Appendix 2: Selected standard electrode potentials | 附录2：部分标准电极电势

📚 Appendix 2: Selected standard electrode potentials | 附录2：部分标准电极电势

In Cambridge A-Level Chemistry, standard electrode potentials are the bedrock of electrochemistry. Appendix 2 of the data booklet provides a carefully selected table of E° values that allow you to compare the relative oxidising or reducing power of different species, predict whether a redox reaction is feasible, and calculate the emf of an electrochemical cell under standard conditions. This article unpacks the meaning, measurement, and use of these values so you can confidently apply them in exam calculations and interpretation questions.

在剑桥A-Level化学中，标准电极电势是电化学的基石。数据手册中的附录2提供了一张精心挑选的E°数据表，使你能够比较不同物种的相对氧化或还原能力，预测氧化还原反应是否可行，并计算标准条件下电化学电池的电动势。本文将剖析这些数值的含义、测量方法和使用技巧，帮助你自信地应对考试中的计算和分析题。

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1. Introduction to Standard Electrode Potentials | 标准电极电势简介

A standard electrode potential, E°, is the voltage measured when a half-cell is connected to a standard hydrogen electrode under standard conditions (298 K, 1.0 mol dm⁻³ ion concentration, 100 kPa pressure for gases). It reflects the tendency of a species to be reduced: the more positive the E° value, the stronger the oxidising agent (and the more easily the species gains electrons). Conversely, a very negative E° indicates a strong reducing agent.

标准电极电势E°是指在标准条件下（298 K、离子浓度1.0 mol dm⁻³、气体压强100 kPa），将某个半电池与标准氢电极相连时测得的电压。它反映了物种被还原的倾向：E°值越正，氧化剂越强（该物种越容易得电子）。反之，E°值越负，则代表还原剂越强。

→ Red. By convention, all half-cell potentials are measured relative to the standard hydrogen electrode, which is assigned a potential of exactly 0.00 V.

这些电势始终以还原电势的形式给出，即半反应方程写为 Ox + n e⁻ → Red 的形式。按照惯例，所有半电池的电势均相对于标准氢电极测量，标准氢电极的电势被指定为精确的0.00 V。

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2. The Standard Hydrogen Electrode (SHE) | 标准氢电极

The SHE is the universal reference point for electrode potentials. It consists of a platinum electrode coated with finely divided platinum (platinum black) immersed in a solution containing H⁺ ions at 1.0 mol dm⁻³, with hydrogen gas bubbled through at 100 kPa. The half-reaction is 2H⁺(aq) + 2e⁻ ⇌ H₂(g), and its E° is defined as 0.00 V at all temperatures.

标准氢电极是所有电极电势的通用参考点。它由一根镀有铂黑（细粉状铂）的铂电极构成，浸入含有1.0 mol dm⁻³ H⁺离子的溶液中，并通入压强为100 kPa的氢气。其半反应为2H⁺(aq) + 2e⁻ ⇌ H₂(g)，在所有温度下其E°都定义为0.00 V。

In practice, the SHE is difficult to set up, so secondary reference electrodes (such as the silver–silver chloride electrode) are often used in the laboratory. However, all E° values in Appendix 2 are still reported relative to the SHE, allowing consistent comparisons.

实际操作中，标准氢电极较难搭建，因此实验室常使用二级参比电极（例如银–氯化银电极）。但附录2中的所有E°值仍以标准氢电极为基准报告，从而确保比较的一致性。

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3. Definition of Standard Electrode Potential E° | 标准电极电势E°的定义

The standard electrode potential of a half-cell is defined as the emf of a cell composed of that half-cell as the right-hand electrode and the standard hydrogen electrode as the left-hand electrode, with all species at unit activity (approximated by 1 mol dm⁻³ for solutions, 100 kPa for gases) and a temperature of 298 K. The sign of E° is the sign of the right-hand electrode as it is connected to the voltmeter.

某一半电池的标准电极电势定义为：将该半电池作为右侧电极，标准氢电极作为左侧电极，组成一个电池，在溶液离子活度为1（近似为1 mol dm⁻³）、气体压强为100 kPa、温度为298 K的条件下，测得的电池电动势。E°的符号即是右侧电极在电压表上显示的极性符号。

It is crucial to remember that E° is an intensive property: it does not depend on the amount of substance. Doubling the half-equation does not change the E° value. The data in Appendix 2 list E° for the reduction process, which allows direct comparison of oxidising strengths.

