Review AP Chem Unit 7 to build fluency with equilibrium constants, ICE tables, Le Chatelier's principle, and solubility equilibria. These concepts connect reversible reactions, Q vs. K reasoning, and Ksp calculations into one unified framework that appears throughout the rest of the course.
Use the topic guides, key terms, and practice questions available here to work through every concept from dynamic equilibrium to the common-ion effect.
Chemical equilibrium describes a reversible system where the forward and reverse reactions proceed at equal rates, so concentrations stay constant even though both reactions keep happening. Unit 7 builds from that conceptual foundation through quantitative tools like K expressions, ICE tables, and Ksp calculations, then ties everything together with Le Chatelier's principle.
Equilibrium is a dynamic state where forward and reverse reaction rates are equal. The equilibrium constant K describes the ratio of product to reactant concentrations at equilibrium. Comparing Q to K tells you which direction a reaction will shift. Le Chatelier's principle predicts how a system responds to concentration, pressure, or temperature changes. Ksp extends these ideas to the dissolution of slightly soluble salts.
The equilibrium constant expression Kc = [C]^c[D]^d / [A]^a[B]^b (or Kp using partial pressures) describes the system at equilibrium. Q has the same form but applies at any moment. When Q is less than K the reaction shifts toward products; when Q is greater than K it shifts toward reactants. Pure solids and liquids are always omitted from both expressions.
An ICE table organizes Initial concentrations, the Change in terms of a variable x, and the Equilibrium concentrations. You substitute the equilibrium row into the K expression and solve for x. When K is very small, the small-x approximation often avoids a quadratic. Always check whether the approximation is valid.
Adding a reactant, removing a product, decreasing volume in a gas-phase reaction, or changing temperature all disturb equilibrium. The system shifts to partially counteract the stress. Temperature changes are unique because they change the value of K itself, not just Q. Concentration and pressure changes shift Q without changing K.
Every equilibrium concept in Unit 7 follows from one idea: a reversible system reaches a state where opposing processes occur at equal rates. K quantifies that state, Q measures where the system is at any moment, Le Chatelier's principle describes how the system returns to balance after a disturbance, and Ksp applies the same logic to dissolving salts. Understanding this unifying idea lets you reason through any equilibrium problem rather than memorizing separate rules.
Equilibrium is a dynamic state where forward and reverse reaction rates are equal. Concentration, partial pressure, and rate-vs-time graphs all level off at equilibrium. Examples include evaporation/condensation, dissolution/precipitation, and acid-base proton transfer.
The direction a reversible reaction proceeds depends on whether the forward or reverse rate is larger. Net conversion of reactants to products occurs when the forward rate exceeds the reverse rate, and vice versa. Equilibrium is reached when the two rates are equal.
Q and K share the same mathematical form: products over reactants, each raised to its stoichiometric coefficient. Pure solids and liquids are excluded. Comparing Q to K predicts the direction of shift: Q < K means shift toward products, Q > K means shift toward reactants.
Substitute measured equilibrium concentrations or partial pressures into the K expression to calculate Kc or Kp. Stoichiometric coefficients become exponents. K is reported as a unitless number.
K >> 1 means the equilibrium mixture is mostly products. K << 1 means mostly reactants. K near 1 means significant amounts of both. K magnitude describes composition, not reaction speed.
Reversing a reaction inverts K. Multiplying coefficients by c raises K to the power c. Adding reactions multiplies their K values. These same rules apply to Q.
Use an ICE table with initial concentrations, changes expressed as multiples of x, and equilibrium values substituted into the K expression. Compare Q to K first to set the direction of x. Use the small-x approximation when K is very small and verify with the 5% rule.
Particulate diagrams show the relative numbers of reactant and product particles at equilibrium. More product particles indicate a larger K. Atom conservation must be maintained in any particle diagram.
A system at equilibrium shifts to partially counteract a stress. Stresses include adding or removing a species, changing volume or pressure in a gas-phase system, changing temperature, or diluting the system. Observable properties like color and pH can signal the direction of shift.
Any disturbance changes Q relative to K. Concentration and pressure changes affect Q only; temperature changes affect K itself. In both cases, the system shifts until Q equals K again.
