Redox Reactions

1. Oxidation and Reduction Processes

Redox reactions are chemical processes where oxidation and reduction occur simultaneously.

(a) Oxidation and Reduction Definitions

Addition and Removal of Oxygen and Hydrogen

• Oxidation: Gain of oxygen or loss of hydrogen in a substance.
  Example: $2Mg + O_2 \to 2MgO$ (Magnesium oxidized to Magnesium oxide).
• Reduction: Loss of oxygen or gain of hydrogen in a substance.
  Example: $CuO + H_2 \to Cu + H_2O$ (Copper oxide reduced to Copper).

Loss and Gain of Electrons

• Oxidation: Loss of electrons by an atom or ion.
  Example: $Na \to Na^+ + e^-$.
• Reduction: Gain of electrons by an atom or ion.
  Example: $Cl_2 + 2e^- \to 2Cl^-$.

Change in Oxidation Numbers

• Oxidation: Increase in oxidation number.
  Example: $Fe^{2+} \to Fe^{3+} + e^-$.
• Reduction: Decrease in oxidation number.
  Example: $MnO_4^- \to Mn^{2+}$.

2. Oxidizing and Reducing Agents

Agents that facilitate oxidation or reduction in a chemical reaction.

Description

(a) Addition and Removal of Oxygen and Hydrogen

• Oxidizing Agent: Adds oxygen or removes hydrogen (e.g., $KMnO_4$).
• Reducing Agent: Removes oxygen or adds hydrogen (e.g., $H_2$).

(b) Loss and Gain of Electrons

• Oxidizing Agent: Gains electrons (e.g., $Cl_2 \to 2Cl^-$).
• Reducing Agent: Loses electrons (e.g., $Zn \to Zn^{2+}$).

(c) Change in Oxidation Numbers

• Oxidizing Agent: Causes increase in oxidation state of another substance.
• Reducing Agent: Causes decrease in oxidation state of another substance.

3. Balancing Redox Equations

(a) Ion, Electron, or Change in Oxidation Number

1. Assign oxidation numbers to all elements.
2. Identify changes in oxidation numbers.
3. Balance using electron gain/loss.

(b) Half-Reactions and Overall Reaction

• Split into oxidation and reduction half-reactions.
• Balance atoms and charges separately.
• Combine half-reactions ensuring electron counts match.

Example:

Overall: $Zn + Cu^{2+} \to Zn^{2+} + Cu$.

4. Electrochemical Cells

(i) Standard Electrode Potential

• The potential of an electrode compared to the Standard Hydrogen Electrode (SHE) ($E^\circ = 0$ V).
• Measured using a metal/metal ion system under standard conditions ($25^\circ C, 1 \, M, 1 \, atm$).

(ii) Drawing Cell Diagram and Writing Cell Notation

Diagram

• Illustrate a Daniell Cell with Zn and Cu electrodes immersed in their respective solutions.

Cell Notation

(iii) Electromotive Force (e.m.f)

• $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$.

(iv) Applications

1. Batteries: Lead-acid battery, dry cell.
2. Fuel cells: Generating energy in spacecraft.
3. Electrochemical sensors.

5. Electrolysis

(i) Electrolytic Cells

• Devices converting electrical energy into chemical energy.

(ii) Principles

1. Ion migration under an electric field.
2. Redox reactions at electrodes (anode and cathode).

(iii) Factors Influencing Discharge

1. Ion concentration.
2. Electrode material.
3. Nature of electrolyte.

(iv) Faraday's Laws

• First Law: Mass of substance deposited is proportional to charge passed.
  $m = Z \cdot Q$.
• Second Law: Amounts of substances deposited are proportional to their equivalent weights.

(v) Applications

1. Electroplating.
2. Purification of metals.
3. Extraction of aluminum from bauxite.

(vi) Corrosion

• A redox process leading to deterioration of metals (e.g., rusting of iron).
• Prevention:
  • Galvanizing.
  • Cathodic protection.
  • Coatings (paint, grease).

6. Oxidation Numbers

Determination Rules

1. Free element: $0$.
2. Monoatomic ion: Charge of ion.
3. Oxygen: $-2$ (except peroxides).
4. Hydrogen: $+1$ (except hydrides).

Comparison of Electrolytic and Electrochemical Cells

This structured note provides a comprehensive understanding of redox reactions and their applications in both theoretical and practical contexts.