📚 Acids and Bases Key Points for IB and OCR Sciences | IB OCR 科学：酸与碱 考点精讲
Acids and bases form one of the most fundamental frameworks in chemistry, underlying countless reactions in the laboratory, industry, and living organisms. This revision guide brings together essential concepts from both IB Chemistry and OCR A-level Chemistry, highlighting definitions, calculations, equilibria, titrations, and real-world contexts to help you master every exam question.
酸与碱是化学中最基本的理论框架之一，支撑着实验室、工业和生物体内的无数反应。本文结合 IB 化学与 OCR A-level 化学的核心考点，系统梳理定义、计算、平衡、滴定和实际应用，助你全面攻克酸碱类考题。
1. Brønsted-Lowry Theory of Acids and Bases | 布朗斯特-劳里酸碱理论
According to the Brønsted-Lowry theory, an acid is a proton (H⁺) donor, and a base is a proton acceptor. This definition expands the classical Arrhenius concept to include non-aqueous systems and ions like NH₄⁺ and CO₃²⁻.
根据布朗斯特-劳里理论，酸是质子（H⁺）的给予体，碱是质子的接受体。这一定义扩展了经典的阿伦尼乌斯概念，适用于非水体系以及 NH₄⁺、CO₃²⁻ 等离子。
When hydrogen chloride gas dissolves in water, HCl donates a proton to H₂O, forming H₃O⁺ and Cl⁻. Water acts as a base in this reaction. Similarly, ammonia (NH₃) accepts a proton from water to form NH₄⁺ and OH⁻, demonstrating that water can also act as an acid.
当氯化氢气体溶于水时，HCl 向 H₂O 提供一个质子，生成 H₃O⁺ 和 Cl⁻，此时水作为碱。同样，氨（NH₃）从水中接受质子形成 NH₄⁺ 和 OH⁻，说明水也可作为酸。
2. Lewis Acids and Bases | 路易斯酸碱理论
The Lewis theory offers an even broader view: a Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. This definition includes species such as BF₃ and metal cations that do not necessarily involve protons.
路易斯理论给出了更宽泛的定义：路易斯酸是电子对接受体，路易斯碱是电子对给予体。该定义涵盖了 BF₃ 和金属阳离子等不一定涉及质子的物种。
Many organic reactions and transition metal complex formations are best described using Lewis acid-base interactions. For example, the reaction between NH₃ (electron-pair donor) and BCl₃ (electron-pair acceptor) forms a coordinate bond, producing H₃N–BCl₃.
许多有机反应和过渡金属配合物的形成最适合用路易斯酸碱作用来描述。例如 NH₃（电子对给予体）与 BCl₃（电子对接受体）反应，生成配位键产物 H₃N–BCl₃。
3. Conjugate Acid-Base Pairs | 共轭酸碱对
Every Brønsted-Lowry acid has a conjugate base, formed after donating a proton; every base has a conjugate acid, formed after accepting a proton. The acid-base pair differ by a single proton, H⁺.
每个布朗斯特酸都有一个共轭碱，即酸给出质子后形成的物种；每个碱都有一个共轭酸，即碱接受质子后形成的物种。酸碱对之间仅相差一个质子 H⁺。
In the equilibrium CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺, acetic acid and acetate ion form one conjugate pair, while H₂O and H₃O⁺ form the other. The strength of an acid is inversely related to the strength of its conjugate base.
在平衡 CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺ 中，乙酸和乙酸根离子是一对共轭酸碱对，而 H₂O 和 H₃O⁺ 是另一对。酸的强度与其共轭碱的强度成反比。
4. The pH Scale and Strong Acids/Bases | pH 标度与强酸强碱
pH is defined as the negative logarithm to base 10 of the hydrogen ion concentration: pH = –log₁₀[H⁺]. In aqueous solutions at 25 °C, a neutral solution has pH = 7.00, acidic solutions have pH < 7, and basic solutions have pH > 7.
pH 定义为氢离子浓度的负常用对数：pH = –log₁₀[H⁺]。在 25 °C 的水溶液中，中性溶液 pH = 7.00，酸性溶液 pH < 7，碱性溶液 pH > 7。
Strong acids such as HCl, HNO₃, and H₂SO₄ fully dissociate in water, so [H⁺] equals the initial concentration of the acid multiplied by the number of protons released per formula unit. For a 0.10 mol dm⁻³ HCl solution, [H⁺] = 0.10 mol dm⁻³, giving pH = 1.00.
强酸（如 HCl、HNO₃ 和 H₂SO₄）在水中完全电离，因此 [H⁺] 等于酸的初始浓度乘以每个分子释放的质子数。对于 0.10 mol dm⁻³ HCl 溶液，[H⁺] = 0.10 mol dm⁻³，pH = 1.00。
Similarly, strong bases like NaOH and KOH release OH⁻ fully. The pOH (= –log[OH⁻]) and the relationship pH + pOH = 14.00 at 25 °C allow easy calculation of pH for strong bases.
