Equilibrium Shifts and Dissolution: Applying Le Chatelier to Solubility in AP Chemistry

Le Chatelier’s Principle

A rule predicting that when an equilibrium system is stressed (by changes in concentration, pressure/volume, or temperature), it shifts in the direction that partially counteracts the stress and establishes a new equilibrium.

Dynamic Equilibrium

A state where forward and reverse reactions continue to occur but at equal rates, so macroscopic concentrations remain constant.

Equilibrium Shift (Right/Left)

A change in the equilibrium mixture’s composition (more products = shift right; more reactants = shift left); it does not mean the reaction goes to completion.

Stress (Equilibrium Disturbance)

A change imposed on an equilibrium system, typically involving concentration, pressure/volume (gases), or temperature, which drives the system to re-establish equilibrium.

Concentration Stress

A disturbance caused by adding/removing reactants or products; the system shifts to consume what was added or replace what was removed.

Pressure/Volume Stress (Gas-Phase)

A disturbance in gas equilibria where decreasing volume (increasing pressure) shifts equilibrium toward the side with fewer moles of gas; increasing volume shifts toward more moles of gas.

Inert Gas at Constant Volume

Adding an inert gas while keeping volume constant increases total pressure but does not change reactant/product partial pressures, so it does not shift equilibrium.

Endothermic Forward Reaction

A reaction where heat acts like a reactant; increasing temperature shifts equilibrium toward products (to consume added heat).

Exothermic Forward Reaction

A reaction where heat acts like a product; increasing temperature shifts equilibrium toward reactants (to consume the added heat by favoring the reverse direction).

Equilibrium Constant (K)

A temperature-dependent constant that equals the equilibrium value of the product-to-reactant ratio (with coefficients as exponents) for a given reaction.

Temperature and K

The idea that only changes in temperature change the equilibrium constant K; concentration/pressure/catalysts do not change K.

Catalyst (Equilibrium Context)

A substance that lowers activation energy and speeds up both forward and reverse reactions equally, helping equilibrium be reached faster without changing K or the equilibrium composition.

Reaction Quotient (Q)

An expression with the same form as K but using current (not necessarily equilibrium) concentrations or partial pressures to assess the system’s direction of change.

Q vs. K Comparison

Rule for predicting direction: if Q < K, net reaction proceeds forward (shift right); if Q > K, net reaction proceeds backward (shift left); if Q = K, the system is at equilibrium.

Net Reaction Direction

The overall (macroscopic) direction a system proceeds after a disturbance, determined by whether Q is less than, greater than, or equal to K.

Solubility Equilibrium

An equilibrium involving a sparingly soluble ionic solid dissolving into its ions in water, with undissolved solid typically present at equilibrium.

Solubility Product Constant (Ksp)

The equilibrium constant for dissolution of a sparingly soluble ionic solid, written using only aqueous ion concentrations raised to stoichiometric powers (solids are omitted).

Omitting Solids and Liquids in K Expressions

The rule that pure solids and pure liquids are not included in equilibrium constant expressions because their effective concentrations are constant.

Molar Solubility (s)

The number of moles of a solid that dissolve per liter to form a saturated solution; used with stoichiometry to relate ion concentrations to Ksp.

Ion Product (Qsp)

A reaction quotient for a dissolution/precipitation system, computed like Ksp but using current ion concentrations to predict dissolving vs. precipitating.

Precipitation Condition (Qsp > Ksp)

If Qsp exceeds Ksp, ions are too concentrated and precipitation occurs until Qsp is reduced back to Ksp.

Unsaturated Condition (Qsp < Ksp)

If Qsp is less than Ksp, the solution can dissolve more solid (net dissolution occurs) until equilibrium is reached.

Common-Ion Effect

The decrease in solubility of an ionic solid when a soluble compound containing a common ion is added, shifting the dissolution equilibrium left.

Selective Precipitation

A method where, upon adding an ion (like Cl−) to a mixture of cations, the salt that reaches Qsp = Ksp at the lowest added-ion concentration precipitates first.

Gibbs Free Energy–Equilibrium Links

Key relationships: ΔG = ΔG° + RT ln Q and ΔG° = −RT ln K; these connect Q vs. K comparisons to spontaneity (ΔG < 0 forward when Q < K).