📚 Acids and Bases: Key Concepts for IB & WJEC Science | 酸与碱：IB与WJEC科学考点精讲
Acids and bases are fundamental to chemistry, from the laboratory to living systems. This revision guide walks you through the essential theories, properties, and calculations you will face in IB and WJEC science exams. We break everything down into clear, bilingual explanations to boost your confidence.
酸与碱是化学的基础，从实验室到生命体系都离不开它们。本复习指南带你梳理IB与WJEC科学考试中必须掌握的理论、性质与计算。我们用清晰的双语讲解帮你夯实每一个考点，轻松拿分。
1. Defining Acids and Bases | 酸与碱的定义
An acid is a substance that releases hydrogen ions (H⁺) when dissolved in water, while a base is a substance that releases hydroxide ions (OH⁻). In the Brønsted–Lowry framework, an acid is a proton donor and a base is a proton acceptor.
酸是溶于水时释放出氢离子(H⁺)的物质，而碱是释放出氢氧根离子(OH⁻)的物质。在Brønsted–Lowry理论中，酸是质子给予体，碱是质子接受体。
The simplest test for an acid is its sour taste and its ability to turn blue litmus paper red. Bases feel slippery and turn red litmus paper blue. However, in the lab we never taste chemicals – we use indicators instead.
最简单的鉴别方法是：酸有酸味，能使蓝色石蕊试纸变红；碱有滑腻感，能使红色石蕊试纸变蓝。但实验室严禁品尝化学品，我们只使用指示剂。
2. Arrhenius vs Brønsted–Lowry Theories | 阿伦尼乌斯与Brønsted–Lowry理论
Arrhenius defined acids as substances that produce H⁺ in water and bases as those that produce OH⁻. This model explains many reactions but is limited to aqueous solutions. For example, HCl → H⁺ + Cl⁻ and NaOH → Na⁺ + OH⁻.
阿伦尼乌斯把酸定义为在水中产生H⁺的物质，碱定义为产生OH⁻的物质。这个模型能解释许多反应，但仅限于水溶液。例如，HCl → H⁺ + Cl⁻，NaOH → Na⁺ + OH⁻。
Brønsted–Lowry theory broadens the definition: an acid is a proton (H⁺) donor and a base is a proton acceptor. This allows us to describe reactions in non‑aqueous solvents and explains the role of species like ammonia (NH₃), which acts as a base by accepting a proton to become NH₄⁺.
Brønsted–Lowry理论拓展了定义：酸是质子(H⁺)给予体，碱是质子接受体。这让我们能够描述非水溶剂中的反应，并解释了氨(NH₃)等物种的作用——它接受一个质子变成NH₄⁺。
Conjugate acid‑base pairs are central to this theory. When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid. For example, in the reaction HCl + H₂O → H₃O⁺ + Cl⁻, HCl/Cl⁻ and H₃O⁺/H₂O are acid‑base conjugate pairs.
共轭酸碱对是该理论的核心。酸失去质子后形成其共轭碱；碱得到质子后形成其共轭酸。例如反应 HCl + H₂O → H₃O⁺ + Cl⁻中，HCl/Cl⁻和H₃O⁺/H₂O都是共轭酸碱对。
3. Common Properties of Acids | 酸的常见性质
Acids have a pH less than 7. They react with metals (above hydrogen in the reactivity series) to produce a salt and hydrogen gas. The general equation is: metal + acid → salt + hydrogen. Example: Mg + 2HCl → MgCl₂ + H₂.
酸的pH小于7。它们能与金属活动性顺序中氢以上的金属反应，生成盐和氢气。通式为：金属 + 酸 → 盐 + 氢气。实例：Mg + 2HCl → MgCl₂ + H₂。
Acids neutralise bases to form salt and water. They also react with carbonates to produce a salt, water and carbon dioxide: acid + carbonate → salt + water + CO₂. This effervescence is a classic test for carbonates.
酸能中和碱，生成盐和水。它们还能与碳酸盐反应生成盐、水和二氧化碳：酸 + 碳酸盐 → 盐 + 水 + CO₂。这种冒泡现象是检验碳酸盐的经典方法。
In aqueous solutions, strong acids ionise completely, while weak acids only partially dissociate. This affects conductivity and the rate of reactions.