务必记住，E°是一个强度性质，与物质的量无关。将半反应方程乘以2并不会改变E°数值。附录2中的数据列出的是还原过程的E°，因此能够直接比较氧化能力的强弱。

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4. Measuring Standard Electrode Potentials | 测量标准电极电势

ter. The voltmeter reading gives the cell emf. Under standard conditions, if the Zn electrode is connected to the negative terminal of the voltmeter, the measured voltage will be -0.76 V, meaning E°(Zn²⁺/Zn) = -0.76 V.

测量未知E°时，需将含有待测电对（如Zn²⁺/Zn）的半电池通过盐桥和高阻抗电压表与标准氢电极相连。电压表的读数即为电池电动势。在标准条件下，若锌电极与电压表的负极相连，则测得的电压为-0.76 V，即E°(Zn²⁺/Zn) = -0.76 V。

The arrangement ensures that the potential difference is measured under virtually zero current flow, so the measurement reflects the equilibrium potential rather than a potential affected by reaction kinetics. A salt bridge (often filter paper soaked in KNO₃) completes the circuit while preventing the two solutions from mixing directly.

这样的布置确保在几乎无电流通过的情况下测量电位差，因此所测得的是平衡电位，而非受动力学影响的电位。盐桥（通常是浸有KNO₃的滤纸）既能连通电路，又能防止两种溶液直接混合。

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5. Selected Standard Electrode Potentials Table | 部分标准电极电势表

Appendix 2 presents a curated list of electrode potentials commonly encountered in A-Level chemistry. The table below reproduces a selection of these values, all measured at 298 K. Notice that the half-reactions are written as reduction processes, and the E° values range from very negative (strong reducing agents) to very positive (strong oxidising agents).

附录2给出了一张A-Level化学中常见的精选电极电势表。下表复现了其中一部分数值，均在298 K下测得。请注意，半反应均写成还原过程，E°值从极负（强还原剂）一直跨越到极正（强氧化剂）。

Half-reaction (reduction)	E° / V
Li⁺ + e⁻ ⇌ Li	−3.04
K⁺ + e⁻ ⇌ K	−2.92
Ca²⁺ + 2e⁻ ⇌ Ca	−2.87
Na⁺ + e⁻ ⇌ Na	−2.71
Mg²⁺ + 2e⁻ ⇌ Mg	−2.37
Al³⁺ + 3e⁻ ⇌ Al	−1.66
Zn²⁺ + 2e⁻ ⇌ Zn	−0.76
Fe²⁺ + 2e⁻ ⇌ Fe	−0.44
2H⁺ + 2e⁻ ⇌ H₂	0.00
Cu²⁺ + 2e⁻ ⇌ Cu	+0.34
I₂ + 2e⁻ ⇌ 2I⁻	+0.54
Fe³⁺ + e⁻ ⇌ Fe²⁺	+0.77
Ag⁺ + e⁻ ⇌ Ag	+0.80
Br₂ + 2e⁻ ⇌ 2Br⁻	+1.07
Cl₂ + 2e⁻ ⇌ 2Cl⁻	+1.36
MnO₄⁻ + 8H⁺ + 5e⁻ ⇌ Mn²⁺ + 4H₂O	+1.51

Notice that reactive metals like lithium and potassium have very negative E° values, indicating they are easily oxidised. Halogens, on the other hand, have positive E° values, increasing down the group, making fluorine the strongest oxidising agent in its series (not shown above due to its extreme reactivity, but its E° is +2.87 V).

注意，锂、钾等活泼金属具有很负的E°值，表明它们容易被氧化。而卤素则具有正的E°值，并且沿着族往下递增，使得氟成为该系列中最强的氧化剂（由于活性过强上表中未列出，但其E°为+2.87 V）。

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6. Using E° Values to Predict Redox Reactions | 利用E°值预测氧化还原反应

quations and subtracting the more negative (or less positive) E° from the more positive E°: E°cell = E°(reduced species) – positive E° values tend to undergo reduction themselves while oxidising the other species.

