Welcome to Chemical Equilibrium
Welcome to one of the most important concepts in physical and industrial chemistry! In chemical manufacturing, making a product is not just about whether a reaction can happen, but how much product you get and how quickly you can make it. In this chapter, you will learn why many reactions never go to \(100\%\) completion, how chemical systems respond to changes, and how industrial chemists balance yield, speed, and safety in processes like the synthesis of ammonia.
Don't worry if this topic feels a bit abstract at first. We will break down every concept step-by-step with real-world analogies and clear exam-focused explanations.
---1. Reversible Reactions, Closed Systems, and Dynamic Equilibrium
What is a Reversible Reaction?
In many reactions you have met before, reactants turn into products and the reaction stops when the reactants run out (an irreversible reaction). However, in a reversible reaction, the products can react together to reform the original reactants.
We represent a reversible reaction using the equilibrium symbol: \(\rightleftharpoons\)
Forward reaction: Reactants \(\rightarrow\) Products
Reverse (backward) reaction: Products \(\rightarrow\) Reactants
What is a Closed System?
A closed system is a system where matter (reactants and products) cannot enter or leave, but energy (heat) can be freely exchanged with the surroundings. For example, a sealed flask is a closed system. If the flask is left open (an open system), gases could escape, and the reverse reaction would not be able to happen at the same rate.
Crucial Rule: Dynamic equilibrium can ONLY be established in a closed system.
Defining Dynamic Equilibrium
Imagine walking up a "down" escalator. If you walk up at the exact same speed the escalator is moving down, your overall position does not change, even though both you and the escalator are constantly moving. This is dynamic equilibrium!
To score full marks on CCEA exam questions, you must state both essential features of dynamic equilibrium:
1. The rate of the forward reaction equals the rate of the reverse reaction.
2. The concentrations of reactants and products remain constant (all macroscopic properties, such as colour, pressure, and temperature, stay unchanged).
Examiner Warning: A very common mistake is saying that the concentrations of reactants and products are equal. They are not necessarily equal to each other! They are simply constant (they stop changing over time).
Key Takeaway: Dynamic equilibrium is an active balance. Reactions are still taking place in both directions at identical speeds, which keeps the amounts of reactants and products constant in a closed system.
---2. Le Chatelier’s Principle
What happens when you disturb a system that is in dynamic equilibrium? A French chemist named Henri Le Chatelier figured out how systems respond.
Statement of Le Chatelier’s Principle
"If a system at dynamic equilibrium is subjected to a change in conditions (concentration, temperature, or pressure), the position of equilibrium will shift in the direction that opposes the change."
Memory Trick: Think of equilibrium as stubborn. Whatever you do to it, it tries to do the exact opposite to undo your change!
A. Effect of Changing Concentration
When you change the concentration of a substance in the reaction vessel:
• Increasing the concentration of a reactant: The system opposes this by using up the added reactant. The equilibrium shifts to the right (towards the products).
• Decreasing/removing the concentration of a product: The system opposes this by replacing the lost product. The equilibrium shifts to the right (towards the products).
• Increasing the concentration of a product: The system opposes this by consuming the added product. The equilibrium shifts to the left (towards the reactants).
B. Effect of Changing Pressure (Gases Only)
Pressure is caused by gas particles colliding with the walls of the container. More gas molecules mean higher pressure.
• Increasing total pressure: The system opposes this by trying to lower the pressure. It shifts the equilibrium towards the side with fewer moles of gas.
• Decreasing total pressure: The system opposes this by trying to increase the pressure. It shifts the equilibrium towards the side with more moles of gas.
• Equal moles of gas on both sides? A change in pressure has no effect on the position of equilibrium.
Step-by-Step Method for Pressure Questions:
Step 1: Count the total moles of gas on the left-hand side of the balanced equation.
Step 2: Count the total moles of gas on the right-hand side.
Step 3: If pressure increases, shift to the side with the smaller number of gas moles.
C. Effect of Changing Temperature
Every reversible reaction is exothermic in one direction (\(\Delta H < 0\), releases heat) and endothermic in the opposite direction (\(\Delta H > 0\), absorbs heat).
• Increasing temperature: The system opposes this by trying to cool down and absorb the added heat. The equilibrium shifts in the endothermic direction (\(\Delta H > 0\)).
• Decreasing temperature: The system opposes this by trying to warm up and release heat. The equilibrium shifts in the exothermic direction (\(\Delta H < 0\)).
D. Effect of a Catalyst
This is one of the most frequently tested concepts in CCEA AS 3 examinations:
• A catalyst provides an alternative reaction pathway with a lower activation energy (\(E_a\)).
• It increases the rate of the forward and reverse reactions equally.
• Therefore, a catalyst does NOT change the position of equilibrium and does NOT change the equilibrium yield.
• Its only role is to allow dynamic equilibrium to be reached faster.
Key Takeaway: Changing concentration, pressure, or temperature shifts the position of equilibrium to oppose the change. Adding a catalyst speeds up both directions equally and has zero effect on the yield.
