AP®︎ Chemistry: Unit 7 Flashcards
Master key terms and definitions for Unit 7 of AP Chemistry – Equilibrium to help you prep for quizzes and the AP exam.
Forward Reaction and Reverse Reaction
The forward direction forms products from reactants; the reverse direction reforms reactants from products.
The forward direction forms products from reactants; the reverse direction reforms reactants from products.
The symbol showing a reaction is reversible and can establish equilibrium.
Examples include evaporation-condensation, dissolution-precipitation, gas absorption-desorption, acid-base, and redox reactions.
A system where matter does not enter or leave, allowing equilibrium to be established.
At equilibrium, rates are equal while concentrations remain constant but are not usually equal.
Graphs show concentrations leveling off and forward and reverse rates becoming equal over time.
A reaction that can proceed in both the forward and reverse directions under the same conditions.
A state in which forward and reverse reactions continue at equal rates, so the amounts or concentrations of reactants and products remain constant.
Rates become equal over time, while reactant and product concentrations level off and remain constant.
Whichever reaction rate is greater determines the net change toward products or reactants.
The size of the equilibrium constant shows which side is favored at equilibrium: K > 1 means product-favored, K < 1 means reactant-favored, and K = 1 means neither side is strongly favored; larger K values indicate equilibrium lies farther toward products.
They are omitted because their effective concentrations remain constant during the reaction.
Compare the reaction quotient, Q, to the equilibrium constant, K, to predict direction: if Q < K, the reaction shifts toward products; if Q > K, it shifts toward reactants; if Q = K, the system is at equilibrium.
Write Q or K as products over reactants, with each concentration or partial pressure raised to its stoichiometric coefficient.
Substitute equilibrium concentrations or partial pressures into the correct expression and solve for K.
Include gases and aqueous species; omit pure solids and pure liquids from the expression.
Divide moles by solution volume in liters to get equilibrium molarity for substitution into Kc.
Find a gas partial pressure with PA = XA times Ptotal before using it in Kp.
The equilibrium constant is treated as having no units in AP Chemistry.
Kc uses equilibrium molar concentrations, and Kp uses equilibrium partial pressures of gases; both are interpreted the same way when judging whether products or reactants are favored.
Q and K have the same mathematical form, but Q uses current concentrations or partial pressures, whereas K uses equilibrium values only; at equilibrium, Q = K.
The reaction with the larger K lies farther to the product side at equilibrium.
A very large value means nearly complete reaction, while a very small value means little reaction.
Product and reactant terms in the equilibrium expression are equal in magnitude.
Manipulate each step as needed, then combine their K values to match the overall reaction.
Because they have the same mathematical form, inversion, exponentiation, and multiplication rules apply to both.
Reversing inverts K, scaling coefficients raises K to that power, and adding reactions multiplies K values.
A setup listing initial, change, and equilibrium amounts to solve for unknown equilibrium concentrations.
Initial lists starting amounts, change uses stoichiometric x values, and equilibrium combines initial plus change.
Species change by coefficient multiples of x, with reactants negative and products positive.
Substitute equilibrium concentrations or partial pressures from the table into the K expression.
Use the K equation after substitution, then solve algebraically for the unknown change variable.
Use molar concentrations for Kc problems and partial pressures for Kp problems.
Use provided starting amounts, and assign zero to species not initially present unless stated otherwise.
Treat a ± x term as constant when x causes less than 5% error.
A diagram that represents matter as discrete atoms, ions, or molecules.
Reactant and product particles remain present in constant relative amounts with no net change.
Right shift increases product particles; left shift increases reactant particles between two representations.
Particle numbers change according to reaction coefficients when reactants convert to products or reverse.
The total number of each type of atom stays constant as particles rearrange.
Larger product-to-reactant particle ratios indicate larger K values at the same temperature.
A system at equilibrium shifts to oppose an imposed change and reestablish equilibrium.
An external change such as concentration, temperature, pressure, volume, or dilution disturbance.
Adding a species shifts away from it; removing one shifts toward it.
If another reaction consumes a component, equilibrium shifts to replace the removed substance.
Increasing temperature favors the endothermic direction; decreasing temperature favors the exothermic direction.
In endothermic reactions heat acts like a reactant; in exothermic reactions it acts like a product.
Adding a nonreactive gas at constant volume does not change equilibrium position.
Adding solvent lowers concentrations, shifting equilibrium toward the side with more dissolved particles.
A catalyst speeds both directions equally, so equilibrium position does not change.
Equilibrium shifts can change measurable properties such as pH, color intensity, or temperature.
Pressure or volume changes shift gas equilibria toward fewer or more moles, unless gas moles are equal.
Changing temperature changes the equilibrium constant, K; concentration and pressure changes do not change K at constant temperature.
Adding products increases Q, while adding reactants decreases Q, with K unchanged at constant temperature.
A stress makes Q differ from K, so the system shifts until they are equal again.
Uses partial pressures and stoichiometric exponents to predict pressure-driven shifts in gas equilibria.
The equilibrium constant for a sparingly soluble ionic compound dissolving into its ions.
A dynamic equilibrium between an undissolved salt and its dissolved ions in a saturated solution.
A solution containing the maximum dissolved solute possible at equilibrium under given conditions.
The number of moles of a substance that dissolve per liter of solution.
Include only dissolved ion concentrations, each raised to its stoichiometric coefficient; omit solids.
Ion coefficients in the dissolution equation determine concentration relationships used in Ksp calculations.
Larger Ksp usually means greater solubility, but comparisons require matching dissolution stoichiometry.
Salts with Ksp greater than 1 are generally soluble, while smaller values indicate limited solubility.
Use dissolution stoichiometry to relate ion concentrations and solubility in equilibrium calculations.
The maximum amount of solute that dissolves in a given amount of solvent.
A shared dissolved ion shifts dissolution toward solid, reducing a salt’s molar solubility.
Use the initial common-ion concentration in the Ksp expression to find the reduced solubility.
