Key Concepts of Equilibrium Constants

Equilibrium constants (K) are crucial in understanding chemical reactions at balance. They show the ratio of products to reactants, helping predict how changes in conditions affect a reaction's direction and extent, which is key in Honors Chemistry.

1. Definition of equilibrium constant (K)

   • K is a numerical value that expresses the ratio of the concentrations of products to reactants at equilibrium.
   • It is specific to a given reaction at a specific temperature.
   • A large K value indicates a reaction that favors products, while a small K value indicates a reaction that favors reactants.

2. Relationship between K and reaction quotient (Q)

   • Q is calculated using the same expression as K but for non-equilibrium conditions.
   • If Q < K, the reaction will proceed forward to reach equilibrium.
   • If Q > K, the reaction will shift backward to reach equilibrium.

3. Equilibrium constant expressions for homogeneous and heterogeneous reactions

   • Homogeneous reactions involve reactants and products in the same phase (gas, liquid, or solid).
   • Heterogeneous reactions involve reactants and products in different phases; only gases and aqueous solutions are included in the K expression.
   • The equilibrium constant expression is derived from the balanced chemical equation.

4. Units of equilibrium constants

   • K is unitless for reactions where the total number of moles of gaseous products equals the total number of moles of gaseous reactants.
   • For reactions with different numbers of moles, K may have units (e.g., Molarity, atm) depending on the reaction's stoichiometry.

5. Relationship between K and Gibbs free energy (ΔG)

   • The relationship is given by the equation ΔG = ΔG° + RT ln(Q), where ΔG° is the standard Gibbs free energy change.
   • At equilibrium, ΔG = 0, leading to the equation ΔG° = -RT ln(K).
   • A negative ΔG° indicates a spontaneous reaction favoring products at equilibrium.

6. Le Chatelier's Principle and its effect on equilibrium

   • This principle states that if a system at equilibrium is disturbed, the system will shift to counteract the disturbance.
   • Changes in concentration, pressure, or temperature can shift the position of equilibrium.
   • The direction of the shift depends on the nature of the disturbance.

7. Calculation of equilibrium concentrations

   • Use an ICE table (Initial, Change, Equilibrium) to organize initial concentrations, changes during the reaction, and equilibrium concentrations.
   • Set up the equilibrium expression using the equilibrium concentrations to solve for unknowns.
   • Ensure that the stoichiometry of the balanced equation is applied correctly.

8. Kc vs. Kp (concentration vs. pressure-based equilibrium constants)

   • Kc is used for reactions expressed in terms of molarity (concentration).
   • Kp is used for reactions expressed in terms of partial pressures of gases.
   • The relationship between Kc and Kp is given by the equation Kp = Kc(RT)^(Δn), where Δn is the change in moles of gas.

9. Temperature dependence of equilibrium constants

   • K is temperature-dependent; changing the temperature will change the value of K.
   • For exothermic reactions, increasing temperature decreases K, while for endothermic reactions, increasing temperature increases K.
   • The van 't Hoff equation can be used to quantify the effect of temperature on K.

10. Equilibrium constants for reversible reactions

    • For a reversible reaction, the equilibrium constant for the forward reaction (Kf) is the inverse of the equilibrium constant for the reverse reaction (Kr), Kf = 1/Kr.
    • The overall equilibrium constant for a series of reactions can be calculated by multiplying the individual equilibrium constants.
    • The stoichiometry of the balanced equation affects the value of K; coefficients become exponents in the K expression.