Unit 6: Oxidation and Reduction

Unit 6: Oxidation and Reduction (5 Teaching Hours)

1. General and Electronic Concept of Oxidation and Reduction

• Oxidation: The process in which an atom, ion, or molecule loses electrons, resulting in an increase in oxidation state.
  • Example: $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$
• Reduction: The process in which an atom, ion, or molecule gains electrons, resulting in a decrease in oxidation state.
  • Example: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$
• Redox Reaction: A reaction where oxidation and reduction occur simultaneously.
  • Example: $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$

2. Oxidation Number and Rules for Assigning Oxidation Number

• Oxidation Number (Oxidation State): A number assigned to an element in a compound that represents the number of electrons lost or gained by the atom.
• Rules for Assigning Oxidation Numbers:
  1. The oxidation number of an atom in its elemental form (e.g., $\text{O}_2$, $\text{N}_2$) is 0.
  2. For monoatomic ions, the oxidation number is equal to the charge on the ion (e.g., $\text{Na}^+ = +1$, $\text{Cl}^- = -1$).
  3. Hydrogen has an oxidation number of +1 when bonded to non-metals and -1 when bonded to metals.
  4. Oxygen has an oxidation number of -2 in most compounds, except in peroxides (where it is -1) and when bonded to fluorine (where it is +2).
  5. The sum of oxidation numbers in a neutral compound is 0, and in a polyatomic ion, it equals the charge of the ion.

3. Balancing Redox Reactions

• By Oxidation Number Method:
  • Identify the oxidation and reduction half-reactions.
  • Balance the changes in oxidation numbers by adjusting coefficients.
  • Ensure that the total increase and decrease in oxidation numbers are equal.
• By Ion-Electron (Half Reaction) Method:
  1. Separate the redox reaction into two half-reactions (oxidation and reduction).
  2. Balance all elements except hydrogen and oxygen.
  3. Balance oxygen atoms by adding water ($\text{H}_2\text{O}$) molecules.
  4. Balance hydrogen atoms by adding hydrogen ions ($\text{H}^+$) in acidic medium or hydroxide ions ($\text{OH}^-$) in basic medium.
  5. Balance charges by adding electrons ($e^-$).
  6. Combine the half-reactions and ensure the total number of electrons is the same in both half-reactions.

4. Electrolysis

• Electrolysis: The process by which electrical energy is used to drive a non-spontaneous chemical reaction.
  • Qualitative Aspect: Electrolysis occurs in an electrolytic cell where the positive ions move to the cathode (reduction occurs), and the negative ions move to the anode (oxidation occurs).
    • Example: Electrolysis of $\text{NaCl(aq)}$ produces $\text{H}_2$ gas at the cathode and $\text{Cl}_2$ gas at the anode.
  • Quantitative Aspect: Governed by Faraday’s Laws of Electrolysis:
    1. First Law: The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity (charge) passed through the electrolyte. $m = Z \times Q \quad \text{where} \, Q = I \times t$ $m$ = mass of substance, $Z$ = electrochemical equivalent, $Q$ = charge (in coulombs), $I$ = current, and $t$ = time.
    2. Second Law: When the same quantity of electricity is passed through different electrolytes, the mass of the substances produced is proportional to their equivalent weights. $\frac{m_1}{m_2} = \frac{E_1}{E_2} \quad \text{where} \, E_1 \, \text{and} \, E_2 \, \text{are equivalent weights.}$