在水溶液中，强酸完全电离，而弱酸仅部分解离。这会影响导电性和反应速率。
4. Common Properties of Bases and Alkalis | 碱与可溶碱的常见性质
A base is any substance that neutralises an acid. An alkali is a soluble base that releases OH⁻ ions in water. All alkalis are bases, but not all bases are alkalis. For example, copper(II) oxide (CuO) is a base but not an alkali because it is insoluble.
碱是任何能中和酸的物质。可溶碱(alkali)是在水中能释放OH⁻离子的碱。所有可溶碱都是碱，但并非所有碱都可溶。例如氧化铜(CuO)是碱，但不溶于水，因此不是可溶碱。
Alkaline solutions feel soapy and turn red litmus blue. They react with acids to form salt and water, and with ammonium salts to release ammonia gas. A classic test: warm the base with ammonium chloride – the pungent smell of NH₃ confirms the presence of a base.
碱性溶液有滑腻感，使红色石蕊变蓝。它们与酸反应生成盐和水，与铵盐反应释放出氨气。经典检验方法：将碱与氯化铵共热，刺鼻的NH₃气味即可证明碱的存在。
Common alkalis include NaOH (sodium hydroxide), KOH (potassium hydroxide), and Ca(OH)₂ (calcium hydroxide). Weak bases like ammonia solution (NH₃) only partially dissociate in water.
常见可溶碱包括NaOH(氢氧化钠)、KOH(氢氧化钾)和Ca(OH)₂(氢氧化钙)。弱碱如氨水(NH₃)在水中只部分解离。
5. The pH Scale: Measuring Acidity | pH标度：衡量酸碱性
The pH scale runs from 0 to 14. A solution with pH < 7 is acidic, pH = 7 is neutral, and pH > 7 is basic (alkaline). Each whole pH step represents a tenfold change in H⁺ concentration: a solution of pH 3 has 10 times more H⁺ ions than one of pH 4.
pH标度从0到14。pH < 7为酸性，pH = 7为中性，pH > 7为碱性。每一个整数pH步长代表H⁺浓度10倍的变化：pH 3的溶液H⁺浓度是pH 4溶液的10倍。
pH is measured using universal indicator, pH meters, or data loggers. Universal indicator is a mixture of dyes that turns from red (strong acid) through green (neutral) to violet (strong base). Litmus only tells you acidic or basic, not the pH value.
pH可以用通用指示剂、pH计或数据记录器测量。通用指示剂是一种混合染料，颜色从红(强酸)到绿(中性)再到紫(强碱)。石蕊只能指示酸碱性，不能给出具体pH值。
The formula pH = −log₁₀[H⁺] is used in IB calculations. Although you do not need to calculate pH from concentrations in WJEC GCSE, you should understand the logarithmic nature of the scale.
IB计算使用公式 pH = −log₁₀[H⁺]。尽管WJEC GCSE不要求用浓度计算pH，但你仍需理解标度的对数特性。
6. Strong and Weak Acids | 强酸与弱酸
A strong acid completely ionises in water. Examples: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃). For HCl, the equation is HCl → H⁺ + Cl⁻; essentially all molecules dissociate.
强酸在水中完全电离。例如盐酸(HCl)、硫酸(H₂SO₄)和硝酸(HNO₃)。对HCl：HCl → H⁺ + Cl⁻，几乎全部分子解离。
A weak acid only partially dissociates, setting up an equilibrium. Ethanoic acid (CH₃COOH) is a typical weak acid: CH₃COOH ⇌ H⁺ + CH₃COO⁻. At any moment, most molecules remain un‑ionised.
弱酸仅部分解离，存在平衡。乙酸(CH₃COOH)是典型的弱酸：CH₃COOH ⇌ H⁺ + CH₃COO⁻。在任何时刻，大部分分子仍保持未电离态。
This difference has practical consequences: at the same concentration, a strong acid has a lower pH (higher [H⁺]), reacts faster, and conducts electricity better than a weak acid. In titrations, the equivalence point pH differs sharply.
