Core Knowledge Review of Chemical Reaction Principles | 化学反应原理核心知识梳理

📚 Core Knowledge Review of Chemical Reaction Principles | 化学反应原理核心知识梳理

Chemical reaction principles unify thermochemistry, kinetics, equilibrium, acid-base behaviour, and electrochemistry into one logical framework. This article organises the core concepts and exam-facing formulas for A-Level and equivalent chemistry revision.

化学反应原理将热化学、动力学、平衡、酸碱行为与电化学统一于一个逻辑框架之中。本文面向A-Level及同等学力化学备考，系统梳理核心概念与高频考点公式。

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1. Enthalpy Changes and Thermochemistry | 焓变与热化学

Enthalpy (H) is the total heat content of a system at constant pressure. The enthalpy change, ΔH, equals the heat exchanged with the surroundings when pressure is constant.

焓（H）是体系在恒压条件下的总热含量。焓变ΔH等于恒压条件下体系与环境交换的热量。

Exothermic reactions release heat and have ΔH < 0; endothermic reactions absorb heat and have ΔH > 0. Combustion and neutralisation are typically exothermic; thermal decomposition is usually endothermic.

放热反应释放热量，ΔH < 0；吸热反应吸收热量，ΔH > 0。燃烧与中和反应通常放热，热分解通常吸热。

The standard enthalpy change of formation, ΔHf°, is the enthalpy change when one mole of a compound forms from its elements in their standard states under standard conditions.

标准摩尔生成焓（ΔHf°）是指在标准条件下，由稳定单质生成1 mol化合物的焓变。

C(s) + O₂(g) → CO₂(g), ΔH = −394 kJ mol⁻¹

The enthalpy change can be calculated from bond enthalpies: ΔH = sum of bonds broken − sum of bonds formed.

焓变可由键能计算：ΔH = 断裂键能总和 − 形成键能总和。

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2. Hess’s Law | 盖斯定律

Hess’s Law states that the enthalpy change of a reaction depends only on the initial and final states, not on the reaction pathway. Therefore, ΔH for a multi-step reaction equals the sum of ΔH for each step.

盖斯定律指出：反应的焓变只取决于始态和终态，与反应途径无关。因此，分步反应的总焓变等于各步焓变之和。

This enables calculation of ΔH for reactions that cannot be measured directly, such as the formation of carbon monoxide from carbon and oxygen, by using formation or combustion data.

借此可以计算无法直接测定的反应焓变，例如利用生成焓或燃烧焓数据求一氧化碳的生成焓。

ΔH_reaction = ΣΔHf°(products) − ΣΔHf°(reactants)

When applying Hess cycles, ensure reaction equations are correctly reversed (which flips the sign of ΔH) and multiplied (which scales ΔH by the same factor).

应用盖斯循环时，注意方程式反转要改变ΔH符号，方程式乘以系数时ΔH也要乘以相同倍数。

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3. Reaction Rates and Factors Affecting Rate | 反应速率及其影响因素

The rate of a reaction is the change in concentration of a reactant or product per unit time, usually expressed in mol dm⁻³ s⁻¹.

反应速率是指单位时间内反应物或产物浓度的变化量，常用单位mol dm⁻³ s⁻¹表示。

Factors affecting rate include concentration, pressure (for gases), temperature, surface area, and catalysts. Increasing concentration or pressure increases collision frequency; raising temperature increases both collision frequency and the fraction of energetic collisions.

影响反应速率的因素包括浓度、压强（对气体）、温度、接触面积和催化剂。增大浓度或压强可提高碰撞频率；升温既提高碰撞频率，又提高有效碰撞的比例。

Instantaneous rate is obtained from the gradient of a concentration–time graph; the initial rate method uses the tangent at t = 0 for kinetics studies.

瞬时速率可通过浓度–时间图像的切线斜率求得；初始速率法在动力学研究中以t = 0处的切线斜率作为初速率。

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4. Rate Laws and Activation Energy | 速率方程与活化能

For a reaction aA + bB → cC + dD, the rate law is rate = k[A]ᵐ[B]ⁿ, where k is the rate constant and m, n are reaction orders that must be determined experimentally.

对于反应 aA + bB → cC + dD，速率方程为 rate = k[A]ᵐ[B]ⁿ，其中k为速率常数，m、n为反应级数，须由实验测定。

The overall order is m + n. Zero-order, first-order, and second-order reactions show characteristic concentration–time relationships and half-life patterns.

总级数为 m + n。零级、一级和二级反应呈现出特征性的浓度–时间关系与半衰期规律。

For a first-order reaction, rate = k[A] and t₁/₂ = ln 2 / k, independent of initial concentration.

一级反应满足 rate = k[A]，且半衰期 t₁/₂ = ln 2 / k，与初始浓度无关。

k = A e^(−Ea/RT)

The Arrhenius equation relates the rate constant to activation energy Ea. Plotting ln k against 1/T gives a straight line with gradient −Ea/R.

阿伦尼乌斯方程将速率常数与活化能Ea联系起来。作ln k对

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