📚 Combined 121: Acid-Base Equilibria for Edexcel A-Level Chemistry | 综合复习121：Edexcel A-Level化学酸碱平衡
Acid-base equilibria lie at the heart of quantitative chemistry, linking equilibrium constants, pH calculations, buffer design and titration analysis. This combined revision guide unpacks every core idea from Edexcel Topic 12, from Brønsted–Lowry definitions to buffer pH derivations, giving you a structured path to exam confidence.
酸碱平衡是定量化学的核心，它将平衡常数、pH 计算、缓冲液设计和滴定分析紧密联系在一起。这篇综合复习指南拆解了 Edexcel 专题 12 的每一个核心概念，从布朗斯特-劳里定义直到缓冲液 pH 推导，为你铺设一条通向考试自信的系统路径。
1. Brønsted–Lowry Acids and Bases | 布朗斯特-劳里酸碱理论
A Brønsted–Lowry acid is a proton (H⁺) donor, and a base is a proton acceptor. This definition extends acid-base behaviour beyond aqueous solutions, allowing species like gaseous HCl and NH₃ to be classified without requiring water as the medium.
布朗斯特-劳里酸是质子 (H⁺) 给体，碱是质子受体。这一定义将酸碱行为拓展到水溶液之外，即使没有水作为介质，气态的 HCl 和 NH₃ 也可以被分类。
When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid. The strength of an acid is linked to the stability of its conjugate base.
酸给出质子后形成其共轭碱；碱接受质子后形成其共轭酸。酸的强度与其共轭碱的稳定性密切相关。
2. Conjugate Acid-Base Pairs | 共轭酸碱对
A conjugate acid-base pair consists of two species that differ by a single proton. For example, CH₃COOH and CH₃COO⁻ are a pair; NH₄⁺ and NH₃ are another.
共轭酸碱对由两个仅相差一个质子的物种组成。例如 CH₃COOH 与 CH₃COO⁻ 是一对；NH₄⁺ 与 NH₃ 是另一对。
In any acid-base equilibrium, both forward and reverse reactions involve proton transfer between conjugate pairs. The equilibrium position indicates which side has the weaker acid and base — equilibrium favours the side with the weaker acid.
在任何酸碱平衡中，正逆反应都是在共轭对之间进行质子转移。平衡位置可以指出哪一侧的酸和碱更弱——平衡倾向于弱酸一侧。
3. The pH Scale and [H⁺] | pH 标度与 [H⁺]
pH is defined as the negative logarithm to base 10 of the hydrogen ion concentration:
pH 定义为氢离子浓度的以 10 为底的负对数：
pH = –lg [H⁺]
Similarly, [H⁺] can be recovered using:
同理，可通过下式求回 [H⁺]：
[H⁺] = 10⁻ᵖᴴ
ten-fold change in [H⁺]. Neutral water at 298 K has [H⁺] = 1.0 × 10⁻⁷ mol dm⁻³, so pH = 7.
一个 pH 单位的变化对应 [H⁺] 的十倍变化。298 K 时中性水的 [H⁺] = 1.0 × 10⁻⁷ mol dm⁻³，因此 pH = 7。
4. Strong Acids and Bases: Complete Dissociation | 强酸和强碱：完全电离
sociate fully in water. For a monoprotic strong acid of concentration c, [H⁺] = c and pH = –lg c.
强酸如 HCl、HNO₃ 和 H₂SO₄ (第一步电离) 在水中完全电离。对于浓度 c 的一元强酸，[H⁺] = c，pH = –lg c。
er, Kₑᵧ = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K, the pH can be calculated from [OH⁻].
