Most chemical reactions don't go to completion. They reach a state where reactants and products coexist. That state is equilibrium. Unit 7 is about understanding how and why reactions reach equilibrium, what it means for the reaction to be at equilibrium, and how to calculate and predict equilibrium concentrations. You'll work with the equilibrium constant, Le Châtelier's principle, and solubility equilibria.
🎯 What You Need to Know for the Exam
Unit 7 makes up about 7-9% of the AP Chemistry exam. Focus your energy on these priorities:
- Reversible reactions and equilibrium: what it means when a reaction reaches equilibrium, and how to represent it with equations
- The equilibrium constant (K): what it tells you about a reaction, how to calculate it from experimental data, and what large or small values mean
- Reaction quotient (Q) and predicting direction: how to use Q to determine which way a reaction will shift to reach equilibrium
- Calculating equilibrium concentrations: using initial conditions, the equilibrium constant, and ICE tables to find unknown concentrations
- Le Châtelier's principle: predicting how a system responds to stresses like concentration, temperature, and pressure changes
- Solubility equilibria and Ksp: calculating molar solubility and predicting precipitation using the solubility product constant
What's in this review:
- 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
- Common-Ion Effect
- Study Tips for Unit 7
- Summary, Review Questions & Practice
Topic 7.1: Introduction to Equilibrium
Most chemical and physical processes are reversible. Water evaporates and condenses. Gases dissolve in and escape from solutions. Salts dissolve in water and then precipitate out. In all these cases, both the forward and reverse processes are happening at the same time.
When a reversible reaction first begins, the forward reaction dominates and products form. But as products accumulate, the reverse reaction starts to pick up speed. Eventually, the forward and reverse reactions proceed at the same rate. When this happens, the concentrations of all reactants and products stay constant. This is equilibrium. The key insight is that equilibrium is dynamic. The reactions don't stop. Both directions continue, but they balance each other out.
You can see equilibrium being established by plotting concentration or pressure versus time. At first, reactant concentration drops and product concentration rises. Then both curves flatten out and stay constant. That flat region is equilibrium. Graphs of reaction rate versus time show both forward and reverse rates. At equilibrium, these lines meet.
Key concepts to know:
- Reversible reaction: A reaction where both reactants and products can coexist. Processes can proceed in both directions.
- Equilibrium: The state where forward and reverse reaction rates are equal, and concentrations (or partial pressures) remain constant over time.
- Dynamic equilibrium: Reactants and products continue to react, but no net change occurs because forward and reverse rates are equal.
- Observable change: At equilibrium, you observe no change in concentration, color, pressure, or other measurable properties.
⚠ Watch out for:
Students often think that at equilibrium, the reaction has stopped. It hasn't. Both forward and reverse reactions continue at equal rates. Also, "equilibrium" doesn't mean equal concentrations of reactants and products. It means equal rates of reaction in both directions.
Topic 7.2: Direction of Reversible Reactions
The direction a reaction proceeds depends on comparing the rates of the forward and reverse reactions. If the forward reaction is faster, reactants convert to products and the reaction moves right. If the reverse reaction is faster, products convert back to reactants and the reaction moves left. Equilibrium occurs when these rates are equal.
This is simpler than it sounds. You don't need to calculate the actual rates. You just need to know how to compare them using the reaction quotient (Q) and the equilibrium constant (K). For now, focus on the principle: whichever direction has the faster rate is the direction the reaction will proceed.
Key concepts to know:
- Forward reaction rate: The rate at which reactants form products.
- Reverse reaction rate: The rate at which products reform reactants.
- Net reaction direction: When forward rate exceeds reverse rate, there's net conversion of reactants to products. When reverse rate exceeds forward rate, there's net conversion of products to reactants.
- Equilibrium condition: Reached when forward rate equals reverse rate, resulting in no net change in concentrations.
⚠ Watch out for:
Don't confuse "equal rates" with "equal concentrations." At equilibrium, forward and reverse rates are equal, but that doesn't mean [reactants] = [products]. The equilibrium concentrations depend on the equilibrium constant, not on equal amounts.
