Chapter 6: Chemical Equilibrium

Vocabulary practice flashcards covering fundamental chemical equilibrium concepts, equilibrium expressions, gaseous equilibria, Le Chatelier's principle, and ICE table operations.

Dynamic Equilibrium

The state reached by a chemical reaction when the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in system composition.

Vapor Pressure

The pressure exerted by a vapor in dynamic equilibrium with its condensed phase in a closed container, which depends solely on temperature.

Equilibrium Constant ($K_c$)

A unitless number at a given temperature determined by the ratio of the equilibrium molar concentrations of products to reactants, with each raised to the power of its stoichiometric coefficient.

Equilibrium Constant Expression

For the general reaction $a\text{A} + b\text{B} + \text{\dots} \rightleftharpoons g\text{G} + h\text{H} + \text{\dots}$, the mathematical relation defined as $K_c = \frac{[\text{G}]^g [\text{H}]^h \dots}{[\text{A}]^a [\text{B}]^b \dots}$, where square brackets represent unitless equilibrium molar concentrations.

Equilibrium Constant in Terms of Partial Pressures ($K_p$)

An equilibrium constant expression defined for gaseous reactions using the equilibrium partial pressures (in $\text{atm}$) of reactants and products: $K_p = \frac{P_{\text{G}}^g P_{\text{H}}^h \dots}{P_{\text{A}}^a P_{\text{B}}^b \dots}$.

Relation Between $K_p$ and $K_c$

The equation relating gas-phase equilibrium constants: $K_p = K_c(RT)^{\Delta n_{\text{gas}}}$, where $\Delta n_{\text{gas}} = (g + h + \text{\dots}) - (a + b + \text{\dots})$ represents the change in moles of gas.

Heterogeneous Equilibria Rule

The principle that the equilibrium position does not depend on the amounts of pure solids or pure liquids; species with state symbols $(s)$ or $(l)$ are omitted from the equilibrium constant expression.

Reaction Quotient ($Q$)

A ratio calculated using the identical algebraic expression as the equilibrium constant, but evaluated using initial concentrations ($[\text{A}]_0$, $[\text{B}]_0$, etc.) to establish which direction a reaction will proceed to attain equilibrium.

Direction of Net Change when $Q_c < K_c$

The condition where the reactant concentration is too high, causing the net reaction to proceed to the right (toward products).

Direction of Net Change when $Q_c > K_c$

The condition where the product concentration is too high, causing the net reaction to proceed to the left (toward reactants).

Le Chatelier's Principle

The rule that when an equilibrium system is subjected to a disturbance in temperature, pressure, or concentration of a reacting species, it responds by establishing a new equilibrium that partially offsets the impact of the change.

Effect of Volume Reduction on Gas Equilibrium

A reduction in volume causes the system to undergo a net shift in the direction that produces fewer moles of gas.

Effect of Volume Increase on Gas Equilibrium

An increase in volume causes the system to undergo a net shift in the direction that produces more moles of gas.

Effect of Temperature on $K$ and $Q$

Changing the temperature alters the numerical value of the equilibrium constant $K$, while leaving the initial value of $Q$ unaffected.

Effect of Raising Temperature on Equilibrium

Increasing the temperature shifts the equilibrium position in the direction of the endothermic reaction (absorbing heat).

Effect of Lowering Temperature on Equilibrium

Decreasing the temperature shifts the equilibrium position in the direction of the exothermic reaction (releasing heat).

Catalyst Effect on Equilibrium

A catalyst increases the rate at which equilibrium is reached but has no effect on the condition or position of equilibrium, meaning Le Chatelier's principle does not apply.

ICE Table

A tabular accounting method used to calculate unknown equilibrium values by tracking Initial quantities, Changes, and Equilibrium quantities.

Inversion Rule for Equilibrium Constants

Reversing the direction of a chemical equation causes the inversion of its equilibrium constant: $K_{\text{reverse}} = \frac{1}{K_{\text{forward}}}$.

Coefficient Multiplier Rule for Equilibrium Constants

Multiplying the stoichiometric coefficients of a chemical reaction by a common factor $x$ raises the equilibrium constant to the power of $x$ ($K_{\text{new}} = K^x$).

Rule for Combining Equilibrium Constants

When individual chemical equations are added together to form an overall reaction, their equilibrium constants are multiplied together: $K_{\text{overall}} = K_1 \times K_2 \times K_3 \dots$.