Redox (reduction–oxidation) reactions are chemical reactions involving the transfer of electrons or changes in oxidation number (oxidation state). This chapter explores how to identify redox reactions, determine oxidation numbers, balance equations, and understand key redox terminology.
1. Calculating Oxidation Numbers of Elements in Compounds and Ions
Oxidation Number (Oxidation State):
A number assigned to an atom in a compound or ion that represents the number of electrons lost or gained by that atom compared to its elemental state.
Rules for Assigning Oxidation Numbers:
| Rule | Explanation | Example |
|---|---|---|
| 1 | Elements in their standard state have oxidation number 0 | O₂, H₂, Na all = 0 |
| 2 | Group 1 metals = +1 | Na⁺ in NaCl |
| 3 | Group 2 metals = +2 | Ca²⁺ in CaCl₂ |
| 4 | Fluorine = –1 in compounds | F in HF |
| 5 | Hydrogen = +1, but –1 in metal hydrides | H in HCl = +1; in NaH = –1 |
| 6 | Oxygen = –2, except in peroxides (–1) and OF₂ (+2) | O in H₂O = –2; O in H₂O₂ = –1 |
| 7 | Sum of oxidation numbers in a neutral compound = 0 | H₂O: 2(+1) + (–2) = 0 |
| 8 | In ions, sum = charge of ion | SO₄²⁻: S + 4(–2) = –2 ⇒ S = +6 |
2. Using Changes in Oxidation Numbers to Help Balance Chemical Equations
When elements change oxidation numbers in a reaction, electrons are either lost (oxidation) or gained (reduction).
Steps to Balance Redox Equations Using Oxidation Numbers:
- Assign oxidation numbers to all atoms.
- Identify which atoms are oxidised and reduced.
- Calculate the change in oxidation numbers.
- Balance the number of electrons gained and lost.
- Balance the rest of the atoms (including oxygen and hydrogen).
- Add state symbols if required.
Example:
Balance the reaction:
Fe²⁺ + Cr₂O₇²⁻ + H⁺ → Fe³⁺ + Cr³⁺ + H₂O
- Fe²⁺ → Fe³⁺ (oxidation, +1)
- Cr⁶⁺ → Cr³⁺ (reduction, –3 × 2 = –6)
Balance electrons:
- Multiply Fe²⁺ by 6 and Cr₂O₇²⁻ by 1 → 6 Fe²⁺ + Cr₂O₇²⁻ + 14 H⁺ → 6 Fe³⁺ + 2 Cr³⁺ + 7 H₂O
3. Explaining and Using Redox, Oxidation, Reduction, and Disproportionation
Redox Reaction:
A chemical reaction where both oxidation and reduction occur simultaneously.
Oxidation:
- Loss of electrons
- Increase in oxidation number
- E.g., Zn → Zn²⁺ + 2e⁻
Reduction:
- Gain of electrons
- Decrease in oxidation number
- E.g., Cu²⁺ + 2e⁻ → Cu
Disproportionation:
A single substance undergoes both oxidation and reduction.
Example:
Cl₂ + H₂O → HCl + HClO
- Cl in Cl₂ is 0
- In HCl: –1 (reduction)
- In HClO: +1 (oxidation)
4. Oxidising and Reducing Agents
Oxidising Agent (Oxidant):
- Accepts electrons
- Gets reduced
- Causes another substance to be oxidised
Example:
In Cu + HNO₃ → Cu(NO₃)₂ + NO₂ + H₂O
HNO₃ acts as oxidising agent (NO₃⁻ → NO₂)
Reducing Agent (Reductant):
- Donates electrons
- Gets oxidised
- Causes another substance to be reduced
Example:
Zn + Cu²⁺ → Zn²⁺ + Cu
Zn is the reducing agent
5. Using Roman Numerals to Indicate Oxidation Numbers
Roman numerals are used in the names of compounds to indicate the oxidation number of elements, particularly for transition metals and elements with multiple oxidation states.
Examples:
- Iron(II) sulfate: FeSO₄ → Fe²⁺
- Iron(III) chloride: FeCl₃ → Fe³⁺
- Sulfur(IV) oxide: SO₂ → S = +4
- Sulfur(VI) oxide: SO₃ → S = +6
These notations help clearly distinguish compounds with the same elements but different oxidation states.
Summary Table: Key Terms
| Term | Meaning |
|---|---|
| Oxidation | Loss of electrons / increase in oxidation number |
| Reduction | Gain of electrons / decrease in oxidation number |
| Redox | Simultaneous oxidation and reduction |
| Oxidising agent | Accepts electrons, gets reduced |
| Reducing agent | Donates electrons, gets oxidised |
| Disproportionation | One species is both oxidised and reduced |
| Oxidation Number | Indicator of the degree of oxidation of an atom |
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