    • ZnCl2\mathrm{ZnCl_2}: +2+2(−1)=0+2 + 2(-1) = 0.
    • HBr\mathrm{HBr}: +1+(−1)=0+1 + (-1) = 0.
  • (d) Sum in a polyatomic ion = overall ionic charge.
    • SO42−\mathrm{SO_4^{2-}}: S +6+6, each O −2-2 → +6+4(−2)=−2+6 + 4(-2) = -2 (matches ion charge).

Using Oxidation Numbers to Confirm Redox (Extended)

  • If oxidation number decreases (more negative) ➔ electrons gained ➔ reduction.
  • If oxidation number increases (more positive) ➔ electrons lost ➔ oxidation.
  • Worked example: Cl<em>2+2KI  →  2KCl+I</em>2\mathrm{Cl<em>2 + 2KI \;\rightarrow\; 2KCl + I</em>2}
    1. Assign numbers: Cl 00 (reactant) → −1-1 (in Cl−\mathrm{Cl^-}); I −1-1 → 00.
    2. Changes:
    • Cl: 0→−10 \rightarrow -1 (reduced).
    • I: −1→0-1 \rightarrow 0 (oxidised).

Colour Tests for Redox Agents (Extended)

  • Acidified aqueous KMnO4\mathrm{KMnO_4} [manganate(VII)] – tests for reducing agents.
    • Contains Mn7+\mathrm{Mn^{7+}} (purple).
    • When reduced to Mn2+\mathrm{Mn^{2+}} the solution turns purple → colourless.
  • Aqueous KI\mathrm{KI} – tests for oxidising agents.
    • Contains I−\mathrm{I^-} (colourless).
    • Oxidised to I2\mathrm{I_2} → solution turns colourless → brown.
    • With starch indicator: brown I2\mathrm{I_2} forms blue-black complex.

Oxidising & Reducing Agents (Extended)

  • Oxidising agent
    • Causes oxidation by accepting electrons ➔ itself is reduced.
    • Often a non-metal or positive ion.
    • Example half-equation: Br<em>2+2e−  →  2Br−\mathrm{Br<em>2 + 2e^- \;\rightarrow\; 2Br^-} (Br2 reduced; acted as OA).
  • Reducing agent
    • Causes reduction by donating electrons ➔ itself is oxidised.
    • Often a metal or negative ion.
    • Example: K  →  K++e−\mathrm{K \;\rightarrow\; K^+ + e^-} (K oxidised; acted as RA).

Identifying OA & RA in Equations (Extended)

  • Half-equation method OR oxidation-number method.
    • Example: Cl<em>2+2KI  →  2KCl+I</em>2\mathrm{Cl<em>2 + 2KI \;\rightarrow\; 2KCl + I</em>2}
    • I−\mathrm{I^-} loses e−e^- ➔ reducing agent.
    • Cl2\mathrm{Cl_2} gains e−e^- ➔ oxidising agent.
    • Example: Zn+CuSO<em>4  →  ZnSO</em>4+Cu\mathrm{Zn + CuSO<em>4 \;\rightarrow\; ZnSO</em>4 + Cu}
    • Zn: 0→+20 \rightarrow +2 (oxidised) ➔ reducing agent.
    • Cu2+\mathrm{Cu^{2+}}: +2→0+2 \rightarrow 0 (reduced) ➔ oxidising agent.