IB Edexcel Chemistry: Redox Essentials | IB Edexcel 化学：氧化还原考点精讲

📚 IB Edexcel Chemistry: Redox Essentials | IB Edexcel 化学：氧化还原考点精讲

Redox reactions form the backbone of electrochemistry and appear throughout IB and Edexcel chemistry syllabuses. Mastering oxidation numbers, half-equations, cell potentials, and electrolysis predictions is essential for high marks.

氧化还原反应是电化学的主干，贯穿 IB 和 Edexcel 化学课程。掌握氧化数、半反应、电池电势和电解产物预测是赢得高分的关键。

1. What is Redox? | 什么是氧化还原？

A redox reaction involves the simultaneous transfer of electrons between species. Oxidation is the loss of electrons; reduction is the gain of electrons. The mnemonic ‘OIL RIG’ (Oxidation Is Loss, Reduction Is Gain) is widely used.

氧化还原反应涉及物种之间电子的同时转移。氧化是失去电子，还原是得到电子。记忆口诀“OIL RIG”（氧化失电子，还原得电子）被广泛使用。

In terms of oxidation numbers, oxidation is an increase in oxidation number, while reduction is a decrease.

从氧化数的角度来看，氧化意味着氧化数升高，还原则意味着氧化数降低。

These two processes always occur together; you cannot have oxidation without reduction.

这两个过程总是同时发生；不可能只有氧化而没有还原。

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2. Oxidation Numbers: Rules and Applications | 氧化数：规则与应用

Oxidation numbers are bookkeeping tools to track electron distribution. The key rules are:

氧化数是追踪电子分布的工具。关键规则如下：

An atom in its elemental form has oxidation number 0 (e.g., O₂, Na, S₈).

单质中的原子氧化数为 0（如 O₂、Na、S₈）。

For a monatomic ion, the oxidation number equals the ion’s charge (e.g., Na⁺ = +1, Cl⁻ = −1).

对于单原子离子，氧化数等于离子所带电荷（例如 Na⁺ = +1, Cl⁻ = −1）。

Oxygen normally has oxidation number −2, except in peroxides where it is −1 and in OF₂ where it is +2.

氧的氧化数通常为 −2，但在过氧化物中为 −1，在 OF₂ 中为 +2。

Hydrogen has +1 when bonded to non-metals and −1 when bonded to metals (hydrides).

氢与非金属结合时氧化数为 +1，与金属结合时（氢化物）为 −1。

The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion it equals the ion’s charge.

中性化合物中氧化数总和为 0；多原子离子中氧化数总和等于离子所带电荷。

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3. Identifying Redox Reactions | 识别氧化还原反应

A reaction is redox if any atom changes its oxidation number. Disproportionation is a special redox where the same element is simultaneously oxidized and reduced (e.g., 2 H₂O₂ → 2 H₂O + O₂; oxygen goes from −1 to −2 and 0).

只要任一原子的氧化数发生变化，该反应就是氧化还原反应。歧化反应是一种特殊的氧化还原过程，同一元素同时被氧化和还原（例如 2 H₂O₂ → 2 H₂O + O₂；氧从 −1 变为 −2 和 0）。

Many reactions such as acid-base or precipitation are not redox because oxidation numbers remain unchanged.

许多反应，如酸碱反应或沉淀反应，并非氧化还原反应，因为氧化数保持不变。

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4. Oxidizing and Reducing Agents | 氧化剂与还原剂

An oxidizing agent (oxidant) accepts electrons and is itself reduced. A reducing agent (reductant) donates electrons and is itself oxidized.

氧化剂接受电子，自身被还原。还原剂提供电子，自身被氧化。

Common oxidizing agents include MnO₄⁻ (in acidic medium), Cr₂O₇²⁻, and halogens. Common reducing agents include metals such as Zn, Fe²⁺, and I⁻.

常见氧化剂包括酸性条件下的 MnO₄⁻、Cr₂O₇²⁻ 和卤素。常见还原剂包括金属如 Zn、Fe²⁺ 和 I⁻。

The strength of an oxidizing agent increases with more positive standard reduction potential.

氧化剂的强度随着标准还原电势更正而增强。

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5. Balancing Redox Equations: Half-Reaction Method | 半反应法配平

In acidic solution: split into oxidation and reduction half-equations. Balance all atoms except O and H; add H₂O to balance O; add H⁺ to balance H; add electrons (e⁻) to balance charge. Multiply half-equations so electrons cancel, then add and simplify.