强碱如 NaOH 和 KOH 每摩尔释放一摩尔 OH⁻。利用水的离子积 Kₑᵧ = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ (298 K)，即可由 [OH⁻] 计算 pH。
| Common Strong Acids | Common Strong Bases |
| HCl, HBr, HI, HNO₃, H₂SO₄ (first H⁺), HClO₄ | NaOH, KOH, Ba(OH)₂ |
常见强酸：HCl、HBr、HI、HNO₃、H₂SO₄ (第一个 H⁺)、HClO₄；常见强碱：NaOH、KOH、Ba(OH)₂。
5. Weak Acids and the Acid Dissociation Constant Kₐ | 弱酸和酸解离常数 Kₐ
Weak acids such as CH₃COOH partially dissociate, establishing an equilibrium:
弱酸如 CH₃COOH 部分电离，建立平衡：
HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)
The acid dissociation constant Kₐ is given by:
酸解离常数 Kₐ 表示为：
Kₐ = [H⁺][A⁻] / [HA]
Kₐ has units of mol dm⁻³ and its magnitude indicates acid strength — a larger Kₐ means a stronger acid. Because the degree of dissociation is often very small, approximations can simplify calculations.
Kₐ 的单位是 mol dm⁻³，其大小反映酸的强度——Kₐ 越大，酸越强。因为解离度通常很小，常用近似来简化计算。
6. Calculating pH of Weak Acids | 弱酸 pH 计算
ssuming [H⁺] ≈ [A⁻] and [HA] at equilibrium ≈ c, the equilibrium expression simplifies to:
对于初始浓度 c 的一元弱酸 HA，假设 [H⁺] ≈ [A⁻] 且平衡时 [HA] ≈ c，平衡表达式简化为：
Kₐ ≈ [H⁺]² / c → [H⁺] = √Kₐ × c
Then pH is obtained from pH = – oximation is valid when the acid is weak enough and the concentration is not extremely low — typically when c / Kₐ > 100.
随后由 pH = –lg[H⁺] 求 pH。当酸足够弱且浓度不是极低时（通常 c/Kₐ > 100），该近似成立。
7. pKₐ and Its Relationship to Kₐ | pKₐ 及其与 Kₐ 的关系
pKₐ is defined analogously to pH:
pKₐ 的定义与 pH 类似：
pKₐ = –lg Kₐ
while chloroacetic acid has pKₐ = 2.87, reflecting the electron-withdrawing effect of Cl.
pKₐ 越小，酸越强。pKₐ 值可以快速比较酸强度：CH₃COOH 的 pKₐ = 4.76，而氯乙酸的 pKₐ = 2.87，这反映了 Cl 的吸电子效应。
For a given weak acid, when [HA] = [A⁻], pH = pKₐ — a crucial point on titration curves and the basis of buffer pH calculation.
对于给定弱酸，当 [HA] = [A⁻] 时，pH = pKₐ——这是滴定曲线上的关键点，也是缓冲液 pH 计算的基础。
8. Buffer Solutions: What They Are and How They Work | 缓冲溶液：原理与作用
A buffer solution resists changes in pH upon addition of small amounts of acid or base. It consists of a weak acid and its conjugate base in significant concentrations, or a weak base and its conjugate acid.
缓冲溶液能够抵抗少量外加酸或碱引起的 pH 变化。它由高浓度的弱酸及其共轭碱，或弱碱及其共轭酸组成。
nearly constant. Blood pH (≈7.4) is maintained by carbonic acid/hydrogencarbonate buffers.
弱酸中和外加的 OH⁻，共轭碱中和外加的 H⁺，使 [H⁺]/[OH⁻] 比值近乎稳定。血液 pH (≈7.4) 由碳酸/碳酸氢盐缓冲系维持。
9. Calculating pH of Buffer Solutions | 缓冲溶液 pH 计算
For an acidic buffer made from a weak acid HA and its salt NaA, the Henderson–Hasselbalch equation is the standard tool:
对于由弱酸 HA 和其盐 NaA 组成的酸性缓冲液，亨德森-哈塞尔巴尔赫方程是标准工具：
pH = pKₐ + lg ([A⁻] / [HA])
Here [A⁻] is the concentration of the conjugate base (usually from the salt, which fully dissociates) and [HA] is the concentration of the weak acid. The equation can be derived from the Kₐ expression by taking negative logs.
