IB Edexcel Chemistry: Chemical Reactions Key Points | IB Edexcel 化学：化学反应考点精讲

📚 IB Edexcel Chemistry: Chemical Reactions Key Points | IB Edexcel 化学：化学反应考点精讲

Chemical reactions lie at the heart of all chemical science, from the rusting of iron to the metabolic pathways in living cells. This article consolidates the essential principles, definitions, and calculation methods required for IB and Edexcel Chemistry examinations, covering reaction kinetics, energetics, equilibrium, redox, acid-base behaviour, and organic reaction types. Each section interweaves conceptual explanation with worked examples to ensure you can apply your knowledge confidently in both structured and multiple-choice questions.

化学反应是化学科学的核心，从铁的生锈到活细胞中的代谢途径无一不涉及。本文整合了 IB 与 Edexcel 化学考试所要求的基本原理、定义和计算方法，涵盖反应动力学、能量学、平衡、氧化还原、酸碱行为以及有机反应类型。每个部分都将概念讲解与实例交织在一起，确保你能够在结构化试题和选择题中自如运用所学知识。

1. Reaction Rates and Collision Theory | 反应速率与碰撞理论

The rate of a chemical reaction is defined as the change in concentration of a reactant or product per unit time. For a reaction A → B, rate = –Δ[A]/Δt = Δ[B]/Δt. Collision theory states that for a reaction to occur, particles must collide with sufficient kinetic energy (equal to or greater than the activation energy, Eₐ) and with the correct orientation.

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。对于反应 A → B，速率 = –Δ[A]/Δt = Δ[B]/Δt。碰撞理论认为，反应发生的条件是粒子必须发生碰撞，且碰撞动能足够（等于或大于活化能 Eₐ）并具有正确的取向。

The Maxwell–Boltzmann distribution shows the spread of kinetic energies among particles at a given temperature. Only the fraction of particles with energy ≥ Eₐ can lead to successful collisions. This fraction is represented by the area under the curve to the right of Eₐ. Increasing temperature shifts the distribution to higher energies and flattens the curve, greatly enlarging the proportion of particles that exceed Eₐ.

麦克斯韦-玻尔兹曼分布展示了给定温度下粒子动能的范围。只有能量 ≥ Eₐ 的那部分粒子才能导致有效碰撞，这一比例由曲线下 Eₐ 右侧的面积表示。升高温度使分布向高能方向移动并使曲线变得平缓，从而显著增大超过 Eₐ 的粒子比例。

• Rate ∝ collision frequency × fraction of effective collisions.
• 速率正比于碰撞频率 × 有效碰撞分数。

2. Factors Affecting Reaction Rate | 影响反应速率的因素

Concentration (or pressure for gases): Increasing concentration increases the number of particles per unit volume, raising collision frequency. For a simple reaction A + B → products, rate = k[A]ᵐ[B]ⁿ, where m and n are the orders of reaction determined experimentally, not from stoichiometry.

浓度（或气体压强）：增大浓度会提高单位体积内的粒子数，从而增加碰撞频率。对于简单反应 A + B → 产物，速率 = k[A]ᵐ[B]ⁿ，其中 m 和 n 是反应的级数，由实验确定而非由化学计量数决定。

Temperature: As predicted by the Arrhenius equation, k = Ae^(–Eₐ/RT), a small rise in temperature produces a disproportionately large increase in rate because the exponential term is highly sensitive to T. A 10°C rise often doubles the rate near room temperature.

温度：根据阿伦尼乌斯方程 k = Ae^(–Eₐ/RT) 的预测，温度小幅升高会使速率不成比例地大增，因为指数项对 T 非常敏感。室温附近温度每升高 10°C 速率常常加倍。

Surface area: For solid reactants, grinding increases surface area, exposing more particles to collisions and thus accelerating heterogeneous reactions.

表面积：对于固体反应物，研磨可增大表面积，使更多粒子暴露于碰撞之下，从而加快非均相反应。

Catalysts: A catalyst provides an alternative reaction pathway with a lower activation energy, increasing the fraction of successful collisions without being consumed. Enzymes are biological catalysts with extraordinary specificity.