在酸性溶液中：拆分为氧化和还原半反应。配平除 O 和 H 以外的所有原子；加 H₂O 配平 O；加 H⁺ 配平 H；加电子 (e⁻) 配平电荷。将半反应乘以适当倍数使电子抵消，然后相加并化简。

Example: MnO₄⁻ + 5 Fe²⁺ + 8 H⁺ → Mn²⁺ + 5 Fe³⁺ + 4 H₂O. The manganese reduction half-equation is MnO₄⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H₂O; iron oxidation is Fe²⁺ → Fe³⁺ + e⁻.

示例：MnO₄⁻ + 5 Fe²⁺ + 8 H⁺ → Mn²⁺ + 5 Fe³⁺ + 4 H₂O。锰的还原半反应为 MnO₄⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H₂O；铁的氧化为 Fe²⁺ → Fe³⁺ + e⁻。

In basic solution, after balancing as if in acid, add OH⁻ to both sides to neutralise H⁺, forming water.

在碱性溶液中，先假设酸性条件配平，然后在两边加上 OH⁻ 中和 H⁺，生成水。

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6. Electrochemical Cells | 电化学电池

A galvanic (voltaic) cell converts chemical energy into electrical energy through spontaneous redox. It consists of two half-cells connected by a salt bridge, with electrons flowing through an external circuit.

原电池通过自发的氧化还原反应将化学能转化为电能。它由两个半电池通过盐桥连接组成，电子经外电路流动。

Oxidation occurs at the anode (negative in galvanic cell), and reduction occurs at the cathode (positive). The salt bridge maintains electrical neutrality by allowing ion migration.

氧化发生在阳极（原电池中为负极），还原发生在阴极（正极）。盐桥通过允许离子迁移维持电中性。

The cell diagram notation is: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). A single vertical line represents phase boundary, double line represents salt bridge.

电池图示记法为：Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)。单竖线表示相界面，双竖线表示盐桥。

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7. Standard Electrode Potentials (E°) | 标准电极电势

The standard electrode potential is the voltage of a half-cell relative to the standard hydrogen electrode (SHE), measured under standard conditions: 298 K, 1 mol dm⁻³ ion concentration, 1 atm for gases.

标准电极电势是半电池相对于标准氢电极（SHE）的电压，在标准条件下测量：298 K、离子浓度 1 mol dm⁻³、气体压强 1 atm。

By convention, the SHE has E° = 0.00 V. Half-equations are always written as reduction, but the actual direction depends on the other half-cell.

按照惯例，SHE 的 E° = 0.00 V。半反应通常写作还原反应，但实际进行的方向取决于另一个半电池。

E° values are intensive properties; they do not change when the half-equation is multiplied by a factor.

E° 值是强度性质；当半反应乘以某个倍数时，其数值不变。

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8. The Electrochemical Series and Predicting Reactions | 电化学序列与预测反应

The electrochemical series arranges half-cells by increasing reduction potential. A species higher up the series is a stronger reducing agent (more negative E°), while one lower down is a stronger oxidizing agent (more positive E°).

电化学序列按还原电势递增排列半电池。位于序列上方的物种是较强的还原剂（E° 更负），位于下方的物种是较强的氧化剂（E° 更正）。

The standard cell potential is calculated as:

E°cell = E°cathode – E°anode

where both E° are reduction potentials. A positive E°cell indicates a spontaneous reaction.

标准电池电势计算公式为：E°cell = E°cathode – E°anode，其中两者均为还原电势。正的 E°cell 表示反应自发进行。

You can predict whether a metal will displace another from a solution: e.g., Zn (E° = -0.76 V) reduces Cu²⁺ (E° = +0.34 V) because E°cell = +0.34 – (-0.76) = +1.10 V > 0.

你可以预测一种金属能否从溶液中置换另一种金属：例如 Zn (E° = -0.76 V) 可还原 Cu²⁺ (E° = +0.34 V)，因为 E°cell = +0.34 – (-0.76) = +1.10 V > 0。

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9. Nernst Equation and Concentration Effects | 能斯特方程与浓度效应

Under non-standard conditions, the cell potential deviates from E°. The Nernst equation (for 298 K) is often simplified to:

E = E° – (0.0592 V / n) log Q

where n is the number of electrons transferred and Q is the reaction quotient.

