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

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

Acid-base chemistry is a central topic in IB Chemistry, bridging the microscopic behavior of molecules and ions with measurable macroscopic properties such as pH. This article provides a structured review of the key concepts, from classical definitions to quantitative equilibrium calculations.

酸碱化学是IB化学的核心内容，它将分子和离子的微观行为与pH等可观测的宏观性质联系起来。本文系统梳理了从经典定义到定量平衡计算的关键知识点，帮助你在考试中准确应用。

1. The Evolution of Acid-Base Definitions | 1. 酸碱定义的演变

In the 19th century, Arrhenius defined acids as substances that produce H⁺ ions in aqueous solution, and bases as substances that produce OH⁻ ions. This works well for simple aqueous systems but fails for reactions in non-aqueous solvents, such as liquid ammonia.

19世纪，阿伦尼乌斯将酸定义为在水中产生H⁺的物质，碱定义为产生OH⁻的物质。这适用于简单水溶液体系，但在非水溶剂（如液氨）中则失效。

The Brønsted-Lowry definition broadened the idea: an acid is a proton (H⁺) donor, and a base is a proton acceptor. This allows acids and bases to be identified in any solvent or even in the gas phase.

布朗斯特-劳里定义扩展了这一概念：酸是质子供体，碱是质子受体。这使得我们可以在任何溶剂甚至气相中识别酸碱。

Lewis theory goes further: an acid is an electron pair acceptor, and a base is an electron pair donor. This includes many compounds without hydrogen, such as BF₃ and transition metal ions.

路易斯理论更进一步：酸是电子对接受体，碱是电子对给予体。该定义涵盖了不含氢的化合物，如BF₃和过渡金属离子。

Theory	Acid	Base	Example
Arrhenius	Produces H⁺ in water	Produces OH⁻ in water	HCl, NaOH
Brønsted-Lowry	Proton donor	Proton acceptor	NH₄⁺, H₂O
Lewis	Electron pair acceptor	Electron pair donor	BF₃, Fe³⁺

2. Conjugate Acid-Base Pairs | 2. 共轭酸碱对

In the Brønsted-Lowry model, when an acid donates a proton, the remaining species is its conjugate base. Conversely, when a base accepts a proton, the resulting species is its conjugate acid. The pair differs by exactly one proton.

在布朗斯特-劳里模型中，当酸给出一个质子后，剩余物种就是它的共轭碱；反之，碱接受一个质子后，生成的物种就是它的共轭酸。该对物质恰好相差一个质子。

For example, the equilibrium between ethanoic acid and water can be written as:

例如，乙酸与水的平衡可以写成：

CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺

Here, CH₃COOH and CH₃COO⁻ are a conjugate acid-base pair, while H₃O⁺ and H₂O are another pair. Water acts as a base in this reaction, accepting a proton.

这里，CH₃COOH和CH₃COO⁻是一对共轭酸碱，而H₃O⁺和H₂O是另一对。水在该反应中作为碱，接受质子。

It is crucial to identify conjugate pairs correctly in equilibrium equations, as they are central to interpreting acid-base strength relationships.

在平衡方程中正确识别共轭对至关重要，因为它们是理解酸碱强度关系的核心。

3. Autoionization of Water and the pH Scale | 3. 水的自电离与pH标度

Pure water self-ionizes to a very small extent:

纯水会发生极微弱的自电离：

2H₂O ⇌ H₃O⁺ + OH⁻

The equilibrium constant for water, Kw, at 25 °C is 1.0 × 10⁻¹⁴. Since the reaction produces equal concentrations of H₃O⁺ and OH⁻, each equals 1.0 × 10⁻⁷ mol dm⁻³.

水的离子积Kw在25 °C时为1.0 × 10⁻¹⁴。由于反应生成等量的H₃O⁺和OH⁻，两者浓度均为1.0 × 10⁻⁷ mol dm⁻³。

The pH scale is defined as the negative base-10 logarithm of the hydrogen ion concentration:

pH标度定义为氢离子浓度的负以10为底的对数：

pH = -log₁₀[H⁺]

Similarly, pOH = -log₁₀[OH⁻], and at 25 °C, pH + pOH = 14. A neutral solution has pH = 7; acidic solutions have pH < 7; basic solutions have pH > 7.

