Chemical Equilibrium Le Chateliers Principle

Introduction to Reversible Reactions

Many chemical reactions do not go to completion. Instead, they are reversible — the products can react to re-form the original reactants. A reversible reaction is denoted by the ⇌ symbol, indicating that the reaction proceeds in both the forward and backward directions simultaneously.

许多化学反应并不会进行到底。相反，它们是可逆的——产物可以重新反应生成原始反应物。可逆反应用 ⇌ 符号表示，表明反应同时向正向和反向进行。

Consider the reaction between nitrogen and hydrogen to form ammonia: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). This is the Haber process, one of the most important industrial reactions. At the same time ammonia is being formed (forward reaction), it is also decomposing back into nitrogen and hydrogen (backward reaction).

以氮气和氢气生成氨的反应为例：N₂(g) + 3H₂(g) ⇌ 2NH₃(g)。这就是哈伯法，最重要的工业反应之一。在氨被生成的同时（正向反应），它也在分解回氮气和氢气（逆向反应）。

What is Dynamic Equilibrium?

When a reversible reaction reaches a state where the rate of the forward reaction equals the rate of the backward reaction, we say the system has reached dynamic equilibrium. The key word here is “dynamic” — the reactions have NOT stopped. Both forward and backward reactions continue at equal rates, so there is no net change in the concentrations of reactants and products.

当可逆反应达到正向反应速率等于逆向反应速率的状态时，我们称系统达到了动态平衡。关键词是”动态”——反应并没有停止。正向和逆向反应以相等的速率继续进行，因此反应物和产物的浓度没有净变化。

For a system to be at dynamic equilibrium, two conditions must be met:

系统要达到动态平衡，必须满足两个条件：

• Closed system: No matter can enter or leave the system. If gases escape, equilibrium cannot be maintained. 封闭系统：物质不能进入或离开系统。如果气体逸出，平衡就无法维持。
• Constant temperature: Equilibrium is temperature-dependent; changing the temperature shifts the position of equilibrium. 恒温：平衡依赖于温度；改变温度会移动平衡位置。

At equilibrium, the macroscopic properties — such as colour, pressure, and concentration — remain constant. However, at the molecular level, particles are continuously colliding and reacting. This is the essence of dynamic equilibrium: constant change at the microscopic level with no observable change at the macroscopic level.

在平衡状态下，宏观性质——如颜色、压力和浓度——保持不变。然而，在分子层面，粒子在持续碰撞和反应。这就是动态平衡的本质：微观层面的持续变化而宏观层面没有可观察到的变化。

Le Chatelier’s Principle: The Core Concept

Le Chatelier’s Principle states that if a system at dynamic equilibrium is subjected to a change in conditions, the position of equilibrium shifts to oppose that change. In simpler terms: the system “does the opposite” of whatever you do to it.

勒夏特列原理指出，如果一个处于动态平衡的系统受到条件变化的影响，平衡位置会移动以对抗这种变化。简单来说：系统会对你施加的变化做出”相反”的响应。

This principle, formulated by French chemist Henri Louis Le Chatelier in 1884, is a powerful predictive tool. It allows chemists to forecast how an equilibrium mixture will respond to changes in concentration, pressure, and temperature — without needing to know the detailed kinetics of the reaction.

这个原理由法国化学家亨利·路易·勒夏特列于1884年提出，是一个强大的预测工具。它使化学家能够预测平衡混合物如何响应浓度、压力和温度的变化——而无需了解反应的详细动力学。

Effect of Concentration Changes

If you increase the concentration of a reactant, the system will shift equilibrium to the right (towards products) to consume the excess reactant. Conversely, if you increase the concentration of a product, equilibrium shifts to the left (towards reactants).

如果你增加反应物的浓度，系统会将平衡向右移动（朝向产物）以消耗多余的反应物。相反，如果你增加产物的浓度，平衡会向左移动（朝向反应物）。

Worked Example: Consider the equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq). The FeSCN²⁺ ion is blood-red in colour. If you add more Fe³⁺ ions to the system, the solution becomes a deeper red. Why? The equilibrium shifts right to use up the added Fe³⁺, producing more of the red FeSCN²⁺ complex. If you then add SCN⁻, the colour intensifies further — confirming Le Chatelier’s prediction.

