AP Chemistry Unit 7: Building Blocks of Chemical Equilibrium

Chemical equilibrium

A state in a reversible reaction where the forward reaction rate equals the reverse reaction rate, so reactant/product concentrations (or partial pressures) stop changing over time.

Dynamic equilibrium

Equilibrium in which reactions continue in both directions at equal rates; macroscopic properties (concentration, color, pressure) remain constant even though molecules keep reacting.

Reversible reaction

A reaction that can proceed in both the forward (reactants → products) and reverse (products → reactants) directions.

Forward reaction rate

The speed at which reactants are converted into products; depends on reactant availability and conditions.

Reverse reaction rate

The speed at which products are converted back into reactants; increases as product concentrations build.

Equilibrium does NOT mean equal amounts

At equilibrium, reactant and product amounts do not have to be equal; equilibrium only requires equal forward and reverse rates.

Equilibrium does NOT mean the reaction stops

At equilibrium, molecular collisions and conversions continue; there is just no net change because forward and reverse rates match.

Equilibrium position

The relative amounts (equilibrium concentrations/partial pressures) of reactants and products at equilibrium at a given temperature.

Shift right

A net forward reaction (net formation of products) that occurs after a disturbance until equilibrium is re-established.

Shift left

A net reverse reaction (net formation of reactants) that occurs after a disturbance until equilibrium is re-established.

Equilibrium constant (K)

A temperature-dependent number that equals the ratio of products to reactants at equilibrium, with each term raised to stoichiometric coefficients.

Kc

Equilibrium constant written using equilibrium molar concentrations: Kc = ([products]^coefficients)/([reactants]^coefficients).

Kp

Equilibrium constant written using equilibrium partial pressures for gases: Kp = ((Pproducts)^coefficients)/((Preactants)^coefficients).

Reaction quotient (Q)

A “snapshot” ratio with the same form as K, calculated using current (not necessarily equilibrium) concentrations or partial pressures.

Q vs. K direction rule

Compare Q to K to predict net direction: Q < K → forward (toward products); Q > K → reverse (toward reactants); Q = K → at equilibrium (no net reaction).

Large K (products favored)

If K is very large, products predominate at equilibrium (though some reactants may remain).

Small K (reactants favored)

If K is very small, reactants predominate at equilibrium (though some products may be present).

Omitting solids and liquids in K and Q

Pure solids and pure liquids are not included in equilibrium expressions because their effective concentration (activity) is essentially constant.

Heterogeneous equilibrium

An equilibrium involving more than one phase (e.g., solid + gas), where pure solids/liquids are omitted from K and Q expressions.

Partial pressure (Pi)

The pressure contributed by a single gas in a mixture; used in Kp and Qp expressions (e.g., PCO2).

Δn (for Kp/Kc conversions)

Moles of gaseous products minus moles of gaseous reactants from the balanced equation; count only gases.

Kp = Kc(RT)^Δn

Relationship between Kp and Kc for gas-phase equilibria (R = ideal gas constant, T in kelvin, Δn counts gaseous moles only).

Reversing a reaction (effect on K)

If you reverse a balanced equation, the new equilibrium constant is the reciprocal: Krev = 1/K.

Scaling coefficients (effect on K)

If you multiply all coefficients in the balanced equation by n, the new equilibrium constant becomes Knew = K^n.

Adding reactions (effect on K)

If you add reactions to obtain an overall reaction, multiply their equilibrium constants: Koverall = K1 × K2 (after any needed reversing/scaling).