这一差异有实际影响：相同浓度下，强酸的pH更低([H⁺]更高)，反应更快，导电性更好。在滴定中，两者的等当点pH截然不同。
7. Strong and Weak Bases | 强碱与弱碱
Strong bases dissociate completely in water to release OH⁻. Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are strong bases. Equation: NaOH → Na⁺ + OH⁻.
强碱在水中完全解离释放OH⁻。氢氧化钠(NaOH)和氢氧化钾(KOH)是强碱。方程式：NaOH → Na⁺ + OH⁻。
Weak bases like ammonia (NH₃) accept protons but only partially ionise in water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. The equilibrium lies far to the left, so the concentration of OH⁻ is much lower than that of a strong base at the same concentration.
弱碱如氨(NH₃)能接受质子，但在水中仅部分电离：NH₃ + H₂O ⇌ NH₄⁺ + OH⁻。平衡偏向左方，因此OH⁻浓度远比同浓度的强碱低。
Again, electrical conductivity and reaction rates are lower for weak bases. When planning neutralisation, remember that you need the same amount of base in moles, regardless of strength, because the reaction is 1:1 with H⁺.
同样，弱碱的导电性和反应速率较低。设计中和反应时，记住无论碱的强弱，所需碱的摩尔数相同，因为与H⁺是1:1反应的。
8. Neutralisation and Ionic Equations | 中和反应与离子方程式
Neutralisation is the reaction of an acid with a base to produce a salt and water. The general word equation: acid + base → salt + water. For example, HCl + NaOH → NaCl + H₂O.
中和反应是酸与碱反应生成盐和水的过程。通式：酸 + 碱 → 盐 + 水。实例：HCl + NaOH → NaCl + H₂O。
The essential ionic process is the same for all strong acid–strong base reactions: H⁺ (from the acid) + OH⁻ (from the base) → H₂O. This is the net ionic equation. Spectator ions (like Na⁺ and Cl⁻) remain unchanged in solution.
核心的离子过程对所有强酸–强碱反应都相同：H⁺(来自酸) + OH⁻(来自碱) → H₂O。这就是净离子方程式。旁观离子(如Na⁺和Cl⁻)在溶液中保持不变。
Neutralisation is exothermic; the temperature rises. This can be followed with a thermometer in a calorimetry experiment. The enthalpy change of neutralisation for strong acids and bases is about −57 kJ mol⁻¹.
中和反应放热，温度升高。这可在量热实验中用温度计追踪。强酸强碱的中和焓变约为−57 kJ mol⁻¹。
9. Acid–Base Titration | 酸碱滴定
Titration is used to determine the concentration of an unknown acid or base. A known volume of one solution is measured with a pipette, and the other solution is added from a burette until the indicator just changes colour – this is the end point.
滴定用于测定未知酸或碱的浓度。用移液管量取一定体积的一种溶液，从滴定管中逐滴加入另一种溶液直到指示剂恰好变色——这就是终点。
Common indicators: phenolphthalein (colourless in acid, pink in alkali) or methyl orange (red in acid, yellow in alkali). The choice depends on the strength of the acid and base. For a strong acid–strong base titration, either works.
常用指示剂：酚酞(酸中无色，碱中粉红)或甲基橙(酸中红，碱中黄)。选择取决于酸碱的强弱。强酸–强碱滴定两者均可。
IB students need to perform titration calculations: using M₁V₁ = M₂V₂ for monoprotic acids, or accounting for stoichiometry (e.g., H₂SO₄ releases 2 H⁺). Accuracy in reading the meniscus and repeating until concordant results are vital.
IB学生需进行滴定计算：对于一元酸使用M₁V₁ = M₂V₂，或考虑化学计量(如H₂SO₄释放2个H⁺)。准确读取弯月面并重复至结果一致是得分关键。
10. Salts and Their Preparation | 盐及其制备
A salt is the ionic product of neutralisation, consisting of a positive ion from the base and a negative ion from the acid. Table salt, NaCl, is formed from NaOH and HCl. Soluble salts can be prepared by titration, then crystallised by evaporation.
