📚 IB WJEC Chemistry: Last-Minute Revision Notes | IB WJEC 化学：考前冲刺笔记
Whether you are sitting the IB Diploma Chemistry exam or the WJEC A‑level Chemistry paper, the final weeks before the test demand focused, efficient revision. These notes distil the core topics – from mole calculations to organic reaction pathways – into concise, bilingual bullet points that have been aligned with both syllabuses. Use them to plug gaps, reinforce conceptual understanding, and build confidence for the big day.
无论你是参加 IB 文凭化学考试还是 WJEC A‑level 化学考试，考前几周都需要高效、聚焦的复习。这份笔记将摩尔计算、有机反应路径等核心主题提炼成简洁的双语要点，全部与两大课程大纲对齐。用它来查漏补缺、强化概念理解，为考试日积蓄信心。
1. Stoichiometry and Mole Calculations | 化学计量与摩尔计算
The mole (mol) is the SI unit for amount of substance. One mole contains exactly 6.022 × 10²³ elementary entities (Avogadro’s constant, NA). The molar mass (M) is the mass per mole of a substance, usually expressed in g mol−¹. The number of moles n is found by dividing the mass (m) by the molar mass: n = m/M. For gases at STP (273 K, 100 kPa), one mole occupies 22.7 dm³. For solutions, n = concentration (c) × volume (V).
摩尔（mol）是物质的量的 SI 单位。1 mol 恰好包含 6.022×10²³ 个基本单元（阿伏伽德罗常数 NA）。摩尔质量（M）是每摩尔物质的质量，通常用 g mol−¹ 表示。物质的量 n 可由质量 m 除以摩尔质量求得：n = m/M。在标准状况（STP, 273 K、100 kPa）下，1 mol 气体的体积为 22.7 dm³。对于溶液，n = 浓度 c × 体积 V。
n = m / M n = cV n = V(gas) / 22.7
In empirical formula calculations, convert percentage composition to mass, then to moles, and divide by the smallest mole value to obtain the simplest whole‑number ratio. For titration problems, use n = cV to find the unknown concentration; remember to account for the stoichiometric ratio from the balanced equation. Always identify the limiting reagent by comparing the initial moles to the stoichiometric requirement – it decides the theoretical yield.
计算实验式时，先将质量分数转化为质量，再转化为摩尔数，除以最小的摩尔值得到最简整数比。在滴定问题中，使用 n = cV 求未知浓度，务必根据配平的方程式引入化学计量比。总是通过对比起始摩尔数和化学计量要求的摩尔数来确定限制试剂——它决定了理论产率。
2. Atomic Structure | 原子结构
Atoms consist of a dense nucleus containing protons (Z) and neutrons (N), surrounded by electrons arranged in quantised energy levels. The atomic number Z defines the element; mass number A = Z + N. Isotopes have the same Z but different A. Electron configuration is written using s, p, d, f orbitals: e.g., Fe (Z = 26) is 1s²2s²2p&sup6;3s²3p&sup6;4s²3d&sup6; or [Ar]4s²3d&sup6;. Note the 4s orbital fills before 3d, but loses electrons first when forming ions (4s→3d).
原子由一个包含质子（Z）和中子（N）的致密原子核以及按量子化能级排布的电子构成。原子序数 Z 决定元素种类，质量数 A = Z + N。同位素的 Z 相同而 A 不同。电子排布用 s、p、d、f 轨道表示：如 Fe（Z=26）为 1s²2s²2p&sup6;3s²3p&sup6;4s²3d&sup6; 或 [Ar]4s²3d&sup6;。注意 4s 轨道比 3d 先填充，但在形成离子时会先失去 4s 电子（4s→3d）。
The electromagnetic spectrum links electron transitions to emission/absorption lines. For hydrogen, the energy of the nth level is proportional to −1/n²; the convergence limit in the Lyman series gives the ionisation energy. Successive ionisation energies provide evidence for electron shells and sub‑shells: large jumps occur when an electron is removed from a new, closer‑to‑the‑nucleus shell.
