Acid and Base Definitions | 酸和碱的定义

📚 Acid and Base Definitions | 酸和碱的定义

Understanding exactly what qualifies a substance as an acid or a base is fundamental to predicting reactivity, interpreting equilibria and solving problems in IB chemistry. Over time, chemists have developed progressively broader models — from aqueous H⁺ and OH⁻ to proton transfer and finally to electron pair donation — each refining our ability to explain chemical behaviour. This article unpacks the three key definitions required by the IB syllabus: Arrhenius, Brønsted–Lowry and Lewis, and then links them to conjugate pairs, amphiprotic species, strength, the water equilibrium and the pH scale.

精确理解什么物质属于酸、什么物质属于碱，是预测反应活性、解释平衡和解决 IB 化学问题的基石。化学家们先后建立了逐步扩展的模型——从水溶液中的 H⁺ 与 OH⁻，到质子转移，再到电子对给予——每一次都细化了我们对化学行为的解释能力。本文将梳理 IB 大纲所要求的三种核心定义：阿伦尼乌斯、布朗斯特-劳里和路易斯，并进一步关联共轭酸碱对、两性物质、强弱性质、水的电离平衡及 pH 标度。

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1. Why Definitions Matter | 定义的重要性

Acids and bases appear in topics ranging from energetics to organic chemistry and biochemistry. A clear definition allows you to identify the role of a species in any reaction, to write correct equations and to link structure to acidity. IB expects you not only to state the three historical definitions but also to apply them in unfamiliar contexts, such as describing the behaviour of transition metal complexes or non-aqueous solvents.

从能量学到有机化学乃至生物化学，酸与碱几乎贯穿所有主题。清晰的定义能帮助你在任何反应中识别物种的角色，写出正确的方程式，并把结构与酸性关联起来。IB 不仅要求你陈述三种历史定义，还要求你在陌生情境中应用它们，比如描述过渡金属配合物或非水溶剂的行为。

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2. The Arrhenius Definition | 阿伦尼乌斯定义

Svante Arrhenius proposed that an acid is a substance that dissociates in water to produce hydrogen ions, H⁺, while a base dissociates to produce hydroxide ions, OH⁻. For example:

斯万特·阿伦尼乌斯提出，酸是在水中离解产生氢离子 H⁺ 的物质，而碱是在水中离解产生氢氧根离子 OH⁻ 的物质。例如：

HCl → H⁺ + Cl⁻ NaOH → Na⁺ + OH⁻

This model correctly explains many common laboratory acids and bases. However, it is limited to aqueous solutions and cannot explain why ammonia (NH₃) behaves as a base even though it contains no OH⁻ group, nor can it describe gaseous acid–base reactions. The Arrhenius definition therefore serves as a starting point but has been largely superseded.

该模型正确地解释了许多常见的实验室酸碱。但它仅限于水溶液体系，既无法说明不含 OH⁻ 的氨（NH₃）为何显碱性，也无法描述气相中的酸碱反应。因此，阿伦尼乌斯定义只是一个起点，如今已基本被更广义的理论取代。

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3. The Brønsted–Lowry Definition | 布朗斯特-劳里定义

In 1923, Johannes Brønsted and Thomas Lowry independently defined an acid as a proton (H⁺) donor and a base as a proton acceptor. This theory removes the need for OH⁻ and extends acid–base chemistry to any solvent or even the gas phase. A classic Brønsted–Lowry acid–base reaction is the ionisation of hydrogen chloride in water:

1923 年，约翰内斯·布朗斯特和托马斯·劳里各自独立地定义了酸为质子（H⁺）的给予体，碱为质子的接受体。这一理论免去了对 OH⁻ 的依赖，并将酸碱化学推广到任何溶剂乃至气相。一个经典的布朗斯特-劳里酸碱反应是氯化氢在水中的电离：

HCl + H₂O → H₃O⁺ + Cl⁻

Here, HCl donates a proton to H₂O, so HCl is the Brønsted–Lowry acid and H₂O is the base. Note that water can act as a base in this reaction, but as you will see later, it can also act as an acid — a property known as amphiprotic behaviour. The Brønsted–Lowry definition is the main working model used throughout the IB diploma programme.