在非标准条件下，电池电势会偏离 E°。能斯特方程（298 K 下）通常简化为：E = E° – (0.0592 V / n) log Q，其中 n 为转移电子数，Q 为反应商。

Increasing the concentration of reactants or decreasing products shifts E to more positive values, driving the reaction forward more strongly.

增加反应物浓度或减少产物浓度会使 E 更正，从而更强烈地驱动反应正向进行。

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10. Electrolysis: Principles and Predictions | 电解：原理与预测

Electrolysis uses an external power source to force a non-spontaneous redox reaction. Reduction occurs at the cathode (negative electrode) and oxidation at the anode (positive electrode).

电解利用外部电源迫使非自发的氧化还原反应发生。阴极（负极）发生还原，阳极（正极）发生氧化。

In molten salts, the cation is reduced at the cathode and the anion oxidized at the anode (e.g., 2 Cl⁻ → Cl₂ + 2 e⁻).

在熔融盐中，阳离子在阴极被还原，阴离子在阳极被氧化（例如 2 Cl⁻ → Cl₂ + 2 e⁻）。

In aqueous solutions, water may be oxidized or reduced instead. To predict products, compare the reduction potentials of the possible reactions and consider overpotential effects for gases like O₂ or Cl₂.

在水溶液中，水也可能被氧化或还原。要预测产物，需要比较可能反应的还原电势，并考虑氧气或氯气等气体的超电位效应。

Quantitative electrolysis uses Faraday’s laws: Q = I × t, and n(e⁻) = Q / F, where F = 96500 C mol⁻¹.

定量电解使用法拉第定律：Q = I × t，n(e⁻) = Q / F，其中 F = 96500 C mol⁻¹。

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11. Redox Titrations | 氧化还原滴定

Manganate(VII) titrations with Fe²⁺ are classic. Purple MnO₄⁻ is reduced to nearly colourless Mn²⁺ in acidic solution, so the endpoint is a faint pink without an external indicator.

高锰酸根与 Fe²⁺ 的滴定是经典实验。在酸性溶液中，紫色的 MnO₄⁻ 被还原为近乎无色的 Mn²⁺，因此终点呈淡粉色，无需外加指示剂。

The reaction stoichiometry: 5 Fe²⁺ + MnO₄⁻ + 8 H⁺ → 5 Fe³⁺ + Mn²⁺ + 4 H₂O. One mole of MnO₄⁻ reacts with five moles of Fe²⁺.

反应计量关系：5 Fe²⁺ + MnO₄⁻ + 8 H⁺ → 5 Fe³⁺ + Mn²⁺ + 4 H₂O。1 mol MnO₄⁻ 与 5 mol Fe²⁺ 反应。

Iodine-thiosulfate titrations are another common example: 2 S₂O₃²⁻ + I₂ → S₄O₆²⁻ + 2 I⁻. Starch is used as an indicator, added near the endpoint.

碘-硫代硫酸盐滴定是另一个常见例子：2 S₂O₃²⁻ + I₂ → S₄O₆²⁻ + 2 I⁻。淀粉用作指示剂，接近终点时加入。

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12. Exam Tips and Common Mistakes | 考试技巧与常见错误

Always write the oxidation number sign before the number (+2 not 2+), distinguishing charge of ions (2+). Many students confuse the two.

始终将氧化数的符号写在数字前面（+2 而不是 2+），以区别于离子电荷 (2+)。许多学生会混淆两者。

When calculating E°cell, use E°cathode – E°anode without flipping signs. Both are reduction potentials taken directly from the data booklet.

计算 E°cell 时，直接用 E°cathode – E°anode，不要改变符号。两者均为直接从数据手册中获取的还原电势。

In half-equation balancing, double-check that both mass and charge are balanced. Never add extra species like O₂ unless it appears in the overall reaction.

配平半反应时，要仔细核对质量与电荷是否均配平。切勿加入 O₂ 等额外物种，除非它们出现在总反应中。

For electrolysis, remember that water can compete: in aqueous NaCl, Cl₂ is produced at the anode, not O₂, due to high overpotential of O₂ on common electrodes.

对于电解，记住水可能参与竞争：在 NaCl 水溶液中，阳极产生 Cl₂ 而非 O₂，因为 O₂ 在常见电极上的超电位较高。

Use titration data carefully: n(Fe²⁺) = 5 × n(MnO₄⁻). Forgetting the ratio is a frequent error.

使用滴定数据时要小心：n(Fe²⁺) = 5 × n(MnO₄⁻)。忘记这一比例是常见错误。

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