示例：考虑平衡 Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq)。FeSCN²⁺ 离子呈血红色。如果你向系统中加入更多 Fe³⁺ 离子，溶液颜色变深。为什么？平衡向右移动以消耗加入的 Fe³⁺，生成更多红色的 FeSCN²⁺ 配合物。如果再加入 SCN⁻，颜色会进一步加深——这验证了勒夏特列原理的预测。

In industrial processes, removing a product as it forms is a common strategy to drive equilibrium towards completion. This shifts equilibrium continuously to the right, maximising yield.

在工业过程中，在产物生成时将其移除是推动平衡向完成方向移动的常用策略。这使平衡持续向右移动，最大化产率。

Effect of Pressure Changes

Pressure changes only affect equilibria involving gases, and only when there is a difference in the number of gas molecules on each side of the equation. If you increase the pressure, equilibrium shifts to the side with fewer gas molecules to reduce the pressure. If you decrease the pressure, equilibrium shifts to the side with more gas molecules.

压力变化只影响涉及气体的平衡，且仅在方程式两边气体分子数量有差异时。如果增加压力，平衡向气体分子较少的一侧移动以降低压力。如果降低压力，平衡向气体分子较多的一侧移动。

Example: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)

• Left side: 2 + 1 = 3 gas molecules. 左侧：2 + 1 = 3 个气体分子。
• Right side: 2 gas molecules. 右侧：2 个气体分子。
• Increasing pressure shifts equilibrium RIGHT (fewer molecules). 增加压力使平衡向右移动（更少的分子）。
• Decreasing pressure shifts equilibrium LEFT (more molecules). 降低压力使平衡向左移动（更多的分子）。

If the number of gas molecules is equal on both sides — for example, H₂(g) + I₂(g) ⇌ 2HI(g), where both sides have 2 molecules — then pressure changes have no effect on the position of equilibrium. This is a common exam trap.

如果两边气体分子数量相等——例如 H₂(g) + I₂(g) ⇌ 2HI(g)，两边都有2个分子——那么压力变化对平衡位置没有影响。这是一个常见的考试陷阱。

Effect of Temperature Changes

Temperature is the only factor that changes the value of the equilibrium constant Kc. To predict the direction of shift, you must know whether the forward reaction is exothermic or endothermic.

温度是唯一改变平衡常数 Kc 值的因素。要预测移动方向，你必须知道正向反应是放热还是吸热。

• Exothermic forward reaction (ΔH negative): Increasing temperature shifts equilibrium LEFT. Decreasing temperature shifts equilibrium RIGHT. 正向放热反应（ΔH 为负）：升高温度使平衡向左移动。降低温度使平衡向右移动。
• Endothermic forward reaction (ΔH positive): Increasing temperature shifts equilibrium RIGHT. Decreasing temperature shifts equilibrium LEFT. 正向吸热反应（ΔH 为正）：升高温度使平衡向右移动。降低温度使平衡向左移动。

The logic: if temperature increases, the system shifts to absorb the extra heat — favouring the endothermic direction. If temperature decreases, the system shifts to release heat — favouring the exothermic direction.

逻辑是：如果温度升高，系统会移动以吸收额外的热量——有利于吸热方向。如果温度降低，系统会移动以释放热量——有利于放热方向。

Key example — the Haber Process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = -92 kJ mol⁻¹. The forward reaction is exothermic.

关键示例——哈伯法：N₂(g) + 3H₂(g) ⇌ 2NH₃(g)，ΔH = -92 kJ mol⁻¹。正向反应是放热的。

• Low temperature favours ammonia production (exothermic forward direction). 低温有利于氨的生成（放热正向）。
• BUT low temperature also means a slower rate of reaction! 但低温也意味着反应速率更慢！

This is the classic yield vs. rate compromise in industrial chemistry. The Haber process uses a compromise temperature of around 450°C — high enough for a reasonable rate, low enough for an acceptable yield. An iron catalyst is also used to speed up the rate without affecting the position of equilibrium.

这就是工业化学中经典的产率与速率的权衡。哈伯法使用约450°C的折中温度——足够高以获得合理的速率，足够低以获得可接受的产率。还使用铁催化剂来加快速率而不影响平衡位置。

Effect of Catalysts on Equilibrium

A catalyst speeds up both the forward and backward reactions equally. It provides an alternative reaction pathway with a lower activation energy for both directions. Because the rates of both reactions increase by the same factor, a catalyst does NOT change the position of equilibrium. It simply allows the system to reach equilibrium faster.

