IB Chemistry: Core Concepts of Acid-Base Theories | IB化学：酸碱理论核心知识梳理

📚 IB Chemistry: Core Concepts of Acid-Base Theories | IB化学：酸碱理论核心知识梳理

Acid-base chemistry is one of the most central topics in IB Chemistry, bridging stoichiometry, equilibrium, and organic reaction mechanisms. This article provides a systematic review of the core acid-base theories, from the classic Arrhenius model to the more general Brønsted-Lowry and Lewis frameworks, and highlights the key definitions, conjugate pairs, amphiprotic species, and acid strength concepts you need for your exams.

酸碱化学是IB化学中最核心的主题之一，它将化学计量、化学平衡和有机反应机理紧密联系起来。本文系统梳理了酸碱理论的核心内容，从经典的阿伦尼乌斯模型到更广泛的布朗斯特-劳里和路易斯框架，并重点讲解考试中必需的的定义、共轭对、两性物种和酸强度概念。

1. Arrhenius Theory | 阿伦尼乌斯理论

The Arrhenius theory defines acids as substances that produce H⁺ (or H₃O⁺) ions in aqueous solution, and bases as substances that produce OH⁻ ions in aqueous solution. For example, HCl dissociates to give H⁺ and Cl⁻, while NaOH dissociates to give Na⁺ and OH⁻.

阿伦尼乌斯理论将酸定义为在水中产生H⁺（或H₃O⁺）的物质，将碱定义为在水中产生OH⁻的物质。例如，HCl解离产生H⁺和Cl⁻，而NaOH解离产生Na⁺和OH⁻。

HA → H⁺ + A⁻ BOH → B⁺ + OH⁻

This theory is limited because it only applies to aqueous solutions and cannot explain the basicity of substances like NH₃, which does not contain OH⁻ ions but still acts as a base in water.

该理论仅适用于水溶液，无法解释像NH₃这类不含OH⁻但仍在水中显碱性的物质，因此具有很大的局限性。

2. Brønsted-Lowry Theory | 布朗斯特-劳里理论

The Brønsted-Lowry theory expands the definition: an acid is a proton (H⁺) donor, and a base is a proton acceptor. This theory works in any solvent and explains the behavior of NH₃ in water: NH₃ accepts a proton from H₂O to form NH₄⁺ and OH⁻.

布朗斯特-劳里理论扩展了定义：酸是质子的给予体，碱是质子的接受体。该理论适用于任何溶剂，并解释了NH₃在水中的行为：NH₃从H₂O接受一个质子生成NH₄⁺和OH⁻。

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

In this reaction, H₂O acts as an acid because it donates a proton, while NH₃ acts as a base because it accepts a proton. The reverse reaction is also acid-base: NH₄⁺ donates a proton and OH⁻ accepts it.

在该反应中，H₂O作为酸给出质子，NH₃作为碱接受质子。逆反应同样是酸碱反应：NH₄⁺给出质子，OH⁻接受质子。

3. Conjugate Acid-Base Pairs | 共轭酸碱对

A conjugate acid-base pair consists of two species that differ by one proton (H⁺). The acid has one more H⁺ than its conjugate base; the base has one fewer H⁺ than its conjugate acid.

共轭酸碱对由相差一个质子（H⁺）的两个物种构成。酸比其共轭碱多一个H⁺；碱比其共轭酸少一个H⁺。

HA ⇌ H⁺ + A⁻ (HA / A⁻ is a conjugate pair)

Common examples include: HCl/Cl⁻, NH₄⁺/NH₃, H₂O/OH⁻, and H₃O⁺/H₂O. In every acid-base equilibrium, there are two conjugate pairs.

常见例子有：HCl/Cl⁻、NH₄⁺/NH₃、H₂O/OH⁻和H₃O⁺/H₂O。在每个酸碱平衡中都存在两个共轭对。

For the equilibrium CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺, the pairs are CH₃COOH/CH₃COO⁻ and H₃O⁺/H₂O.

对于平衡CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺，其共轭对分别是CH₃COOH/CH₃COO⁻和H₃O⁺/H₂O。

4. Amphiprotic Species | 两性物种

An amphiprotic species can act either as a proton donor (acid) or as a proton acceptor (base), depending on the reaction conditions. Water is the most important amphiprotic substance: it can accept a proton to form H₃O⁺, or donate a proton to form OH⁻.

两性物种既能作为质子给予体（酸），也能作为质子接受体（碱），具体取决于反应条件。水是最重要的两性物质：它既可以接受质子形成H₃O⁺，也可以给出质子形成OH⁻。

Other common amphiprotic species include HCO₃⁻, HSO₄⁻, H₂PO₄⁻, and HPO₄²⁻. For example, HCO₃⁻ reacts with HCl as a base: HCO₃⁻ + H⁺ → H₂CO₃, but reacts with NaOH as an acid: HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O.

