14-3

Chapter 14: Chemical Equilibrium

Introduction to Chemical Equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions in a closed system are equal, resulting in constant concentrations of reactants and products. The position of equilibrium can be affected by changes in concentration, pressure, volume, and temperature.

• Dynamic Equilibrium: Both forward and reverse reactions continue to occur, but there is no net change in concentrations.

• Equilibrium Constant (K): Quantifies the ratio of product and reactant concentrations at equilibrium. For gases, partial pressures are often used ($K_p$).

• Reaction Quotient (Q): Used to predict the direction in which a reaction will proceed to reach equilibrium.

Review of Calculating Equilibrium Constants

The equilibrium constant expression depends on the balanced chemical equation. For a general reaction:

$aA + bB \rightleftharpoons cC + dD$

The equilibrium constant ($K_c$ for concentrations, $K_p$ for partial pressures) is:

$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

$K_p = \frac{(P_C)^c (P_D)^d}{(P_A)^a (P_B)^b}$

• Solids and pure liquids are not included in the equilibrium expression.

• Only concentrations (mol/L) or partial pressures (atm) of gases and aqueous species are used.

Le Châtelier’s Principle

Definition and Statement

Le Châtelier’s Principle states that if a system at equilibrium is disturbed by a change in concentration, pressure, volume, or temperature, the system will shift its equilibrium position to counteract the disturbance.

Any change to a chemical reaction at equilibrium causes the reaction to proceed in the direction that reduces the effect of the change.

Effects of Concentration Changes

Changing the concentration of reactants or products will shift the equilibrium to restore balance.

• Adding Reactant: Shifts equilibrium toward products (right).

• Adding Product: Shifts equilibrium toward reactants (left).

• Removing Reactant: Shifts equilibrium toward reactants (left).

• Removing Product: Shifts equilibrium toward products (right).

Example: For $N_2O_4(g) \rightleftharpoons 2 NO_2(g)$, adding more $NO_2$ increases its concentration, causing the system to shift left to produce more $N_2O_4$.

Effects of Volume and Pressure Changes

For gaseous equilibria, changes in volume or pressure affect the equilibrium position depending on the number of moles of gas on each side of the equation.

• Decreasing Volume (Increasing Pressure): Shifts equilibrium toward the side with fewer moles of gas.

• Increasing Volume (Decreasing Pressure): Shifts equilibrium toward the side with more moles of gas.

• If both sides have equal moles of gas: No shift occurs.

• Adding an inert gas at constant volume: No effect on equilibrium position.

Example: For $N_2(g) + 3 H_2(g) \rightleftharpoons 2 NH_3(g)$, decreasing the volume shifts equilibrium to the right (toward fewer moles of gas).

Change	Effect on Equilibrium
Decrease volume	Shift to side with fewer moles of gas
Increase volume	Shift to side with more moles of gas
Add inert gas (constant V)	No change

Effects of Temperature Changes

Temperature changes affect equilibrium depending on whether the reaction is exothermic or endothermic.

• Endothermic Reaction ($\Delta H > 0$): Heat is a reactant. Increasing temperature shifts equilibrium toward products.

• Exothermic Reaction ($\Delta H < 0$): Heat is a product. Increasing temperature shifts equilibrium toward reactants.

Example: For $CO(g) + 3 H_2(g) \rightleftharpoons CH_4(g) + H_2O(g)$, $\Delta H = -206.2$ kJ/mol (exothermic). Increasing temperature decreases $K$ and shifts equilibrium to the left.

Temperature (K)	Equilibrium Constant, K
298	4.3 × 107
800	1.38 × 102
1000	2.54 × 101
1200	3.92

Additional info: As temperature increases for an exothermic reaction, the equilibrium constant decreases, favoring reactants.

Endothermic vs. Exothermic Reactions

• Endothermic: System absorbs heat ($\Delta H > 0$). Heat is a reactant.

• Exothermic: System releases heat ($\Delta H < 0$). Heat is a product.

Example Equations:

Endothermic: $A + B + \text{heat} \rightarrow C + D$

Exothermic: $A + B \rightarrow C + D + \text{heat}$

Application of Le Châtelier’s Principle: Worked Examples

• CaCO3(s) → CaO(s) + CO2(g): Adding CaO(s) does not affect equilibrium; $K = P_{CO_2}$.

• N2(g) + 3 H2(g) → 2 NH3(g): $K_p = \frac{P_{NH_3}^2}{P_{N_2} P_{H_2}^3}$

Concept Checks and Practice Questions

• What can be done to increase the partial pressure of N2O4?

  • Add more NO2 to the container.

  • Increase the pressure by compressing the mixture.

• How could you increase the concentration of products for $N_2(g) + 3 H_2(g) \rightleftharpoons 2 NH_3(g)$?

  • Increase volume, add N2, or increase temperature.

Summary Table: Effects of Changes on Equilibrium

Change	Direction of Shift
Add reactant	Toward products
Add product	Toward reactants
Remove reactant	Toward reactants
Remove product	Toward products
Decrease volume	Toward fewer moles of gas
Increase volume	Toward more moles of gas
Increase temperature (endothermic)	Toward products
Increase temperature (exothermic)	Toward reactants

Additional info:

• Adding an inert gas at constant volume does not affect the equilibrium position, but at constant pressure, it can change the volume and thus shift equilibrium.

• Le Châtelier’s Principle is a qualitative tool; quantitative predictions require calculation of $Q$ and $K$.