盐是中和的离子产物，由来自碱的正离子与来自酸的负离子组成。食盐NaCl便是由NaOH和HCl生成。可溶性盐可通过滴定制备，再蒸发结晶。
Insoluble salts are made by precipitation: mixing two solutions containing the required ions. For example, BaCl₂ + Na₂SO₄ → BaSO₄ (s) + 2NaCl. The insoluble salt is filtered, washed and dried.
不溶性盐通过沉淀法制备：混合两种含有目标离子的溶液。例如BaCl₂ + Na₂SO₄ → BaSO₄(s) + 2NaCl。将沉淀过滤、洗涤、干燥即可。
Naming salts follows a pattern: hydro‑chlor‑ic acid → chlor‑ide; nitr‑ic acid → nitr‑ate; sulfur‑ic acid → sulf‑ate. Always refer to the acid’s parentage.
盐的命名规则：氢‑氯‑酸 → 氯‑化物；硝‑酸 → 硝‑酸盐；硫‑酸 → 硫‑酸盐。始终要追溯酸的来源。
11. Indicators and Their Colours | 指示剂及其变色
Litmus is a simple indicator: red in pH < 5, blue in pH > 8. Phenolphthalein changes at pH 8.2 (colourless to pink); methyl orange changes at pH 3.1–4.4 (red to yellow).
石蕊是最简单的指示剂：pH < 5 呈红色，pH > 8 呈蓝色。酚酞变色点在pH 8.2(无色变粉红)；甲基橙在pH 3.1–4.4变色(红变黄)。
Universal indicator gives a full spectrum: red (pH 0–3), orange (pH 3–5), yellow (pH 5–6), green (pH 7), blue (pH 8–11), violet (pH 11–14). It is essential for estimating pH in the lab.
通用指示剂呈现全色谱：红(pH 0–3)、橙(pH 3–5)、黄(pH 5–6)、绿(pH 7)、蓝(pH 8–11)、紫(pH 11–14)。它在实验室估测pH时不可或缺。
For IB internal assessments, you might need to justify your choice of indicator based on the equivalence point pH of a particular titration. For a weak acid–strong base titration, phenolphthalein is more suitable; for a strong acid–weak base, methyl orange works better.
在IB内部评估中，你可能需要根据特定滴定的等当点pH来论证指示剂的选择。弱酸‑强碱滴定更适合用酚酞；强酸‑弱碱滴定甲基橙更佳。
12. Exam Tips and Common Pitfalls | 考试技巧与常见误区
Always write state symbols (s, l, g, aq) in equations when asked. In WJEC, omitting them can lose marks. For IB, state symbols are expected in all chemical equations unless otherwise stated.
只要题目要求，务必在方程式中标注状态符号(s, l, g, aq)。WJEC中遗漏会扣分。IB所有化学方程式除非另有说明，否则都要标示。
Distinguish clearly between ‘strength’ and ‘concentration’. A strong acid is fully ionised; a concentrated acid simply has a lot of acid dissolved. You can have a dilute strong acid or a concentrated weak acid.
清晰区分“强度”与“浓度”。强酸是完全电离的；浓酸只是溶解了大量的酸。你可以有稀的强酸，也可以有浓的弱酸。
In neutralisation, “salt” does not always mean table salt; it refers to any ionic compound formed. Don’t forget that ammonia gas turns damp red litmus paper blue – a key test for ammonium salts.
在中和反应中，“盐”不总是指食盐，而是泛指生成的任何离子型化合物。别忘了，氨气能使湿润的红色石蕊试纸变蓝——这是检验铵盐的关键。
Finally, practice writing ionic and net ionic equations. Start by writing the full balanced equation, split aqueous ionic compounds into ions, then cancel spectator ions. This skill is vital for both IB and WJEC higher‑tier questions.
最后，多练习书写离子方程式和净离子方程式。先写出完整的配平方程式，将可溶离子型化合物拆分为离子，再消去旁观离子。这项技能对IB和WJEC高阶题目至关重要。
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