📚 Year 13 Edexcel Chemistry: Formula & Theorem Quick Reference | 爱德思Year 13化学：公式定理速查手册
A comprehensive quick reference covering the key formulas, equations, and theorems for Year 13 Edexcel Chemistry. This guide consolidates essential numerical relationships from kinetics, equilibrium, thermodynamics, electrochemistry, and inorganic chemistry to support your revision.
本速查手册涵盖爱德思Year 13化学核心公式、方程与定理，汇总了动力学、化学平衡、热力学、电化学和无机化学中的重要定量关系，助力你的复习备考。
1. Rate Laws and Reaction Orders | 速率方程与反应级数
For a reaction aA + bB → products, the rate equation relates the rate to the concentrations of reactants raised to some powers. The exponents are the orders with respect to each reactant, and the overall order is their sum.
对于反应 aA + bB → 产物，速率方程将反应速率与各反应物浓度的幂次方联系起来。各浓度项的指数就是该反应物的反应级数，总反应级数为各指数之和。
rate = k [A]ᵐ [B]ⁿ
where k is the rate constant, m is the order with respect to A, and n is the order with respect to B. The values of m and n must be determined experimentally and are not necessarily equal to the stoichiometric coefficients a and b.
式中 k 为速率常数，m 为对 A 的反应级数，n 为对 B 的反应级数。m 和 n 必须通过实验测定，不一定等于计量系数 a 和 b。
Common orders: zero order (rate independent of [A], m = 0); first order (rate ∝ [A], m = 1); second order (rate ∝ [A]² or rate ∝ [A][B], overall order = 2). The units of k depend on the overall order: for overall order 0, k has units mol dm⁻³ s⁻¹; order 1, s⁻¹; order 2, dm³ mol⁻¹ s⁻¹.
常见级数：零级（速率与[A]无关，m = 0）；一级（速率 ∝ [A]，m = 1）；二级（速率 ∝ [A]² 或速率 ∝ [A][B]，总级数为2）。k 的单位取决于总反应级数：总级数为0时，k 的单位为 mol dm⁻³ s⁻¹；总级数为1时，单位为 s⁻¹；总级数为2时，单位为 dm³ mol⁻¹ s⁻¹。
2. The Arrhenius Equation | 阿伦尼乌斯方程
The Arrhenius equation describes how the rate constant k changes with temperature. It shows that a higher temperature or a lower activation energy increases the rate constant exponentially.
阿伦尼乌斯方程描述了速率常数 k 如何随温度变化。它表明，更高的温度或更低的活化能会使速率常数呈指数级增大。
k = A exp(−Eₐ/(RT))
where A is the pre-exponential factor, Eₐ is the activation energy (J mol⁻¹), R is the gas constant (8.31 J K⁻¹ mol⁻¹), and T is the absolute temperature (K). Taking natural logarithms gives a linear form:
式中 A 为指前因子，Eₐ 为活化能（J mol⁻¹），R 为气体常数（8.31 J K⁻¹ mol⁻¹），T 为绝对温度（K）。取自然对数后得到线性形式：
ln k = ln A − Eₐ/(RT)
A plot of ln k against 1/T yields a straight line with slope = −Eₐ/R and intercept = ln A. This is used to calculate Eₐ from experimental data. Remember that the activation energy Eₐ is always positive.
以 ln k 对 1/T 作图得到一条直线，其斜率 = −Eₐ/R，截距 = ln A。此法常用于由实验数据计算活化能。注意活化能 Eₐ 恒为正值。
3. Equilibrium Constants: Kc and Kp | 平衡常数：Kc 与 Kp
The equilibrium constant Kc expresses the ratio of product to reactant concentrations at equilibrium, each raised to the power of its stoichiometric coefficient. It is influenced only by temperature.
平衡常数 Kc 表示在平衡状态下生成物浓度乘积与反应物浓度乘积之比，各物质浓度以其计量系数为指数。Kc 只受温度影响。
For a general reaction aA + bB ⇌ cC + dD:
对于一般反应 aA + bB ⇌ cC + dD：
Kc = [C]ᶜ [D]ᵈ / ([A]ᵃ [B]ᵇ)
For gaseous equilibria, Kp is used, with partial pressures in place of concentrations. The relationship between Kp and Kc is Kp = Kc (RT)^Δn, where Δn = (c+d) − (a+b) for gases only, R = 8.31 J K⁻¹ mol⁻¹ but in pressure terms often use 0.0831 L bar K⁻¹ mol⁻¹ as appropriate.
对于气相平衡，使用 Kp，以分压代替浓度。Kp 与 Kc 的关系为 Kp = Kc (RT)^Δn，其中 Δn = (c+d) − (a+b) 仅针对气体物质，R = 8.31 J K⁻¹ mol⁻¹，但在压强计算中常使用 0.0831 L bar K⁻¹ mol⁻¹ 等适配单位。
A large Kc (≫1) indicates the equilibrium lies far to the right; a small Kc (≪1) indicates it lies to the left. Kc has units that depend on the stoichiometry, but in Edexcel units are often omitted when quoting a value.
