Definition: Le Chatelier's Principle states that when a system at equilibrium is disturbed, it will shift in a direction that counteracts the change.

Example: Temperature changes in exothermic reactions shift equilibrium opposite to the temperature change direction.

Highlight: Pressure changes affect equilibrium systems with different numbers of gas moles differently.

Vocabulary: Spontaneity (ΔG) determines the direction of equilibrium shift, with negative values indicating spontaneous processes.

Quote: "If external factors (pressure, temperature, or concentration) are modified in an equilibrium system, the system evolves in a direction that opposes the introduced modification."

Definition: Solubility equilibrium exists between a solid substance and its dissolved ions in a saturated solution.

Example: For a dissolution AB$s$ ⇌ A+(aq) + B-(aq), the solubility product constant Kps = A+$$B-.

Highlight: The common ion effect reduces solubility when an ion common to the dissolving salt is added to the solution.

Vocabulary: Kps (solubility product constant) represents the product of ionic concentrations in a saturated solution.

Quote: "When adding a common ion, the equilibrium shifts to form precipitate."

Definition: Chemical equilibrium is a state where forward and reverse reaction rates are equal, represented through equilibrium constants KC and KP.

Vocabulary: KC represents concentration-based equilibrium constant, while KP deals with partial pressures.

Example: For a reaction N₂O₄ + 2NO, the equilibrium expressions exclude solids, water, and liquids from calculations.

Highlight: The relationship between KC and KP is given by KP = KC(RT)^Δn, where Δn represents the change in number of moles.

Quote: "Neither solids, nor H₂O, nor liquids are included in equilibrium constant expressions."