IB Chemistry How to Solve: IB Chemistry – Equilibrium (Kc, Kp, Le Chatelier, Acid-Base, Buffers)

How to Solve: IB Chemistry – Equilibrium (Kc, Kp, Le Chatelier, Acid-Base, Buffers)

Complete Guide

Introduction

"Mastering equilibrium lets you predict how reactions shift under stress—like how your blood buffers pH when you sprint, or how industries maximise ammonia yield. On the IB exam, this topic is worth 15-20% of Paper 2, and missing one step can cost you 5+ marks. Let’s break it down so you never lose those marks again."

WHAT YOU NEED TO KNOW FIRST

1. Dynamic equilibrium: Forward and reverse reactions occur at equal rates; concentrations remain constant.
2. Stoichiometry: Balancing equations and mole ratios.
3. Basic algebra: Rearranging equations and solving for unknowns.

KEY TERMS & FORMULAS

1. Equilibrium Constants (Kc and Kp)

2. Le Chatelier’s Principle

3. Acid-Base Equilibria

4. Buffers

STEP-BY-STEP METHOD

1. Calculating Kc or Kp

Step 1: Write the balanced equation. Step 2: Write the K expression (products over reactants, raised to coefficients). Step 3: Substitute equilibrium concentrations (Kc) or partial pressures (Kp). Step 4: Solve for the unknown (K or a concentration/pressure). Step 5: Check units (Kc has no units; Kp uses atm or Pa).

2. Predicting Shifts (Le Chatelier’s Principle)

Step 1: Identify the stress (concentration, pressure, temperature). Step 2: Determine if the reaction is exo/endothermic (if temperature change). Step 3: Predict the shift (left/right) to oppose the change. Step 4: State the effect on yield (e.g., "more product formed").

3. Calculating pH, pKa, or Buffer pH

Step 1: Write the dissociation equation (e.g., HA ⇌ H⁺ + A⁻). Step 2: Write the Ka or Kb expression. Step 3: Substitute known values (initial concentrations, Ka/Kb). Step 4: Solve for [H⁺] or [OH⁻]. Step 5: Calculate pH = -log[H⁺] or use Henderson-Hasselbalch for buffers.

WORKED EXAMPLES

Example 1 – Basic (Kc Calculation)

Question: For the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), equilibrium concentrations are: [N₂] = 0.50 mol/dm³, [H₂] = 1.50 mol/dm³, [NH₃] = 0.20 mol/dm³. Calculate Kc.

Step 1: Write the Kc expression: Kc = [NH₃]² / ([N₂][H₂]³)

Step 2: Substitute values: Kc = (0.20)² / (0.50 × (1.50)³)

Step 3: Calculate: - (0.20)² = 0.04 - (1.50)³ = 3.375 - Denominator = 0.50 × 3.375 = 1.6875 - Kc = 0.04 / 1.6875 = 0.0237 (no units)

What we did and why: We used the Kc formula, substituted equilibrium concentrations, and solved algebraically. No shortcuts—always write the expression first!

Example 2 – Medium (Le Chatelier’s Principle)

Question: For the reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = -198 kJ/mol, predict the effect of: a) Increasing pressure b) Increasing temperature c) Adding a catalyst

Step 1: Identify the stress: - a) Pressure ↑ - b) Temperature ↑ - c) Catalyst added

Step 2: Predict shifts: - a) Pressure ↑: Shifts to side with fewer moles of gas (right, 2 moles vs. 3 moles). - b) Temperature ↑: Shifts to absorb heat (left, since forward reaction is exothermic). - c) Catalyst: No shift (speeds up both directions equally).

What we did and why: We applied Le Chatelier’s Principle by analysing each stress separately. Always check if the reaction is exo/endothermic for temperature changes!

Example 3 – Exam-Style (Buffer pH Calculation)

Question: A buffer contains 0.10 mol/dm³ CH₃COOH (Ka = 1.8 × 10⁻⁵) and 0.20 mol/dm³ CH₃COO⁻. Calculate its pH.

Step 1: Write the Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA])

Step 2: Calculate pKa: pKa = -log(1.8 × 10⁻⁵) = 4.74

Step 3: Substitute values: pH = 4.74 + log(0.20 / 0.10)

Step 4: Solve: - log(2) = 0.30 - pH = 4.74 + 0.30 = 5.04

What we did and why: We used the Henderson-Hasselbalch equation because it’s a buffer. Never use Ka directly for buffers—this formula is faster!

COMMON MISTAKES

EXAM TRAPS

1-MINUTE RECAP

"Here’s the night-before cheat sheet:
1. Kc/Kp: Products over reactants, raised to coefficients. Only equilibrium values!
2. Le Chatelier: Concentration? Oppose the change. Pressure? Fewer gas moles. Temperature? Absorb heat. Catalyst? No shift.
3. Acid-base: Ka = [H⁺][A⁻]/[HA]. pH = -log[H⁺]. Buffers? Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]).
4. Common mistakes: Forgetting coefficients, using initial concentrations, ignoring units.
5. Exam traps: Q vs. K, pressure units, temperature effects on K.

Write out one Kc and one buffer problem tonight. You’ve got this!