Kc 远大于1表明平衡强烈偏向产物一侧；Kc 远小于1表明平衡偏向反应物一侧。Kc 的单位取决于计量关系，但在爱德思考试中引用数值时往往省略单位。
4. Acid Dissociation Constant and pH | 酸解离常数与pH
The strength of a weak acid is measured by its acid dissociation constant Ka. For a weak acid HA dissociating as HA ⇌ H⁺ + A⁻, the equilibrium expression is:
弱酸的强度用酸解离常数 Ka 衡量。对于弱酸 HA 的解离 HA ⇌ H⁺ + A⁻，平衡表达式为：
Ka = [H⁺][A⁻] / [HA]
The logarithmic form gives pKa = −log₁₀Ka. The higher the Ka, the stronger the acid and the smaller the pKa. For water, the ionic product Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K, and pKw = 14.
对数形式为 pKa = −log₁₀Ka。Ka 越大，酸性越强，pKa 越小。水的离子积 Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶（298 K），pKw = 14。
pH is defined as pH = −log₁₀[H⁺]. For a strong monoprotic acid, [H⁺] equals the acid concentration. For a weak acid, [H⁺] is found by solving Ka = [H⁺]² / (c − [H⁺]) or, when dissociation is very small, by [H⁺] ≈ √(Ka c).
pH 定义为 pH = −log₁₀[H⁺]。对于强的一元酸，[H⁺] 等于酸的浓度。对于弱酸，需通过解 Ka = [H⁺]² / (c − [H⁺]) 求出 [H⁺]；若解离度极小，可用 [H⁺] ≈ √(Ka c) 近似。
5. Buffer Solutions: The Henderson-Hasselbalch Equation | 缓冲溶液：亨德森-哈塞尔巴尔赫方程
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 (or a weak base and its conjugate acid). The pH of an acidic buffer can be calculated using the Henderson-Hasselbalch equation.
缓冲溶液能抵抗少量酸或碱加入时引起的 pH 变化。它由弱酸及其共轭碱（或弱碱及其共轭酸）组成。酸性缓冲溶液的 pH 可用亨德森-哈塞尔巴尔赫方程计算。
pH = pKa + log₁₀([A⁻] / [HA])
where [A⁻] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. This equation assumes that the concentrations of A⁻ and HA at equilibrium are essentially equal to their initial concentrations, which holds when the acid is weak and the concentrations are reasonably high.
式中 [A⁻] 为共轭碱浓度，[HA] 为弱酸浓度。该方程假设 A⁻ 和 HA 在平衡时的浓度基本上等于其初始浓度，这在酸较弱且浓度较高时成立。
When preparing a buffer, choose a weak acid with a pKa close to the desired pH. The buffer capacity is greatest when [A⁻] = [HA], i.e., pH = pKa. Dilution does not change the pH of a buffer appreciably because the ratio [A⁻]/[HA] remains constant.
制备缓冲溶液时，应选择 pKa 接近目标 pH 的弱酸。当 [A⁻] = [HA] 时，即 pH = pKa，缓冲容量最大。稀释不会显著改变缓冲液的 pH，因为 [A⁻]/[HA] 比值保持不变。
6. Gibbs Free Energy and Spontaneity | 吉布斯自由能与自发性
The Gibbs free energy change ΔG determines whether a reaction is thermodynamically feasible at constant temperature and pressure. The fundamental relationship combines enthalpy and entropy changes.
吉布斯自由能变 ΔG 用于判断在恒温恒压下反应是否热力学可行。其基本关系式综合了焓变和熵变。
ΔG = ΔH − TΔS
where ΔH is the enthalpy change, T is the absolute temperature in kelvin, and ΔS is the entropy change. A reaction is spontaneous (feasible) when ΔG < 0. If ΔG > 0, the reaction is not feasible under those conditions. When ΔG = 0, the system is at equilibrium.
式中 ΔH 为焓变，T 为开尔文绝对温度，ΔS 为熵变。当 ΔG < 0 时，反应自发（可行）；若 ΔG > 0，则在此条件下反应不可行；ΔG = 0 时系统处于平衡状态。
Note that a negative ΔH (exothermic) favours a negative ΔG, while a positive ΔS favours a negative ΔG at high temperatures. The interplay explains why some endothermic reactions occur spontaneously at high T.
注意，ΔH 为负（放热）有利于 ΔG 为负，而 ΔS 为正时高温有利于 ΔG 为负。二者协同作用解释了为何某些吸热反应在高温下能自发进行。
7. Gibbs Free Energy and Equilibrium Constant | 吉布斯自由能与平衡常数
The standard Gibbs free energy change ΔG° is directly related to the equilibrium constant K. This provides a link between
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