催化剂：催化剂提供一条活化能较低的替代反应路径，在不被消耗的情况下提高有效碰撞分数。酶是具有极高专一性的生物催化剂。

3. Chemical Equilibrium and Le Chatelier’s Principle | 化学平衡与勒夏特列原理

A dynamic equilibrium is established in a closed system when the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. The equilibrium constant Kc or Kp has a fixed value at a given temperature. For aA + bB ⇌ cC + dD, Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ (units omitted in IB expressions).

当正逆反应速率相等且反应物和产物浓度保持恒定时，在封闭体系中建立起动态平衡。平衡常数 Kc 或 Kp 在给定温度下有固定值。对于 aA + bB ⇌ cC + dD，Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ（IB 表达式省略单位）。

Le Chatelier’s principle states that if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the equilibrium position shifts to counteract the imposed change. Adding a reactant shifts the equilibrium to the right; increasing temperature favours the endothermic direction; increasing pressure (by volume reduction) shifts the equilibrium towards the side with fewer gas molecules.

勒夏特列原理指出，若平衡体系受到浓度、温度或压强的改变，平衡位置将移动以抵消外加的变化。增加反应物使平衡向右移动；升高温度有利于吸热方向；增大压强（通过缩小体积）使平衡向气体分子数较少的一侧移动。

Importantly, catalysts do not alter the equilibrium position or the value of K; they merely speed up the attainment of equilibrium.

重要的是，催化剂不会改变平衡位置或 K 值，它们只能加快达到平衡的速度。

4. Enthalpy Changes (ΔH) and Energy Profiles | 焓变与能量曲线

Enthalpy change ΔH is the heat transferred under constant pressure. Exothermic reactions release energy to the surroundings (ΔH < 0), while endothermic reactions absorb energy (ΔH > 0). Standard enthalpy changes refer to measurements at 298 K and 100 kPa.

焓变 ΔH 是恒压条件下传递的热量。放热反应向环境释放能量（ΔH < 0），吸热反应则吸收能量（ΔH > 0）。标准焓变指的是在 298 K 和 100 kPa 下的测量值。

An energy profile diagram plots enthalpy against the reaction coordinate. The peak corresponds to the transition state, and the energy gap between reactants and the peak is Eₐ. For an exothermic reaction, the products sit at a lower enthalpy than the reactants; the enthalpy change is negative.

能量曲线图以焓对反应进程作图。最高点对应过渡态，反应物与最高点之间的能量差为 Eₐ。对于放热反应，产物焓值低于反应物，焓变为负值。

Standard enthalpy of combustion (ΔH°c) is the enthalpy change when one mole of a substance is completely burned in oxygen. Standard enthalpy of formation (ΔH°f) is the enthalpy change when one mole of a compound is formed from its elements in their standard states. By definition, ΔH°f of an element is zero.

标准燃烧焓（ΔH°c）是指 1 mol 物质在氧气中完全燃烧时的焓变。标准生成焓（ΔH°f）是指由标准状态下的元素生成 1 mol 化合物时的焓变。根据定义，元素的 ΔH°f 为零。

5. Hess’s Law and Bond Enthalpies | 赫斯定律与键能

Hess’s Law states that the total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same. It allows the calculation of ΔH for reactions that cannot be measured directly by combining known enthalpy changes of related reactions.

赫斯定律指出，只要始态和终态相同，反应的总焓变与所经途径无关。这使我们能够通过组合已知相关反应的焓变来计算难以直接测量的反应的 ΔH。

ΔH° reaction = ΣΔH°f(products) – ΣΔH°f(reactants)

Bond enthalpy (average) is the energy required to break one mole of a specific bond in the gaseous state, averaged over a range of compounds. The enthalpy change of a reaction can be approximated using bond enthalpies:

键能（平均键能）是指气态下断裂 1 mol 特定化学键所需的能量，是多种化合物中的平均值。可利用键能近似计算反应焓变：

ΔH ≈ Σ (bond enthalpies broken) – Σ (bond enthalpies formed)

This method is less accurate than using formation enthalpies because average bond enthalpies ignore variations due to molecular environment. Remember: breaking bonds requires energy (+), making bonds releases energy (–).

