Redox Reactions
FoundationA redox reaction is essentially electron transfer: oxidation is the loss of electrons (oxidation number rising) and reduction the gain of electrons (oxidation number falling). The two always occur together; an oxidizing agent is itself reduced while a reducing agent is itself oxidised.
- Oxidation = loss of electrons, oxidation number rises; reduction = gain of electrons, number falls.
- Oxidation and reduction always occur together; electrons lost equal electrons gained.
- An oxidizing agent oxidises another and is itself reduced; a reducing agent reduces another and is itself oxidised.
文體小知識：科學知識說明文——由燃燒與生鏽說起，說明氧化還原的電子轉移本質、氧化數規則，以及判斷氧化劑與還原劑的方法。
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1A magnesium ribbon burns in oxygen with a dazzling white flame; an iron gate rusts slowly in humid air. These seemingly different changes are essentially the same kind of reaction — redox reactions. 'Oxidation' originally meant combining with oxygen and 'reduction' meant having oxygen removed, but chemists found that the more fundamental definition lies in the transfer of electrons, which is the core idea of this lesson.
2In terms of electron transfer: oxidation is the loss of electrons and reduction is the gain of electrons. When magnesium burns, each atom loses two electrons and becomes a magnesium ion — oxidised; the oxygen molecules gain electrons — reduced. Loss and gain must come as a pair, so oxidation and reduction always happen at the same time, on different species; hence the combined name, redox reactions.
3Since electrons are invisible, we judge redox by oxidation numbers. The oxidation number of an element in its free state is zero; that of a monatomic ion equals its charge; oxygen is usually −2 in compounds; hydrogen is +1 when bonded to non-metals. The species whose oxidation number rises is oxidised and the one whose number falls is reduced: in burning magnesium, magnesium goes from 0 to +2 while oxygen falls from 0 to −2.
4A species that oxidises another while being reduced itself is an oxidizing agent — oxygen and potassium permanganate are common examples. A species that reduces another while being oxidised itself is a reducing agent, such as carbon, carbon monoxide and hydrogen. In metallurgy, carbon monoxide reduces iron from its ore while being oxidised to carbon dioxide — the classic role of a reducing agent.
5Redox is everywhere: the body oxidises glucose for energy, food spoils, metals rust and bleaches remove stains — all involve electron transfer. Once you can judge redox confidently, the next lesson shows how such reactions turn chemical energy into electrical energy in chemical cells, and how an electric current can drive them backwards in electrolysis.
1Summary 重點摘要
A redox reaction is essentially electron transfer: oxidation is the loss of electrons (oxidation number rising) and reduction the gain of electrons (oxidation number falling). The two always occur together; an oxidizing agent is itself reduced while a reducing agent is itself oxidised.
2Vocabulary cards Vocabulary
A reaction involving oxidation and reduction together; in essence, electron transfer.
The loss of electrons; the oxidation number rises.
The gain of electrons; the oxidation number falls.
A species that oxidises another and is itself reduced, e.g. oxygen, potassium permanganate.
A species that reduces another and is itself oxidised, e.g. carbon, carbon monoxide, hydrogen.
A number showing the charged state of an atom in a species; its rise or fall identifies oxidation and reduction.
An equation showing only the oxidation or the reduction half, balanced with electrons.
An equation showing only the ions and molecules actually taking part in the reaction.
The movement of electrons from one species to another — the essence of redox.
3Step by step Steps
- Step 1: Meeting redox through burning and rustingRedox is inseparable from daily life: a magnesium ribbon burns with a white flame, an iron gate rusts and flakes, and a gas stove turns fuel into carbon dioxide and water — all redox reactions. For example, when iron rusts, it reacts with oxygen and water in the air: the iron is oxidised, the oxygen is reduced, and the rust is mainly iron oxide. Early definitions saw oxidation as combining with oxygen and reduction as losing it — intuitive, but unable to cover reactions without oxygen; hence the modern electron-transfer definition, the starting point of this lesson.
- Step 2: Oxidation loses electrons, reduction gains themThe modern definition is compact: oxidation is the loss of electrons, reduction the gain. For example, place zinc granules in copper(II) sulfate solution: each zinc atom loses two electrons and enters the solution as a zinc ion, while copper ions gain two electrons and deposit as copper atoms. Zinc is oxidised, the copper ions are reduced, and electrons flow from zinc to the ions. The mnemonic is 'the loser of electrons is oxidised, the gainer reduced'. Electrons are only handed on, never destroyed, so total electrons lost must equal total electrons gained in every transfer.
