Chemical Equilibrium (Ice Method)

CHEMICAL EQUILIBRIUM (ICE METHOD) Introduction • Chemical equilibrium occurs when opposing reactions are proceeding at equal rates. • The rate at which the products are formed from the reactants equals the rate at which the reactants are formed from the products. • As a result, concentrations cease to change, making the reaction appear to be stopped. How fast a reaction reaches this equilibrium state is a matter of kinetics. • An equilibrium state results when a reaction is reversible At equilibrium the concentrations of reactants and products is still changing, however, the rate of the forward reaction (kf) is equal to the rate of the reverse reaction (kr) in what is described as a dynamic equilibrium such that no change in their concentrations is observed. Thus, for equilibrium to occur, neither reactants nor products can escape from the system. • The law of mass action states the ratio of forward and reverse processes is described by the equilibrium constant KC which can be calculated using a knowledge of the equilibrium concentrations of reactants and products. The equilibrium constant expression depends only on the stoichiometry of the reaction, not on the mechanism. Objectives and Success Criteria • Mastering the application of the ICE table methodology to equilibrium problems. • Accurate solutions to problems involving reactant and product concentrations and equilibrium constants. Prerequisites • Stoichiometry • Kinetics 1 MODEL 1: The ICE Table A worked example: Initially 1.50 moles of N2(g) and 3.50 moles of H2(g) were added to a 1 L container at 700 °C. As a result of the reaction the equilibrium concentration of NH3(g) became 0.540 M. What is the value of the equilibrium constant for this reaction at the given temperature of 700 °C. N2(g) I. Write the Initial concentrations of reactants and products: C. Write the Change in concentration due to reaction using the given reaction stoichiometric coefficients: E. Write the reactant and product concentrations at Equilibrium. + 1.50 mol L‐1 −x 1.50 mol L‐1 –x 3H2(g) 2NH3 3.50 mol L‐1 0 mol L‐1 −3x +2x 3.50 mol L‐1 −3x +2x mol L‐1 We are now set to solve for the equilibrium constant KC using the equilibrium equation: eqn. 1 where the reactant and product concentrations should be expressed at equilibrium. The problem tells us that the equilibrium concentration of NH3 is 0.540 M, thus we can solve for the unknown ‘x’ [NH3]eq = +2x = 0.540 M x = 0.270 M we can now solve for KC 2 Key Questions Concentration (M) (i) In the above reaction we can monitor the change in concentration of reactants over time (just as we discussed when dealing with kinetics) and we can plot the data as follows: 3 2 1 a b c d Time Using the worked example from Model 1 label each data plot as either [H2], [N2] or [NH3]. Why does one data plot show an initial positive slope whereas the other two data plots show initial negative slopes? Why does the uppermost plot have a steeper initial slope than the middle plot. At which time a, b, c or d is an equilibrium state reached? (ii) The equilibrium constant for this reaction was calculated as shown below. What is Kc for the reverse reaction? 3 EXERCISES 1. Initially, 1.0 mol of NO(g) and 1 mol of Cl2(g) were added to a 1 L container. As a result of the reaction the equilibrium concentration of NOCl(g) became 0.96 M. Using the RICE table methodology determine the value of the equilibrium constant KC for this reaction. I. Write the Initial concentrations of reactants and products: C. Write the Change in concentration due to reaction using the given reaction stoichiometric coefficients: E. Write the reactant and product concentrations at Equilibrium. Solve for KC : Answer: KC = 1.11 x 103 4 PROBLEMS 1. If a 10.00 L flask at 500 K is filled with a 0.30 mole of hydrogen and 0.30 mole of iodine, what are the equilibrium concentrations of the three gases? The equilibrium constant KC = 45.0. The relevant reaction is Answer: [H2] = [I2] = 0.0069 M; [HI] = 0.0462 M 5 2. Nitrogen dioxide is introduced into a flask at a pressure of one atmosphere (1.00 atm). After some time it dimerizes to produce N2O4. a) What is the equilibrium partial pressure of N2O4 at a temperature where the equilibrium constant for dimerization is KC = 3.33 ? Answer: P (N2O4) = 0.32 atm b) Calculate the total pressure in the flask. Answer: PTotal = 0.68 atm c) Explain why the total pressure is more than, less than, or still equal to the initial pressure (1.00 atm). 6 3. Gaseous carbon dioxide is partially decomposed according to the following equation. An initial pressure of 1.00 atm of CO2 is placed in a closed container at 2500 K, and 2.1 % of the molecules decompose. Determine the equilibrium constant Kp at this temperature. Answer: Kp = 4.83 x 10‐6 7