这种方法不如使用生成焓准确，因为平均键能忽略了分子环境引起的差异。记住：断键吸热（+），成键放热（–）。

6. Redox Reactions and Oxidation Numbers | 氧化还原反应与氧化数

Oxidation is the loss of electrons, and reduction is the gain of electrons (OIL RIG). A redox reaction involves both processes simultaneously. The oxidation number (or oxidation state) is a book-keeping tool to track electron transfer. Key rules: free elements have oxidation number 0; H is +1 (except in metal hydrides where it is –1); O is –2 (except in peroxides where it is –1, and in OF₂ where it is +2); the sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, the sum equals the ion’s charge.

氧化是失去电子，还原是得到电子（OIL RIG）。氧化还原反应同时包含这两个过程。氧化数（或氧化态）是追踪电子转移的记帐工具。关键规则：游离态元素氧化数为 0；H 为 +1（金属氢化物中为 –1 除外）；O 为 –2（过氧化物中为 –1，OF₂ 中为 +2 除外）；中性化合物中各元素氧化数之和为 0；多原子离子中氧化数之和等于离子电荷。

A disproportionation reaction is a special redox process in which the same element is simultaneously oxidised and reduced, such as the decomposition of hydrogen peroxide: 2H₂O₂ → 2H₂O + O₂, where oxygen goes from –1 to both –2 (reduction) and 0 (oxidation).

歧化反应是一种特殊的氧化还原过程，其中同一元素同时被氧化和被还原，例如过氧化氢的分解：2H₂O₂ → 2H₂O + O₂，氧的氧化数从 –1 变为 –2（还原）和 0（氧化）。

Redox titrations, like those involving manganate(VII) and iron(II), require balancing half-equations in acidic medium. The end point is detected by a colour change without an external indicator.

氧化还原滴定，如涉及高锰酸根(VII)和铁(II)的滴定，需要在酸性介质中书写并配平半反应式。终点可通过颜色变化指示，无需外加指示剂。

7. Electrochemical Cells | 电化学电池

A voltaic (galvanic) cell converts chemical energy into electrical energy through spontaneous redox reactions. The cell consists of two half-cells connected by a salt bridge. The half-cell with the more negative standard electrode potential (E°) is the anode, where oxidation occurs; the half-cell with the more positive E° is the cathode, where reduction occurs. The standard cell potential E°cell = E°cathode – E°anode must be positive for a spontaneous reaction.

伏打（原）电池通过自发的氧化还原反应将化学能转化为电能。电池由两个通过盐桥连接的半电池组成。标准电极电势（E°）较负的半电池为阳极，发生氧化；E° 较正的半电池为阴极，发生还原。标准电池电动势 E°cell = E°阴极 – E°阳极，自发反应时该值必须为正。

Electrolytic cells, by contrast, use an external power source to drive non-spontaneous reactions. Here the anode is positive and the cathode is negative. Faraday’s laws relate the quantity of electricity passed to the amount of substance deposited or dissolved.

与之相反，电解池使用外部电源驱动非自发反应。此时阳极接正极，阴极接负极。法拉第定律关联了通过的电量与沉积或溶解物质的量。

The reactivity series and the ease of discharge can predict products in electrolysis of aqueous solutions. For example, in the electrolysis of concentrated NaCl(aq), Cl₂ is produced at the anode and H₂ at the cathode because Cl⁻ is oxidised more readily than OH⁻, and H₂O is reduced more readily than Na⁺.

利用金属活动性顺序和离子放电难易度可以预测水溶液电解的产物。例如，电解浓 NaCl(aq) 时，阳极产生 Cl₂，阴极产生 H₂，因为 Cl⁻ 比 OH⁻ 更容易被氧化，而 H₂O 比 Na⁺ 更容易被还原。

8. Acid-Base Reactions (Brønsted-Lowry) | 酸碱反应（布朗斯特-劳里理论）

According to the Brønsted-Lowry theory, an acid is a proton (H⁺) donor and a base is a proton acceptor. A conjugate acid-base pair differs by one proton. Strong acids and bases fully dissociate in water, whereas weak acids and bases exist in equilibrium with only partial dissociation. The acid dissociation constant Ka = [H⁺][A⁻]/[HA] quantifies the strength of a weak acid; pKa = –log₁₀Ka.

