Reversible Reactions Explained: How Chemical Equilibrium Works

Reversible Reactions Explained: How Chemical Equilibrium Works

What a reversible reaction is

A reversible reaction can proceed both forward (reactants → products) and backward (products → reactants). In a closed system, both directions occur simultaneously.

Dynamic equilibrium

• Definition: Dynamic equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction.
• Consequence: Concentrations (or partial pressures) of reactants and products remain constant over time, though reactions continue at the molecular level.

Equilibrium constant

• For a general reaction aA + bB ⇌ cC + dD, the equilibrium constant (Kc for concentrations, Kp for pressures) is:

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  K = [C]^c [D]^d / ([A]^a [B]^b) 

• Meaning: K quantifies the ratio of product to reactant activities at equilibrium. Large K → products favored; small K → reactants favored.

Reaction quotient (Q)

• Q has the same form as K but uses current concentrations/pressures.
• If Q < K the reaction proceeds forward; if Q > K it proceeds in reverse; if Q = K the system is at equilibrium.

Le Châtelier’s Principle (predicting shifts)

If an equilibrium is disturbed, it shifts to oppose the disturbance:

• Concentration: Adding a reactant shifts equilibrium toward products; removing a product shifts toward products.
• Pressure (gaseous systems): Increasing pressure favors the side with fewer moles of gas; decreasing pressure favors the side with more moles.
• Temperature: Raising T favors the endothermic direction (K changes accordingly); lowering T favors the exothermic direction.
• Catalyst: Speeds both directions equally — no change to equilibrium position (only faster attainment).

Thermodynamic view

• Equilibrium corresponds to the minimum Gibbs free energy (ΔG = 0 at equilibrium).
• Relationship: ΔG° = −RT ln K. Temperature changes alter K according to van ’t Hoff relation.

Examples

• N2O4(g) ⇌ 2NO2(g): increasing T shifts right (endothermic dissociation); increasing pressure shifts left (fewer moles).
• N2(g) + 3H2(g) ⇌ 2NH3(g) (Haber process): high pressure favors NH3 formation; lower temperature favors NH3 (but slows rate), so industry uses compromise conditions and a catalyst.

Calculations (brief)

1. Write balanced equation and K expression.
2. Insert initial concentrations/pressures, compute Q.
3. If not at equilibrium, set change = ±x, apply stoichiometry, solve for x using K.
4. Check approximations and units.

Key practical points

• Catalysts speed approach to equilibrium but do not change yields.
• Only temperature changes K; concentration/pressure changes shift position but leave K unchanged.
• For heterogeneous equilibria, pure solids/liquids are omitted from the K expression.

If you want, I can (pick one):

1. solve a specific equilibrium calculation step‑by‑step, or
2. make a short study-sheet or flashcards for exam revision.