Topic 7.3: Reaction Quotient and Equilibrium Constant
The reaction quotient (Q) and the equilibrium constant (K) are expressions that describe the ratio of products to reactants in a reaction. They have the same mathematical form. For a generic reaction aA + bB ⇌ cC + dD, the equilibrium expression is:
K = [C]^c [D]^d / [A]^a [B]^b
The equilibrium constant K uses equilibrium concentrations. The reaction quotient Q uses whatever concentrations are present at any moment in the reaction. The Q expression is identical to the K expression, but Q changes as the reaction proceeds. K is a constant for a given reaction at a given temperature.
An important detail: the expression doesn't include pure solids or pure liquids. If you have a solid reactant or product, you leave it out of the Q and K expressions. The concentration of a solid doesn't change as it dissolves or precipitates, so it doesn't appear.
Key concepts to know:
- Reaction quotient (Q): Calculated using the mass action expression with the current concentrations (or partial pressures). Q varies as the reaction proceeds.
- Equilibrium constant (K): The value of Q when the system is at equilibrium. K is constant at a given temperature.
- Mass action expression: The ratio of product concentrations (raised to stoichiometric coefficients) to reactant concentrations.
- Kc and Kp: Kc uses molar concentrations; Kp uses partial pressures of gases.
- Exclusions: Pure solids and pure liquids are not included in K or Q expressions because their concentrations don't change.
⚠ Watch out for:
Remember the stoichiometric coefficients. If the equation is N2 + 3H2 ⇌ 2NH3, then K = [NH3]^2 / ([N2][H2]^3). The exponents come from the coefficients, not from the subscripts in the formula. Also, don't forget to omit solids and pure liquids from your expression. Note: conversion between Kc and Kp (using Kp = Kc(RT)^Δn) is excluded from the AP Exam. You should know that both forms exist and when each is used, but you won't be asked to convert between them.
Topic 7.4: Calculating the Equilibrium Constant
You calculate the equilibrium constant by measuring the concentrations (or partial pressures) of all species when the system reaches equilibrium, and then plugging those values into the equilibrium expression.
Here's the process: (1) Start with the balanced equation. (2) Write the equilibrium expression. (3) Measure or determine the equilibrium concentrations. (4) Substitute into the expression and calculate K.
The equilibrium constant is always the same for a given reaction at a given temperature, regardless of initial concentrations. Whether you start with pure reactants or a mixture of reactants and products, the equilibrium constant stays the same. This makes K a useful tool for predicting where the system will end up.
Key concepts to know:
- Equilibrium measurements: K is calculated from experimental measurements of equilibrium concentrations or partial pressures.
- Equilibrium expression: Use the balanced equation coefficients as exponents.
- Temperature dependence: K changes with temperature, but remains constant at a given temperature regardless of initial conditions.
⚠ Watch out for:
Make sure you're using equilibrium concentrations, not initial concentrations. If the problem gives you initial conditions and some changes, you need to calculate the equilibrium concentrations first using an ICE table, then plug those into K.
Topic 7.5: Magnitude of the Equilibrium Constant
The size of K tells you something crucial about the reaction. A very large K (like 10^6 or larger) means the reaction goes essentially to completion. Nearly all reactants convert to products, and at equilibrium, products dominate. A very small K (like 10^-6 or smaller) means the reaction barely proceeds at all. The system remains mostly reactants, with only a tiny amount of products formed.
If K is around 1, then reactants and products are present in similar amounts at equilibrium. For intermediate values of K, the position of equilibrium is somewhere between mostly reactants and mostly products.
This is a key conceptual insight: K tells you the direction the reaction favors. Large K favors products. Small K favors reactants. You don't need to do calculations to know the outcome. Just look at the magnitude of K.
Key concepts to know:
- Large K (>>1): Reaction proceeds nearly to completion. Products are favored at equilibrium.
- Small K (<<1): Reaction barely proceeds. Reactants are favored at equilibrium.
- K near 1: Reactants and products are present in comparable amounts at equilibrium.
- Prediction: The magnitude of K lets you predict whether the system will be mostly reactants, mostly products, or a mixture.
⚠ Watch out for:
Don't confuse a large K with a reaction that happens quickly. K tells you where equilibrium is, not how fast you get there. A reaction with K = 10^10 might reach equilibrium in seconds or take years. Speed and position are separate concepts.
Topic 7.6: Properties of the Equilibrium Constant
The equilibrium constant has algebraic properties that make it powerful. When you reverse a reaction, K becomes its reciprocal (1/K). If K for the forward reaction is 100, then K for the reverse reaction is 1/100 = 0.01. When you multiply the coefficients of a reaction by a factor, you raise K to that same power. If you double the coefficients, you square K. If you divide by 2, you take the square root of K.
