📚 A-Level OCR Chemistry: Formula Summary Handbook | A-Level OCR 化学：公式汇总手册
This handbook brings together the essential equations and expressions you will encounter across the A-Level OCR Chemistry (H432) specification. Each formula is presented with a concise explanation of its terms and practical context, ensuring you can apply them confidently in calculations, data analysis and exam questions. Use this as your go-to revision companion for quantitative chemistry.
本手册汇集了 A-Level OCR 化学（H432）课程中所有核心公式和表达式。每个公式都配有术语解析与应用情境的简要说明，帮助你在计算、数据分析和考试中自信地运用。将这份资料作为你定量化学复习的首选伴侣。
1. The Mole and Molar Mass | 摩尔与摩尔质量
The Avogadro constant, L = 6.02 × 10²³ mol⁻¹, defines the number of particles in one mole.
物质的量（单位：摩尔 mol）是连接质量、粒子数和体积的核心概念。阿伏加德罗常数 L = 6.02 × 10²³ mol⁻¹ 定义了一摩尔物质所含的粒子数。
n = m / M
where n = amount (mol), m = mass (g), M = molar mass (g mol⁻¹).
其中 n = 物质的量，m = 质量，M = 摩尔质量。
n = N / L
where N = number of particles (atoms, molecules, ions).
其中 N = 粒子数。
For solutions: n = c × V, where c = concentration (mol dm⁻³), V = volume (dm³).
对于溶液：n = c × V，c = 浓度（mol dm⁻³），V = 体积（dm³）。
Always convert cm³ to dm³ by dividing by 1000.
务必记得将 cm³ 除以 1000 换算为 dm³。
2. Ideal Gas Equation | 理想气体状态方程
The ideal gas law combines pressure, volume, temperature and amount of gas. It works best at low pressure and high temperature.
理想气体定律将气体的压力、体积、温度与物质的量联系起来，在低压高温下最为适用。
pV = nRT
V = volume (m³), n = amount (mol), R = gas constant 8.31 J mol⁻¹ K⁻¹, T = temperature (K).
p = 压力（Pa），V = 体积（m³），R = 气体常数 8.31 J mol⁻¹ K⁻¹，T = 温度（K）。
Temperature must be in kelvin: T(K) = θ(°C) + 273.
温度必须使用开尔文：T(K) = θ(°C) + 273。
Common conversions: 1 atm = 101 325 Pa, 1 dm³ = 1 × 10⁻³ m³, 1 cm³ = 1 × 10⁻⁶ m³.
常见换算：1 atm = 101 325 Pa，1 dm³ = 1 × 10⁻³ m³，1 cm³ = 1 × 10⁻⁶ m³。
3. Empirical and Molecular Formulae | 经验式与分子式
Empirical formula gives the simplest whole-number ratio of atoms in a compound; molecular formula shows the actual number of atoms in one molecule.
经验式表示化合物中原子的最简整数比；分子式则表示一个分子中原子的实际数目。
To find empirical formula from % composition:
由元素质量分数求经验式的步骤：
Divide % obtain moles → divide each by the smallest mole value → multiply to obtain whole numbers.
将每种元素的质量分数除以相对原子质量得到物质的量 → 除以最小物质的量 → 乘上适当倍数得到最简整数比。
lecular formula = n × empirical formula, where n = Mᵣ (compound) / Mᵣ (empirical formula).
分子式 = n × 经验式，其中 n = 化合物相对分子质量 / 经验式相对质量。
n = Mᵣ(molecule) / Mᵣ(empirical unit)
4. Enthalpy Changes (ΔH) | 焓变 (ΔH)
Enthalpy change is the heat energy transferred at constant pressure. Standard enthalpy changes are measured under 100 kPa and a stated temperature, usually 298 K.
焓变是恒压条件下传递的热量。标准焓变在 100 kPa 和指定温度（通常 298 K）下测定。
ΔH = H(products) – H(reactants)
Exothermic reactions: ΔH is negative; endothermic: ΔH is positive.
放热反应：ΔH 为负值；吸热反应：ΔH 为正值。
Key experimental equation using calorimetry:
量热实验的关键公式：
q = mcΔT
c = specific heat capacity (4.18 J g⁻¹ K⁻¹ for water), ΔT = temperature change (K or °C).
q = 热量（J），m = 溶液质量（g），c = 比热容（水为 4.18 J g⁻¹ K⁻¹），ΔT = 温度变化。
Then ΔH = – moles of limiting reactant. The negative sign reflects the enthalpy change of the system.
