Welcome to Chemical Equilibrium!
Welcome to one of the most interesting and important topics in GCSE Chemistry! In everyday life, many processes only go in one direction — once you bake a cake or burn a piece of paper, you cannot un-bake the cake or un-burn the ash back into paper. But did you know that many chemical reactions can go both forwards and backwards? In this chapter, you will learn how these "two-way" reactions work, what happens when they balance out, and how industrial chemists control them to make everyday products like fertilizers.
Don't worry if this topic feels a bit strange at first! We will break it down into easy, step-by-step pieces with everyday analogies to help you master every single concept.
1. Reversible Reactions
Most chemical reactions you have studied so far are irreversible. That means the reactants turn completely into products, and the reaction stops when reactants run out.
However, some reactions are reversible. In a reversible reaction, the products can react together to reform the original reactants.
The Reversible Arrow Symbol
In chemical equations, we represent a reversible reaction using a special double arrow symbol: \( \rightleftharpoons \)
• The top arrow pointing right (\( \rightarrow \)) represents the forward reaction.
• The bottom arrow pointing left (\( \leftarrow \)) represents the reverse (or backward) reaction.
Key Examples of Reversible Reactions
Example 1: Dehydration of Hydrated Copper(II) Sulfate
Hydrated copper(II) sulfate is bright blue and contains water trapped in its crystals. When you heat it, water is driven off, leaving behind a white powder called anhydrous copper(II) sulfate:
\( \text{CuSO}_4 \cdot 5\text{H}_2\text{O (s)} \rightleftharpoons \text{CuSO}_4\text{ (s)} + 5\text{H}_2\text{O (l)} \)
• Forward reaction (endothermic): Blue crystals \( \rightarrow \) White powder + Water (requires heat)
• Reverse reaction (exothermic): White powder + Water \( \rightarrow \) Blue crystals (releases heat)
Example 2: Heating Ammonium Chloride
When solid ammonium chloride is heated, it breaks down into ammonia gas and hydrogen chloride gas. As the gases cool down higher up the test tube, they react together to reform solid white ammonium chloride:
\( \text{NH}_4\text{Cl (s)} \rightleftharpoons \text{NH}_3\text{ (g)} + \text{HCl (g)} \)
Key Takeaway: If a reaction is endothermic in the forward direction, it is always equally exothermic in the reverse direction, and vice versa!
2. Dynamic Equilibrium
Imagine walking up a "down" escalator. If you walk up at the exact same speed that the escalator is moving down, what happens? To someone watching, you stay in the exact same spot! You are still moving, and the stairs are still moving, but there is no overall change in your position. This is the perfect picture of dynamic equilibrium.
What is Dynamic Equilibrium in Chemistry?
When a reversible reaction takes place in a closed system, the forward and reverse reactions happen at the exact same time. Eventually, they reach a state of balance called dynamic equilibrium.
The Crucial Features of Dynamic Equilibrium
To get full marks in an exam, remember these essential features:
1. Closed System: No substances (reactants or products) can enter or leave the reaction container.
2. Equal Rates: The rate of the forward reaction is equal to the rate of the reverse reaction.
3. Constant Concentrations: The amounts (concentrations) of reactants and products remain constant (they do not have to be equal to each other, but they stop changing).
4. Macroscopic properties remain unchanged: Observable properties like colour, temperature, and pressure stay constant.
Common Mistake to Avoid:
Do not say that the concentrations of reactants and products are equal! They are usually different; the key is that their amounts remain constant because they are being made and used up at the exact same rate.
Key Takeaway: Dynamic equilibrium means the forward rate equals the reverse rate in a closed system, so concentrations stay constant.
3. Changing the Position of Equilibrium (Le Chatelier's Principle)
The position of equilibrium tells us whether there are more products or more reactants at equilibrium:
• If equilibrium lies to the right, there are more products.
• If equilibrium lies to the left, there are more reactants.
Le Chatelier's Principle
"If a change in conditions is applied to a system at equilibrium, the system will adjust itself to oppose that change."
Think of equilibrium like a stubborn teenager: whatever change you try to make, the system does the exact opposite to undo it!
Rule 1: Effect of Temperature
Every reversible reaction is exothermic in one direction and endothermic in the other.
• If you increase the temperature: The system tries to cool down by favouring the endothermic direction (which absorbs heat).
• If you decrease the temperature: The system tries to warm up by favouring the exothermic direction (which releases heat).
