Rucete ✏ SAT Chemistry In a Nutshell
10. Chemical Equilibrium
This chapter introduces the dynamic nature of equilibrium, how to write equilibrium expressions, and how to use Le Châtelier’s Principle to predict shifts. It also covers acid-base equilibrium, solubility product constants, and Gibbs free energy—crucial concepts for understanding reversible reactions.
Reversible Reactions and Equilibrium
- A reversible reaction is one where reactants and products continuously interconvert:
A + B ⇌ C + D - Equilibrium is reached when the forward and reverse reaction rates are equal.
- At equilibrium, the reaction is dynamic, not static—molecules continue reacting, but concentrations remain constant.
Law of Mass Action and Equilibrium Constant (Keq)
- The rate of a reaction is proportional to the product of the concentrations of the reactants.
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Example:
A₂ + B₂ ⇌ 2AB
Forward rate: R = k₁[A₂][B₂]
Reverse rate: R′ = k₂[AB]² -
At equilibrium:
k₁[A₂][B₂] = k₂[AB]²
→ Keq = k₁/k₂ = [AB]² / [A₂][B₂] -
For a general reaction:
aA + bB ⇌ cC + dD
→ Keq = [C]^c[D]^d / [A]^a[B]^b -
Keq is constant at a given temperature.
- Large Keq → mostly products
- Small Keq → mostly reactants
Sample Problem: Solving for Equilibrium Concentrations
- Given initial amounts and Keq, set up variables for change in concentration (use “x”) and solve.
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Example:
H₂ + I₂ ⇌ 2HI Keq = 45.9
Start: [H₂] = [I₂] = 3.00 mol/L
At equilibrium: [H₂] = [I₂] = 3.00 − x, [HI] = 2x
Plug into Keq and solve: Keq = (2x)² / (3.00 − x)²
Le Châtelier’s Principle
- If a system at equilibrium is disturbed, it shifts to counteract the change and restore equilibrium.
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Concentration changes:
- Adding more reactants → shifts right
- Removing products → shifts right
- Adding products → shifts left
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Pressure changes (only for gases):
- Increase pressure → shifts to the side with fewer moles of gas
- Decrease pressure → shifts to the side with more moles of gas
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Temperature changes:
- For endothermic reactions (ΔH > 0), heat is a reactant → increase T = shift right
- For exothermic reactions (ΔH < 0), heat is a product → increase T = shift left
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Catalysts:
- Affect rate of reaction but do not change equilibrium position or Keq
Heterogeneous Equilibrium
- Includes solids and liquids, but their concentrations do not appear in the Keq expression.
- Only gases and aqueous species are included in equilibrium calculations.
- Example:
CaCO₃(s) ⇌ CaO(s) + CO₂(g)
Keq = [CO₂]
Acid-Base Equilibrium (Ka and Kb)
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Weak acids and bases partially ionize and have equilibrium constants:
- Ka = [H⁺][A⁻] / [HA]
- Kb = [OH⁻][BH⁺] / [B]
- Strong acids/bases dissociate completely → no Ka or Kb is used.
- If Ka is large → stronger acid; if Ka is small → weaker acid
pH and pOH
- pH = −log[H⁺], pOH = −log[OH⁻]
- pH + pOH = 14 at 25°C
- pH < 7 → acidic, pH = 7 → neutral, pH > 7 → basic
- For weak acids, you may need to use Ka to find [H⁺] and then calculate pH.
Solubility Product Constant (Ksp)
- Ksp = equilibrium constant for dissolving slightly soluble salts.
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Example:
AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)
Ksp = [Ag⁺][Cl⁻] - A small Ksp means low solubility.
- Adding a common ion (already present in the solution) shifts equilibrium left → decreases solubility.
Gibbs Free Energy (ΔG)
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Determines whether a reaction is spontaneous:
- ΔG < 0 → spontaneous
- ΔG > 0 → non-spontaneous
- ΔG = 0 → equilibrium
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Related to enthalpy and entropy:
ΔG = ΔH − TΔS -
You can also calculate ΔG using equilibrium constant:
ΔG = −RT ln Keq
In a Nutshell
Chemical equilibrium is a dynamic balance between forward and reverse reactions. The position of equilibrium can be described using Keq and predicted with Le Châtelier’s Principle. Related ideas like Ka, Kb, Ksp, and ΔG show how equilibrium applies to acid-base systems, solubility, and spontaneity. These concepts form a foundation for understanding reversible chemical systems.