Dissolution of a slightly soluble salt is modeled as an equilibrium with constant Ksp. Molar solubility is calculated from Ksp using an ICE table, but the math depends on the stoichiometry of the dissolution equation. Comparing Qsp to Ksp predicts precipitation.
A salt dissolves less in a solution already containing one of its ions. The common ion shifts the dissolution equilibrium back toward the solid, reducing molar solubility. Ksp itself does not change; only the equilibrium position shifts.
Review AP Chemistry free energy of dissolution, including entropy, enthalpy, Gibbs free energy, Delta G = Delta H - T Delta S, salt dissolution, solvent reorganization, ion-solvent attractions, thermodynamic favorability, and Ksp.
This snapshot uses Fiveable practice activity to show where students tend to miss questions and which review moves are worth prioritizing first.
Across 8.9k multiple-choice practice attempts for this unit.
Practice activity included in this snapshot.
Across 25 scored free-response attempts for this unit.
Review Common-Ion Effect with attention to how the concept appears in AP-style source and evidence questions.
Review Reaction Quotient and Equilibrium Constant with attention to how the concept appears in AP-style source and evidence questions.
Review Magnitude of the Equilibrium Constant with attention to how the concept appears in AP-style source and evidence questions.
Review Direction of Reversible Reactions with attention to how the concept appears in AP-style source and evidence questions.
Equilibrium is reached when the forward and reverse reaction rates become equal. At that point, concentrations and partial pressures stop changing, but both reactions continue. On a concentration-vs-time or rate-vs-time graph, equilibrium appears where the curves level off. The direction a reversible reaction proceeds depends entirely on which rate is currently larger.
Sketch a concentration-vs-time graph for a reversible reaction and label the point where equilibrium is established. Explain why the concentrations are constant but the reaction has not stopped.
The reaction quotient Q and equilibrium constant K share the same mathematical form: products over reactants, each raised to its stoichiometric coefficient. For aA + bB = cC + dD, Kc = [C]^c[D]^d / [A]^a[B]^b and Kp uses partial pressures. Pure solids and pure liquids are excluded from both expressions. To calculate K experimentally, substitute measured equilibrium concentrations or partial pressures directly into the expression.
Write the Kc expression for 2SO2(g) + O2(g) = 2SO3(g). Then write the Kp expression. Explain why you would omit a solid like CaCO3 if it appeared in the reaction.
The size of K tells you the composition of the equilibrium mixture. K much greater than 1 means the mixture is mostly products; K much less than 1 means mostly reactants. K values also follow predictable algebraic rules when reactions are manipulated: reversing a reaction inverts K, multiplying coefficients by a factor c raises K to the power c, and adding reactions together multiplies their K values. These rules apply equally to Q.
Given K1 for A = B and K2 for B = C, write the K expression for 2A = 2C. Show each algebraic step.
| Manipulation | Effect on K |
|---|---|
| Reverse the reaction | K becomes 1/K |
| Multiply coefficients by factor c | K becomes K^c |
| Add two reactions | K becomes K1 x K2 |
| No change to reaction | K stays the same |
ICE tables are the standard method for finding equilibrium concentrations when you know initial conditions and K. Set up Initial, Change (as +x or -x based on stoichiometry), and Equilibrium rows, then substitute into the K expression. When K is very small, the small-x approximation simplifies the algebra. Particulate diagrams represent equilibrium visually: the ratio of product particles to reactant particles reflects the magnitude of K.
For N2(g) + 3H2(g) = 2NH3(g) with given initial concentrations and Kc, set up the ICE table and identify whether you need the quadratic formula or can use the small-x approximation.
Le Chatelier's principle states that a system at equilibrium shifts to partially counteract any applied stress. Adding a reactant or removing a product shifts the reaction toward products; the reverse is also true. Decreasing volume in a gas-phase reaction increases pressure and shifts toward the side with fewer moles of gas. Temperature changes are fundamentally different: they change the value of K itself. For an endothermic reaction, increasing temperature increases K; for an exothermic reaction, increasing temperature decreases K. Every shift can be explained quantitatively by noting that the stress changes Q relative to K, and the system moves to restore Q = K.
