- Define and explain Chemical Equilibria in your own words
- Use key terms such as dynamic equilibrium accurately
- Apply what you have learned to new examples and questions
- Avoid the common mistakes learners make with this topic
This lesson focuses on Chemical Equilibria: apply Le Chatelier's principle and calculate equilibrium constants.
Apply Le Chatelier's principle and calculate equilibrium constants.
Key ideas
Equilibria respond to change
At dynamic equilibrium the forward and backward rates are equal, so concentrations stay constant even though both reactions continue. Le Chatelier's principle predicts the response to change: raising the temperature shifts equilibrium in the endothermic direction, and raising pressure favours the side with fewer moles of gas. Kc quantifies the position: a large Kc means the equilibrium lies far to the right.
Enthalpy changes are measured — or found with Hess's law
In a simple calorimetry experiment, the heat released warms a known mass of water: q = mcΔT, with c = 4.18 J/g/K for water. Dividing q by the moles of fuel burned gives ΔH in kJ/mol — but heat escapes to the apparatus and air, so experimental values are always less exothermic than data-book values. When a change cannot be measured directly, Hess's law comes to the rescue: the total enthalpy change is independent of the route, so you can add and subtract known enthalpy changes around an energy cycle to find the unknown one.
Key term — dynamic equilibrium: The state where forward and backward reactions occur at equal rates in a closed system, so concentrations stay constant.
For N₂ + 3H₂ ⇌ 2NH₃, equilibrium concentrations are [N₂] = 0.50, [H₂] = 1.20, [NH₃] = 0.80 mol/dm³. Calculate Kc.
Kc = 0.80² ÷ (0.50 × 1.20³) = 0.64 ÷ 0.864 = 0.741, units mol⁻² dm⁶.
Answer: Kc = 0.80² ÷ (0.50 × 1.20³) = 0.64 ÷ 0.864 = 0.741, units mol⁻² dm⁶.
- Saying a catalyst changes the position of equilibrium A catalyst speeds up both forward and backward reactions equally, so equilibrium is reached faster but Kc and the final concentrations do not change.
- Forgetting the minus sign on exothermic ΔH Exothermic enthalpy changes are always negative — heat leaving the system. Writing +209 kJ/mol for a combustion is simply wrong, however good the arithmetic.
Practice
High pressure shifts equilibrium to the right (fewer moles of gas: 4 → 2), raising yield. Low temperature would favour the exothermic forward reaction but make it too slow, so a moderate temperature with an iron catalyst is the compromise.
ΔHc = −25.0 ÷ 0.0200 = −1250 kJ/mol.
ΔHf = (−394) + 2(−286) − (−890) = −394 − 572 + 890 = −76 kJ/mol.
Particles move faster, colliding more often, but the key effect is that many more collisions have energy at or above the activation energy, so far more collisions are successful.
Quick check
Which of these best defines "dynamic equilibrium"?
A Maxwell–Boltzmann distribution is drawn at a higher temperature. Describe two changes to the curve.
- Chemical Equilibria: apply Le Chatelier's principle and calculate equilibrium constants.
- Equilibria respond to change: At dynamic equilibrium the forward and backward rates are equal, so concentrations stay constant even though both reactions continue.
- Kc: The equilibrium constant: for aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d ÷ ([A]^a[B]^b).
- Watch out for: saying a catalyst changes the position of equilibrium