根据布朗斯特-劳里理论，酸是质子（H⁺）给予体，碱是质子接受体。共轭酸碱对之间相差一个质子。强酸和强碱在水中完全解离，而弱酸和弱碱仅部分解离，存在平衡。酸解离常数 Ka = [H⁺][A⁻]/[HA] 可量化弱酸的强度；pKa = –log₁₀Ka。

A neutralisation reaction involves the transfer of a proton from an acid to a base, forming a salt and water. The pH curve for a strong acid-strong base titration shows a steep vertical section around pH 7. For a weak acid with a strong base, the equivalence point lies above pH 7 due to the hydrolysing anion.

中和反应涉及质子从酸转移到碱，生成盐和水。强酸-强碱滴定的 pH 曲线在 pH 7 附近出现陡峭的垂直区段。弱酸与强碱滴定时，由于水解阴离子的存在，等当点位于 pH > 7 区域。

Buffer solutions resist changes in pH upon addition of small amounts of acid or base. They consist of a weak acid and its conjugate base (e.g., CH₃COOH/CH₃COO⁻) or a weak base and its conjugate acid. The Henderson–Hasselbalch equation, pH = pKa + log₁₀([A⁻]/[HA]), is used for buffer calculations.

缓冲溶液能在加入少量酸或碱时抵抗 pH 变化。它们由弱酸及其共轭碱（如 CH₃COOH/CH₃COO⁻）或弱碱及其共轭酸组成。亨德森-哈塞尔巴尔赫方程 pH = pKa + log₁₀([A⁻]/[HA]) 可用于缓冲液的计算。

9. Organic Reactions: Substitution, Addition, Elimination | 有机反应：取代、加成、消除

Organic chemistry is dominated by a few reaction mechanisms that will be tested repeatedly. Free-radical substitution occurs with alkanes and halogens in the presence of UV light, proceeding through initiation, propagation, and termination steps. Electrophilic addition occurs with alkenes: the C=C double bond attacks an electrophile such as HBr or Br₂, forming a carbocation intermediate that quickly combines with a nucleophile.

有机化学以少数几种反应机理为主导，这些机理会在考试中反复出现。自由基取代反应发生在烷烃与卤素之间，需紫外光引发，经历链引发、链增长和链终止步骤。亲电加成反应发生在烯烃中：C=C 双键进攻亲电试剂如 HBr 或 Br₂，形成碳正离子中间体，后者迅速与亲核试剂结合。

Nucleophilic substitution can follow either an SN1 (two-step, via planar carbocation, common for tertiary haloalkanes) or an SN2 (one-step, backside attack, favoured for primary haloalkanes) pathway. The rate equations reflect these mechanisms: SN1 rate = k[haloalkane]; SN2 rate = k[haloalkane][nucleophile].

亲核取代可遵循 SN1（两步，经平面碳正离子，常见于叔卤代烷）或 SN2（一步，背面进攻，伯卤代烷有利）路径。速率方程反映了这些机理：SN1 速率 = k[卤代烷]；SN2 速率 = k[卤代烷][亲核试剂]。

Elimination reactions produce alkenes from haloalkanes or alcohols. In the E2 mechanism, a strong base abstracts a proton while the leaving group departs, requiring antiperiplanar geometry. Saytzeff’s rule predicts the more substituted alkene is the major product.

消除反应由卤代烷或醇生成烯烃。在 E2 机理中，强碱夺取质子的同时离去基团离去，需要反式共平面构型。扎伊采夫规则预测取代基更多的烯烃为主要产物。

10. Stoichiometry and Limiting Reactants | 化学计量学与限量试剂

Stoichiometry is the quantitative relationship between reactants and products in a balanced equation. The mole ratio determines the theoretical yield. The limiting reactant is the substance that is completely consumed first, thus determining the maximum amount of product formed. The other reactants are in excess.

化学计量学是平衡方程中反应物与产物间的定量关系。物质的量之比决定了理论产率。限量试剂是首先被完全消耗的物质，从而决定了产物的最大生成量。其他反应物则过量。

Percentage yield = (actual yield / theoretical yield) × 100. Reasons for yields below 100% include incomplete reactions, side reactions, and losses during purification. Atom economy = (molar mass of desired product / sum of molar masses of all reactants) × 100, which is a key green chemistry metric.