For the exothermic reaction 2NO2(g) = N2O4(g), predict the direction of shift when (a) NO2 is added, (b) volume is decreased, and (c) temperature is increased. Justify each answer using Q vs. K reasoning.
| Stress | Effect on Q | Effect on K | Direction of shift |
|---|---|---|---|
| Add reactant | Q decreases | No change | Toward products |
| Remove product | Q decreases | No change | Toward products |
| Decrease volume (gas) | Q changes based on moles of gas | No change | Toward fewer moles of gas |
| Increase temperature (endothermic) | No immediate change | K increases | Toward products |
| Increase temperature (exothermic) | No immediate change | K decreases | Toward reactants |
The dissolution of a slightly soluble salt is a reversible process with its own equilibrium constant, Ksp, called the solubility product. For a salt MxAy dissolving into M^y+ and A^x- ions, the Ksp expression includes only the ion concentrations raised to their stoichiometric coefficients; the solid is omitted. Molar solubility s can be calculated from Ksp using an ICE table, but the relationship between s and Ksp depends on the stoichiometry of the dissolution. Comparing the ion product Qsp to Ksp predicts whether a precipitate will form.
Write the Ksp expression for Ca3(PO4)2. Set up an ICE table and express Ksp in terms of molar solubility s. Identify how the stoichiometry affects the relationship between s and Ksp.
When a salt is dissolved in a solution that already contains one of its ions, the solubility of the salt decreases. This is the common-ion effect. It follows directly from Le Chatelier's principle: the extra ion shifts the dissolution equilibrium back toward the solid. Quantitatively, place the initial concentration of the common ion in the ICE table and solve for the new molar solubility. The Ksp value itself does not change; only the solubility changes.
Calculate the molar solubility of AgCl in 0.10 M NaCl solution given Ksp = 1.8 x 10^-10. Compare your answer to the molar solubility in pure water and explain the difference.
Try stimulus-based AP practice questions and written prompts after you review the notes.
A timed, AP-style test built only from unit 7 questions — multiple choice plus free response, scored with AI feedback. Use the standard format or build your own.
| Term | Definition |
|---|---|
| Chemical Equilibrium | A dynamic state in which the forward and reverse reactions occur at equal rates, so the concentrations of all species remain constant over time. |
| Dynamic Equilibrium | The condition where forward and reverse reaction rates are equal and both reactions continue, resulting in no net change in concentrations. |
| Reversible Reaction | A reaction that can proceed in both the forward and reverse directions, represented by a double arrow and capable of reaching equilibrium. |
| Equilibrium Constant Expression | The ratio of product concentrations to reactant concentrations at equilibrium, each raised to its stoichiometric coefficient; pure solids and liquids are excluded. |
| K Value | The numerical value of the equilibrium constant; K >> 1 indicates a product-favored equilibrium, K << 1 indicates a reactant-favored equilibrium. |
| Partial Pressure | The pressure exerted by a single gas in a mixture; used in Kp expressions for gas-phase equilibria. |
| Particulate Diagram | A molecular-level drawing showing the relative numbers of reactant and product particles at equilibrium; the particle ratio reflects the magnitude of K. |
| Molar Solubility | The number of moles of a slightly soluble salt that dissolve per liter of solution to form a saturated solution; calculated from Ksp using an ICE table. |
| Solubility product constant (Ksp) | The equilibrium constant for a slightly soluble salt dissolving into its constituent ions; equals the product of ion concentrations each raised to their stoichiometric coefficients. |
| Precipitation | The formation of an insoluble solid when the ion product Qsp exceeds Ksp, driving the dissolution equilibrium back toward the solid. |
| Stoichiometric Coefficients | The numbers in a balanced equation that become exponents in the K expression and determine how molar solubility relates to ion concentrations in Ksp problems. |
| Dissociation | The separation of an ionic compound into its constituent ions when dissolved in water; the basis for writing Ksp expressions. |
Pure solids and pure liquids have concentrations that do not depend on amount, so they are always excluded from Kc, Kp, and Ksp expressions. A common error is including the dissolving solid in a Ksp expression or including water in an aqueous equilibrium.