电磁波谱将电子跃迁与发射/吸收谱线联系起来。对氢原子，第 n 能级的能量与 −1/n² 成正比；莱曼系的收敛极限对应电离能。逐级电离能数据为电子层和子壳层结构提供了证据：当从更靠近原子核的内层移走电子时，会出现电离能的巨大跳跃。
3. Bonding and Intermolecular Forces | 化学键与分子间作用力
Three principal types of strong bonding are ionic, covalent, and metallic. Ionic bonding arises from electrostatic attraction between oppositely charged ions, forming giant lattices with high melting points. Covalent bonding involves the sharing of electron pairs; simple molecular substances (CO&sub2;, H&sub2;O) have low melting points, while giant covalent networks (diamond, SiO&sub2;) are extremely hard and high‑melting. Metallic bonding is a lattice of cations immersed in a sea of delocalised electrons, explaining malleability and electrical conductivity.
三类主要的强相互作用是离子键、共价键和金属键。离子键源于正负离子之间的静电吸引，形成高熔点的巨型晶格。共价键通过共享电子对实现；简单分子（CO&sub2;、H&sub2;O）熔点低，而巨型共价网络（金刚石、SiO&sub2;）硬度极高、熔点极高。金属键是阳离子点阵沉浸在离域电子“海洋”中，这解释了金属的延展性和导电性。
Molecular shape is determined by VSEPR theory: electron pairs around a central atom repel each other and adopt the arrangement that minimises repulsion. For example, CH&sub4; is tetrahedral (bond angle 109.5°), NH&sub3; is trigonal pyramidal (107°), and H&sub2;O is bent (104.5°). The presence of lone pairs distorts bond angles because lone‑pair–bond‑pair repulsion is stronger than bond‑pair–bond‑pair repulsion.
分子形状由 VSEPR 理论决定：中心原子周围的电子对互相排斥，采取使排斥力最小的排列方式。例如 CH&sub4; 是四面体形（键角 109.5°），NH&sub3; 是三角锥形（107°），H&sub2;O 是 V 形（104.5°）。孤对电子的存在会压缩键角，因为孤对-键对排斥力大于键对-键对排斥力。
Polarity arises when bonds are polar (due to electronegativity differences) and the molecular geometry does not cancel the dipole vectors. Intermolecular forces (IMFs) dictate physical properties: London forces (all molecules), dipole‑dipole interactions (polar molecules), and hydrogen bonding (molecules with H attached to N, O, or F). Hydrogen bonding significantly raises boiling points, as seen in H&sub2;O and NH&sub3;.
当键具有极性（由电负性差异所致）且分子几何结构无法使偶极矢量抵消时，分子便呈现极性。分子间作用力（IMFs）决定物理性质，包括伦敦力（所有分子）、偶极‑偶极作用（极性分子）和氢键（H 与 N、O、F 相连的分子）。氢键显著提高沸点，如水（H&sub2;O）和氨（NH&sub3;）所示。
4. Energetics and Thermochemistry | 能量学与热化学
The standard enthalpy change ΔH° is measured under standard conditions (100 kPa, 298 K). Exothermic reactions release heat (ΔH negative); endothermic reactions absorb heat (ΔH positive). Hess’s Law states that the enthalpy change for a reaction is independent of the pathway. It is applied by combining known enthalpy changes of formation or combustion to find an unknown ΔH.
标准焓变 ΔH° 是在标准条件（100 kPa，298 K）下测定的。放热反应释放热量（ΔH 为负值），吸热反应吸收热量（ΔH 为正值）。赫斯定律表明，反应的焓变与途径无关。通过组合已知的生成焓或燃烧焓数据，可以求得未知的 ΔH。
Average bond enthalpies allow estimation of ΔH by considering bonds broken (endothermic) minus bonds formed (exothermic). Born‑Haber cycles (IB HL / WJEC) link lattice enthalpy, ionisation energies, electron affinity, and atomisation enthalpy through a closed energy loop. The Gibbs free energy change, ΔG = ΔH − TΔS, determines reaction feasibility: a reaction is spontaneous at constant temperature and pressure when ΔG < 0. Entropy (S) measures the dispersal of energy; ΔS° can be calculated from standard molar entropy values.
