Redox Principles (OIL RIG)
- Oxidation: Loss of electrons (\(\text{OIL}\)), increase in oxidation state, or addition of oxygen.
- Reduction: Gain of electrons (\(\text{RIG}\)), decrease in oxidation state, or loss of oxygen / gain of hydrogen.
- Oxidising agent: Accepts electrons, oxidises another species, and is itself reduced.
- Reducing agent: Donates electrons, reduces another species, and is itself oxidised.
RIG
Reduction Is Gain (of electrons)
Oxidising agent = accepts electrons (is itself reduced). Reducing agent = donates electrons (is itself oxidised).
Three Ways to Define Redox
| Perspective | Oxidation | Reduction |
|---|---|---|
| Electrons | Loss of e⁻ | Gain of e⁻ |
| Oxidation state | Increase in ox. Number | Decrease in ox. Number |
| Oxygen / Hydrogen | Gain of O / Loss of H | Loss of O / Gain of H |
Oxidation States (Numbers)
An oxidation state is a formal charge assigned to an atom, assuming all bonds are completely ionic. It lets us track electron transfer in complex molecules.
8 Rules for Assigning Oxidation States
- Free uncombined elements = 0 (e.g. \(\text{O}_2, \text{P}_4, \text{Na}\)).
- Simple monoatomic ions = charge on ion (e.g. \(\text{Fe}^{3+} = +3, \text{Cl}^- = -1\)).
- Fluorine in all compounds = -1.
- Group 1 metals = +1; Group 2 metals = +2; Aluminium = +3.
- Hydrogen in compounds = +1 (except metal hydrides like \(\text{NaH}\) where it is -1).
- Oxygen in compounds = -2 (except peroxides like \(\text{H}_2\text{O}_2\) where -1, and in \(\text{OF}_2\) where +2).
- Sum of oxidation states in neutral molecule = 0.
- Sum of oxidation states in polyatomic ion = overall charge on ion.
Oxidation State of Mn in Permanganate
Problem: Find the oxidation state of Mn in \(\text{MnO}_4^-\).
\[\text{OS(Mn)} + 4(-2) = -1 \implies \text{OS(Mn)} - 8 = -1 \implies \text{OS(Mn)} = \mathbf{+7}\]Identifying Redox Reactions
Assign oxidation states to every element on both sides. If any element's oxidation number changes, it is a redox reaction. If no elements change, it is not redox (e.g. Acid-base neutralisation).
Half-Equations
Half-equations show oxidation and reduction separately. Every redox reaction can be split into two half-equations.
Half-Equations: Mg + CuSO₄
Overall: \(\text{Mg}(\text{s}) + \text{CuSO}_4(\text{aq}) \rightarrow \text{MgSO}_4(\text{aq}) + \text{Cu}(\text{s})\)
- Oxidation half-equation: \(\text{Mg}(\text{s}) \rightarrow \text{Mg}^{2+}(\text{aq}) + 2\text{e}^-\) (Mg is reducing agent)
- Reduction half-equation: \(\text{Cu}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Cu}(\text{s})\) (\(\text{Cu}^{2+}\) is oxidising agent)
Balancing Complex Half-Equations (Acidic Conditions)
- Balance elements other than H and O.
- Balance oxygen atoms by adding \(\text{H}_2\text{O}\).
- Balance hydrogen atoms by adding \(\text{H}^+\).
- Balance charge by adding electrons (\(\text{e}^-\)) to the more positive side.
- Multiply half-equations to equate electron counts, then add and cancel spectators.
Balancing the Permanganate Half-Equation
1. \(\text{MnO}_4^- \rightarrow \text{Mn}^{2+}\)
2. Balance O: \(\text{MnO}_4^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}\)
3. Balance H: \(\text{MnO}_4^- + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}\)
4. Balance charge (\(+7 \rightarrow +2\)): \(\mathbf{\text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}}\)
Disproportionation
A reaction where a single element is simultaneously oxidised and reduced.
Example: Cl₂ + H₂O → HClO + HCl
Cl goes from 0 (in Cl₂) to both +1 (in HClO) and −1 (in HCl).
Common Redox Notation and Naming Errors
- Charge vs Oxidation State: Charge is written \(2+\) or \(2-\); oxidation state is written \(+2\) or \(-2\) with the sign first.
- Per-atom calculations: In \(\text{Cr}_2\text{O}_7^{2-}\), each Cr is \(+6\), not \(+12\).
- Naming the whole agent: When asked to identify the oxidising agent, name \(\text{CuSO}_4\) or \(\text{Cu}^{2+}\), not just "Cu".
- Stock Roman numerals: Iron(III) chloride = \(\text{FeCl}_3\); Manganese(IV) oxide = \(\text{MnO}_2\).
Identifying Agents in Ionic Equations
In the reaction \(2\text{Fe}^{2+} + \text{Cl}_2 \rightarrow 2\text{Fe}^{3+} + 2\text{Cl}^-\), \(\text{Fe}^{2+}\) increases from \(+2 \rightarrow +3\) (oxidised, reducing agent), and \(\text{Cl}_2\) decreases from \(0 \rightarrow -1\) (reduced, oxidising agent).
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