Equilibrium Constants

(a) Effect of Temperature on Kp and Kc for Exothermic and Endothermic Reactions

The equilibrium constant (Kp for gases, Kc for concentrations) depends on temperature.

Exothermic reactions (ΔH < 0):

• Increasing temperature decreases Kp/Kc.

• The equilibrium shifts left (towards reactants).

• Example: Combustion reactions.

Endothermic reactions (ΔH > 0):

• Increasing temperature increases Kp/Kc.

• The equilibrium shifts right (towards products).

• Example: Thermal decomposition of calcium carbonate.

Changes in temperature alter the value of Kp/Kc, unlike changes in concentration or pressure.

(b) Calculating Kp, Kc, and Equilibrium Quantities

Kc (Equilibrium Constant in Terms of Concentration)

General formula for a reaction:

aA + bB ⇌ cC + dD

Kc = [C]^c [D]^d / [A]^a [B]^b

Reactants / Products

Steps to calculate Kc:

1. Write the balanced equation.

2. Use the ICE table (Initial, Change, Equilibrium).

3. Substitute equilibrium concentrations into the Kc expression.

Kp (Equilibrium Constant in Terms of Partial Pressure)

For gases: Kp = (P_C)^c (P_D)^d / (P_A)^a (P_B)^b

Relationship between Kp and Kc:

K_p = K_c (RT)^Δn
Where Δn = (moles of gaseous products - moles of gaseous reactants).

Finding Equilibrium Quantities

• Use initial moles and reaction stoichiometry to set up an ICE table.

• Express changes in terms of x.

• Solve for x using Kc or Kp.

• Determine final equilibrium concentrations or partial pressures.

(c) Significance of the Magnitude of K and Its Relation to Equilibrium Position

K >> 1 (large equilibrium constant)

• Reaction favors products.

• Equilibrium lies far to the right.

• Almost complete conversion of reactants.

K << 1 (small equilibrium constant)

• Reaction favors reactants.

• Equilibrium lies far to the left.

• Very little product is formed.

K ≈ 1 (moderate value)

• Both reactants and products are present in significant amounts.

• Equilibrium is balanced between reactants and products.

Key Takeaways

• Larger K → More product at equilibrium.

• Smaller K → More reactant at equilibrium.

• K determines the extent of the reaction but not the rate (reaction kinetics are separate).