📚 AS Chemistry: Acid-Base Theories Explained | AS 化学：酸碱理论 考点精讲
Acid-base chemistry is a cornerstone of A-level Chemistry, linking concepts of bonding, equilibria, and reaction mechanisms. Understanding the historical development of acid-base theories—from Arrhenius to Lewis—equips you with the tools to explain a wide range of chemical behaviour, whether in aqueous solution or beyond. This revision guide covers all essential theories, definitions, conjugate pairs, pH calculations, and key constants required for AS level, with clear explanations and worked examples.
酸碱化学是 A-level 化学的基石，它将化学键、平衡和反应机理等概念联系起来。理解酸碱理论的发展历程——从阿伦尼乌斯到路易斯——能帮助你解释水溶液乃至非水体系中的众多化学行为。这份复习指南全面梳理了 AS 阶段必须掌握的酸碱理论、定义、共轭对、pH 计算和关键常数，并配有清晰的解释与例题。
1. Introduction to Acids and Bases | 酸碱简介
Acids and bases are encountered everywhere, from laboratory reagents to biological systems. Over time, chemists have proposed several theories to define what constitutes an acid or a base. Each successive theory built upon the limitations of its predecessors, extending the scope of reactions that can be rationalised. At AS level, you are expected to distinguish between the Arrhenius, Brønsted-Lowry, and Lewis definitions, and apply them to predict products and explain reactivity.
酸和碱无处不在，从实验室试剂到生物体系都有它们的身影。随着化学的发展，科学家们提出了多种理论来定义什么是酸、什么是碱。每一种新理论都是在克服前一种理论的局限性的基础上建立的，从而能够解释更广泛的反应。在 AS 阶段，你需要区分阿伦尼乌斯、布朗斯特-劳里和路易斯三种酸碱定义，并能运用它们预测产物、解释反应活性。
2. The Arrhenius Theory | 阿伦尼乌斯酸碱理论
Arrhenius defined an acid as a substance that dissociates in water to produce hydrogen ions, H⁺. For example, hydrogen chloride gas dissolves in water to form hydrochloric acid, which is fully dissociated into H⁺ and Cl⁻ ions.
阿伦尼乌斯将酸定义为在水溶液中离解产生氢离子 H⁺ 的物质。例如，氯化氢气体溶于水形成盐酸，完全离解为 H⁺ 和 Cl⁻ 离子。
HCl(g) + aq → H⁺(aq) + Cl⁻(aq)
Similarly, an Arrhenius base dissociates in water to yield hydroxide ions, OH⁻. Sodium hydroxide is a classic example:
类似地，阿伦尼乌斯碱是指在水中离解产生氢氧根离子 OH⁻ 的物质。氢氧化钠就是一个典型例子：
NaOH(s) + aq → Na⁺(aq) + OH⁻(aq)
The Arrhenius theory successfully describes many neutralisation reactions as H⁺ + OH⁻ → H₂O, but it is limited to aqueous systems and cannot explain basic behaviour of substances like ammonia (NH₃) that lack OH⁻ in their formula.
阿伦尼乌斯理论成功地将许多中和反应描述为 H⁺ + OH⁻ → H₂O，但它只适用于水溶液体系，无法解释氨 (NH₃) 这类分子式不含 OH⁻ 却表现出碱性的物质。
3. The Brønsted-Lowry Theory | 布朗斯特-劳里酸碱理论
In 1923, Brønsted and Lowry independently proposed a more general definition: an acid is a proton (H⁺) donor, and a base is a proton acceptor. This proton-transfer model no longer requires the presence of water, although aqueous solutions are still commonly used.
1923 年，布朗斯特和劳里各自独立地提出了一个更普适的定义：酸是质子 (H⁺) 的给体，碱是质子的受体。这个质子传递模型不再要求必须有水存在，尽管水溶液仍然是最常见的情况。
When hydrogen chloride gas reacts with ammonia gas, HCl donates a proton to NH₃, forming ammonium chloride. HCl is the Brønsted-Lowry acid, and NH₃ is the base.
当氯化氢气体与氨气反应时，HCl 把质子给了 NH₃，生成氯化铵。在此反应中，HCl 是布朗斯特-劳里酸，NH₃ 是碱。
HCl(g) + NH₃(g) → NH₄⁺Cl⁻(s)
In water, the reaction between HCl and H₂O is also viewed as a proton transfer: HCl donates a proton to H₂O, producing the hydronium ion H₃O⁺ and Cl⁻. Water acts as a base here.
