Reversible Reactions and Equilibrium

Not every reaction goes only one way. In a reversible reaction the products can react together to re-form the original reactants. If a reversible reaction is left in a closed system, it settles into a special balanced state called a dynamic equilibrium, and a rule called Le Chatelier's principle lets you predict how changing the conditions will shift that balance. This is the climax of Topic C5 of your OCR Gateway Combined Science course, and it explains, among other things, how the ammonia in fertilisers is manufactured. This lesson defines reversible reactions and dynamic equilibrium, and (for Higher tier) applies Le Chatelier's principle to changes in temperature, pressure and concentration.

By the end of this lesson you should be able to recognise reversible reactions and use the $\rightleftharpoons$⇌ symbol, explain what dynamic equilibrium means, and (Higher tier) use Le Chatelier's principle to predict the effect of changing temperature, pressure or concentration.

This lesson moves from AO1 (recalling what a reversible reaction and dynamic equilibrium are) to AO3, where you apply Le Chatelier's principle to predict and justify how a change in conditions shifts the position of equilibrium.

Reversible reactions are the final piece of the "controlling reactions" story. Many reactions go essentially to completion, but a great many important ones — including the reaction that makes the fertilisers feeding 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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Reversible Reactions

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

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

The top arrow represents the forward reaction (reactants turning into products) and the bottom arrow the backward reaction (products turning back into reactants).

A familiar example is the heating of hydrated copper(II) sulfate, the blue crystals. Heating drives off water to leave anhydrous copper(II) sulfate, a 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 — it takes in energy; the backward change (adding water, white → blue) is exothermic — it gives energy out. This is a general rule for reversible reactions: if the forward reaction is endothermic, the backward reaction is exothermic, by the same amount (and vice versa). That fact matters when you predict the effect of temperature.

Another simple example is ammonium chloride, a white solid that decomposes on heating into two colourless gases and recombines on cooling:

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

Heated at the bottom of a tube, the white solid seems to vanish as it breaks apart into ammonia and hydrogen chloride; higher up, where it is cooler, the two gases recombine and deposit white ammonium chloride again — a striking demonstration that the products can re-form the reactants.

Exam Tip: The $\rightleftharpoons$⇌ arrow always means the reaction is reversible. Remember the energy rule: if the forward reaction is endothermic, the reverse is exothermic. The hydrated/anhydrous copper(II) sulfate colour change (blue ⇌ white) is a common example and is also a chemical test for water.

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Dynamic Equilibrium

If a reversible reaction takes place in a closed system — a container that lets nothing enter or leave — it eventually reaches a dynamic equilibrium. At equilibrium:

• the forward and backward reactions are both still happening (this is why it is called "dynamic", not "stopped");
• they are happening at exactly the same rate;
• so the concentrations of the reactants and products stay constant — they no longer change with time.

The word "equilibrium" does not mean the amounts of reactants and products are equal. It means they are constant, because the two opposing reactions are perfectly balanced. There may be far more product than reactant, or far more reactant than product; what matters is that the amounts are no longer changing. The actual position of equilibrium — whether it favours the products or the reactants — depends on the conditions.

Exam Tip: Define dynamic equilibrium precisely: in a closed system, the forward and backward reactions occur at the same rate, so the concentrations stay constant. Stress "constant, not equal" — a very common misconception is to think equilibrium means equal amounts of reactants and products, but it only means unchanging amounts.

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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 will shift to oppose (counteract) that change. In other words, the system responds so as to partly cancel out whatever you did to it. This one rule lets you predict the effect of changing temperature, pressure or concentration.

Change you make	The equilibrium shifts…
Increase the temperature	in the endothermic direction (to take in the added heat)
Decrease the temperature	in the exothermic direction (to release heat)
Increase the pressure (gases)	to the side with fewer gas molecules
Decrease the pressure (gases)	to the side with more gas molecules
Increase a concentration	away from the substance you added (to use it up)

Working through the logic each time is more reliable than memorising the table. If you heat a system, it shifts in the direction that absorbs heat (endothermic), because that opposes the temperature rise. If you squeeze a gas equilibrium (raise the pressure), it shifts to the side with fewer gas molecules, because fewer molecules take up less space and so relieve the pressure. If you add more of a substance, the equilibrium shifts to use it up, moving away from the side you added to.

A catalyst is the exception: it does not change the position of equilibrium at all. A catalyst speeds up the forward and backward reactions equally, so equilibrium is simply reached faster — the amounts of reactants and products at equilibrium are exactly the same as they would have been without it.

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. A catalyst changes only the speed, never the position.

flowchart TD
  A["A condition is changed at equilibrium"] --> B{"What changed?"}
  B -->|Heat added| C["Shift in the endothermic direction"]
  B -->|Pressure raised| D["Shift to the side with fewer gas molecules"]
  B -->|Concentration raised| E["Shift away from what was added"]
  B -->|Catalyst added| F["No shift: equilibrium reached faster only"]

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Higher Tier: Applying Le Chatelier to Ammonia Production

Higher tier only: A good way to see Le Chatelier's principle in action is the industrial manufacture of ammonia ($\text{NH}_3$NH3​), used to make fertilisers, from nitrogen (from the air) and hydrogen (from natural gas). The reaction is reversible, 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)

The conditions are chosen as a compromise between getting a good yield, getting it quickly, and keeping costs down: