Reversible Reactions and Equilibrium
Understand how reversible reactions reach equilibrium and respond to changes
Dynamic Balance
Forward ⇌ Backward
Most reactions we study go to completion—reactants fully convert to products. But reversible reactions can proceed in both directions. Products can reform the original reactants. We show this with the ⇌ symbol instead of a single arrow.
Examples include: the thermal decomposition of ammonium chloride (NH₄Cl ⇌ NH₃ + HCl), hydrated copper sulfate losing water when heated (CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O), and the Haber process for making ammonia (N₂ + 3H₂ ⇌ 2NH₃).
In a closed system (nothing enters or leaves), a reversible reaction eventually reaches equilibrium. At this point, the rate of the forward reaction equals the rate of the backward reaction. Reactions haven't stopped—they continue in both directions at equal rates.
At equilibrium, the concentrations of reactants and products remain constant (but not necessarily equal). This is called dynamic equilibrium—"dynamic" because reactions are still happening, "equilibrium" because there's no net change.
Le Chatelier's Principle states: if a system at equilibrium is disturbed, it will shift to oppose the change and restore equilibrium. This helps us predict what happens when we change conditions.
Temperature: If you increase temperature, equilibrium shifts in the endothermic direction (to absorb the extra heat). For exothermic reactions, this means shifting backwards, reducing product yield.
Pressure: Increasing pressure shifts equilibrium towards the side with fewer gas molecules (to reduce pressure). For N₂ + 3H₂ ⇌ 2NH₃, high pressure favours products (2 moles) over reactants (4 moles).
Concentration: Adding more of a substance shifts equilibrium away from that substance. Adding N₂ to the Haber process shifts equilibrium right, making more NH₃.
Key Exam Point
Catalysts do NOT affect equilibrium position—they only help equilibrium be reached faster. Both forward and backward rates increase equally. Yield stays the same, but you get there quicker.
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
ΔH = -92 kJ/mol (exothermic forward reaction)
Industrial: 450°C (compromise)
Industrial: 200 atm (compromise)
50%
Reactants (N₂ + H₂)
50%
Products (NH₃)
Le Chatelier's Principle
If a system at equilibrium is disturbed, it will shift to oppose the change. Increase temperature? System shifts to absorb heat (endothermic direction). Increase pressure? System shifts to reduce pressure (fewer gas molecules).
Question:
The Contact Process produces sulfur trioxide: 2SO₂ + O₂ ⇌ 2SO₃ (ΔH = -196 kJ/mol). Explain why the industrial process uses 450°C rather than a lower temperature, even though lower temperatures give higher yields.
Answer:
The forward reaction is exothermic (ΔH negative). According to Le Chatelier's Principle, lower temperatures would shift equilibrium to the right, giving a higher yield of SO₃.
However, at low temperatures the rate of reaction is very slow—it would take too long to reach equilibrium to be economically practical.
450°C is a compromise temperature: the yield is still acceptable (about 96% with catalyst), but the rate is fast enough for industrial production. The catalyst (V₂O₅) also helps achieve equilibrium faster.
What symbol represents a reversible reaction?