在水溶液中，HCl 与 H₂O 的反应同样可看作质子传递：HCl 将质子给予 H₂O，生成水合氢离子 H₃O⁺ 和 Cl⁻。这里水充当了碱的角色。
HCl + H₂O → H₃O⁺ + Cl⁻
Brønsted-Lowry theory elegantly explains why aqueous ammonia is basic: NH₃ accepts a proton from water, leaving OH⁻ ions in solution.
布朗斯特-劳里理论很自然地解释了为什么氨水呈碱性：NH₃ 从水分子中接受一个质子，使得溶液中留下了 OH⁻ 离子。
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
4. Conjugate Acid-Base Pairs | 共轭酸碱对
A central concept in Brønsted-Lowry theory is that every acid has a conjugate base, formed after it donates a proton, and every base has a conjugate acid, formed after it accepts a proton. The pair differs by exactly one H⁺.
布朗斯特-劳里理论的一个核心概念是：每种酸都有一个对应的共轭碱（酸给出质子后形成），每种碱都有一个对应的共轭酸（碱接受质子后形成）。一个共轭酸碱对之间只相差一个 H⁺。
For the reaction HCl + H₂O → H₃O⁺ + Cl⁻, the conjugate pairs are HCl / Cl⁻ and H₃O⁺ / H₂O. The stronger the acid, the weaker its conjugate base. Strong acids like HCl have negligible conjugate base strength, whereas weak acids like ethanoic acid produce relatively stronger conjugate bases (ethanoate ion).
在反应 HCl + H₂O → H₃O⁺ + Cl⁻ 中，共轭酸碱对是 HCl / Cl⁻ 和 H₃O⁺ / H₂O。酸越强，其共轭碱就越弱。像 HCl 这样的强酸，其共轭碱 Cl⁻ 几乎无碱性；而弱酸如乙酸，其共轭碱乙酸根离子的碱性就相对较强。
Water is an amphiprotic solvent – it can act as both an acid and a base depending on the reaction partner. In the presence of a stronger base, water donates a proton; in the presence of a stronger acid, it accepts one.
水是一种两性溶剂——它既能作酸也能作碱，取决于反应对象。遇到更强的碱时，水给出质子；遇到更强的酸时，水接受质子。
H₂O + H₂O ⇌ H₃O⁺ + OH⁻
5. Amphoteric Substances | 两性物质
Some species can react as either a Brønsted-Lowry acid or a base, depending on the conditions. These are called amphoteric (or amphiprotic when referring specifically to proton transfer). Water is the most common example, but many metal oxides and hydroxides also display amphoteric character.
有些物质既可作为布朗斯特-劳里酸，也可作为其碱，这取决于反应条件。它们被称为两性物质（在水溶液中涉及质子转移时也常称为两性电解质）。水是最常见的例子，但许多金属氧化物和氢氧化物也表现出两性特征。
Aluminium hydroxide, Al(OH)₃, reacts with acids to form Al³⁺ salts, and with bases to form aluminate ions [Al(OH)₄]⁻. Amino acids, with both an amine group (–NH₂) and a carboxyl group (–COOH), are excellent biological examples of amphoteric behaviour.
氢氧化铝 Al(OH)₃ 既能与酸反应生成铝盐，也能与碱反应生成铝酸根离子 [Al(OH)₄]⁻。氨基酸含有一个氨基 (–NH₂) 和一个羧基 (–COOH)，是体现两性行为的绝佳生物例子。
6. Lewis Acid-Base Theory | 路易斯酸碱理论
Gilbert N. Lewis proposed an even broader theory based on electron pair donation and acceptance. A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor. This definition does not require hydrogen, protons, or even an aqueous environment.
吉尔伯特·路易斯提出了一个更为宽泛的、基于电子对给予和接受的理论。路易斯酸是电子对受体，路易斯碱是电子对给体。这个定义完全不要求有氢原子、质子，甚至不需要水溶液环境。
A classic example is the reaction between boron trifluoride, BF₃ (electron deficient, Lewis acid), and ammonia, NH₃ (lone pair donor, Lewis base), to form a coordinate bond.
经典例子是三氟化硼 BF₃（缺电子，路易斯酸）和氨 NH₃（孤对电子给体，路易斯碱）反应，生成配位键化合物。
BF₃ + :NH₃ → F₃B–NH₃
Many cations, such as H⁺, Cu²⁺, and Fe³⁺, act as Lewis acids by accepting electron pairs from ligands. This theory unifies classic acid-base reactions with complex ion formation and organic mechanisms.
很多阳离子，如 H⁺、Cu²⁺ 和 Fe³⁺，都能作为路易斯酸接受来自配体的电子对。路易斯理论将传统的酸碱反应与配位离子形成、有机反应机理统一了起来。
| Theory | Acid | Base | Scope |
|---|---|---|---|
| Arrhenius | H⁺ producer in H₂O | OH⁻ producer in H₂O | Aqueous only |
| Brønsted-Lowry | Proton donor | Proton acceptor | Proton transfer, often aqueous |
| Lewis | Electron pair acceptor | Electron pair donor | Broadest (any phase) |
7. Strong and Weak Acids and Bases | 强酸强碱与弱酸弱碱
A strong acid is fully ionised in aqueous solution. Common strong acids include HCl, H₂SO₄ (first ionisation), and HNO₃. Because the dissociation goes to completion, we use a single arrow → in equations.
强酸在水溶液中完全电离。常见的强酸包括 HCl、H₂SO₄（一级电离）和 HNO₃。由于电离进行到底，方程式中使用单向箭头 →。
A weak acid partially dissociates, establishing an equilibrium. Ethanoic acid, CH₃COOH, is typical: only about 1% of molecules donate a proton in a 0.1 mol dm⁻³ solution. The equilibrium is represented with a reversible arrow ⇌.
弱酸仅部分电离，建立起平衡。乙酸 CH₃COOH 是一个典型：在 0.1 mol dm⁻³ 溶液中只有约 1% 的分子给出质子。其平衡方程式用可逆箭头 ⇌ 表示。
CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq)
Strong bases, such as NaOH and KOH, fully dissociate to release OH⁻. Weak bases like ammonia undergo partial protonation, leaving a relatively low concentration of OH⁻. Importantly, strength refers to degree of ionisation, not concentration – a concentrated weak acid may still have a lower pH than a dilute strong acid, but the terminology must be used precisely.
强碱如 NaOH 和 KOH 完全离解释放出 OH⁻。弱碱如氨仅部分质子化，溶液中 OH⁻ 浓度相对较低。需注意的是，强弱指的是电离程度，而非浓度——高浓度的弱酸仍可能比稀强酸的 pH 低，但在使用术语时必须严谨区分。
8. The Ionic Product of Water, Kw | 水的离子积 Kw
Water undergoes slight autoionisation. The equilibrium constant for this process is called the ionic product of water, Kw, and at 298 K it has the value 1.0 × 10⁻¹⁴ mol² dm⁻⁶.
水分子会微弱的自电离。这一过程的平衡常数称为水的离子积 Kw，在 298 K 时其值为 1.0 × 10⁻¹⁴ mol² dm⁻⁶。
2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)
Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ (at 298 K)
Since the concentration of water is essentially constant, it is incorporated into Kw. In pure water, [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol dm⁻³, giving a neutral pH of 7. Kw is temperature dependent; it increases with rising temperature, meaning the pH of neutral water decreases at higher temperatures despite equal concentrations of H₃O⁺ and OH⁻.
由于水的浓度几乎不变，它被并入 Kw 值中。在纯水中，[H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol dm⁻³，中性 pH = 7。Kw 随温度变化；温度升高时 Kw 增大，这意味着尽管 H₃O⁺ 和 OH⁻ 浓度仍然相等，中性水的 pH 值却会降低。
9. pH and pOH | pH 和 pOH
The pH scale is a convenient way to express hydronium ion concentration. pH is defined as the negative logarithm to base 10 of [H₃O⁺]:
pH 标度是表达水合氢离子浓度的一种便捷方式。pH 定义为 [H₃O⁺] 的以10为底的负对数：
pH = –log₁₀[H₃O⁺]
Similarly, pOH = –log₁₀[OH⁻]. At 298 K, pH + pOH = 14. For a 0.01 mol dm⁻³ solution of HCl (strong acid, complete ionisation), [H₃O⁺] = 0.01 mol dm⁻³, so pH = 2. If the solution were 0.05 mol dm⁻³ NaOH, [OH⁻] = 0.05 mol dm⁻³, pOH ≈ 1.30, and pH = 12.70.
类似地，pOH = –log₁₀[OH⁻]。298 K 时，pH + pOH = 14。对于 0.01 mol dm⁻³ 的 HCl 溶液（强酸，完全电离），[H₃O⁺] = 0.01 mol dm⁻³，pH = 2。如果是 0.05 mol dm⁻³ NaOH 溶液，[OH⁻] = 0.05 mol dm⁻³，pOH ≈ 1.30，pH = 12.70。
Always remember that a change of one pH unit corresponds to a tenfold change in [H₃O⁺]. The logarithmic nature is essential when comparing acid strengths or working out dilutions.