催化剂同等地加速正向和逆向反应。它为两个方向提供了具有更低活化能的替代反应路径。因为两个反应的速率以相同的倍数增加，催化剂不改变平衡位置。它只是让系统更快达到平衡。

This is a very common exam question. Students often mistakenly think that a catalyst increases the yield of products. It does not. It only reduces the time needed to reach equilibrium. The equilibrium yield is determined by thermodynamics, not kinetics.

这是一个非常常见的考题。学生经常错误地认为催化剂会增加产物产率。它不会。它只减少达到平衡所需的时间。平衡产率由热力学决定，而不是动力学。

The Equilibrium Constant Kc

For a general reaction: aA + bB ⇌ cC + dD, the equilibrium constant Kc is defined as:

对于一般反应：aA + bB ⇌ cC + dD，平衡常数 Kc 定义为：

Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ

Where square brackets denote equilibrium concentrations in mol dm⁻³. The concentrations of solids and pure liquids are omitted from the expression because their concentrations are effectively constant.

其中方括号表示平衡浓度，单位为 mol dm⁻³。固体和纯液体的浓度从表达式中省略，因为它们的浓度实际上是常数。

Important facts about Kc:

关于 Kc 的重要事实：

• Kc is temperature-dependent only. Changing concentration or pressure does not change Kc. Kc 仅依赖于温度。改变浓度或压力不会改变 Kc。
• A large Kc (much greater than 1) means the equilibrium lies to the right — products are favoured. 大的 Kc（远大于1）意味着平衡偏右——产物占优势。
• A small Kc (much less than 1) means the equilibrium lies to the left — reactants are favoured. 小的 Kc（远小于1）意味着平衡偏左——反应物占优势。
• Kc has no units when the total number of concentration terms is equal in numerator and denominator. Otherwise, units depend on the stoichiometry. 当分子和分母中浓度项的总数相等时，Kc 没有单位。否则，单位取决于化学计量数。

Calculating Kc from experimental data is a standard A-Level skill. You need the equilibrium concentrations of all species. If initial concentrations and one equilibrium concentration are given, use an ICE table (Initial, Change, Equilibrium) to determine the rest.

从实验数据计算 Kc 是一项标准的 A-Level 技能。你需要所有物种的平衡浓度。如果给出了初始浓度和一个平衡浓度，使用ICE 表（初始、变化、平衡）来确定其余浓度。

Industrial Application: The Haber Process

The Haber process for ammonia synthesis is the perfect case study for applying equilibrium principles. Here is how Le Chatelier’s Principle guides the choice of conditions:

哈伯法合成氨是应用平衡原理的完美案例研究。以下是勒夏特列原理如何指导条件选择的：

• Pressure — 200 atm: The forward reaction reduces gas molecules (4 → 2). High pressure shifts equilibrium right, increasing yield. However, extremely high pressures are expensive and require stronger equipment. 200 atm is the economic compromise. 压力——200 atm：正向反应减少气体分子（4 → 2）。高压使平衡向右移动，增加产率。然而，极高的压力成本高昂且需要更强的设备。200 atm 是经济折中。
• Temperature — 450°C: The forward reaction is exothermic. Low temperature favours higher yield, but would make the reaction too slow. 450°C is the compromise — reasonable yield with an acceptable rate. 温度——450°C：正向反应是放热的。低温有利于更高的产率，但会使反应太慢。450°C 是折中——合理的产率和可接受的速率。
• Catalyst — Iron: Speeds up the reaction without affecting equilibrium position, making the compromise temperature viable. 催化剂——铁：加速反应而不影响平衡位置，使折中温度变得可行。
• Continuous removal of NH₃: Ammonia is liquefied and removed as it forms, continually shifting equilibrium to the right. 持续移除 NH₃：氨在生成时被液化并移除，持续使平衡向右移动。

This process produces approximately 150 million tonnes of ammonia annually, primarily for fertilisers. It is estimated that without the Haber process, the world could only support about half its current population.

这个过程每年生产约1.5亿吨氨，主要用于化肥。据估计，没有哈伯法，世界只能养活大约一半的现有人口。

The Contact Process: Sulfuric Acid Production

Another important industrial equilibrium is the Contact Process: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), ΔH = -197 kJ mol⁻¹. Sulfuric acid is the most widely produced chemical globally.