类似地，强碱如 NaOH 和 KOH 完全释放 OH⁻。利用 pOH = –log[OH⁻] 及 25 °C 下 pH + pOH = 14.00 的关系，可轻松计算强碱的 pH。
5. Weak Acids, Ka and pKa | 弱酸、Ka 与 pKa
Weak acids such as ethanoic acid (CH₃COOH) only partially dissociate in water. The equilibrium constant for a weak acid is the acid dissociation constant, Ka. For a generic weak acid HA: HA + H₂O ⇌ H₃O⁺ + A⁻, Ka = [H₃O⁺][A⁻]/[HA], where water concentration is treated as constant.
弱酸如乙酸（CH₃COOH）在水中仅部分电离。弱酸的电离平衡常数称为酸解离常数 Ka。对于一般弱酸 HA：HA + H₂O ⇌ H₃O⁺ + A⁻，Ka = [H₃O⁺][A⁻]/[HA]，水的浓度视为常数。
The pKa value, defined as –log₁₀(Ka), gives a direct measure of acid strength: the smaller the pKa, the stronger the weak acid. Typical Ka values are very small, so using pKa numbers (e.g., ethanoic acid pKa = 4.76) is more convenient.
pKa 定义为 –log₁₀(Ka)，可直接衡量酸强度：pKa 越小，弱酸越强。典型的 Ka 值非常小，使用 pKa（如乙酸 pKa = 4.76）更为方便。
Calculating pH of a weak acid solution involves an ICE table and the approximation [H⁺] ≈ √(Ka × [HA]₀), provided the acid is less than 5 % dissociated. This approximation is valid when [HA]₀ / Ka > 100.
计算弱酸溶液的 pH 需使用 ICE 表格及近似 [H⁺] ≈ √(Ka × [HA]₀)，前提是电离度小于 5 %。当 [HA]₀ / Ka > 100 时该近似成立。
6. Weak Bases, Kb and pKb | 弱碱、Kb 与 pKb
Weak bases such as ammonia (NH₃) accept protons partially, establishing an equilibrium with a base dissociation constant, Kb. For the reaction NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, Kb = [NH₄⁺][OH⁻]/[NH₃].
弱碱如氨（NH₃）部分接受质子，建立平衡，其碱解离常数为 Kb。对于反应 NH₃ + H₂O ⇌ NH₄⁺ + OH⁻，Kb = [NH₄⁺][OH⁻]/[NH₃]。
Analogous to pKa, pKb = –log₁₀(Kb). For a conjugate acid-base pair at 25 °C, the useful relationship is pKa + pKb = 14.00. This allows you to convert between the acid and base strength of a conjugate system.
与 pKa 类似，pKb = –log₁₀(Kb)。对于共轭酸碱对，在 25 °C 下存在有用关系 pKa + pKb = 14.00，可实现共轭体系酸碱强度的相互转换。
The pH of a weak base solution can be found by calculating pOH first using [OH⁻] ≈ √(Kb × [B]₀), then converting to pH at 25 °C.
计算弱碱溶液的 pH 时，可先利用 [OH⁻] ≈ √(Kb × [B]₀) 求出 pOH，再于 25 °C 下转换为 pH。
7. Autoionization of Water and Kw | 水的自耦电离与 Kw
Pure water undergoes autoionization: 2H₂O ⇌ H₃O⁺ + OH⁻. The equilibrium constant Kw = [H₃O⁺][OH⁻] has a value of 1.0 × 10⁻¹⁴ at 25 °C. This relation is fundamental to all aqueous acid-base calculations.
纯水发生自耦电离：2H₂O ⇌ H₃O⁺ + OH⁻，平衡常数 Kw = [H₃O⁺][OH⁻]，在 25 °C 时为 1.0 × 10⁻¹⁴。这一关系是所有水溶液酸碱计算的基础。
Kw changes with temperature; it increases as temperature rises, making the neutral point shift. For example, at 40 °C, Kw ≈ 2.9 × 10⁻¹⁴, so a neutral solution has pH ≈ 6.77, not 7.
Kw 随温度变化，温度升高 Kw 增大，中性点随之移动。例如 40 °C 时 Kw ≈ 2.9 × 10⁻¹⁴，中性溶液的 pH ≈ 6.77，不再是 7。
8. Buffer Solutions | 缓冲溶液
A buffer solution resists changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (or a weak base and its conjugate acid) in appreciable concentrations.
缓冲溶液能在加入少量酸或碱时抵抗 pH 的变化。它由浓度可观的弱酸及其共轭碱（或弱碱及其共轭酸）组成。
The Henderson-Hasselbalch equation for an acidic buffer is:
pH = pKa + log₁₀([A⁻]/[HA]). This equation is extremely powerful for designing buffers and predicting pH after addition of small amounts of H⁺ or OH⁻.
