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High Yield Topic

Redox Reactions

Chemistry Unit 4
20 min read
IAT Advanced

1. Classical Idea of Redox

Originally, redox reactions were defined based on the addition or removal of oxygen and hydrogen.

Process Classic Definition Example
Oxidation Addition of O2 or removal of H2. 2Mg + O2 → 2MgO
Reduction Removal of O2 or addition of H2. CuO + H2 → Cu + H2O

Extended Classical View

  • Oxidation: Also involves addition of an electronegative element (e.g., Cl, F) or removal of an electropositive element.
  • Reduction: Also involves addition of an electropositive element or removal of an electronegative element.

2. Electron Transfer Theory

The modern and most useful definition: Redox is a transfer of electrons from one species to another.

Memory Hack: OIL RIG

Oxidation Is Loss (of e-) | Reduction Is Gain (of e-)

Identifying Agents

  • Reducing Agent: Donor of electrons (gets oxidized).
  • Oxidizing Agent: Acceptor of electrons (gets reduced).

3. Oxidation Number (O.N.)

A formal charge assigned to an atom based on a set of rules, assuming total electron transfer in bonds. It is a bookkeeping tool.

Rules for Assigning Oxidation States
1. Free State: O.N. of elements = 0 (e.g., O2, Na, P4).
2. Oxygen: Usually -2. Exceptions: Peroxides (-1, H2O2), Superoxides (-1/2, KO2), with F (+2, OF2).
3. Hydrogen: +1 with non-metals, -1 with active metals (Metal Hydrides).
4. Halogens: F is always -1. Cl, Br, I are -1 unless bonded to O or F.
5. Sum Rule: ∑ O.N. = 0 (neutral molecule) or Charge of Ion.

Paradox of Fractional O.N.

Some compounds show fractional average O.N. (e.g., C3O2 has carbon with +4/3). This is because different atoms of the same element have different oxidation states within the molecule.

  • Fe3O4: Mixture of FeO (+2) and Fe2O3 (+3). Average = +8/3.
  • Br3O8: Terminal Br are +6, Central Br is +4.

4. Electrode Processes

Electrochemical cells utilize redox reactions to generate or utilize electrical energy.

Galvanic (Voltaic) Cells

Converts Chemical Energy → Electrical Energy (Spontaneous, ΔG < 0).

  • Anode: Oxidation occurs here (Negative pole).
  • Cathode: Reduction occurs here (Positive pole).
  • Salt Bridge: Maintains electrical neutrality and completes the circuit.
E°cell = E°cathode - E°anode
Using Standard Reduction Potentials.

Daniel Cell Example

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Zn is oxidized (Anode), Cu2+ is reduced (Cathode).

Quick Revision Flashcards

Reducing Agent

The substance that causes reduction by losing electrons (getting oxidized itself).

Disproportionation

A reaction where the same element is simultaneously oxidized and reduced (e.g., 2H2O2 → 2H2O + O2).

Anode

The electrode where oxidation takes place. In a galvanic cell, it is the negative terminal.

5. Common Mistakes

  • Confusing OA and RA: Remember, the Oxidizing Agent is the one being Reduced.
  • Hydrogen with Metals: Hydrogen is NOT always +1. In LiH, NaH, CaH2, it is -1.
  • Fractional O.N.: Don't panic if you get a fraction; it usually represents an average of different atoms.
  • Ionic Product vs Ksp: (Wait, that's Equilibrium). For Redox: E°cell signs. Spontaneous reactions MUST have a positive E°cell.

6. Practice Mock Test

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Take a quick 15-question assessment specifically designed for Redox Reactions.

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End of Chapter

Redox Reactions

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