AP Chemistry Unit 7 Review: Equilibrium

Every Unit 7 topic, the Q-vs-K reasoning the exam actually rewards, and the one stress that really changes K.

Unit 7 asks what happens when a reversible reaction stops changing on the outside while staying busy on the inside. It's 7–9% of the multiple-choice section across 12 topics and 13–15 class periods -- the longest pacing of any unit at that weighting -- and it introduces the reaction quotient, ICE tables, and Le Châtelier's principle that Unit 8's acid-base chemistry depends on entirely. Here's what's actually in it, topic by topic, verified against the current CED.

Unit 7 at a Glance

Unit 7 (Equilibrium) builds directly on the reversible elementary steps introduced in Unit 6 and, further back, the reaction-rate reasoning from Unit 5. See the full AP Chemistry units breakdown for how it fits alongside the other 8 units, and the Course and Exam Description for the full framework.

Introduction to Equilibrium

A reaction reaches dynamic equilibrium when the forward and reverse reaction rates become equal. Both reactions keep happening -- nothing stops -- but because they happen at the same rate, the concentrations of every species stop changing. Those equilibrium concentrations are not necessarily equal to each other; only the forward and reverse rates are equal.

Direction of Reversible Reactions

A reversible reaction reaches the same equilibrium state -- the same value of K, at a given temperature -- whether it's approached starting from pure reactants, pure products, or any mixture in between.

Reaction Quotient and Equilibrium Constant

For a general reaction, the equilibrium constant expression puts products over reactants, each raised to its stoichiometric coefficient: for aA + bB ⇌ cC + dD, K = [C]^c[D]^d / [A]^a[B]^b.

Critical rule: only gases and aqueous species appear in a K expression. Pure solids and pure liquids are omitted entirely, since their concentration (technically, activity) doesn't change as the reaction proceeds -- by convention it's treated as 1. This is the heterogeneous-equilibrium rule, and leaving a solid or liquid out of the expression on purpose is correct, not an oversight.

Calculating the Equilibrium Constant

Given a table of measured equilibrium concentrations (or partial pressures), K is calculated by substituting directly into the expression above. K_c (using concentrations) and K_p (using partial pressures) describe the same equilibrium and are related through the ideal gas law, but K_c is the form used throughout the rest of this unit.

Magnitude of the Equilibrium Constant

K's size tells you where the equilibrium mixture sits, not how fast it gets there. K ≫ 1 means products are strongly favored at equilibrium; K ≪ 1 means reactants are strongly favored; K ≈ 1 means significant amounts of both reactants and products are present at equilibrium.

Properties of the Equilibrium Constant

Because K comes directly from the balanced equation, manipulating the equation manipulates K in a predictable way: reversing a reaction takes the reciprocal, K_reverse = 1/K_forward; scaling every coefficient by n raises K to the nth power, K_scaled = Kⁿ; and adding two reactions together multiplies their K values, K_combined = K₁ × K₂ -- the same additive logic as Hess's Law, just multiplicative instead of additive because K is exponential where ΔH is linear.

Calculating Equilibrium Concentrations

An ICE table (Initial, Change, Equilibrium) is the standard way to organize an equilibrium calculation: write starting concentrations, the change each undergoes in terms of a single unknown x (scaled by stoichiometric coefficients), and add them to get equilibrium concentrations in terms of x. Substituting into the K expression gives an equation to solve for x.

When K is small (products barely form) and the initial concentration is much larger than K, x is often negligible compared to the initial concentration -- letting you simplify (initial − x) ≈ initial and avoid the quadratic formula entirely. Always check this assumption with the 5% rule: if x divided by the initial concentration is less than 5%, the approximation is valid. If it's 5% or more, the approximation isn't valid and the full quadratic formula must be used instead.

Representations of Equilibrium

Equilibrium shows up in two common graph types. A concentration-vs-time graph shows reactant concentrations decreasing and product concentrations increasing until every curve goes flat at equilibrium. A rate-vs-time graph shows the forward rate decreasing and the reverse rate increasing until they meet at a single, nonzero, equal value -- this second graph is the direct picture of "dynamic equilibrium": the rates become equal, they never become zero.