When you add two reactions together, you multiply their K values. This is why K is useful: it lets you combine reactions algebraically and find the K for the overall reaction.
These properties follow directly from the mathematical form of the equilibrium expression. They're not arbitrary. Understanding them helps you manipulate equilibrium problems efficiently.
Key concepts to know:
- Reaction reversal: If K for A ⇌ B is 100, then K for B ⇌ A is 1/100.
- Coefficient multiplication: If you multiply all coefficients by a factor c, raise K to the power c.
- Adding reactions: When two reactions are added, Koverall = K1 × K2.
- Algebraic validity: All valid algebraic manipulations of K expressions work with Q expressions too.
⚠ Watch out for:
Be careful when multiplying reactions. If you reverse one and keep the other forward, you reverse the K of one reaction and multiply. For example, if Reaction 1: K1 = 10 and Reaction 2 (reversed): 1/K2 = 0.5, then the overall K = (1/10) × (0.5) = 0.05. Track which reactions you're reversing.
Topic 7.7: Calculating Equilibrium Concentrations
This is where the equilibrium constant becomes a practical tool. You're given initial concentrations and K, and you need to find the equilibrium concentrations. The standard method is an ICE table: Initial, Change, Equilibrium.
Start by writing the balanced equation and the equilibrium expression. Fill in the initial concentrations. Use stoichiometry to determine how the concentrations change as the reaction proceeds. If you know that concentration of A drops by x, then concentrations of B and C change by -x or +x depending on stoichiometry. Finally, calculate the equilibrium concentrations and use them to solve for any unknown.
The key comparison is Q versus K. If Q < K, the reaction will shift right (forward) to make more products. If Q > K, the reaction will shift left (reverse) to make more reactants. If Q = K, the system is at equilibrium and no shift occurs.
Key concepts to know:
- ICE table: Initial concentrations, Change in concentrations, Equilibrium concentrations.
- Stoichiometric relationships: If 1 mole of A reacts, then 2 moles of B are produced and 3 moles of C are consumed, based on coefficients.
- Q < K: Reaction shifts forward (right), forming more products.
- Q > K: Reaction shifts backward (left), forming more reactants.
- Q = K: System is at equilibrium. No net shift occurs.
⚠ Watch out for:
Track signs carefully. If a reactant decreases, its change is negative. If a product increases, its change is positive. Use the stoichiometry to make sure the changes are proportional to the coefficients. The most common mistake is getting the stoichiometric relationships wrong in the ICE table.
Topic 7.8: Representations of Equilibrium
You can describe an equilibrium system using words, equations, or particulate diagrams. A particulate representation shows reactant and product particles before and at equilibrium. The relative numbers of each type of particle give you a visual sense of the equilibrium position.
For example, if you have a container with N2O4 ⇌ 2NO2, a particulate diagram before equilibrium might show many N2O4 molecules and few NO2 molecules. At equilibrium, the diagram shows fewer N2O4 molecules and more NO2 molecules, depending on K.
These diagrams help you see that equilibrium is about the relative amounts present, not absolute numbers. A large K means more product particles relative to reactant particles at equilibrium. A small K means mostly reactant particles.
Key concepts to know:
- Particulate models: Show the relative numbers of reactant and product particles before and at equilibrium.
- K interpretation: A higher product-to-reactant particle ratio indicates a larger K.
- Visual equilibrium: Diagrams help illustrate that equilibrium is dynamic, not static.
⚠ Watch out for:
When interpreting particulate diagrams, count carefully. The diagram shows relative particle numbers, and you can use these to calculate K by determining the concentrations represented by the particle counts.
Topic 7.9: Introduction to Le Châtelier's Principle
Le Châtelier's principle says that when a system at equilibrium is disturbed, it will shift to counteract the disturbance. If you add a reactant, the system will consume some of it by shifting right. If you remove a product, the system will produce more by shifting right. If you increase temperature, the system will respond based on whether the reaction is exothermic or endothermic.
The stresses you can apply are: adding or removing species, changing temperature, changing volume or pressure (for gas systems), or diluting the solution. The system responds by shifting left or right to minimize the effect of the stress. It never fully cancels the stress, but it reduces it.
