Reversible Reactions and Equilibrium

Not every reaction goes only one way. In a reversible reaction the products can react to re-form the reactants, and if the reaction happens in a closed system it settles into a dynamic equilibrium where the forward and backward changes balance. This idea — and how changing the conditions shifts the balance — is the climax of Topic C5 of OCR Gateway Science A, and it explains how an industrial process like the manufacture of ammonia is run. This lesson recaps catalysts, defines reversible reactions and dynamic equilibrium, and (for Higher tier) applies Le Chatelier's principle to the Haber process.

By the end of this lesson you should be able to describe what a catalyst does, recognise reversible reactions, explain dynamic equilibrium, and (Higher tier) use Le Chatelier's principle to predict the effect of changing conditions, applied to the Haber process.

Reversible reactions are the final piece of the "controlling reactions" story. Many of the reactions met so far go essentially to completion, but a great many important reactions — including the one that makes the fertilisers that feed much of the world — do not. Understanding why, and how the conditions can be tuned to favour the products, is one of the most powerful ideas in the whole of GCSE chemistry.

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Catalysts (recap)

A catalyst speeds up a reaction by providing an alternative reaction pathway with a lower activation energy, so a greater proportion of collisions are successful. A catalyst is not used up in the reaction — it is unchanged at the end and can be used again. Enzymes are biological catalysts: protein molecules that catalyse the reactions in living things.

Catalysts are valuable in industry because they let a reaction run faster (and often at a lower temperature), saving energy and money, without being consumed.

Exam Tip: A catalyst lowers the activation energy and is not used up. Enzymes are biological catalysts. A catalyst does not change the products or the overall energy change — only the rate.

Because a catalyst is not consumed, only a small amount is needed to process a large quantity of reactant — the same catalyst particles take part again and again. This is why catalysts, even expensive ones, are economical to use in industry: the cost is spread over a huge amount of product. It also explains why a catalyst does not appear in the overall balanced equation for a reaction — it is present at the start and the end unchanged, so it is conventionally written above the arrow rather than as a reactant or product.

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Reversible Reactions

In a reversible reaction the products can react together to re-form the original reactants. We show this with a special double arrow ($\rightleftharpoons$⇌) instead of a single arrow:

$\text{reactants} \rightleftharpoons \text{products}$reactants⇌products

A familiar example is the heating of hydrated copper(II) sulfate (blue crystals). Heating drives off water to leave anhydrous copper(II) sulfate (white powder); adding water back turns it blue again:

$\text{hydrated copper sulfate} \rightleftharpoons \text{anhydrous copper sulfate} + \text{water}$hydrated copper sulfate⇌anhydrous copper sulfate+water $\text{(blue)} \rightleftharpoons \text{(white)}$(blue)⇌(white)

The forward change (heating, blue → white) is endothermic; the backward change (adding water, white → blue) is exothermic. This colour change is used as a test for water. Another example is ammonium chloride, which decomposes on heating and recombines on cooling:

$\text{NH}_4\text{Cl} \rightleftharpoons \text{NH}_3 + \text{HCl}$NH4​Cl⇌NH3​+HCl

Exam Tip: The double arrow $\rightleftharpoons$⇌ means the reaction is reversible. If the forward reaction is endothermic, the reverse is exothermic (and vice versa), by the same amount — a point that matters for predicting the effect of temperature.

The copper(II) sulfate colour change is a useful chemical test for the presence of water: if a liquid is added to white anhydrous copper(II) sulfate and the solid turns blue, the liquid contains water. (To test that the water is pure you would also check that it boils at exactly $100 \text{ °C}$100 °C.) The ammonium chloride example is a striking demonstration of reversibility too: heated at the bottom of a test tube, the white solid disappears as it decomposes into two colourless gases, which then recombine higher up the tube where it is cooler, depositing white ammonium chloride again. Both examples show the same key feature — the products can re-form the reactants — which is what the double arrow represents.

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Dynamic Equilibrium

If a reversible reaction takes place in a closed system (nothing can enter or leave), it reaches a dynamic equilibrium. At equilibrium:

• the forward and backward reactions are still happening (it is "dynamic", not stopped);
• they occur at exactly the same rate;
• so the concentrations of reactants and products stay constant (they do not change with time).

"Equilibrium" does not mean the amounts of reactants and products are equal — it means they are constant, because the two opposing reactions are balanced. The actual position of equilibrium (whether there is more product or more reactant) depends on the conditions.

Exam Tip: Define dynamic equilibrium as: in a closed system, the forward and backward reactions occur at the same rate, so the concentrations stay constant. Stress "constant, not equal" — this is the distinction examiners test most.

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Higher Tier: Le Chatelier's Principle

Higher tier only: Le Chatelier's principle states that if a condition of a system at equilibrium is changed, the equilibrium shifts to oppose (counteract) the change. This lets us predict the effect of changing temperature, pressure or concentration:

Change	Equilibrium shifts…
Increase temperature	in the endothermic direction (to absorb the added heat)
Decrease temperature	in the exothermic direction
Increase pressure (gases)	to the side with fewer gas molecules
Decrease pressure (gases)	to the side with more gas molecules
Increase a concentration	away from the substance added (using it up)

A catalyst does not change the position of equilibrium — it speeds up the forward and backward reactions equally, so equilibrium is reached faster but the amounts of reactants and products at equilibrium are unchanged.

Exam Tip: Le Chatelier in one line: the equilibrium shifts to oppose the change you make. For temperature, shift in the endothermic direction when heated; for pressure, shift to the side with fewer gas molecules when squeezed; for concentration, shift away from what you add.

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Higher Tier: The Haber Process

Higher tier only: The Haber process makes ammonia ($\text{NH}_3$NH3​) for fertilisers from nitrogen (from the air) and hydrogen (from natural gas). It is a reversible reaction, and the forward reaction is exothermic:

$\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3 \quad (\text{forward exothermic})$N2​+3H2​⇌2NH3​(forward exothermic)