然后 ΔH = –q / n，n = 限制试剂的物质的量。负号表示体系的焓变。
5. Hess's Law and Enthalpy Cycles | 盖斯定律与焓循环
Hess's Law: The total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.
盖斯定律：只要始态和终态相同，反应的总焓变与途径无关。
ΔH(direct) = ΔH(route A) + ΔH(route B) (if multiple steps)
For combustion cycles: ΔHᶠ (formation) or ΔHᶜ (combustion) are used to construct cycles.
利用燃烧或生成焓构建循环：例如 ΔH⦵(reaction) = ΣΔH⦵f(products) – ΣΔH⦵f(reactants)。
ΔH⦵ = ΣΔH⦵f(products) – ΣΔH⦵f(reactants)
Or using combustion data: ΔH⦵ = ΣΔH⦵c(reactants) – ΣΔH⦵c(products).
或用燃烧数据：ΔH⦵ = ΣΔH⦵c(反应物) – ΣΔH⦵c(生成物)。
Always balance equations and pay attention to the sign of each ΔH value.
务必配平方程式并注意每个 ΔH 值的符号。
6. Equilibrium Constant Kc | 平衡常数 Kc
For a reversible reaction aA + bB ⇌ cC + dD at a given temperature, the equilibrium constant in terms of concentration is Kc.
对于可逆反应 aA + bB ⇌ cC + dD，在指定温度下，以浓度表示的平衡常数为 Kc。
Kc = [C]ᶜ [D]ᵈ / ([A]ᵃ [B]ᵇ)
Only species in the gaseous or aqueous phase appear in the expression; solids and pure liquids are omitted.
只有气态或水溶液中的物种出现在表达式中；固体和纯液体省略。
Kc is constant only at a fixed temperature. If Kc >> 1, equilibrium lies to the right; if Kc << 1, equilibrium lies to the left.
Kc 仅在温度恒定时为常数。Kc >> 1 表明平衡趋向产物；Kc << 1 则平衡偏向反应物。
7. Acid Dissociation Constant Ka and pH | 酸解离常数 Ka 与 pH
For a weak acid HA ⇌ H⁺ + A⁻, the acid dissociation constant Ka measures the strength of the acid.
对于弱酸 HA ⇌ H⁺ + A⁻，酸解离常数 Ka 用于衡量酸的强度。
Ka = [H⁺][A⁻] / [HA]
The higher the Ka, the stronger the weak acid. pKa = –log₁₀ Ka.
Ka 越大，弱酸越强。pKa = –log₁₀ Ka。
pH is defined as:
pH = –log₁₀ [H⁺]
] = [acid]; for a weak acid, assuming [H⁺] = [A⁻] and [HA] ≈ initial concentration, [H⁺] = √Ka × [HA].
强一元酸：[H⁺] = 酸的浓度；对于弱酸，假设 [H⁺] = [A⁻] 且 [HA] ≈ 起始浓度，则 [H⁺] = √Ka × [HA]。
8. Ionic Product of Water Kw | 水的离子积 Kw
Water undergoes slight self-ionisation: 2H₂O ⇌ H₃O⁺ + OH⁻, simplified as H₂O ⇌ H⁺ + OH⁻.
水存在微弱的自解离：2H₂O ⇌ H₃O⁺ + OH⁻，简写为 H₂O ⇌ H⁺ + OH⁻。
Kw = [H⁺][OH⁻]
At 298 K, Kw = 1.0 × 10⁻¹⁴ mol² dm⁻⁶.
在 298 K 时，Kw = 1.0 × 10⁻¹⁴ mol² dm⁻⁶。
pKw = –log₁₀ Kw = 14 at 298 K, and pKw = pH + pOH.
298 K 时 pKw = –log₁₀ Kw = 14，且 pKw = pH + pOH。
ationship allows calculation of pH for strong bases: if [OH⁻] is known, [H⁺] = Kw / [OH⁻].