务必记住，pH 每变化一个单位，[H₃O⁺] 就改变十倍。在进行酸强度比较或稀释计算时，这种对数关系至关重要。
10. Acid Dissociation Constant, Ka and pKa | 酸解离常数 Ka 和 pKa
For a weak acid HA, the equilibrium established in water is:
对于弱酸 HA，在水溶液中建立的平衡为：
HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)
The acid dissociation constant, Ka, is given by:
酸解离常数 Ka 的表达式为：
Ka = [H₃O⁺][A⁻] / [HA]
Ka has units of mol dm⁻³. Its magnitude indicates the strength of the acid: the larger the Ka, the more the equilibrium lies to the right, and the stronger the acid. For convenience, pKa is used: pKa = –log₁₀Ka. A smaller pKa value means a stronger acid.
Ka 的单位是 mol dm⁻³。其数值大小反映了酸的强弱：Ka 越大，平衡越偏向右侧，酸性越强。为方便起见常使用 pKa：pKa = –log₁₀Ka。pKa 值越小，酸性越强。
For ethanoic acid, Ka ≈ 1.8 × 10⁻⁵ mol dm⁻³, so pKa ≈ 4.74. When calculating pH of a weak acid solution, we often assume that [H₃O⁺] = [A⁻] and that the amount of acid dissociated is negligible compared to the initial concentration, leading to the approximate formula [H₃O⁺] ≈ √(Ka × c). These approximations are valid only when the acid is weak and the solution is not extremely dilute.
乙酸的 Ka ≈ 1.8 × 10⁻⁵ mol dm⁻³，故 pKa ≈ 4.74。计算弱酸溶液的 pH 时，通常假设 [H₃O⁺] = [A⁻] 且酸已电离部分与原浓度相比可忽略不计，从而得到近似公式 [H₃O⁺] ≈ √(Ka × c)。这些近似仅在酸很弱且溶液不太稀的情况下成立。
11. Dilution and pH Changes | 稀释与 pH 变化
Diluting a strong acid like HCl reduces [H₃O⁺] predictably: a tenfold dilution raises the pH by one unit. For example, diluting 0.1 mol dm⁻³ HCl (pH 1) to 0.01 mol dm⁻³ gives pH 2. However, further dilution towards extremely low concentrations must consider the contribution of water’s autoionisation, limiting the pH to just below 7.
稀释强酸（如 HCl）时，[H₃O⁺] 会按比例降低：稀释十倍，pH 升高一个单位。例如，将 0.1 mol dm⁻³ HCl (pH 1) 稀释至 0.01 mol dm⁻³，pH 变为 2。然而，当稀释到极低浓度时，必须考虑水的自电离贡献，pH 将趋近但低于 7。
Weak acids behave differently upon dilution. As the solution is diluted, the weak acid equilibrium shifts to the right (Le Chatelier’s principle), increasing the degree of ionisation. Therefore, a tenfold dilution of a weak acid does not raise the pH by a full unit; the change is smaller because more HA dissociates to partially counteract the dilution effect. This concept is a common AS exam question.
弱酸在稀释时表现不同。随着溶液变稀，弱酸电离平衡向右移动（勒夏特列原理），电离度增大。因此，弱酸稀释十倍后 pH 上升不足一个单位；因为更多 HA 分子电离，部分抵消了稀释带来的浓度降低。这是 AS 考试中的常见考点。
12. Summary and Comparison of Theories | 理论总结与比较
The three major acid-base theories provide a progressive framework. Arrhenius is the simplest but limited to aqueous OH⁻ and H⁺. Brønsted-Lowry extends the concept to any proton-transfer system and introduces conjugate pairs. Lewis further generalises to electron pair sharing, covering metal-ligand reactions and organic mechanisms. For most aqueous acid-base problems at AS, Brønsted-Lowry theory is the most useful, but recognising the wider Lewis view reinforces a deeper understanding of reaction chemistry.
三大酸碱理论构成了一个层层递进的体系。阿伦尼乌斯理论最简单，但局限于水溶液中的 OH⁻ 和 H⁺。布朗斯特-劳里理论将概念扩展到任何质子传递体系，并引入了共轭酸碱对。路易斯理论则进一步推广至电子对共享，涵盖了金属-配位体反应和有机机理。在 AS 阶段的大多数水溶液酸碱问题中，布朗斯特-劳里理论最为实用，但理解更宽广的路易斯视角将加深你对化学反应本质的认识。
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