6.2.1Additional lessons
Oxidation and Reduction — IB Chemistry HL Revision Notes & Topical Questions
5 min read 5 parts Lesson 132 of 172
Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution
pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴
Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14
Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.
Section 1 of 5Defining Redox
Redox reactions involve the transfer of electrons, and the two processes cannot happen separately.
Oxidation
- Oxidation The loss of electrons by a species, which makes its oxidation state increase, for example Mg becoming Mg²⁺.
Reduction
- Reduction The gain of electrons by a species, which makes its oxidation state decrease, for example Cu²⁺ becoming Cu.
Redox
- Redox Oxidation and reduction always occur together because electrons lost by one species must be gained by another.
Oxygen and hydrogen
- Oxygen and hydrogen In organic chemistry gain of oxygen or loss of hydrogen is oxidation, while loss of oxygen or gain of hydrogen is reduction.
Section 2 of 5Oxidation States
Oxidation States
Uncombined elements
- Uncombined elements Atoms in an element on its own, such as Fe, O₂ or S₈, always have an oxidation state of zero.
Simple ions
- Simple ions The oxidation state of a monatomic ion equals its charge, so Al³⁺ is +3 and Cl⁻ is −1.
Oxygen and hydrogen
- Oxygen and hydrogen Oxygen is usually −2 and hydrogen is usually +1, with exceptions such as peroxides and metal hydrides.
Totals
- Totals The oxidation states in a neutral compound sum to zero, and in a polyatomic ion they sum to the ion charge.
Roman numerals
- Roman numerals A numeral in a name such as iron(III) oxide states the oxidation state of the metal as +3.
Section 3 of 5Oxidising and Reducing Agents
1
Identify which species changes oxidation state by comparing every element on the left with the same element on the right.
2
The species whose oxidation state rises has been oxidised, and it is the reducing agent because it supplies the electrons.
3
The species whose oxidation state falls has been reduced, and it is the oxidising agent because it takes the electrons.
4
Check that the total electrons lost equal the total electrons gained, which confirms the reaction is a true redox process.
5
Remember that the agent is the substance causing the change, so the agent is always the opposite of what happens to it.
| Term | Electrons | Oxidation state | Example |
|---|---|---|---|
| Oxidising agent | Gains electrons | Decreases | MnO₄⁻, Cl₂ |
| Reducing agent | Loses electrons | Increases | Fe²⁺, Zn |
Section 4 of 5Worked Example
Question Use oxidation states to show that Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) is a redox reaction, and state the oxidising agent.
- Zn changes from 0 in the element to +2 in ZnSO₄, so its oxidation state increases.
- Cu changes from +2 in CuSO₄ to 0 in the element, so its oxidation state decreases.
- Zn is oxidised and Cu²⁺ is reduced, so Zn is the reducing agent and Cu²⁺ is the oxidising agent.
Answer Zinc is oxidised, copper(II) is reduced, and the oxidising agent is Cu²⁺(aq).
Section 5 of 5Exam Checklist & Tips
6.2.1 – Key Points for the Exam
- Oxidation Loss of electrons, or an increase in oxidation state.
- Reduction Gain of electrons, or a decrease in oxidation state.
- Oxidising agent A species that accepts electrons and is itself reduced.
- Reducing agent A species that donates electrons and is itself oxidised.
- Oxidation state The charge an atom would have if all its bonding electrons were assigned to the more electronegative atom.
Common Mistakes
- Saying the oxidising agent is oxidised; it is the species that is reduced.
- Giving oxidation states as 2+ instead of +2; the sign comes first for oxidation states.
- Forgetting that elements in their standard state, such as O₂, have oxidation state zero.
- Assuming oxygen is always −2, which is wrong in peroxides such as H₂O₂ where it is −1.
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