由此可计算强碱的 pH：若已知 [OH⁻]，则 [H⁺] = Kw / [OH⁻]。
9. Buffer Solutions | 缓冲溶液
A buffer solution resists changes in pH when small amounts of acid or base are added. It usually contains a weak acid and its conjugate base.
缓冲溶液能在加入少量酸或碱时抵抗 pH 变化，通常由弱酸及其共轭碱组成。
The pH of an acidic buffer can be estimated using the Henderson–Hasselbalch equation:
酸性缓冲液的 pH 可用 Henderson–Hasselbalch 方程估算：
pH = pKa + log₁₀ ([A⁻] / [HA])
Here [A⁻] is the concentration of the salt (conjugate base) and [HA] is the concentration of the weak acid.
其中 [A⁻] 为盐（共轭碱）的浓度，[HA] 为弱酸的浓度。
[HA], pH = pKa. This means the buffer works most effectively when the ratio is close to 1.
当 [A⁻] = [HA] 时，pH = pKa，此时缓冲效率最高。
Buffers are essential in biochemical systems, such as blood (H₂CO₃/HCO₃⁻ buffer).
缓冲液在生物体系中至关重要，例如血液中的 H₂CO₃/HCO₃⁻ 缓冲对。
10. Rate Equations | 速率方程
f reactants through a rate equation. For a reaction A + B → products, the general form is:
反应速率通过速率方程与反应物浓度联系。对于反应 A + B → 产物，一般形式为：
rate = k [A]ᵐ [B]ⁿ
= rate constant, m, n = orders of reaction with respect to A and B. Overall order = m + n.
k = 速率常数，m, n = 对 A 和 B 的反应级数。总反应级数 = m + n。
Orders can be 0, 1, 2 and must be determined experimentally, not from the stoichiometric equation.
反应级数可为 0, 1, 2，必须通过实验测定，不能从化学计量方程推导。
Units of k depend on overall order:
速率常数 k 的单位取决于总级数：
Zero order: k in mol dm⁻³ s⁻¹
First order: k in s⁻¹
Second order: k in dm³ mol⁻¹ s⁻¹
零级反应：k 单位为 mol dm⁻³ s⁻¹
一级反应：k 单位为 s⁻¹
二级反应：k 单位为 dm³ mol⁻¹ s⁻¹
11. Arrhenius Equation | 阿伦尼乌斯方程
The Arrhenius equation describes how the rate constant k varies with temperature and activation energy.
阿伦尼乌斯方程描述了速率常数 k 随温度和活化能的变化关系。
k = A e–Eₐ / RT
ponential factor, Eₐ = activation energy (J mol⁻¹), R = gas constant, T = temperature (K).
A 为指前因子，Eₐ 为活化能（J mol⁻¹），R 为气体常数，T 为温度（K）。
Taking natural logarithms gives a linear form:
取自然对数后得到线性形式：
ln k = –Eₐ / (RT) + ln A
A graph of ln k against 1/T yields a straight line with gradient = –Eₐ / R and y-intercept = ln A.
绘制 ln k 对 1/T 的图，得到一条直线，斜率为 –Eₐ / R，截距为 ln A。
This allows calculation of Eₐ from experimental data.
由此可从实验数据中计算出活化能 Eₐ。
12. Electrode Potentials and Cell EMF | 电极电势与电池电动势
The standard electrode potential E° measures the tendency of a half-cell to gain electrons under standard conditions. The cell potential (EMF) is determined by connecting two half-cells.
标准电极电势 E° 衡量半电池在标准条件下获得电子的趋势。电池的电动势 (EMF) 由两个半电池连接而成。
E°cell = E°(right-hand electrode) – E°(left-hand electrode)
A positive E°cell indicates a feasible reaction. It is common to write E°cell = E°cathode – E°anode, where the cathode is where reduction occurs.
E°cell 为正表示反应可行。通常写作 E°cell = E°cathode – E°anode，阴极发生还原反应。
The relationship between free energy and cell potential is:
吉布斯自由能与电池电势的关系为：
ΔG° = – n F E°cell
n = number of moles of electrons transferred, F = Faraday constant (96 500 C mol⁻¹).
n = 转移电子的物质的量，F = 法拉第常数（96 500 C mol⁻¹）。
A feasible reaction has negative ΔG°, meaning E°cell must be positive.
可行的反应 ΔG° 为负，意味着 E°cell 必须为正。
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