如果计算出的电池电动势E°cell为正值，则该氧化还原反应在热力学上是可行的。E°cell可以通过组合两个半反应方程，用较正的E°减去较负（或较不正）的E°得到：E°cell = E°(被还原物种) – E°(被氧化物种)。E°值更正的物种倾向于自身被还原，并氧化另一方。

For example, will zinc metal reduce copper(II) ions? The relevant half-reactions are: Zn²⁺ + 2e⁻ ⇌ Zn (E° = −0.76 V) and Cu²⁺ + 2e⁻ ⇌ = +0.34 V). Zinc must be oxidised, so we reverse its half-equation. Then E°cell = +0.34 V – ce the reaction Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) is feasible under standard conditions.

例如，锌金属能否还原铜(II)离子？相关的半反应为：Zn²⁺ + 2e⁻ ⇌ Zn (E° = −0.76 V) 和 Cu²⁺ + 2e⁻ ⇌ Cu (E° = +0.34 V)。锌必须被氧化，因此将其半反应反转。那么E°cell = +0.34 V – (−0.76 V) = +1.10 V，该值为正。因此在标准条件下，反应Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)是可行的。

A common exam trick is to ask whether a reaction is feasible, but the rate may be so slow that no observable change occurs despite a positive E°cell. This highlights the distinction between thermodynamics and kinetics.

考试中常见的陷阱是要求判断反应是否可行，但即便E°cell为正值，反应速率也可能极慢，导致观察不到任何变化。这凸显了热力学与动力学之间的区别。

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7. Calculating Cell EMF (E°cell) | 计算电池电动势（E°电池）

To calculate the standard cell emf, write the two reduction half-equations as they appear in Appendix 2. Identify which species is reduced (higher E°) and which is oxidised (lower E°). The cell emf is then:

E°cell = E°(cathode) – E°(anode)

要计算标准电池电动势，先按附录2的形式写出两个还原半反应。确定哪个物种被还原（E°较高者）和哪个被氧化（E°较低者）。然后电池电动势为：

E°cell = E°(阴极) – E°(阳极)

In a galvanic cell, the cathode is the electrode where reduction occurs (positive terminal), and the anode is where oxidation occurs (negative terminal). The sign of the voltmeter reading will match the sign of the right-hand electrode if the cell is constructed with the cathode on the right.

在原电池中，阴极是发生还原反应的电极（正极），阳极是发生氧化反应的电极（负极）。如果构建电池时阴极置于右侧，电压表读数的符号将与右侧电极的符号一致。

When balancing the overall equation, do not multiply the E° values by any coefficients. The cell emf remains the same because E° is an intensive property. This is a frequent source of error in exam answers.

在配平总反应方程时，切勿给E°值乘以任何系数。电池电动势保持不变，因为E°是强度性质。这是考试答案中常见的错误来源。

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8. Relationship with Gibbs Free Energy | 与吉布斯自由能的关系

The thermodynamic feasibility indicated by E°cell is directly linked to the change in Gibbs free energy, ΔG°, through the equation:

ΔG° = –n F E°cell

E°cell所指示的热力学可行性与吉布斯自由能的变化ΔG°直接相关，方程如下：

ΔG° = –n F E°cell

where n is the number of moles of electrons transferred in the balanced redox equation, and F is the Faraday constant (approximately 96 500 C mol⁻¹). A positive E°cell gives a negative ΔG°, confirming that the reaction is spontaneous under standard conditions.

其中n是配平的氧化还原方程中转移电子的物质的量，F为法拉第常数（约为96 500 C mol⁻¹）。正的E°cell导致ΔG°为负，证实该反应在标准条件下是自发的。

This relationship also allows the calculation of an unknown E° from thermochemical data, and it explains why some reactions with positive E°cell still appear not to proceed: a large activation energy may mean the reaction is under kinetic control, and ΔG° alone does not guarantee a fast rate.