This principle is powerful because it lets you predict what will happen without calculations. You just need to think about which direction the system will shift to reduce the stress.
Key concepts to know:
- Le Châtelier's principle: A system at equilibrium shifts to counteract an applied stress.
- Types of stresses: Concentration changes (add/remove species), temperature changes, volume/pressure changes, dilution.
- System response: Shifts left (reverse) or right (forward) to minimize the disturbance.
- Observable effects: Changes in pH, temperature, color, or pressure indicate a shift.
⚠ Watch out for:
Le Châtelier's principle tells you which way a system shifts, but it doesn't tell you by how much or how fast. It also doesn't completely cancel the stress. If you add a reactant, some of it will be consumed, but the concentration of that reactant is still higher than before you added it.
Topic 7.10: Reaction Quotient and Le Châtelier's Principle
When you disturb an equilibrium system, Q changes. If the disturbance adds reactants or removes products, Q decreases and becomes less than K. The system shifts right (forward) to bring Q back to K. If the disturbance removes reactants or adds products, Q increases and becomes greater than K. The system shifts left (reverse) to bring Q back down to K.
Temperature is special. Changing temperature actually changes K. For an exothermic reaction, increasing temperature makes K smaller. For an endothermic reaction, increasing temperature makes K larger. The system responds by shifting to establish a new equilibrium with the new K value.
This is the mechanism behind Le Châtelier's principle. The system always tries to return to Q = K. If you change K (by changing temperature), the system shifts until Q again equals the new K.
Key concepts to know:
- Q shifts to K: When disturbed, a system shifts to bring Q back to K.
- Concentration changes: Affect Q but not K. System shifts to restore Q = K.
- Temperature changes: Change K itself. System shifts to establish new equilibrium with new K.
- Disturbance and response: Smaller Q than K means shift right; larger Q than K means shift left.
⚠ Watch out for:
Remember that concentration changes don't affect K, only Q. Temperature changes affect K. This is a critical distinction. A concentration change causes a shift to a new equilibrium with the same K. A temperature change shifts to a new equilibrium with a different K.
Topic 7.11: Introduction to Solubility Equilibria
When a salt dissolves in water, the dissolution process is reversible. The solid dissolves to form ions in solution, and ions precipitate back out to reform the solid. At equilibrium, both processes are happening at equal rates. The system reaches solubility equilibrium.
The solubility product constant (Ksp) is the equilibrium constant for this dissolution process. For a salt like AgCl, the dissolution is AgCl(s) ⇌ Ag+(aq) + Cl-(aq), and Ksp = [Ag+][Cl-]. The Ksp is always written with products (ions in solution) in the numerator and excludes the solid.
Ksp tells you how soluble a salt is. A large Ksp means the salt is very soluble (lots of ions in solution at equilibrium). A small Ksp means the salt is barely soluble (few ions in solution at equilibrium). You can calculate the molar solubility of a salt using Ksp. You can also use Ksp to predict whether a precipitate will form when two solutions are mixed.
Key concepts to know:
- Solubility equilibrium: The salt dissolves and precipitates at equal rates, establishing equilibrium between solid and ions.
- Ksp expression: For a salt MxAy, Ksp = [M^n+]^x[A^m-]^y. Solids and pure liquids are excluded.
- Solubility relationship: Ksp can be used to calculate molar solubility. For AgCl, if solubility is s, then Ksp = s^2.
- Precipitation prediction: If the ionic product Q > Ksp, a precipitate will form.
⚠ Watch out for:
Remember that Ksp only describes the equilibrium between the solid and the saturated solution. It doesn't tell you anything about the dissolution rate or how long it takes to reach equilibrium. Also, be careful with the exponents in the Ksp expression. They come from the stoichiometry of the dissolution, not from the subscripts in the formula.
Topic 7.12: Common-Ion Effect
The common-ion effect describes what happens when you add an ion that's already present in a solubility equilibrium. For example, if you have a saturated solution of AgCl, adding more Cl- ions will cause some AgCl to precipitate out and the solubility of AgCl to decrease.
This is an application of Le Châtelier's principle to solubility equilibria. When you increase the concentration of one of the ions (the common ion), Q temporarily exceeds Ksp, so the system shifts left and more precipitate forms. The solubility decreases because some of the salt must precipitate to restore the Ksp equilibrium.