平均键焓可用于估算 ΔH，即断裂键吸收的能量减去形成键释放的能量。波恩‑哈伯循环（IB HL / WJEC）将晶格焓、电离能、电子亲和能和原子化焓通过闭合能量循环联系起来。吉布斯自由能变 ΔG = ΔH − TΔS 决定反应的可行性：恒温恒压下，当 ΔG < 0 时反应可自发进行。熵（S）衡量能量的分散程度；ΔS° 可从标准摩尔熵值计算得出。
5. Chemical Kinetics | 化学动力学
The rate of a chemical reaction is defined as the change in concentration of a reactant or product per unit time: rate = −Δ[R]/Δt = Δ[P]/Δt. For the general rate equation rate = k[A]¹[B]², the reaction is first order in A, second order in B, and third order overall. The rate constant k is temperature‑dependent and has units that vary with the overall order.
化学反应速率定义为单位时间内反应物或产物浓度的变化量：rate = −Δ[R]/Δt = Δ[P]/Δt。对于通式 rate = k[A]¹[B]²，反应对 A 为一级，对 B 为二级，总级数为三级。速率常数 k 随温度变化，其单位取决于总级数。
Collision theory requires particles to collide with sufficient energy (E ≥ Ea) and correct orientation for a reaction to occur. The temperature dependence of k is described by the Arrhenius equation: k = A exp(−Ea/RT), where A is the pre‑exponential factor. A Maxwell‑Boltzmann distribution plot shows that increasing temperature broadens the distribution and greatly increases the fraction of particles with energy exceeding Ea. Catalysts provide an alternative pathway with a lower activation energy, without being consumed.
碰撞理论要求粒子碰撞时具有足够能量（E ≥ Ea）和正确的取向，反应才能发生。k 对温度的依赖性由阿伦尼乌斯方程描述：k = A exp(−Ea/RT)，其中 A 为指前因子。麦克斯韦‑玻尔兹曼分布图表明，升高温度使分布曲线变宽，并显著增加能量超过 Ea 的粒子比例。催化剂通过提供活化能更低的替代路径而起作用，本身不被消耗。
6. Chemical Equilibrium | 化学平衡
Dynamic equilibrium is established when the rates of the forward and reverse reactions are equal, so the concentrations of all species remain constant. The equilibrium constant Kc (for solutions) is expressed as Kc = [C]c[D]d / [A]a[B]b for the reaction aA + bB ⇄ cC + dD. Solids and pure liquids are omitted from the expression. Kc is temperature‑dependent; a change in temperature shifts the equilibrium position and alters the value of Kc in accordance with Le Chatelier’s principle.
当正、逆反应速率相等时，体系达到动态平衡，各物种的浓度保持恒定。对于反应 aA + bB ⇄ cC + dD，平衡常数 Kc（用于溶液）表示为 Kc = [C]¹[D]² / [A]¹[B]²（按计量数），固体和纯液体不写入表达式。Kc 依赖于温度；根据勒夏特列原理，温度改变会使平衡位置移动并改变 Kc 的数值。
Le Chatelier’s principle predicts how a system at equilibrium responds to disturbances: adding a reactant shifts equilibrium to the right; increasing pressure shifts equilibrium toward the side with fewer gas moles; increasing temperature favours the endothermic direction. For gaseous equilibria, Kp is used, where partial pressures replace concentrations. The reaction quotient Q provides a snapshot of the system; if Q < K, the forward reaction is favoured, and if Q > K, the reverse reaction is favoured.