另一个重要的工业平衡是接触法：2SO₂(g) + O₂(g) ⇌ 2SO₃(g)，ΔH = -197 kJ mol⁻¹。硫酸是全球产量最大的化学品。

• Pressure — 1-2 atm: Although the forward reaction reduces gas molecules (3 → 2) and high pressure would favour SO₃ production, the equilibrium already lies far to the right at atmospheric pressure. The extra cost of high-pressure equipment is not justified by the small yield increase. 压力——1-2 atm：虽然正向反应减少气体分子（3 → 2），高压有利于 SO₃ 生产，但在大气压下平衡已经大幅偏右。高压设备的额外成本无法由小幅产率提升来证明合理。
• Temperature — 450°C: Similar compromise to the Haber process — balancing yield against rate for this exothermic reaction. 温度——450°C：与哈伯法类似的折中——在这种放热反应中平衡产率与速率。
• Catalyst — Vanadium(V) oxide (V₂O₅): Increases the rate of attainment of equilibrium. 催化剂——五氧化二钒（V₂O₅）：增加达到平衡的速率。

Notice how both the Haber and Contact processes use similar temperatures (around 450°C). This is because both forward reactions are exothermic with similar activation energies. The pattern is common: exothermic industrial equilibria operate at moderate temperatures (compromise), while endothermic industrial processes often operate at very high temperatures (both rate and yield benefit from high temperature).

注意哈伯法和接触法都使用相似的温度（约450°C）。这是因为两个正向反应都是放热的，具有相似的活化能。这种模式很常见：放热的工业平衡在中等温度下运行（折中），而吸热的工业过程通常在非常高的温度下运行（高温对速率和产率都有利）。

Exam Tips and Common Mistakes

Here are the most important points to remember for A-Level examinations, along with the most frequent student errors:

以下是 A-Level 考试中最重要的要点以及学生最常见的错误：

1. Pressure has NO effect when gas molecules are equal on both sides. Always count gas molecules before answering. 当两边气体分子相等时，压力没有影响。在回答之前始终数一下气体分子数。
2. Catalysts do NOT change equilibrium position or yield. They only affect the rate at which equilibrium is reached. 催化剂不改变平衡位置或产率。它们只影响达到平衡的速率。
3. Only temperature changes the value of Kc. Concentration and pressure changes do not affect Kc; they only shift the position of equilibrium. 只有温度改变 Kc 的值。浓度和压力变化不影响 Kc；它们只移动平衡位置。
4. Use ICE tables for Kc calculations. Write Initial, Change, and Equilibrium rows clearly. The Change row must reflect the stoichiometric ratios. 用 ICE 表进行 Kc 计算。清楚地写出初始、变化和平衡行。变化行必须反映化学计量比。
5. State symbols matter. Solids and pure liquids are excluded from Kc expressions. Gases and aqueous species are included. 状态符号很重要。固体和纯液体从 Kc 表达式中排除。气体和水相物种包括在内。
6. “Equilibrium shifts” vs. “Rate increases”. These are distinct concepts. Equilibrium shift is about thermodynamics; rate is about kinetics. “平衡移动”与”速率增加”。这些是不同的概念。平衡移动关乎热力学；速率关乎动力学。
7. When explaining using Le Chatelier, always state the OPPOSING effect. E.g., “Increasing temperature shifts equilibrium in the endothermic direction to absorb the added heat.” 使用勒夏特列原理解释时，始终说明对抗效应。例如，”升高温度使平衡向吸热方向移动，以吸收增加的热量。”

Summary

Chemical equilibrium is a fundamental concept in physical chemistry that describes the balance point of reversible reactions. Le Chatelier’s Principle provides a qualitative framework for predicting how an equilibrium system responds to external changes. Understanding the distinction between thermodynamic control (equilibrium position and Kc) and kinetic control (rate of reaction and catalysts) is essential for success at A-Level and beyond. The industrial applications — the Haber and Contact processes — demonstrate how these principles are applied in the real world to produce chemicals that sustain modern civilisation.

化学平衡是物理化学中的一个基本概念，描述了可逆反应的平衡点。勒夏特列原理提供了一个定性框架，用于预测平衡系统如何响应外部变化。理解热力学控制（平衡位置和 Kc）与动力学控制（反应速率和催化剂）之间的区别对于 A-Level 及更高层次的成功至关重要。工业应用——哈伯法和接触法——展示了这些原理如何在现实世界中应用，以生产维持现代文明的化学品。