---3. Industrial Applications & Compromise Conditions: The Haber Process
In industry, chemical companies want to make as much product as possible (high yield) as fast as possible (high reaction rate) while keeping costs and hazards manageable. Often, the ideal condition for a high equilibrium yield gives a terrible reaction rate, requiring a compromise.
The Haber Process for the Synthesis of Ammonia
Ammonia (\(\text{NH}_3\)) is vital for manufacturing agricultural fertilisers and chemical feedstocks.
\(\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)} \quad \Delta H = -92\text{ kJ mol}^{-1}\)
Let's analyse the forward reaction:
• Gas moles: \(1\text{ mole of }\text{N}_2 + 3\text{ moles of }\text{H}_2 = 4\text{ moles of gas on the left}\). \(2\text{ moles of }\text{NH}_3 = 2\text{ moles of gas on the right}\).
• Enthalpy change: \(\Delta H = -92\text{ kJ mol}^{-1}\) (The forward reaction is exothermic).
1. Choosing the Temperature (\(400\text{ }^\circ\text{C} - 450\text{ }^\circ\text{C}\))
• Equilibrium argument: Because the forward reaction is exothermic, a low temperature shifts the equilibrium to the right, giving a higher equilibrium yield of ammonia.
• Rate argument: At low temperatures, gas particles have low kinetic energy, resulting in infrequent collisions and very few collisions having energy \(\ge E_a\). The reaction rate is far too slow to be commercially viable.
• The Compromise: A moderate temperature of \(400\text{ }^\circ\text{C} - 450\text{ }^\circ\text{C}\) is used. This provides an acceptable equilibrium yield within a reasonable, cost-effective timeframe.
2. Choosing the Pressure (\(200\text{ atm}\) / \(20\text{ MPa}\))
• Equilibrium argument: There are \(4\text{ moles of gas}\) on the left and \(2\text{ moles of gas}\) on the right. A high pressure shifts equilibrium to the side with fewer gas moles (the right), giving a higher equilibrium yield of ammonia.
• Rate argument: High pressure pushes gas particles closer together, increasing collision frequency and therefore increasing the rate of reaction.
• The Cost and Safety Limits: Extremely high pressures require very expensive, thick-walled reaction vessels and pipes, high electricity costs to run powerful compressors, and pose significant safety and explosion risks.
• The Compromise: A compromise pressure of approximately \(200\text{ atm}\) (\(20\text{ MPa}\)) is chosen to maximise yield and rate without excessive capital and maintenance costs.
3. The Catalyst
• An iron (Fe) catalyst (finely divided to provide a large surface area) is used.
• The catalyst speeds up both forward and reverse reactions equally.
• It does not increase the percentage yield, but it allows the process to run at a commercially acceptable rate at the compromise temperature of \(400\text{ }^\circ\text{C} - 450\text{ }^\circ\text{C}\).
4. Separation and Recycling
At each pass over the catalyst bed, only around \(15\%\) of the nitrogen and hydrogen gases are converted into ammonia. To make the process efficient:
1. The gases are cooled so that gaseous ammonia condenses into a liquid and is removed from the reaction vessel. (Removing product shifts the equilibrium further to the right!).
2. The unreacted \(\text{N}_2\) and \(\text{H}_2\) are continuously recycled back over the iron catalyst bed so that nothing is wasted.
Key Takeaway: Industrial conditions represent a carefully calculated balance. The Haber process uses \(400\text{ }^\circ\text{C} - 450\text{ }^\circ\text{C}\), \(200\text{ atm}\), an iron catalyst, and continuous recycling of unreacted gases.
---4. Common Exam Pitfalls & How to Avoid Them
Keep these examiner insights in mind when answering questions on Topic 3.5:
• Pitfall 1: Half-definitions of Dynamic Equilibrium.
Always mention both points: (1) forward rate equals reverse rate, AND (2) concentrations of reactants and products remain constant.
• Pitfall 2: Confusing "Equal" with "Constant".
Never say reactants and products have "equal concentrations". Write "concentrations remain constant".
• Pitfall 3: Claiming catalysts increase yield.
A catalyst never shifts the equilibrium or increases the amount of product formed. It only allows equilibrium to be reached in less time.
• Pitfall 4: Forgetting to count gas moles.
When explaining pressure changes, always state the exact numbers: "There are 4 moles of gas on the left and 2 moles of gas on the right, so increasing pressure shifts equilibrium to the right."
• Pitfall 5: Giving one-sided compromise answers.
When asked why a compromise temperature or pressure is used, clearly state the effect on equilibrium yield AND the effect on reaction rate / economic cost.
Quick Summary Checklist
Before sitting your AS 3 exam, make sure you can:
• State the two conditions for dynamic equilibrium.
• Define a closed system.
• State Le Chatelier’s principle accurately.
• Predict equilibrium shifts when concentration, pressure, or temperature changes.
• Explain the effect of a catalyst on rate and equilibrium position.
• Explain the compromise temperature (\(400\text{ }^\circ\text{C} - 450\text{ }^\circ\text{C}\)), compromise pressure (\(200\text{ atm}\)), catalyst (iron), and recycling in the Haber process.