Only a temperature change alters the value of K. Adding or removing a species, changing pressure, or diluting the solution changes Q but leaves K the same. Saying that adding a reactant increases K is incorrect.
Always compare Q to K before writing the change row. If Q > K, x represents a decrease in products and an increase in reactants. Skipping this step leads to a negative equilibrium concentration, which signals an error.
The approximation that x is negligible compared to the initial concentration is only valid when K is very small relative to the initial concentration. Always verify using the 5% rule: x divided by the initial concentration must be less than 0.05.
For a salt like Ag2CrO4 that dissolves as Ag2CrO4(s) = 2Ag+(aq) + CrO4^2-(aq), [Ag+] = 2s, not s. Forgetting the coefficient leads to a Ksp expression of s^3 instead of the correct 4s^3.
AP Chemistry free-response questions frequently ask you to explain why a system shifts in a particular direction after a stress. A complete answer names the stress, states how it changes Q relative to K, and concludes with the direction of shift. Saying only that Le Chatelier's principle predicts a shift is not sufficient without the Q vs. K justification.
Calculation tasks in Unit 7 often require writing a correct K expression, setting up an ICE table, and solving for an unknown concentration or Ksp. Errors in the K expression, such as including a solid or using the wrong exponent, propagate through the entire calculation. Show each step clearly and include units where appropriate.
The AP Chemistry exam uses multiple representations of equilibrium, including graphs, particulate diagrams, and symbolic equations. You may be asked to identify the equilibrium state from a concentration-vs-time graph, draw a particulate model consistent with a given K value, or connect a particle-level description to a macroscopic observation such as a color change or pH shift.
Given any balanced equation, write Kc or Kp correctly: products over reactants, coefficients as exponents, pure solids and liquids excluded. Practice with heterogeneous equilibria like CaCO3(s) = CaO(s) + CO2(g).
Given initial concentrations or partial pressures, calculate Q, compare it to K, and state whether the reaction shifts toward products, toward reactants, or is already at equilibrium.
Organize initial, change, and equilibrium rows correctly. Choose the sign of x based on Q vs. K. Decide whether to use the small-x approximation or the quadratic formula, and verify the approximation if used.
Reverse a reaction, scale coefficients, or combine reactions and adjust K correctly in each case. Show each step when combining multiple equilibria into an overall K.
Distinguish between concentration stresses (change Q, not K), pressure/volume stresses (affect gas-phase reactions based on moles of gas), and temperature stresses (change K itself). Justify each prediction with Q vs. K reasoning.
Write the dissolution equation, set up an ICE table with s as the variable, and solve. Account for stoichiometry: for CaF2, [F-] = 2s, so Ksp = (s)(2s)^2 = 4s^3.
Place the initial common-ion concentration in the ICE table, solve for the new molar solubility, and explain the reduction using Le Chatelier's principle. Confirm that Ksp is unchanged.
Read the topic guides for 7.1 and 7.2 to build the conceptual foundation, then work through 7.3 and 7.4 to practice writing and calculating K and Q expressions. Focus on getting the form of the expression right before moving to calculations.
Review 7.5 and 7.6 together. Practice interpreting K values qualitatively and manipulating K algebraically by reversing reactions, scaling coefficients, and combining equilibria. Use the comparison table in the review notes as a quick reference.
Work through 7.7 and 7.8 by setting up ICE tables for a variety of reactions. Practice both the small-x approximation and the quadratic approach. Sketch particulate diagrams and check that your particle ratios match the K value.
Review 7.9 and 7.10 together. For each stress type, concentration, pressure/volume, and temperature, practice predicting the shift and justifying it with Q vs. K reasoning. Use the comparison table in the review notes to check your logic.
Work through 7.11 and 7.12 by writing Ksp expressions, calculating molar solubility with ICE tables, and solving common-ion problems. Pay close attention to stoichiometry in the dissolution equation. Use the AP score calculator to estimate how your practice performance maps to an AP score.
Open the individual guides for Unit 7 when you want a closer review of one topic.
Use AP-style practice after you review the notes so you can check what you understand.
Practice free-response reasoning and compare your answer with scoring guidance.
Watch past review streams filtered to Unit 7 when you want a video walkthrough.