其他常见的两性物种包括HCO₃⁻、HSO₄⁻、H₂PO₄⁻和HPO₄²⁻。例如，HCO₃⁻与HCl反应时作为碱：HCO₃⁻ + H⁺ → H₂CO₃；但与NaOH反应时作为酸：HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O。

5. Lewis Acid-Base Theory | 路易斯酸碱理论

The Lewis theory is the most general acid-base model. A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. This theory covers reactions that do not involve protons, such as the formation of complex ions.

路易斯理论是最广义的酸碱模型。路易斯酸是电子对的接受体，路易斯碱是电子对的给予体。该理论涵盖了不涉及质子的反应，例如配合离子的形成。

BF₃ + NH₃ → BF₃-NH₃

In this reaction, BF₃ accepts a lone pair from NH₃, so BF₃ is a Lewis acid and NH₃ is a Lewis base. Similarly, H⁺ is a Lewis acid (it accepts an electron pair), and OH⁻, CN⁻, and H₂O are Lewis bases.

在该反应中，BF₃接受来自NH₃的孤对电子，因此BF₃是路易斯酸，NH₃是路易斯碱。类似地，H⁺是路易斯酸（它接受电子对），而OH⁻、CN⁻和H₂O都是路易斯碱。

6. Acid and Base Strength | 酸和碱的强度

Strong acids and bases dissociate completely in aqueous solution. Common strong acids include HCl, HBr, HI, HNO₃, H₂SO₄, and HClO₄. Common strong bases include group 1 and group 2 hydroxides such as NaOH, KOH, and Ca(OH)₂.

强酸和强碱在水中完全解离。常见强酸包括HCl、HBr、HI、HNO₃、H₂SO₄和HClO₄。常见强碱包括第1族和第2族金属氢氧化物，如NaOH、KOH和Ca(OH)₂。

Weak acids and bases only partially dissociate in water. Examples include CH₃COOH, H₂CO₃, and NH₃. The strength of a weak acid is quantified by the acid dissociation constant Kₐ, while the strength of a weak base is quantified by K_b.

弱酸和弱碱在水中仅部分解离，例如CH₃COOH、H₂CO₃和NH₃。弱酸的强度用电离常数Kₐ定量表示，弱碱的强度用K_b定量表示。

CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺ Kₐ = [CH₃COO⁻][H₃O⁺]/[CH₃COOH]

A larger Kₐ indicates a stronger acid. Conversely, a larger K_b indicates a stronger base. For a conjugate pair, Kₐ × K_b = K_w = 1.0 × 10⁻¹⁴ at 25 °C.

Kₐ越大表示酸越强；同理，K_b越大表示碱越强。对于共轭酸碱对，Kₐ × K_b = K_w = 1.0 × 10⁻¹⁴（25 °C时）。

7. The pH Scale and pOH | pH标度和pOH

The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration: pH = -log₁₀[H⁺]. Similarly, pOH = -log₁₀[OH⁻]. At 25 °C, pH + pOH = 14.

溶液的pH定义为氢离子浓度的负对数：pH = -log₁₀[H⁺]。类似地，pOH = -log₁₀[OH⁻]。在25 °C时，pH + pOH = 14。

Solution type	Condition at 25 °C
Acidic	[H⁺] > [OH⁻], pH < 7
Neutral	[H⁺] = [OH⁻], pH = 7
Basic	[H⁺] < [OH⁻], pH > 7

Remember that pH is a logarithmic scale: a decrease in pH by 1 unit corresponds to a ten-fold increase in [H⁺]. This means pH 3 is ten times more acidic than pH 4, and pH 2 is one hundred times more acidic than pH 4.

记住pH是对数标度：pH每降低1个单位，[H⁺]增大10倍。这意味着pH 3的酸性是pH 4的10倍，pH 2的酸性是pH 4的100倍。

8. Strong vs. Weak Acid Calculations | 强酸与弱酸的计算

For a strong acid such as HCl with concentration c, [H⁺] = c, so pH = -log₁₀ c. For a strong base such as NaOH, [OH⁻] = c, so pOH = -log₁₀ c and pH = 14 – pOH.

对于浓度为c的强酸（如HCl），[H⁺] = c，因此pH = -log₁₀ c。对于强碱（如NaOH），[OH⁻] = c，因此pOH = -log₁₀ c，pH = 14 – pOH。

For a weak acid HA, the equilibrium must be considered. If the initial concentration is c and the degree of dissociation is x, then:

对于弱酸HA，需要考虑平衡。若初始浓度为c，解离度为x，则：

Kₐ = x²/(c – x) ≈ x²/c (if x is much smaller than c)

Then [H⁺] = x, and pH = -log₁₀ x. Always check the approximation: if x/c is less than 5%, the approximation is usually considered valid in IB.

此时[H⁺] = x，pH = -log₁₀ x。始终要检验近似条件：在IB课程中，通常认为x/c小于5%时近似成立。

9. The Role of Water in Acid-Base Equilibria | 水在酸碱平衡中的作用

Water undergoes autoionization, producing H₃O⁺ and OH⁻. The ion product constant of water is K_w = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.