勒夏特列原理可预测平衡体系对外界干扰的响应：增加反应物使平衡右移，增大压强使平衡向气体分子数减少的方向移动，升高温度有利于吸热方向。对于气相平衡，则使用 Kp，以分压代替浓度。反应商 Q 能反映某一时刻体系的状况；若 Q < K，则正向反应占优势；若 Q > K，则逆向反应占优势。
7. Acids and Bases | 酸与碱
A Brønsted–Lowry acid is a proton donor; a base is a proton acceptor. Strong acids (HCl, HNO&sub3;, H&sub2;SO&sub4;) dissociate completely in water, while weak acids (CH&sub3;COOH, HF) dissociate only partially, establishing an equilibrium. For a weak acid HA, the acid dissociation constant Ka = [H+][A−]/[HA]. The pH is defined as pH = −log[H+]; for strong monoprotic acids, [H+] equals the acid concentration. For a weak acid, [H+] ≈ √(Ka × [HA]).
布朗斯特‑劳里酸是质子给予体，碱是质子接受体。强酸（HCl、HNO&sub3;、H&sub2;SO&sub4;）在水中完全解离，而弱酸（CH&sub3;COOH、HF）仅部分解离，建立平衡。对弱酸 HA，酸解离常数 Ka = [H+][A−]/[HA]。pH 定义为 pH = −log[H+]；对于一元强酸，[H+] 等于酸浓度。对于弱酸，[H+] ≈ √(Ka × [HA])。
The ionic product of water Kw = [H+][OH−] = 1.0×10−¹&sup4; at 298 K, giving pKw = 14. Strong bases (NaOH, KOH) dissociate fully to give [OH−]; pOH = −log[OH−], and pH = 14 − pOH. Buffer solutions resist changes in pH on adding small amounts of acid or base. They consist of a weak acid and its conjugate base (e.g., CH&sub3;COOH/CH&sub3;COO−). The Henderson–Hasselbalch equation (pH = pKa + log([A−]/[HA])) is used for buffer calculations. Titration curves show the pH change during neutralisation; the equivalence point must be matched with a suitable indicator whose pKin lies within the steep pH jump.
水的离子积 Kw = [H+][OH−] = 1.0×10−¹&sup4;（298 K），故 pKw = 14。强碱（NaOH、KOH）完全解离产生 OH−；pOH = −log[OH−]，且 pH = 14 − pOH。缓冲溶液能在加入少量酸或碱时抵抗 pH 变化，它由弱酸及其共轭碱（如 CH&sub3;COOH/CH&sub3;COO−）构成。亨德森‑哈塞尔巴尔赫公式（pH = pKa + log([A−]/[HA])）用于缓冲溶液计算。滴定曲线展示中和过程中的 pH 变化；等当点需与合适的指示剂匹配，指示剂的 pKin 应落在突跃范围之内。
8. Redox Processes and Electrochemistry | 氧化还原与电化学
Oxidation involves an increase in oxidation number (electron loss); reduction involves a decrease in oxidation number (electron gain). Oxidation numbers are assigned using rules: free elements = 0, monatomic ions = charge, oxygen usually −2 (except in peroxides), hydrogen usually +1 (except in hydrides). Redox equations are balanced by separating into half‑equations and combining them, ensuring electrons lost equal electrons gained.
氧化伴随氧化数的升高（失电子），还原伴随氧化数的降低（得电子）。氧化数的确定规则如下：游离态单质为 0，单原子离子等于其所带电荷，氧通常为 −2（过氧化物除外），氢通常为 +1（金属氢化物除外）。配平氧化还原反应时，先拆分为半反应式，再加以合并，确保失电子总数等于得电子总数。
An electrochemical cell converts chemical energy into electrical energy. The cell potential E°cell = E°cathode − E°anode, using standard electrode potentials measured against the standard hydrogen electrode (SHE). A positive E°cell indicates a spontaneous reaction. In electrolytic cells, an external power source drives non‑spontaneous reactions; the strongest oxidising
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