Open the Fiveable one-page unit review, then explore visual cheatsheets for a quick refresher.
Estimate your broader AP score goal after you review the course and exam format.
Print the unit cheatsheets ($3) or unit exams ($5) for Unit 7. Yours to keep, no subscription needed.
get printablesAP Chem Unit 7 covers 12 topics on equilibrium: Introduction to Equilibrium, Direction of Reversible Reactions, Reaction Quotient and Equilibrium Constant, Calculating the Equilibrium Constant, Magnitude of the Equilibrium Constant, Properties of the Equilibrium Constant, Calculating Equilibrium Concentrations, Representations of Equilibrium, Introduction to Le Châtelier's Principle, Reaction Quotient and Le Châtelier's Principle, Introduction to Solubility Equilibria, and Common-Ion Effect. The unit builds from the concept of dynamic equilibrium up through predicting how systems shift under stress. You'll work with K expressions, ICE tables, Q vs. K comparisons, and Ksp calculations. See AP Chem Unit 7 for topic-by-topic breakdowns.
Unit 7 makes up 7-9% of the AP Chem exam. That weight covers everything from writing equilibrium constant expressions and calculating K to applying Le Châtelier's principle and solving solubility equilibria problems. It's a focused unit, but the concepts show up in calculation-heavy multiple-choice and free-response questions.
The AP Chem Unit 7 progress check includes MCQ and FRQ parts drawn from all 12 equilibrium topics. MCQ questions test your ability to interpret Q vs. K, predict reaction direction, and identify how Le Châtelier's principle applies to concentration, temperature, and volume changes. FRQ questions typically ask you to set up ICE tables, calculate equilibrium concentrations, write Ksp expressions, or explain shifts using Le Châtelier's principle. The progress check pulls heavily from topics like Reaction Quotient and Equilibrium Constant (7.3), Calculating Equilibrium Concentrations (7.7), Introduction to Le Châtelier's Principle (7.9), and Introduction to Solubility Equilibria (7.11). Practicing those topics first gives you the most progress check coverage. Head to AP Chem Unit 7 for matched practice on each topic.
AP Chem Unit 7 FRQs most often come from three areas: calculating equilibrium concentrations using ICE tables, applying Le Châtelier's principle to explain system shifts, and solving solubility equilibria problems with Ksp. Questions usually ask you to show your setup, not just a final answer, so writing out every step of an ICE table or Ksp calculation matters. To practice effectively, work through past College Board FRQs that involve equilibrium constant expressions and reaction quotient comparisons. Focus on explaining your reasoning in full sentences when the prompt says 'justify' or 'explain.' You can find topic-aligned FRQ practice at AP Chem Unit 7.
The best place to find AP Chem Unit 7 practice questions, including multiple-choice and practice test sets, is AP Chem Unit 7. That page organizes practice by topic so you can target equilibrium constant calculations, reaction quotient problems, Le Châtelier's principle scenarios, and solubility equilibria separately. For MCQ practice, look for questions that give you a reaction and ask whether Q is greater than, less than, or equal to K. For a practice test experience, work through a full set of Unit 7 questions timed, then review any ICE table or Ksp problems you missed.
Start AP Chem Unit 7 by building a solid understanding of what equilibrium means physically: a dynamic state where forward and reverse reaction rates are equal. Once that clicks, the math follows more naturally. Here's a practical study order: 1. **Learn K expressions first.** Practice writing equilibrium constant expressions for gases and solutions before touching calculations. 2. **Master the Q vs. K comparison.** Knowing whether the reaction quotient Q is less than, greater than, or equal to K tells you which direction a reaction shifts. This shows up constantly. 3. **Work ICE tables by hand.** Set up and solve at least 10 ICE table problems for Calculating Equilibrium Concentrations (7.7). Speed and accuracy here pay off on FRQs. 4. **Study Le Châtelier's principle with real examples.** Practice predicting shifts for concentration changes, temperature changes, and volume changes separately. 5. **Finish with solubility equilibria.** Ksp problems and the Common-Ion Effect (7.12) build on everything above, so save them for last. Visit AP Chem Unit 7 for topic-by-topic resources that follow this order.