1.1 Oxidation and Reduction
Redox reaction is a chemical reaction where oxidation and reduction occur simultaneously.
Definitions of Oxidation and Reduction
- In terms of Oxygen: Oxidation is gain of oxygen; Reduction is loss of oxygen.
- In terms of Hydrogen: Oxidation is loss of hydrogen; Reduction is gain of hydrogen.
- In terms of Electron Transfer: Oxidation is loss of electrons; Reduction is gain of electrons.
- In terms of Oxidation Number: Oxidation is an increase in oxidation number; Reduction is a decrease in oxidation number.
Oxidising and Reducing Agents
- Oxidising Agent: A substance that oxidises another substance while itself gets reduced (electron acceptor / oxidation number decreases).
- Reducing Agent: A substance that reduces another substance while itself gets oxidised (electron donor / oxidation number increases).
Rules for Determining Oxidation Numbers
- Free elements have an oxidation number of $0$ (e.g., $\text{Zn}$, $\text{O}_2$, $\text{Cl}_2$).
- Simple ions have an oxidation number equal to their charge (e.g., $\text{Cu}^{2+} = +2$, $\text{Cl}^- = -1$).
- Neutral molecules have a sum of oxidation numbers equal to $0$ (e.g., $\text{H}_2\text{O}$, $\text{CO}_2$).
- Polyatomic ions have a sum of oxidation numbers equal to the ion's charge (e.g., $\text{SO}_4^{2-} = -2$).
- Hydrogen in compounds is generally $+1$ (except in metal hydrides like $\text{NaH}$ where it is $-1$).
- Oxygen in compounds is generally $-2$ (except in peroxides like $\text{H}_2\text{O}_2$ where it is $-1$).
Redox Reactions in Aqueous Solutions
Conversion between $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$ ions:
- Oxidation of $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$: Requires an oxidising agent (e.g., acidified $\text{KMnO}_4$ solution, bromine water $\text{Br}_2$). Half-equation: $\text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^-$.
- Reduction of $\text{Fe}^{3+}$ to $\text{Fe}^{2+}$: Requires a reducing agent (e.g., zinc powder $\text{Zn}$). Half-equation: $\text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+}$.
1.2 Standard Electrode Potential
The Standard Electrode Potential ($E^\circ$) is the potential difference produced when a half-cell is connected to the Standard Hydrogen Electrode (SHE) under standard conditions ($25^\circ\text{C}$, $1.0\text{ mol dm}^{-3}$ ion concentration, $1\text{ atm}$ pressure).
Reactivity and $E^\circ$ Values
- A more negative or less positive $E^\circ$ value indicates a stronger reducing agent (greater tendency to lose electrons / undergo oxidation).
- A more positive or less negative $E^\circ$ value indicates a stronger oxdising agent (greater tendency to gain electrons / undergo reduction).
Predicting Redox Reactions Using $E^\circ$
For a spontaneous reaction, $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} > 0$.
1.3 Voltaic Cell (Galvanic Cell)
A voltaic cell converts chemical energy into electrical energy using spontaneous redox reactions.
- Anode (Negative Pole): Metal with a more negative $E^\circ$ value. Oxidation occurs here.
- Cathode (Positive Pole): Metal with a more positive $E^\circ$ value. Reduction occurs here.
- Salt Bridge / Porous Pot: Completes the circuit and maintains electrical neutrality by allowing ion migration.
1.4 Electrolysis of Aqueous Solutions
An electrolytic cell converts electrical energy into chemical energy to drive a non-spontaneous redox reaction.
Factors Affecting Discharge of Ions at Electrodes
- Standard Electrode Potential ($E^\circ$): Ions with more positive $E^\circ$ are selectively discharged at the cathode; ions with less positive/more negative $E^\circ$ are selectively discharged at the anode.
- Concentration of Electrolyte: In a concentrated halide solution (e.g., concentrated $\text{NaCl}$), chloride ions ($\text{Cl}^-$) are selectively discharged at the anode instead of hydroxide ions ($\text{OH}^-$).
- Type of Electrodes: Reactive electrodes (e.g., copper anode in $\text{CuSO}_4$ solution) dissolve into the electrolyte: $\text{Cu} \rightarrow \text{Cu}^{2+} + 2e^-$.
1.5 Extraction of Metal from Its Ore
Metals high in the reactivity series (e.g., $\text{Al}$, $\text{Na}$, $\text{Mg}$) are extracted via electrolysis of molten ores. Metals lower in the series (e.g., $\text{Fe}$, $\text{Zn}$) are extracted by reduction with carbon in a blast furnace.
1.6 Rusting of Iron
Rusting is a redox process requiring both water ($\text{H}_2\text{O}$) and oxygen ($\text{O}_2$).
- Anode (Oxidation): $\text{Fe (s)} \rightarrow \text{Fe}^{2+}\text{(aq)} + 2e^-$
- Cathode (Reduction): $\text{O}_2\text{(g)} + 2\text{H}_2\text{O(l)} + 4e^- \rightarrow 4\text{OH}^-\text{(aq)}$
- Overall Chemical Rust Formula: Hydrated iron(III) oxide, $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$.
- Rust Prevention: Sacrificial protection (attaching a metal with a more negative $E^\circ$ than iron, such as zinc or magnesium), painting, galvanising, or tin plating.