Introduction to Le Châtelier's Principle

A system at equilibrium that experiences a stress -- a change in concentration, volume/pressure, or temperature -- responds by shifting in the direction that partially counteracts the stress. A few conditions narrow when this actually applies: a volume/pressure change only shifts equilibrium when the moles of gas differ between the reactant and product sides (Δn_gas ≠ 0); adding an inert gas at constant volume changes total pressure but changes no concentration, so it does not shift equilibrium at all.

Not a shift: a catalyst speeds up both the forward and reverse reactions equally, so equilibrium is reached faster -- but the equilibrium position (and K) is completely unchanged. This is the same distinction Unit 5 draws between reaction rate and reaction extent: a catalyst is a kinetics tool, not an equilibrium tool.

Reaction Quotient and Le Châtelier's Principle

The rigorous version of Le Châtelier's principle -- and the version the AP Exam actually rewards over memorized "shifts left" / "shifts right" rules -- compares the reaction quotient Q (calculated with current, non-equilibrium concentrations) to K: if Q < K, the system shifts toward products to increase Q back up to K; if Q > K, the system shifts toward reactants to decrease Q back down to K.

Critical rule: concentration, volume, and pressure stresses only shift Q away from K -- the system then shifts to bring Q back to the same, unchanged K. Temperature is the only stress that changes the actual value of K. The easiest way to predict which direction: treat heat as a product in an exothermic reaction (or a reactant in an endothermic one), then apply the same Q-vs-K logic to that "heat" term.

Introduction to Solubility Equilibria

For a sparingly soluble ionic solid dissolving in water, the equilibrium constant is called K_sp (the solubility-product constant), and the same heterogeneous-equilibrium rule applies: the undissolved solid is omitted from the expression. K_sp can be used to calculate molar solubility (how many moles of the solid dissolve per liter), and comparing the reaction quotient Q to K_sp predicts whether a precipitate forms: if Q > K_sp, a precipitate forms; if Q < K_sp, the solution stays unsaturated and no precipitate forms.

Common-Ion Effect

Dissolving a salt in a solution that already contains one of its ions produces less dissolution than dissolving it in pure water -- the common-ion effect. It's Le Châtelier's principle applied directly to a solubility equilibrium: the extra common ion is a concentration stress that shifts the dissolution equilibrium backward, toward the undissolved solid, lowering the salt's molar solubility.

Common Mistakes in Unit 7

• Treating Le Châtelier's principle as memorized "shift left/shift right" rules. Use Q-vs-K reasoning instead -- and remember that a concentration, volume, or pressure change never changes K itself; only temperature does.
• Including a pure solid or liquid in a K expression. Only gases and aqueous species belong in the expression; solids and liquids are omitted by convention.
• Using the small-x approximation without checking the 5% rule. If x turns out to be 5% or more of the initial concentration, the approximation is invalid and the quadratic formula is required instead.
• Confusing a large K with a fast reaction. K's magnitude describes where the equilibrium mixture sits, not how quickly the system gets there -- that's a kinetics question from Unit 5, not an equilibrium one.
• Forgetting to raise each concentration to its stoichiometric coefficient. A missing exponent in a K expression is one of the most common algebra slips on this unit's free-response questions.

How Unit 7 Connects to the Rest of the Course

• K and the rate constant k from Unit 5 share similar notation but are otherwise unrelated -- equilibrium is defined by forward and reverse rates becoming equal, not by the reaction stopping.
• Whether a temperature increase shifts equilibrium toward products or reactants depends on the reaction's exothermic or endothermic sign from Unit 6 -- heat behaves like a reactant or product in the Q-vs-K comparison.
• Unit 8 (Acids and Bases) is equilibrium applied to one specific system: K_a, K_b, and K_sp all share the exact algebraic structure introduced here.

Related Resources

• AP Chemistry Units (All 9)
• AP Chemistry Particulate Diagrams
• AP Chemistry Unit 1 Review
• AP Chemistry Unit 2 Review
• AP Chemistry Unit 3 Review
• AP Chemistry Unit 4 Review
• AP Chemistry Unit 5 Review
• AP Chemistry Unit 6 Review
• AP Chemistry Unit 8 Review
• AP Chemistry Unit 9 Review
• AP Chemistry Course and Exam Description
• AP Chemistry FRQ Calculator

Frequently Asked Questions

Sourced from College Board's official AP Chemistry Course and Exam Description, Effective Fall 2024. This page describes the document's real content; it is not a copy of it and is not affiliated with or endorsed by College Board.