What it covers: The balance point — dynamic equilibrium, the equilibrium constant, ICE tables, Le Châtelier's principle, and solubility equilibria.
Exam weight: About 7–9% of the AP Chemistry exam.
The big question: How can the reaction quotient be compared to the equilibrium constant to predict and explain the behavior of a chemical system at or approaching equilibrium?
Big Ideas covered: Chemical Effects (CE), Transformations (TRA).
Key topics at a glance
Dynamic Equilibrium
Forward rate = reverse rate. Concentrations stop changing, but the reaction never actually stops.
Q vs. K
Q < K → shifts forward (toward products). Q > K → shifts reverse (toward reactants). Q = K → already at equilibrium.
Properties of K
Reversing a reaction takes the reciprocal of K. Multiplying coefficients by n raises K to the nth power.
ICE Tables
Track Initial, Change, Equilibrium concentrations to solve for unknowns using K.
Le Châtelier's Principle
A stress (concentration, volume/pressure, temperature) shifts equilibrium to partially counteract the change.
Solubility Equilibria (Ksp)
Ksp = equilibrium constant for dissolving a slightly soluble solid. Use it to find molar solubility.
Common-Ion Effect
Adding an ion already in the equilibrium decreases molar solubility — it's Le Châtelier's principle applied to Ksp.
Free Energy & K
ΔG° = −RT ln K connects thermodynamic favorability to the size of an equilibrium (or Ksp) constant.
The key terms you must know
Dynamic equilibrium — forward and reverse rates are equal; the reaction hasn't stopped.
Reaction quotient (Q) / equilibrium constant (K) — same expression, but Q applies at any moment, K applies specifically at equilibrium.
Magnitude of K — K >> 1 favors products; K << 1 favors reactants.
ICE table — the standard tool for solving equilibrium concentration problems.
Le Châtelier's principle — equilibrium shifts to partially counteract an applied stress.
Solubility product constant (Ksp) — the equilibrium constant for dissolving a slightly soluble ionic solid.
Common-ion effect — solubility decreases when a shared ion is already present in solution.
Free energy of dissolution — links ΔG° to Ksp via ΔG° = −RT ln K.
Key themes to remember
Equilibrium is dynamic, never static. Reactions are still happening at equilibrium — they just cancel out at the macroscopic level.
Comparing Q to K always tells you the direction of shift. This single comparison underlies almost every prediction in this unit.
Pure solids and liquids don't appear in K or Q expressions. Only aqueous species and gases matter for heterogeneous equilibria.
Le Châtelier's principle is really just "the system re-establishes Q = K." Every stress changes Q relative to K momentarily, and the system shifts to restore equality.
The common-ion effect is Le Châtelier's principle, not a brand-new rule. Adding a shared ion increases Q relative to Ksp, so the system shifts to consume some of the added ion — precipitating more solid and lowering solubility.
Common exam traps
Don't include pure solids or liquids in K expressions. Their "concentration" is constant and folded into K itself.
Adding more of a pure solid or liquid doesn't shift equilibrium. Since they're not in the K expression, changing their amount has no effect on Q.
A catalyst speeds up reaching equilibrium but does NOT change K or shift equilibrium position. It speeds up forward and reverse rates equally.
Temperature changes actually shift K itself, not just Q. Unlike concentration/volume changes, temperature changes alter the value of the equilibrium constant.
Increasing volume (decreasing pressure) shifts equilibrium toward the side with more moles of gas — but only affects gas-phase equilibria with unequal moles of gas on each side.