酸性缓冲溶液的 Henderson-Hasselbalch 方程为：
pH = pKa + log₁₀([A⁻]/[HA])。该方程在缓冲体系的设计及少量 H⁺ 或 OH⁻ 加入后 pH 的预测中功能强大。
Buffers are vital in biological systems (e.g., blood buffered by H₂CO₃/HCO₃⁻) and in industrial processes. In the lab, a buffer can be prepared by partially neutralizing a weak acid with a strong base or by mixing a weak acid and its salt.
缓冲体系在生物系统（如血液的 H₂CO₃/HCO₃⁻ 缓冲对）和工业过程中至关重要。实验室中可通过用强碱部分中和弱酸或混合弱酸及其盐来制备缓冲溶液。
9. Acid-Base Titrations and Curves | 酸碱滴定与滴定曲线
An acid-base titration involves the gradual addition of a titrant of known concentration to an analyte until neutralization is reached. The equivalence point is where the amount of acid equals the amount of base stoichiometrically. The end point is indicated by a colour change of an indicator.
酸碱滴定是将已知浓度的滴定剂逐渐加入待测物中直至中和的过程。化学计量点指酸与碱的物质的量按化学计量数完全反应之时，而滴定终点由指示剂颜色变化指示。
Titration curves plot pH against volume added. Key curve shapes differ for strong acid-strong base, strong acid-weak base, weak acid-strong base, and weak acid-weak base combinations. The vertical region of rapid pH change helps identify the equivalence point.
滴定曲线以 pH 对加入体积作图。强酸-强碱、强酸-弱碱、弱酸-强碱和弱酸-弱碱的曲线形状各有不同。pH 快速变化的陡直区域有助于识别化学计量点。
For a weak acid-strong base titration, the pH at half-neutralization equals the pKa of the weak acid. This point, often marked as the half-equivalence point, is a quick method to determine Ka experimentally.
在弱酸-强碱滴定中，半中和点处溶液的 pH 等于弱酸的 pKa。此点常称作半点化学计量点，是实验测定 Ka 的快捷方法。
10. Indicators | 酸碱指示剂
Acid-base indicators are weak acids or bases that change colour within a specific pH range. The colour change occurs because the protonated and deprotonated forms have different colours, and the ratio shifts around pH = pKIn.
酸碱指示剂本身是弱酸或弱碱，在特定 pH 范围内发生颜色变化。颜色变化源于其质子化与去质子化形式颜色不同，在 pH ≈ pKIn 附近比例发生转变。
For a successful titration, the indicator must have a pH range that falls entirely within the steep portion of the titration curve. Common choices include phenolphthalein (range 8.2–10.0) for strong base-strong acid titrations, and methyl orange (3.1–4.4) for strong acid-weak base titrations.
要获得准确的滴定终点，指示剂的变色范围必须完全落在滴定曲线的陡直段内。常用选择如酚酞（变色范围 8.2–10.0）适用于强碱-强酸滴定，甲基橙（3.1–4.4）适用于强酸-弱碱滴定。
| Indicator | Colour in acid | Colour in base | pH range |
|---|---|---|---|
| Methyl orange | Red | Yellow | 3.1 – 4.4 |
| Bromothymol blue | Yellow | Blue | 6.0 – 7.6 |
| Phenolphthalein | Colourless | Pink | 8.2 – 10.0 |
Using an indicator with a pKIn as close as possible to the equivalence point pH ensures a sharp end point with minimal titration error.
选择 pKIn 尽可能接近化学计量点 pH 的指示剂，可确保终点敏锐且滴定误差极小。
11. Acid-Base Properties of Oxides | 氧化物的酸碱性
Metallic oxides are typically basic, reacting with acids to form salts and water. For example, Na₂O + 2HCl → 2NaCl + H₂O. Non-metallic oxides are usually acidic; they react with bases to form salts and water, e.g., CO₂ + 2NaOH → Na₂CO₃ + H₂O.
金属氧化物通常呈碱性，与酸反应生成盐和水。例如 Na₂O + 2HCl → 2NaCl + H₂O。非金属氧化物通常呈酸性，与碱反应生成盐和水，如 CO₂ + 2NaOH → Na₂CO₃ + H₂O。
Amphoteric oxides, such as Al₂O₃ and ZnO, can react with both acids and bases. Aluminium oxide dissolves in hydrochloric acid to give AlCl₃ and in sodium hydroxide to form NaAlO₂ or Al(OH)₄⁻ complexes.
两性氧化物如 Al₂O₃ 和 ZnO，既能与酸又能与碱反应。氧化铝既溶于盐酸生成 AlCl₃，又溶于氢氧化钠形成 NaAlO₂ 或 Al(OH)₄⁻ 配合物。
Some oxides are neutral, such as CO and NO, showing no acidic or basic behaviour. Understanding these trends is important for predicting reaction outcomes and for environmental issues like acid rain caused by SO₂ and NOₓ.
某些氧化物如 CO 和 NO 为中性氧化物，不显酸性或碱性。掌握这些变化规律对预测反应产物和理解酸雨（由 SO₂、NOₓ 引起）等环境问题至关重要。
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