AP Chemistry Unit 7: Equilibrium - ICE Tables | StarSpark

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

Study Tips for Unit 7

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.

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.