You can calculate the effect using Le Châtelier's principle qualitatively, or you can set up an equilibrium expression with the new ion concentration and calculate the new solubility quantitatively.
Key concepts to know:
- Common-ion effect: Adding an ion that's already present in equilibrium decreases the solubility of the salt.
- Mechanism: The ion concentration increase shifts the equilibrium left, reducing solubility.
- Le Châtelier's application: Adding a common ion is a stress that shifts equilibrium to counteract it.
- Calculation: Use Ksp with the new ion concentration to calculate the new solubility.
⚠ Watch out for:
The common-ion effect only decreases solubility. It doesn't increase it. If you add a non-common ion (an ion not present in the dissolution equilibrium), it might change solubility, but that's a different effect and usually small. Focus on the common-ion effect as shown in the CED.
Study Tips for Unit 7
Unit 7 is mathematical and conceptual. Here are strategies that work:
Master ICE tables. Equilibrium calculations depend on them. Practice until you can set up an ICE table in your sleep. Get the stoichiometry right. Check your work by substituting back into the equilibrium expression.
Understand Q versus K. This is the heartbeat of equilibrium chemistry. Q tells you which direction the reaction will shift. Practice comparing them without calculating exact values.
Practice Le Châtelier with actual systems. Pick real reactions and predict what happens when you change temperature, pressure, or concentration. Use Le Châtelier to check your intuition before doing calculations.
Distinguish concentration changes from temperature changes. Concentration changes shift the system to the same K. Temperature changes shift to a different K. This distinction comes up constantly on the exam.
Use Ksp for precipitation problems. Set up the Ksp expression carefully. Make sure you know whether you're calculating Ksp from solubility, or solubility from Ksp. Both are tested.
StarSpark Practice Prompts:
- "Give me an ICE table problem with a multi-step equilibrium expression and ask me to find equilibrium concentrations"
- "Show me a Le Châtelier scenario with multiple stresses and ask me to predict the net effect"
- "Describe a solubility equilibrium problem where I need to calculate Ksp or solubility with the common-ion effect"
Summary, Review Questions & Practice
You've covered all the topics in Unit 7. Before you move on, test yourself with these scenario-based questions. If you can answer them confidently, you're in great shape for this section of the exam.
Review Questions: Test Yourself
- For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g), you measure equilibrium concentrations of [N2] = 0.1 M, [H2] = 0.3 M, and [NH3] = 0.6 M. Calculate Kc for this reaction at this temperature.
- A system contains H2, I2, and HI at equilibrium. When you add more I2 to the container, explain what happens to Q, how the system responds, and why this is consistent with Le Châtelier's principle.
- For the reaction A ⇌ 2B, you're told that K = 4. If you start with 1 mole of A in a 1 L container, use an ICE table to find the equilibrium concentrations of A and B.
- A solution is saturated with BaCO3. Calculate Ksp for BaCO3 if the molar solubility is 1.0 × 10^-5 M. Then, predict what happens to the solubility if you add more CO3^2- ions to the solution.
- For an exothermic reaction at equilibrium, describe what happens to K, to Q, and to the equilibrium position when you increase the temperature. Explain the mechanism using Q and K.
Want more practice? Paste these questions into StarSpark to generate a full quiz with explanations.
Explore the Full AP Chemistry Study Guide
Unit 7 is where equilibrium becomes your central tool for predicting chemical behavior. The equilibrium constant connects everything: it tells you the direction of reactions, the extent of reactions, how systems respond to stress, and how solids dissolve in water.
Check out the full AP Chemistry study plan to see how this unit connects to the rest of the course.
Other Unit Reviews:
- AP Chemistry Unit 1: Atomic Structure and Properties
- AP Chemistry Unit 2: Compound Structure and Properties
- AP Chemistry Unit 3: Properties of Substances and Mixtures
- AP Chemistry Unit 4: Chemical Reactions
- AP Chemistry Unit 5: Kinetics
- AP Chemistry Unit 6: Thermochemistry
- AP Chemistry Unit 8: Acids and Bases
- AP Chemistry Unit 9: Thermodynamics and Electrochemistry
For official AP Chemistry resources, visit apcentral.collegeboard.org.
This review is aligned with the AP Chemistry Course and Exam Description. AP is a registered trademark of the College Board, which was not involved in the production of this guide.
