When a reversible reaction proceeds in both directions at equal rates, the concentrations of reactants and products remain constant. This is dynamic equilibrium — the reaction hasn't stopped, it's balanced.

aA + bB ⇌ cC + dD → forward rate ← backward rate At equilibrium: r_forward = r_backward

The Equilibrium Constant (Kc):

$$K_c = \frac{[\text{C}]^c[\text{D}]^d}{[\text{A}]^a[\text{B}]^b}$$

"If a system at equilibrium is subjected to a disturbance or stress, the equilibrium shifts in the direction that tends to nullify the effect of that stress."

Think of it as the system's self-correction mechanism.

✅ Key Rule:

• Exothermic reaction (ΔH < 0): ↑T → Kc decreases (favors reactants)
• Endothermic reaction (ΔH > 0): ↑T → Kc increases (favors products)

⚠️ Temperature is the ONLY factor that changes the value of K

📌 Example — Haber Process:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol (Exothermic)

• ↑ Temperature → Equilibrium shifts LEFT → Less NH₃ produced
• Industrially, a compromise temperature (~450°C) is used (rate vs. yield trade-off)

Pressure affects equilibrium only when Δn_gas ≠ 0

$$\Delta n_{gas} = \text{moles of gaseous products} - \text{moles of gaseous reactants}$$

📌 Examples:

Example 1 — Haber Process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) Δn = 2 − (1+3) = −2 → fewer moles on product side ↑ Pressure → shifts RIGHT → more NH₃ ✅

Example 2 — Decomposition of PCl₅: PCl₅(g) ⇌ PCl₃(g) + Cl₂(g) Δn = 2 − 1 = +1 → more moles on product side ↑ Pressure → shifts LEFT → less decomposition

Example 3 — H₂ + I₂ ⇌ 2HI: Δn = 2 − 2 = 0 → Pressure has NO effect

• TECS → Temperature changes Equilibrium Constant; Stress only shifts
• High P → Low n (pressure increase favors fewer gas molecules)
• Heat = Reactant for endothermic reactions (add heat = add reactant = shift right)

✦ Kp = Kc(RT)^Δn ✦ Q < K → forward reaction proceeds ✦ Q > K → reverse reaction proceeds
✦ Q = K → system at equilibrium