水发生自偶电离，产生H₃O⁺和OH⁻。水的离子积常数K_w = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴（25 °C时）。

2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

Because the autoionization of water is endothermic, K_w increases with temperature. At temperatures above 25 °C, pH 7 is no longer neutral; neutrality occurs when pH = pOH, which is less than 7.

由于水的自偶电离是吸热过程，K_w随温度升高而增大。在高于25 °C时，pH 7不再代表中性；中性条件是pH = pOH，此时pH小于7。

10. pH Curves and Indicators | pH曲线与指示剂

A pH titration curve shows how the pH of a solution changes as a titrant is added. There are four main types: strong acid-strong base, strong acid-weak base, weak acid-strong base, and weak acid-weak base.

pH滴定曲线表示随着滴定剂的加入，溶液pH的变化。主要有四种类型：强酸-强碱、强酸-弱碱、弱酸-强碱和弱酸-弱碱。

Strong acid + strong base: equivalence point pH = 7; steep vertical section around the equivalence point.

强酸+强碱：等当点pH = 7；等当点附近有陡峭的垂直段。
Strong acid + weak base: equivalence point pH < 7 (because the conjugate acid of the weak base hydrolyzes).

强酸+弱碱：等当点pH < 7（因为弱碱的共轭酸发生水解）。
Weak acid + strong base: equivalence point pH > 7 (because the conjugate base of the weak acid hydrolyzes).

弱酸+强碱：等当点pH > 7（因为弱酸的共轭碱发生水解）。

An acid-base indicator is a weak acid (HIn) whose colour differs from its conjugate base (In⁻). The end point of a titration is chosen so that the indicator’s colour change range overlaps with the steep part of the pH curve.

酸碱指示剂是一种弱酸（HIn），其颜色与其共轭碱（In⁻）不同。选择指示剂时，应使其变色范围与pH曲线的陡峭部分重叠。

11. Buffer Solutions | 缓冲溶液

A buffer solution resists changes in pH when small amounts of acid or base are added. It typically contains a weak acid and its conjugate base (e.g., CH₃COOH / CH₃COO⁻) or a weak base and its conjugate acid (e.g., NH₃ / NH₄⁺).

缓冲溶液在加入少量酸或碱时能抵抗pH的变化。缓冲溶液通常由弱酸及其共轭碱（如CH₃COOH / CH₃COO⁻）或弱碱及其共轭酸（如NH₃ / NH₄⁺）组成。

The Henderson-Hasselbalch equation is used to calculate the pH of a buffer:

亨德森-哈塞尔巴尔赫方程用于计算缓冲溶液的pH：

pH = pKₐ + log₁₀([A⁻]/[HA])

When [A⁻] = [HA], the buffer has maximum capacity and pH = pKₐ. Adding a small amount of H⁺ shifts the equilibrium to convert A⁻ to HA; adding OH⁻ converts HA to A⁻. As long as the reservoir of HA and A⁻ is not exhausted, pH changes are minimal.

当[A⁻] = [HA]时，缓冲容量最大，pH = pKₐ。加入少量H⁺会使平衡移动，将A⁻转化为HA；加入OH⁻则将HA转化为A⁻。只要HA和A⁻的储备未被耗尽，pH变化就很小。

12. Exam Tips and Common Pitfalls | 考试技巧与常见误区

In IB exams, students often confuse the terms “strong” and “concentrated.” Strength refers to the degree of dissociation, while concentration refers to the amount of solute per unit volume. A dilute strong acid can have a higher pH than a concentrated weak acid.

在IB考试中，学生常常混淆“强”和“浓”的概念。强度指解离程度，而浓度指单位体积内溶质的量。稀的强酸其pH可能大于浓的弱酸。

Always identify conjugate pairs by looking for species that differ by exactly one H⁺.

始终通过寻找相差恰好一个H⁺的物种来识别共轭对。
Remember that amphiprotic species appear in the middle of a polyprotic acid sequence, such as HCO₃⁻.

记住两性物种出现在多元酸序列的中间位置，例如HCO₃⁻。
Do not use Kₐ for strong acids; strong acids have Kₐ values too large to measure and are assumed to dissociate fully.

不要对强酸使用Kₐ；强酸的Kₐ值太大无法测量，通常假定完全解离。
When calculating pH of a weak base, first find [OH⁻], then pOH, then pH using pH = 14 – pOH (at 25 °C).

计算弱碱的pH时，先求[OH⁻]，再求pOH，最后用pH = 14 – pOH（25 °C）得到pH。

Understanding the relationships between Kₐ, K_b, K_w, pH, and pOH is essential for solving IB acid-base problems systematically. Always write the equilibrium first, then apply the appropriate approximation.

理解Kₐ、K_b、K_w、pH和pOH之间的关系，是系统解决IB酸碱问题的关键。务必先写出平衡方程式，再应用适当的近似方法。

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IB Chemistry: Core Concepts of Acid-Base Theories | IB化学：酸碱理论核心知识梳理