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Calculations using Equilibrium Constants in Heterogeneous Equilibria

Learn the concept of activity in chemistry and how it differs from equilibrium constants, especially in heterogeneous equilibria. Discover how to predict reaction directions, use reaction quotients, solve equilibrium problems, and calculate equilibrium pressures and concentrations. Practice problems included.

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Calculations using Equilibrium Constants in Heterogeneous Equilibria

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  1. Lecture #18 Activity, Heterogeneous Equilibria and Calculations using the Equilibrium Constant Chemistry 142 B Autumn Quarter 2004 J. B. Callis, Instructor

  2. The Concept of Activity • The ‘true’ equilibrium constant expression does not simply involve pressure or concentration for a substance. • Instead, it uses the ratio of the equilibrium pressure (or concentration) for a given substance to a reference pressure (or concentration) for that substance. • This ratio is defined as the activity of the substance.

  3. Use of Activities in Equilibrium Expressions

  4. The Difference between K and Kp • In general K and Kp are not equal in value. • Exception: when Dn = 0.

  5. Heterogeneous Equilibria • Homogeneous equilibria take place in one phase. • Heterogeneous equilibria involve more than one phase.

  6. NH4NO2(s) = N2(g) + 2 H2O(g) is an example of a heterogeneous equilibrium The equilibrium constant for this reaction may be expressed as:

  7. What We Can Learn About a Reaction from Its Equilibrium Constant • The tendency of the reaction to occur: • A value of K > 1 favors products • A value of K < 1 favors reactants • However, the value of K says nothing about the speed of the reaction. • Whether a given set of concentrations represents an equilibrium condition. • The equilibrium position that will be reached for a given set of initial conditions.

  8. The Reaction Quotient, Q

  9. Which Direction? • If Q = K, the system is at equilibrium; the concentrations will not change. • If Q > K, the system will shift to the left by converting products back to reactants. • If Q < K, the system will shift to the right by converting reactants to products.

  10. Problem 18-1: Predict Reaction Direction using Q • For the following reaction: • CH4(g) + 2 H2S(g) = CS2(g) + 4 H2(g) • 1.00 mol CH4, 1.00 mol CS2, 2.00 mol H2S and 2.00 mol H2 are mixed in a 250 mL vessel at 960oC. At this temperature, K = 0.036. In which direction will the reaction proceed in order to reach equilibrium? • Solution:

  11. Solution: Predict Reaction Direction using Q (2) Calculating the value of Q: (3) Comparing Q and K:

  12. Solving Equilibrium Problems • Write the balanced equation for the reaction. • Write the equilibrium expression. • List the initial concentrations. • Calculate Q and determine the direction of shift to equilibrium. • Define the change needed to reach equilibrium and define the equilibrium concentrations. • Substitute the equilibrium concentrations into the equilibrium expression and solve for the unknown. • Check the solution by calculating K and making sure it is identical to the original K.

  13. Calculating Equilibrium Pressures and Concentrations – Problem 18-2 Consider the equilibrium: CO(g) + H2O(g) = CO2(g) + H2(g) If 0.250 mol CO and 0.250 mol H2O are placed in a 125 mL flask at 900 K, what is the composition of the equilibrium mixture if K = 1.56. Answer:

  14. Example 18-2: Continued (a) Construct the reaction table:

  15. Example 18-2: Continued (b) Substitute into the equilibrium expression:

  16. Answers to Problems in Lecture 18 • Q = 64. Comparing Q and K: Q > K, so the reaction goes to the left. Therefore reactants increase and products decrease their concentrations.

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