Le Chatelier's Principle & Solubility Product
Explore Le Chatelier's principle for predicting equilibrium shifts and understand solubility product (Ksp) for precipitation reactions.
Key Topics & Instructions
▼- Le Chatelier's Principle: Understand how systems at equilibrium respond to stress.
- Equilibrium Shifts: Study effects of concentration, pressure, and temperature changes.
- Solubility Product: Learn Ksp calculations and precipitation predictions.
- Common Ion Effect: Understand decreased solubility in common ion solutions.
- Equilibrium Simulation: Apply stresses and observe equilibrium shifts.
- Solubility Analysis: Study precipitation and dissolution processes.
- Ksp Calculations: Calculate solubility products and predict precipitation.
- Review the explanations for understanding equilibrium principles.
Experiment 1: Le Chatelier's Principle
Apply different stresses to chemical equilibria and observe how the system responds according to Le Chatelier's principle.
Experiment 2: Solubility Product (Ksp)
Study solubility equilibria, calculate Ksp values, and observe precipitation and common ion effects.
Le Chatelier's Principle states that when a system at equilibrium is subjected to a stress, it will shift in a direction that minimizes the effect of that stress. For chemical equilibria: Concentration changes - increasing reactants shifts toward products, increasing products shifts toward reactants. Pressure changes - increasing pressure shifts toward side with fewer gas moles. Temperature changes - increasing temperature favors endothermic direction. Catalysts - speed up both forward and reverse reactions equally, no equilibrium shift. Solubility Product (Ksp) is the equilibrium constant for dissolution of sparingly soluble salts: MA(s) ⇌ M⁺(aq) + A⁻(aq), Ksp = [M⁺][A⁻]. If ion product > Ksp, precipitation occurs; if < Ksp, dissolution occurs; if = Ksp, equilibrium exists.
Equilibrium and Solubility Principles
Le Chatelier's principle predicts how equilibria respond to disturbances: Concentration changes: N₂ + 3H₂ ⇌ 2NH₃ - adding N₂ shifts right to consume added N₂. Pressure changes: N₂ + 3H₂ ⇌ 2NH₃ (4 gas moles → 2 gas moles) - increasing pressure shifts right toward fewer moles. Temperature changes: N₂ + 3H₂ ⇌ 2NH₃ ΔH = -92 kJ (exothermic) - increasing temperature shifts left (endothermic). Catalyst addition: No shift in equilibrium position, only faster attainment. Real-world applications: Haber process (high pressure, moderate temperature), Contact process (optimal temperature), industrial esterification (remove water to shift equilibrium).
For a general salt: AₐBₑ(s) ⇌ aAⁿ⁺(aq) + eBᵐ⁻(aq), Ksp = [Aⁿ⁺]ᵃ[Bᵐ⁻]ᵉ. AgCl: AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq), Ksp = [Ag⁺][Cl⁻] = 1.8×10⁻¹⁰. BaCO₃: BaCO₃(s) ⇌ Ba²⁺(aq) + CO₃²⁻(aq), Ksp = [Ba²⁺][CO₃²⁻] = 5.1×10⁻⁹. CaSO₄: CaSO₄(s) ⇌ Ca²⁺(aq) + SO₄²⁻(aq), Ksp = [Ca²⁺][SO₄²⁻] = 2.4×10⁻⁵. PbCl₂: PbCl₂(s) ⇌ Pb²⁺(aq) + 2Cl⁻(aq), Ksp = [Pb²⁺][Cl⁻]² = 1.7×10⁻⁵. To calculate solubility: Let s = molar solubility, then express ion concentrations in terms of s, substitute into Ksp expression, solve for s.
The common ion effect decreases solubility of a salt in solutions containing one of its ions. AgCl in pure water: Solubility = √Ksp = √(1.8×10⁻¹⁰) = 1.34×10⁻⁵ M. AgCl in 0.1 M NaCl: [Cl⁻] = 0.1 M, Ksp = [Ag⁺](0.1) = 1.8×10⁻¹⁰, [Ag⁺] = 1.8×10⁻⁹ M, solubility decreases 1000-fold. Applications: Qualitative analysis - precipitation of group ions, water softening - precipitation of Ca²⁺ and Mg²⁺, pharmaceutical formulations - controlled solubility. The effect is significant when common ion concentration is much larger than solubility.
The ion product (Qsp) is calculated like Ksp but using initial concentrations before equilibrium: Qsp < Ksp: Unsaturated, no precipitate, more salt can dissolve. Qsp = Ksp: Saturated, at equilibrium. Qsp > Ksp: Supersaturated, precipitation occurs until Qsp = Ksp. Example: Mix 100 mL of 0.001 M AgNO₃ and 100 mL of 0.001 M NaCl. [Ag⁺] = 0.0005 M, [Cl⁻] = 0.0005 M, Qsp = (5×10⁻⁴)(5×10⁻⁴) = 2.5×10⁻⁷, Ksp(AgCl) = 1.8×10⁻¹⁰, Qsp > Ksp, so precipitation occurs. The amount of precipitate depends on how much Qsp exceeds Ksp.
Solubility of salts containing basic anions increases in acidic solutions: CaCO₃: CaCO₃(s) ⇌ Ca²⁺ + CO₃²⁻, CO₃²⁻ + H⁺ ⇌ HCO₃⁻, HCO₃⁻ + H⁺ ⇌ H₂CO₃ → CO₂ + H₂O. Adding acid removes CO₃²⁻, shifting dissolution right. Mg(OH)₂: Mg(OH)₂(s) ⇌ Mg²⁺ + 2OH⁻, OH⁻ + H⁺ → H₂O. Adding acid removes OH⁻, increasing solubility. Ag₃PO₄: PO₄³⁻ protonates in acid, increasing solubility. Salts with acidic cations: Like Al³⁺, may have decreased solubility in basic solutions due to hydrolysis and precipitation of hydroxides.
Water treatment: Precipitation of heavy metals as hydroxides or sulfides using solubility principles. Analytical chemistry: Qualitative analysis scheme based on selective precipitation. Geology: Formation of mineral deposits through precipitation from saturated solutions. Medicine: Kidney stone formation (calcium oxalate, Ksp = 2.3×10⁻⁹) and prevention. Agriculture: Soil pH adjustment to control nutrient availability. Industrial processes: Precipitation in metallurgical extraction, pharmaceutical manufacturing. Understanding these principles enables control of chemical processes and environmental management.