类似地，pOH = -log₁₀[OH⁻]，且在25 °C下，pH + pOH = 14。中性溶液pH = 7；酸性溶液pH < 7；碱性溶液pH > 7。

Remember that a change of one pH unit corresponds to a tenfold change in [H⁺]. This logarithmic relationship explains why a small pH difference can represent a large concentration change.

注意，pH每变化1个单位，[H⁺]变化10倍。这种对数关系解释了为什么很小的pH差异可能对应很大的浓度变化。

4. Strong vs Weak Acids and Bases | 4. 强酸强碱与弱酸弱碱

Strong acids and bases are completely dissociated in aqueous solution, meaning that the concentration of H⁺ or OH⁻ equals the initial concentration of the acid or base. Examples include HCl, H₂SO₄, HNO₃, NaOH, and KOH.

强酸和强碱在水中完全电离，因此H⁺或OH⁻的浓度等于酸或碱的初始浓度。例如HCl、H₂SO₄、HNO₃、NaOH和KOH。

Weak acids and bases only partially dissociate. For a weak acid HA, the equilibrium lies far to the left:

弱酸和弱碱仅部分电离。对于弱酸HA，平衡强烈向左移动：

HA ⇌ H⁺ + A⁻

Common weak acids include ethanoic acid and carbonic acid; common weak bases include ammonia and methylamine. The position of equilibrium is quantified by the acid dissociation constant, Ka.

常见弱酸包括乙酸和碳酸；常见弱碱包括氨和甲胺。平衡位置用酸解离常数Ka来定量描述。

Do not confuse “strong” with “concentrated”. A strong acid in dilute solution has a high degree of ionisation but a low concentration, while a weak acid in concentrated solution has a low degree of ionisation but a high initial concentration.

不要混淆“强”与“浓”。稀溶液中的强酸电离程度大但浓度可能低；而浓溶液中的弱酸电离程度小但初始浓度可能高。

5. Ka, Kb and the Relationship with Kw | 5. Ka、Kb及其与Kw的关系

For a weak acid, the equilibrium constant expression is:

对于弱酸，其平衡常数表达式为：

Ka = [H⁺][A⁻] / [HA]

For a weak base, B, the hydrolysis reaction and Kb expression are:

对于弱碱B，水解反应和Kb表达式为：

B + H₂O ⇌ BH⁺ + OH⁻, Kb = [BH⁺][OH⁻] / [B]

For a conjugate acid-base pair, KA × KB = KW. This relationship allows us to calculate the Ka of an acid if the Kb of its conjugate base is known, and vice versa.

对于一对共轭酸碱，Ka × Kb = Kw。利用这一关系，如果已知共轭碱的Kb，就能计算对应酸的Ka，反之亦然。

It is often convenient to use pKa and pKb values, defined as pKa = -log₁₀Ka and pKb = -log₁₀Kb. Lower pKa means a stronger acid. At 25 °C, pKa + pKb = 14.

使用pKa和pKb值往往更方便，定义为pKa = -log₁₀Ka，pKb = -log₁₀Kb。pKa越小，酸性越强。在25 °C下，pKa + pKb = 14。

6. pH Calculations Using ICE Tables | 6. 使用ICE表进行pH计算

For strong acids and bases, pH calculations are straightforward because dissociation is complete. For example, 0.01 mol dm⁻³ HCl gives [H⁺] = 0.01 mol dm⁻³, so pH = 2.00.

对于强酸和强碱，由于完全电离，pH计算非常简单。例如，0.01 mol dm⁻³ HCl的[H⁺] = 0.01 mol dm⁻³，因此pH = 2.00。

For weak acids, an ICE (Initial, Change, Equilibrium) table is essential to account for the partial dissociation. Consider a weak acid with concentration c and dissociation constant Ka:

对于弱酸，需要使用ICE（初始、变化、平衡）表格来处理部分电离。设弱酸浓度为c，解离常数为Ka：

	HA	H⁺	A⁻
Initial	c	0	0
Change	-x	+x	+x
Equilibrium	c – x	x	x

Assuming x is negligible compared to c, the expression simplifies to Ka ≈ x²/c, giving x = √(Ka·c). This x equals [H⁺], and pH = -log₁₀x.