- Step 3: Judging by oxidation numbersElectrons are invisible, so oxidation numbers are the safest tool in exams. Basic rules: the oxidation number of an element in its free state is zero; a monatomic ion's number equals its charge; oxygen is usually −2 in compounds; hydrogen is +1 when bonded to non-metals. For example, in magnesium oxide (MgO), oxygen is −2 so magnesium is +2; before the reaction magnesium metal was 0, rising to +2 — oxidised, while oxygen fell from 0 to −2 — reduced. The motto: the species whose oxidation number rises is oxidised, the one whose number falls is reduced.
- Step 4: Writing half equationsHalf equations display the electrons explicitly; the method has three stages: write the species change, balance the atoms, then balance the charges with electrons. For example, the oxidation half equation of magnesium is Mg → Mg²⁺ + 2e⁻ — the right side carries +2 in total while the left is neutral, so the electron count is correct; the reduction of oxygen is O₂ + 4e⁻ → 2O²⁻. Multiplying the two half equations by suitable numbers to equalise the electrons and adding them, with the electrons cancelling, yields the overall equation — a technique used again in chemical cells next lesson.
- Step 5: Telling oxidizing agents from reducing agentsAn oxidizing agent oxidises another species and is itself reduced — its oxidation number falls; a reducing agent reduces another and is itself oxidised — its number rises. Common oxidizing agents include oxygen and acidified potassium permanganate solution; common reducing agents include carbon, carbon monoxide and hydrogen. For example, in the blast furnace, carbon monoxide removes oxygen from iron oxide, reducing the iron while being oxidised to carbon dioxide — so carbon monoxide is the reducing agent and iron oxide the oxidizing agent. Remember that the role of an 'agent' is the opposite of what it undergoes; identify whose oxidation number falls first.
- Step 6: Oxidation and reduction always come togetherWherever electrons are transferred, oxidation and reduction appear as a pair: without something being oxidised, nothing can be reduced. For example, hydrogen burns in chlorine to form hydrogen chloride — no oxygen involved — yet hydrogen's oxidation number rises from 0 to +1 (oxidised) while chlorine falls from 0 to −1 (reduced), so it is still a redox reaction. This shows why 'with or without oxygen' is not a valid criterion. When asked to name the oxidizing agent, first write out the oxidation numbers of all species and only then conclude — clear steps protect your marks.
- Step 7: Redox in life and industryRedox underpins modern life: body cells oxidise glucose step by step, releasing energy to maintain body temperature and activity; protecting steel from rust by oiling, painting or galvanising aims precisely to stop iron being oxidised; bleach oxidises dyes into colourless substances to remove stains. For example, a cut apple browns because substances in the flesh are oxidised by air; a squeeze of lemon juice slows it, as the vitamin C there is oxidised first, acting as a reducing agent. Recognising redox around you gives you plenty to say in data-based extension questions.
4Diagrams Diagrams
5Worked example Example
A magnesium ribbon burns in oxygen to form magnesium oxide (MgO). Use oxidation numbers to identify the species oxidised and reduced, write the two half equations, and combine them into the overall equation.
- Write and balance the equation: 2Mg + O₂ → 2MgO.
- Assign oxidation numbers: magnesium rises from 0 to +2 — oxidised; oxygen falls from 0 to −2 — reduced.
- Write the half equations: oxidation Mg → Mg²⁺ + 2e⁻; reduction O₂ + 4e⁻ → 2O²⁻.
- Multiply the oxidation half equation by 2 to equalise the electrons, then add it to the reduction half to obtain 2Mg + O₂ → 2MgO.
- Answer: Magnesium is oxidised (0 → +2) and oxygen is reduced (0 → −2); the half equations are Mg → Mg²⁺ + 2e⁻ and O₂ + 4e⁻ → 2O²⁻, combining to 2Mg + O₂ → 2MgO, in which magnesium is the reducing agent and oxygen the oxidizing agent.
6Common mistakes Mistakes
7Quick check Quick check
Q1What is the essential difference between oxidation and reduction?
Q2In Zn + CuSO₄ → ZnSO₄ + Cu, which species is the oxidizing agent?
Q3Write the half equation for the oxidation of zinc.
Q4In the blast furnace, carbon monoxide reduces iron oxide to iron. Which species is oxidised?
Q5Why must oxidation and reduction occur together?
8Recap Recap
- Oxidation = loss of electrons, oxidation number rises; reduction = gain of electrons, number falls.
- Oxidation and reduction always occur together; electrons lost equal electrons gained.
- An oxidizing agent oxidises another and is itself reduced; a reducing agent reduces another and is itself oxidised.
- Common oxidizing agents include oxygen and potassium permanganate; common reducing agents include carbon, carbon monoxide and hydrogen.
- Half equations in three steps: write the species change, balance atoms, balance charge with electrons.