在此，HCl 将一个质子给予 H₂O，因此 HCl 是布朗斯特-劳里酸，H₂O 是碱。注意水在这个反应中作为碱，但后面会看到，水同样可以充当酸——这被称为两性行为。布朗斯特-劳里定义是整个 IB 文凭课程中使用的主要工作模型。

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4. Conjugate Acid–Base Pairs | 共轭酸碱对

Every Brønsted–Lowry acid has a conjugate base formed after it donates a proton, and every base has a conjugate acid formed after it accepts a proton. A conjugate pair consists of two species that differ by just one proton. For example, in the reaction above, HCl and Cl⁻ form one conjugate acid–base pair, while H₃O⁺ and H₂O form the other. The table below lists some common conjugate pairs:

每个布朗斯特-劳里酸在给出质子后都形成一个共轭碱，每个碱在得到质子后都形成一个共轭酸。共轭对由两个只相差一个质子的物种组成。例如，在上述反应中，HCl 与 Cl⁻ 组成一个共轭酸碱对，而 H₃O⁺ 与 H₂O 组成另一对。下表列出了一些常见共轭对：

Acid / 酸	Conjugate Base / 共轭碱
HCl	Cl⁻
H₂SO₄	HSO₄⁻
CH₃COOH	CH₃COO⁻
NH₄⁺	NH₃
H₂O	OH⁻
H₃O⁺	H₂O

Recognising conjugate pairs is essential for understanding buffers, acid strength and equilibrium. A key IB point: the stronger the acid, the weaker its conjugate base, and vice versa.

识别共轭对是理解缓冲溶液、酸性强度和平衡的基础。IB 强调的一个关键点是：酸越强，其共轭碱越弱，反之亦然。

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5. Amphiprotic and Amphoteric Species | 两性质子与两性物质

A species that can either donate or accept a proton is said to be amphiprotic. Water is the most familiar example: with HCl it accepts a proton (acting as a base); with NH₃ it donates a proton (acting as an acid). Hydrogencarbonate ion, HCO₃⁻, and hydrogensulfate ion, HSO₄⁻, also display amphiprotic behaviour.

既能给出质子又能接受质子的物种被称为两性质子物质。水是最熟悉的例子：与 HCl 反应时它接受质子（充当碱）；与 NH₃ 反应时它给出质子（充当酸）。碳酸氢根离子 HCO₃⁻ 和硫酸氢根离子 HSO₄⁻ 也表现出两性行为。

The term amphoteric is more general; it describes any substance that can react with both acids and bases, even by mechanisms other than proton transfer. For instance, aluminium oxide, Al₂O₃, is amphoteric because it dissolves in both acidic and alkaline solutions, forming Al³⁺(aq) or [Al(OH)₄]⁻(aq). Under the Brønsted–Lowry definition, Al₂O₃ is not amphiprotic because it does not donate or accept a simple proton; it is instead explained by the Lewis theory.

词语两性（amphoteric）更为广义；它描述任何既能与酸又能与碱反应的物质，哪怕不是通过质子转移机理。例如，氧化铝 Al₂O₃ 具有两性，因为它既能溶于酸也能溶于碱，分别生成 Al³⁺(aq) 或 [Al(OH)₄]⁻(aq)。在布朗斯特-劳里定义中，Al₂O₃ 并非两性质子物质，因为它不涉及简单质子的给予或接受；这需要通过路易斯理论来解释。

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6. The Lewis Definition | 路易斯酸碱定义

Gilbert N. Lewis further generalised acid–base behaviour by shifting the focus from protons to electron pairs. A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor. This definition encompasses all Brønsted–Lowry acids and bases but also includes many reactions that do not involve protons at all.