假设x相对于c可以忽略，则表达式简化为Ka ≈ x²/c，得到x = √(Ka·c)。这里的x就是[H⁺]，因此pH = -log₁₀x。

Always check the “5% rule”: if x/c > 5%, the approximation may be invalid and the quadratic equation must be solved exactly.

务必检查“5%规则”：如果x/c > 5%，近似可能不成立，需要精确求解二次方程。

7. Acid-Base Indicators | 7. 酸碱指示剂

Indicators are weak acids that have different colours in their protonated and deprotonated forms. In a solution, the indicator HIn establishes the equilibrium:

指示剂是弱酸，其质子化和去质子化形式呈现不同颜色。在溶液中，指示剂HIn建立如下平衡：

HIn ⇌ H⁺ + In⁻

When pH is low, the equilibrium lies to the left and the colour of HIn predominates. When pH is high, the equilibrium shifts right and the colour of In⁻ appears. The endpoint occurs near pKa of the indicator.

当pH较低时，平衡向左移动，HIn的颜色占主导。当pH较高时，平衡右移，In⁻的颜色显现。终点出现在指示剂pKa附近。

For a titration, the indicator must have a pH range that overlaps the steep vertical section of the titration curve. For example, phenolphthalein (pKa ≈ 9) is suitable for strong-acid-strong-base and weak-acid-strong-base titrations, whereas methyl orange (pKa ≈ 3.7) is suitable for strong-acid-strong-base and strong-acid-weak-base titrations.

对于滴定，指示剂的变色范围必须与滴定曲线的垂直陡峭段重叠。例如，酚酞（pKa ≈ 9）适用于强酸强碱和弱酸强碱滴定，而甲基橙（pKa ≈ 3.7）适用于强酸强碱和强酸弱碱滴定。

8. Buffer Solutions | 8. 缓冲溶液

A buffer solution resists changes in pH when small amounts of acid or base are added. It is composed of a weak acid and its conjugate base (or a weak base and its conjugate acid) in significant concentrations.

缓冲溶液能够在加入少量酸或碱时抵抗pH变化。它由弱酸及其共轭碱（或弱碱及其共轭酸）以显著浓度组成。

The pH of a buffer made from a weak acid HA and its salt NaA can be calculated using the Henderson-Hasselbalch equation:

由弱酸HA及其盐NaA配制的缓冲溶液，可用亨德森-哈塞尔巴尔赫方程计算pH：

pH = pKa + log₁₀([A⁻] / [HA])

When [A⁻] = [HA], the buffer pH equals pKa. Buffers are most effective when the ratio [A⁻]/[HA] is between 0.1 and 10, corresponding to a pH range of pKa ± 1.

当[A⁻] = [HA]时，缓冲液pH等于pKa。当[A⁻]/[HA]在0.1到10之间时，缓冲效果最佳，对应pH范围为pKa ± 1。

The buffer capacity is the amount of acid or base that can be neutralised before the pH changes appreciably. Adding a small amount of H⁺ shifts the equilibrium to consume A⁻, while adding OH⁻ is neutralised by HA.

缓冲容量是指在pH发生明显变化之前所能中和的酸或碱的量。加入少量H⁺时，平衡移向消耗A⁻；加入OH⁻时，则由HA中和。

9. Titration Curves and Equivalence Points | 9. 滴定曲线与等当点

A titration curve plots pH against the volume of titrant added. The equivalence point is reached when the moles of added titrant are stoichiometrically equal to the moles of analyte in the sample.

滴定曲线表示pH随滴定剂体积变化的图像。等当点是指所加滴定剂的物质的量与样品中待测物质的量按化学计量完全反应时的点。

There are four main types of acid-base titration, each with a characteristic curve shape and pH at the equivalence point:

酸碱滴定主要有四种类型，每种都有其特征曲线形状和等当点pH：

Strong acid + strong base: equivalence pH = 7, steep pH change.
强酸+强碱：等当点pH = 7，pH突变剧烈。
Weak acid + strong base: equivalence pH > 7, buffering region before equivalence.
弱酸+强碱：等当点pH > 7，等当点前有缓冲区域。
Strong acid + weak base: equivalence pH < 7, acidic salt formed.
强酸+弱碱：等当点pH < 7，生成酸性盐。
Weak acid + weak base: equivalence pH depends on Ka and Kb, no clear steep section.
弱酸+弱碱：等当点pH取决于Ka和Kb，无明显突变段。

The midpoint of the vertical section of the curve is for the buffer region where pH = pKa of the weak acid (or pKb of the weak base). Indicators should be selected so that their colour change occurs at the equivalence point.