这一关系也使我们能够根据热化学数据计算未知的E°值，并解释了为什么某些E°cell为正的反应似乎仍不进行：较高的活化能可能意味着反应受动力学控制，仅凭ΔG°并不能保证反应速率很快。

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9. Effect of Concentration and the Nernst Equation | 浓度的影响与能斯特方程

The E° values in Appendix 2 apply strictly under standard conditions. When concentrations, pressures, or temperature deviate from standard values, the electrode potential changes according to the Nernst equation. For a half-reaction a Ox + n e⁻ ⇌ b Red, the Nernst equation at 298 K can be written as:

E = E° – (0.0592 / n) log₁₀ ( [Red]ᵇ / [Ox]ᵃ )

附录2中的E°值严格适用于标准条件。当浓度、压强或温度偏离标准值时，电极电势将根据能斯特方程发生变化。对于半反应 a Ox + n e⁻ ⇌ b Red，298 K下的能斯特方程可写作：

E = E° – (0.0592 / n) log₁₀ ( [Red]ᵇ / [Ox]ᵃ )

Here, 0.0592 is the value of (RT ln 10 / F) at 298 K. This equation shows that increasing the concentration of the oxidised form makes the potential more positive (stronger oxidising agent), while increasing the reduced form makes it more negative.

这里0.0592是298 K下(RT ln 10 / F)的数值。该方程表明，增大氧化型物种的浓度会使电势变得更正（氧化剂更强），而增大还原型物种的浓度则会使电势变得更负。

In a concentration cell, two half-cells with the same chemistry but different ion concentrations can generate a voltage. This concept is used in pH meters and ion-selective electrodes, extending the utility of electrode potentials far beyond Appendix 2.

在浓差电池中，两个化学性质相同但离子浓度不同的半电池可以产生电压。这一概念被应用于pH计和离子选择性电极，使电极电势的用途远远超出了附录2的范畴。

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10. Limitations and Practical Considerations | 局限性和实际考量

While E° values are powerful predictive tools, they have limitations. First, they assume all species are at unit activity, which may not reflect real solution behaviour, particularly for highly charged ions where ion-pairing occurs. Second, kinetic factors can prevent a thermodynamically favourable reaction from occurring at an appreciable rate – a classic example is the reaction between magnesium metal and water, which has a positive E°cell but is very slow at room temperature.

尽管E°值是强有力的预测工具，但它们也有局限性。首先，它们假定所有物种的活度为1，这可能无法反映真实溶液的行为，尤其是对于易发生离子配对的高电荷离子。其次，动力学因素可能阻止热力学上有利的反应以可观速率进行——一个经典例子是金属镁与水的反应，虽然E°cell为正值，但在室温下反应极其缓慢。

Furthermore, the standard hydrogen electrode is a reference that exists by definition; in real measurements, liquid junction potentials and non-standard temperatures introduce small errors. Appendix 2 values are also given for aqueous systems, and many electrode potentials for non-aqueous solvents or molten salts differ significantly.

此外，标准氢电极是一个定义性基准，在实际测量中，液接电位和非标准温度都会引入微小误差。附录2中的数值也仅适用于水溶液体系，许多非水溶剂或熔盐中的电极电势会有很大差异。

In exams, you are expected to use the provided data exactly as they appear, without adjusting for these real-world factors. Recognising these limitations, however, deepens your understanding of the difference between an idealised model and experimental chemistry.

在考试中，你需要完全按照所给数据使用它们，无需针对这些实际因素进行调整。然而，认识到这些局限性可以加深你对理想模型与实验化学之间区别的理解。

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11. Summary of Key Points | 要点总结

eference. A positive E°cell indicates a thermodynamically feasible reaction, and the magnitude of E°cell is proportional to the negative Gibbs free energy change. When using the data, always apply the formula E°cell = E°(cathode) – ode), never multiplying E° by stoichiometric coefficients. Be aware of the Nernst equation' when conditions change, and remember that kinetics can override thermodynamic predictions.

附录2中的标准电极电势根据物种被还原的倾向进行排序。标准氢电极是通用的零基准。正的E°cell表明反应在热力学上可行，且E°cell的大小与负的吉布斯自由能变化成正比。使用数据时，始终采用公式E°cell = E°(阴极) – E°(阳极)，切勿给E°乘以化学计量系数。要了解当条件改变时能斯特方程的作用，并记住动力学可能凌驾于热力学预测之上。

Mastering these principles not only prepares you for direct electrochemistry questions but also supports broader topics such as transition metal redox titrations, corrosion, and industrial electrolysis. Make Appendix 2 your trusted companion in reasoning about electron flow.

掌握这些原理不仅能帮助你应对直接的电化学考题，还能支持更广泛的主题，如过渡金属氧化还原滴定、腐蚀和工业电解。让附录2成为你推断电子流向的可靠伙伴。

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