吉尔伯特·N·路易斯将焦点从质子转移到电子对，进一步推广了酸碱行为。路易斯酸是电子对接受体；路易斯碱是电子对给予体。这一定义涵盖了所有布朗斯特-劳里酸和碱，还包含了大量根本不涉及质子的反应。

A typical Lewis acid–base reaction is the formation of an adduct between boron trifluoride and ammonia:

一个典型的路易斯酸碱反应是三氟化硼与氨形成加合物：

BF₃ + :NH₃ → F₃B←NH₃

BF₃ has an incomplete octet on boron and can accept a lone pair, making it a Lewis acid. NH₃ provides that lone pair, acting as a Lewis base. Metal cations, such as Fe³⁺ or Cu²⁺, are also Lewis acids when they accept lone pairs from ligands like water or ammonia. The Lewis theory is particularly important when studying transition metal complex formation and organic reaction mechanisms.

BF₃ 中的硼原子八隅体不全，可以接受孤对电子，因而是路易斯酸；NH₃ 提供孤对电子，作为路易斯碱。金属阳离子如 Fe³⁺ 或 Cu²⁺ 在从水或氨等配体接受孤对电子时，同样起着路易斯酸的作用。在研究过渡金属配合物的形成和有机反应机理时，路易斯理论尤为重要。

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7. Strong and Weak Acids and Bases | 强酸、弱酸与强碱、弱碱

Strength refers to the extent of ionisation or proton transfer, not concentration. A strong acid is one that completely dissociates in aqueous solution: the equilibrium lies fully to the right. Common strong acids include HCl, HNO₃ and H₂SO₄ (first proton). A weak acid, such as CH₃COOH (ethanoic acid), only partially dissociates, establishing an equilibrium that lies largely to the left.

强度指电离或质子转移的程度，而非浓度。强酸是在水溶液中完全解离的酸：平衡完全移向右侧。常见的强酸包括 HCl、HNO₃ 和 H₂SO₄（第一步质子）。弱酸如 CH₃COOH（乙酸）只能部分解离，平衡主要偏向左侧。

Similarly, strong bases like NaOH, KOH and Ba(OH)₂ dissociate fully to give OH⁻ ions, while weak bases such as NH₃ and CH₃NH₂ accept protons only to a limited extent. The IB also expects you to relate strength to the acid dissociation constant, Kₐ, and the base dissociation constant, K_b Kₐ value indicates a strong acid; a small Kₐ indicates a weak one. The relationship Kₐ × K_b = K_w for a conjugate pair is a powerful tool for linking acid and base strength.

同理，强碱如 NaOH、KOH 和 Ba(OH)₂ 完全解离产生 OH⁻ 离子，而弱碱如 NH₃ 和 CH₃NH₂ 只能有限地接受质子。IB 还要求你将强度与酸解离常数 Kₐ 和碱解离常数 K_b 联系起来。Kₐ 值大表示强酸，Kₐ 值小表示弱酸。对于共轭对，有 Kₐ × K_b = K_w，这一关系是联系酸碱强度的重要工具。

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8. The Ionic Product of Water, Kw | 水的离子积 Kw

Pure water undergoes autoprotolysis, a reaction in which one water molecule donates a proton to another, producing equal amounts of hydronium and hydroxide ions:

纯水会发生自质子转移反应，一个水分子向另一个水分子提供质子，生成等量的水合氢离子和氢氧根离子：

2H₂O ⇌ H₃O⁺ + OH⁻

The equilibrium constant for this process is the ionic product of water, K_w. At 298 K, K_w e the concentration of water in dilute aqueous solution is essentially constant, we write:

这个过程的平衡常数即为水的离子积 K_w。在 298 K 时，K_w = 1.00 × 10⁻¹⁴ (mol dm⁻³)²。由于稀水溶液中水的浓度基本恒定，我们记作：

K_w = [H₃O⁺][OH⁻]

In pure water at this temperature, [H₃O⁺] = [OH⁻] = 1.00 × 10⁻⁷ mol dm⁻³. K_w gher temperature may have [H₃O⁺] > 1.0 × 10⁻⁷ mol dm⁻³, yet still maintain [H₃O⁺] = [OH⁻].