式中 [A⁻] 是共轭碱的浓度（通常来自完全电离的盐），[HA] 是弱酸的浓度。该方程可由 Kₐ 表达式取负对数推导得出。
When preparing a buffer to a specific pH, one chooses an acid with a pKₐ close to the target pH and adjusts the ratio [A⁻]/[HA]. Adding water does not alter the ratio, so buffer pH is independent of dilution.
配制指定 pH 的缓冲液时，选择 pKₐ 接近目标 pH 的酸，并调节 [A⁻]/[HA] 比值。加水不会改变该比值，因此缓冲液 pH 与稀释度无关。
10. Acid–Base Titration Curves | 酸碱滴定曲线
Titration curves plot pH against volume of titrant added. Their shapes reveal equivalence points and buffer regions.
滴定曲线描绘了 pH 随滴定剂加入体积的变化。曲线的形状展现出等当点和缓冲区域。
- Strong acid–Strong base: steep vertical rise at pH 7; negligible buffer region.
- 强酸–强碱：pH 7 处出现近乎垂直的陡升；缓冲区可忽略。
- Weak acid–Strong base: initial pH higher; a buffer region before the equivalence point, which lies above pH 7.
- 弱酸–强碱：起始 pH 较高；等当点前出现缓冲区，等当点 pH > 7。
- Weak base–Strong acid: equivalence point below pH 7; buffer region from the weak base/conjugate acid pair.
- 弱碱–强酸：等当点 pH < 7；弱碱/共轭酸对形成缓冲区。
gion is flatter, where pH changes slowly. At half-neutralisation of a weak acid, pH = pKₐ.
缓冲区较平坦，pH 变化缓慢。弱酸半中和时，pH = pKₐ。
11. Choosing Indicators for Titrations | 滴定指示剂的选择
have different colours. It changes colour over a pH range approximately equal to pKᵢₙ ± 1.
指示剂是一种弱酸或弱碱，其共轭形式具有不同颜色。它在 pH ≈ pKᵢₙ ± 1 的范围内变色。
The indicator must have its pH range lying entirely within the steep part of the titration curve, so that a single drop of titrant causes a visible colour change at the equivalence point.
指示剂的变色范围必须完全落在滴定曲线的陡峭部分内，这样才能在等当点时由一滴滴定剂引起可见颜色变化。
| Indicator | pH range | Typical use |
| Methyl orange | 3.1 – 4.4 | Strong acid–strong base or weak base–strong acid |
| Phenolphthalein | 8.3 – 10.0 | Strong acid–strong base or weak acid–strong base |
甲基橙 pH 范围 3.1–4.4，适用于强酸-强碱或弱碱-强酸滴定；酚酞 pH 范围 8.3–10.0，适用于强酸-强碱或弱酸-强碱滴定。
12. Enthalpy of Neutralisation | 中和焓变
The standard enthalpy change of neutralisation, ΔH⦵ₙₑᵤₜ, is the heat evolved when one mole of water is formed from H⁺(aq) and OH⁻(aq) under standard conditions. For strong acids with strong bases, the value is approximately –57 kJ mol⁻¹, because the net equation is simply:
标准中和焓变 ΔH⦵ₙₑᵤₜ 是在标准条件下由 H⁺(aq) 和 OH⁻(aq) 生成一摩尔水所放出的热量。强酸与强碱中和时，该值约为 –57 kJ mol⁻¹，因为净离子方程式就是：
H⁺(aq) + OH⁻(aq) → H₂O(l)
Weak acids or weak bases give less exothermic values because some energy is absorbed to ionise the undissociated molecules before neutralisation can occur.
弱酸或弱碱的中和焓变数值较小（放热更少），因为部分能量需要先用于电离未解离的分子，然后才能发生中和反应。
Published by TutorHao | Chemistry Revision Series | aleveler.com
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