产率百分数 = (实际产量 / 理论产量) × 100。产率低于 100% 的原因包括反应不完全、副反应以及纯化过程中的损失。原子经济性 = (目标产物的摩尔质量 / 所有反应物摩尔质量之和) × 100，这是一项重要的绿色化学指标。

In titration calculations, using the formula n = cV (mol = concentration × volume in dm³) and the balanced equation allows determination of unknown concentrations. Back titration is employed when the analyte is insoluble or volatile.

在滴定计算中，使用公式 n = cV（摩尔 = 浓度 × 体积，以 dm³ 计）并结合平衡方程式可求出未知浓度。当待测物不溶或易挥发时，则采用返滴定法。

11. Reaction Mechanisms and Rate-Determining Step | 反应机理与速控步

A reaction mechanism is the sequence of elementary steps by which a chemical change occurs. The molecularity of an elementary step is the number of reactant particles involved (unimolecular, bimolecular). The overall rate equation is determined by the slowest step, known as the rate-determining step (RDS).

反应机理是化学变化所经历的一系列基元步骤。基元步骤的分子数指参与反应的粒子数（单分子、双分子）。总反应速率方程由最慢的步骤决定，该步骤称为速控步（RDS）。

If the RDS involves a single molecule A decomposing, the rate = k[A]; if it involves A + B colliding, rate = k[A][B]. Intermediates, produced in one step and consumed in a later step, do not appear in the overall rate equation. Catalysts often appear in a mechanism being consumed in an early step and regenerated in a later step, thereby lowering the activation energy of the RDS.

若速控步涉及单个分子 A 分解，则速率 = k[A]；若涉及 A + B 碰撞，则速率 = k[A][B]。中间体在某一步生成、在后续步骤被消耗，不出现于总速率方程中。催化剂常出现在机理中：在早期步骤中被消耗，随后在后续步骤中再生，从而降低速控步的活化能。

Energy profiles for multi-step reactions have more than one hump; the highest energy barrier corresponds to the RDS. The number of peaks equals the number of elementary steps.

多步反应的能量曲线含有多个波峰；最高的能垒对应速控步。波峰数目等于基元步骤的数目。

12. Catalysis: Homogeneous and Heterogeneous | 催化作用：均相与多相

Homogeneous catalysts are in the same phase as the reactants, usually dissolved in solution. They function by forming an intermediate species that reacts more readily, and are regenerated at the end. An example is the role of Fe²⁺/Fe³⁺ in the iodide-persulfate reaction, or acid catalysis in ester hydrolysis.

均相催化剂与反应物处于同一相，通常溶解在溶液中。它们通过形成更容易反应的中间物种发挥作用，并在反应结束时再生。例如碘离子-过二硫酸根反应中的 Fe²⁺/Fe³⁺，或酯水解中的酸催化。

Heterogeneous catalysts exist in a different phase, commonly solid catalysts with gaseous or liquid reactants. The reaction takes place on the catalyst’s surface where adsorption weakens bonds in the reactant molecules, offering a new pathway with reduced Eₐ. In the Haber process, finely divided iron is used; in the Contact process, V₂O₅ oxidises SO₂ to SO₃. Catalytic converters in cars use platinum, palladium, and rhodium to oxidise CO and unburnt hydrocarbons and reduce NOₓ.

多相催化剂存在于不同相中，通常为固体，而反应物为气体或液体。反应发生在催化剂表面，吸附作用削弱了反应物分子中的化学键，提供了一条降低 Eₐ 的新路径。哈伯法中使用细碎铁粉；接触法中 V₂O₅ 将 SO₂ 氧化为 SO₃。汽车催化转化器使用铂、钯和铑将 CO 和未燃烧烃氧化，并还原 NOₓ。

Enzymes are the most specific catalysts. Their activity is modelled by the lock-and-key and induced-fit hypotheses, and is strongly influenced by temperature and pH. The maximum rate Vmax and the Michaelis constant Km describe the kinetics of enzyme-catalysed reactions.

酶是专一性最强的催化剂。其活性可由锁钥模型和诱导契合假说解释，且极大地受温度和 pH 影响。最大速率 Vmax 和米氏常数 Km 描述了酶催化反应的动力学。

Published by TutorHao | Chemistry Revision Series | aleveler.com

更多咨询请联系16621398022（同微信）