在该温度下的纯水中，[H₃O⁺] = [OH⁻] = 1.00 × 10⁻⁷ mol dm⁻³。K_w 依赖于温度；由于自质子转移反应是吸热的，K_w 随温度升高而增大。因此，在较高温度下具有中性的溶液，其 [H₃O⁺] 可能大于 1.0 × 10⁻⁷ mol dm⁻³，但仍保持 [H₃O⁺] = [OH⁻]。

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9. The pH and pOH Scale | pH 与 pOH 标度

To express the often minute hydronium ion concentration more conveniently, Sorensen introduced the pH scale:

为了方便地表示常极低的水合氢离子浓度，索伦森引入了 pH 标度：

pH = −log₁₀[H₃O⁺]

Analogously, pOH is defined as:

类似地，pOH 定义为：

pOH = −log₁₀[OH⁻]

Taking the negative logarithm of the K_w expression gives the extremely useful relationship valid at 298 K:

对 K_w 表达式取负对数，可得在 298 K 时极其有用的关系式：

pH + pOH = 14.00

0; acidic solutions have pH < 7.00, and basic solutions have pH > 7.00. Because pH is a logarithmic scale, a change of one pH unit corresponds to a tenfold change in [H₃O⁺]. IB students must be comfortable interconverting [H₃O⁺], [OH⁻], pH and pOH for both strong and weak systems.

在 298 K 时，中性溶液的 pH = 7.00；酸性溶液 pH < 7.00，碱性溶液 pH > 7.00。由于 pH 是对数标度，每变化一个 pH 单位对应 [H₃O⁺] 的十倍变化。IB 学生需能熟练地在 [H₃O⁺]、[OH⁻]、pH 和 pOH 之间进行换算，无论针对强还是弱电解质体系。

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10. Calculating pH for Strong Acids and Bases | 强酸与强碱的 pH 计算

onsidered). For example, a 0.050 mol dm⁻³ solution of HCl has [H₃O⁺] = 0.050 mol dm⁻³, so:

对于完全解离的强一元酸，只要初始浓度 cₐ 不是极低（极低时需考虑水的自解离），水合氢离子浓度就等于初始酸浓度。例如，0.050 mol dm⁻³ 的 HCl 溶液中 [H₃O⁺] = 0.050 mol dm⁻³，于是：

pH = −log₁₀(0.050) = 1.30

a(OH)₂, the concentration must be multiplied by the stoichiometric factor: 0.010 mol dm⁻³ Ba(OH)₂ yields [OH⁻] = 0.020 mol dm⁻³.

对于强碱如 NaOH，氢氧根离子浓度等于碱的浓度。0.010 mol dm⁻³ 的 NaOH 溶液给出 [OH⁻] = 0.010 mol dm⁻³，因此 pOH = −log₁₀(0.010) = 2.00，进而在 298 K 下 pH = 14.00 − 2.00 = 12.00。若每个化学式单元能提供多个 OH⁻，如 Ba(OH)₂，则浓度需乘以化学计量因数：0.010 mol dm⁻³ 的 Ba(OH)₂ 给出 [OH⁻] = 0.020 mol dm⁻³。

temperature is 298 K before using pH + pOH = 14.00, because the relationship changes if K_w differs from 1.00 × 10⁻¹⁴.

这些直接的计算是滴定曲线和缓冲体系等更复杂主题的基础。在使用 pH + pOH = 14.00 之前，务请确认温度为 298 K，因为若 K_w 偏离 1.00 × 10⁻¹⁴，该关系式也会变化。

In summary, the historical progression from Arrhenius to Brønsted–Lowry to Lewis equips chemists with an increasingly versatile toolkit for describing acid–base behaviour. Mastering the definitions, recognising conjugate partners and quantifying strength through pH and K_w are all essential skills for success in IB chemistry.

总而言之，从阿伦尼乌斯到布朗斯特-劳里再到路易斯的历史演进，为化学家提供了日益通用的酸碱描述工具包。掌握定义、识别共轭伙伴并通过 pH 和 K_w 量化强度，都是 IB 化学取得成功的必备技能。

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