Rule 2: Effect of Pressure (Gases Only)
Pressure is caused by gas particles colliding with the walls of the container. Count the number of gas molecules (moles of gas) on each side of the equation:
• If you increase the pressure: The system tries to reduce pressure by shifting to the side with fewer moles of gas.
• If you decrease the pressure: The system tries to increase pressure by shifting to the side with more moles of gas.
• If both sides have equal moles of gas: Changing pressure has no effect on the position of equilibrium.
Rule 3: Effect of Concentration
• If you increase the concentration of a reactant: The system shifts to the right to use up the extra reactant and make more product.
• If you remove a product: The system shifts to the right to replace the removed product.
Rule 4: Effect of a Catalyst
• A catalyst speeds up the forward reaction and the backward reaction by the exact same amount.
• Therefore, adding a catalyst does NOT change the position of equilibrium and does not alter the yield.
• It simply helps the reaction reach equilibrium faster.
Quick Summary Box:
• Increase Temp \( \rightarrow \) shifts to endothermic direction (\( +\Delta H \))
• Decrease Temp \( \rightarrow \) shifts to exothermic direction (\( -\Delta H \))
• Increase Pressure \( \rightarrow \) shifts to side with fewer gas molecules
• Decrease Pressure \( \rightarrow \) shifts to side with more gas molecules
• Add Catalyst \( \rightarrow \) no shift in position (reaches equilibrium faster)
4. Real-World Application: The Haber Process
The Haber Process is used industrially to manufacture ammonia (\( \text{NH}_3 \)), which is essential for making agricultural fertilizers to feed the world's growing population.
The Reaction Equation
\( \text{N}_2\text{ (g)} + 3\text{H}_2\text{ (g)} \rightleftharpoons 2\text{NH}_3\text{ (g)} \quad (\Delta H = -92\text{ kJ/mol, exothermic forward}) \)
Raw Materials
• Nitrogen (\( \text{N}_2 \)): Extracted from the air.
• Hydrogen (\( \text{H}_2 \)): Obtained from natural gas (methane).
Industrial Conditions and The "Compromise"
In industry, chemical engineers must balance high yield, fast rate of reaction, and operating costs:
1. Temperature: \( 450^\circ\text{C} \) (A Compromise Temperature)
• Yield reason: Because the forward reaction is exothermic, a low temperature gives the highest yield of ammonia.
• Rate reason: However, at low temperatures, the reaction is far too slow.
• Compromise: \( 450^\circ\text{C} \) gives an acceptable yield in a reasonably fast time.
2. Pressure: \( 200\text{ atmospheres (atm)} \) (A Compromise Pressure)
• Yield reason: There are \( 4 \) moles of gas on the left (\( 1\text{ N}_2 + 3\text{H}_2 \)) and only \( 2 \) moles of gas on the right (\( 2\text{ NH}_3 \)). A high pressure favours the right side, giving a higher yield of ammonia.
• Cost and safety reason: Extremely high pressures require very expensive, thick-walled steel pipes and massive amounts of energy to run compressors, increasing explosion risks.
• Compromise: \( 200\text{ atm} \) is high enough to give a good yield without making equipment dangerously expensive.
3. Catalyst: Iron Catalyst
• The iron catalyst increases the rate of reaction, allowing equilibrium to be reached much faster at lower temperatures.
• It does not affect the yield, but it saves time and reduces energy costs.
4. Recycling Unreacted Gases
• At each pass, only about \( 15\% \) of the gases turn into ammonia.
• The mixture is cooled so ammonia condenses into a liquid and is removed.
• The unreacted nitrogen and hydrogen are recycled back into the reactor so none is wasted!
Key Takeaway: The chosen conditions for the Haber process (\( 450^\circ\text{C} \), \( 200\text{ atm} \), iron catalyst) represent a careful economic compromise between rate, yield, and safety costs.
5. Quick Revision & Common Pitfalls Checklist
Before your exam, make sure you can answer these common questions:
• What is a closed system? A container where no reactants or products can escape or enter.
• What happens to the forward and backward rates at equilibrium? They are completely equal.
• Why does a catalyst not alter the yield? Because it speeds up both the forward and backward reactions at the exact same rate.
• How do you find which side has fewer gas moles? Add up the balancing big numbers in front of all gaseous substances on each side of the equation.
• Why is \( 450^\circ\text{C} \) used in the Haber process instead of room temperature? While lower temperatures give a higher yield, the rate of reaction would be too slow to be commercially viable.