曲线垂直段的中点对应缓冲区域，此时pH等于弱酸的pKa（或弱碱的pKb）。指示剂应选择变色范围与等当点重合的类型。

10. Salt Hydrolysis | 10. 盐的水解

When a salt dissolves in water, its ions may react with water to form acidic or basic solutions. This phenomenon is called hydrolysis. The resulting pH depends on the nature of the parent acid and base.

盐溶于水时，其离子可能与水反应生成酸性或碱性溶液，这种现象称为水解。最终pH取决于盐对应的酸和碱的性质。

Salt of a strong acid and strong base (e.g., NaCl): neutral, pH = 7.
强酸和强碱生成的盐（如NaCl）：中性，pH = 7。
Salt of a weak acid and strong base (e.g., CH₃COONa): basic, pH > 7. The conjugate base reacts with water.
弱酸和强碱生成的盐（如CH₃COONa）：碱性，pH > 7。其共轭碱与水反应。
Salt of a strong acid and weak base (e.g., NH₄Cl): acidic, pH < 7. The conjugate acid donates protons.
强酸和弱碱生成的盐（如NH₄Cl）：酸性，pH < 7。其共轭酸给出质子。
Salt of a weak acid and weak base: pH depends on the relative Ka and Kb values.
弱酸和弱碱生成的盐：pH取决于Ka与Kb的相对大小。

To predict whether a salt solution is acidic, basic, or neutral, compare the Ka of the conjugate acid with the Kb of the conjugate base. The larger one determines the dominant effect.

要预测盐溶液是酸性、碱性还是中性，需比较共轭酸的Ka和共轭碱的Kb。数值较大者决定主导效应。

11. Common Exam Pitfalls | 11. 常见考试陷阱

Students often confuse strong acids with concentrated acids, or assume that pH = 7 always means neutral. Neutrality is temperature-dependent because Kw changes with temperature.

学生常把强酸与浓酸混淆，或认为pH = 7总是表示中性。实际上中性取决于温度，因为Kw会随温度变化。

When doing pH calculations, remember to account for the stoichiometry of polyprotic acids. For example, H₂SO₄ releases up to two protons, while H₃PO₄ releases three. For simplification, IB often treats only the first dissociation as significant for weak polyprotic acids.

进行pH计算时，需注意多元酸的化学计量。例如H₂SO₄最多释放两个质子，H₃PO₄释放三个。为简化，IB通常只考虑弱多元酸的第一步电离。

Another mistake is neglecting the “x is small” approximation check. If Ka is not very small or c is very dilute, the cubic or quadratic equation must be solved. Also, in buffer questions, always identify the volume changes when concentrations are given.

另一个错误是忽略“x很小”近似检验。如果Ka不是很小或c很稀，必须求解二次方程。此外，在缓冲溶液问题中，当给出浓度时，务必考虑体积变化。

12. Summary | 12. 总结

Acid-base chemistry in IB requires a clear understanding of the three theoretical definitions, the concept of conjugate pairs, and the quantitative tools involving Kw, Ka, Kb, pH, and pOH. Mastery of ICE tables and buffer calculations is essential for high marks.

IB酸碱化学要求清晰理解三种定义、共轭对概念，以及涉及Kw、Ka、Kb、pH和pOH的定量工具。掌握ICE表和缓冲溶液计算是获取高分的关键。

Remember that strong acids/bases fully ionise, while weak ones are in equilibrium. Indicators and titration curves rely on the pH changes near the equivalence point. Salt hydrolysis explains the acidic or basic nature of salt solutions.

记住，强酸强碱完全电离，弱酸弱碱处于平衡状态。指示剂和滴定曲线依赖等当点附近的pH变化。盐类水解解释了盐溶液的酸碱性。

By systematically reviewing these core concepts and practicing past-paper questions, you can confidently approach any acid-base problem on the IB exam.

通过系统复习这些核心知识点并练习历年真题，你就能自信应对IB考试中的任何酸碱问题。

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