22. The equilibrium constant for the following reaction is 14.9 2 NO (g) + Cl2(g) 2NOCK(g) Kc = 14.9 A reaction container is found to contain the following concentrations of these three gases: [NO] 0.0500mol/L [Cl₂] 0.100mol/L [NOCI] a. Show that the system is not at equilibrium. 0.100mol/L b. What shift would need to take place in order to establish equilibrium? Justify your answer.

22. The equilibrium constant for the following reaction is 14.9 2 NO (g) + Cl2(g) 2NOCK(g) Kc = 14.9 A reaction container is found to contain the following concentrations of these three gases: [NO] 0.0500mol/L [Cl₂] 0.100mol/L [NOCI] a. Show that the system is not at equilibrium. 0.100mol/L b. What shift would need to take place in order to establish equilibrium? Justify your answer.

Chemistry & Chemical Reactivity

9th Edition

ISBN:9781133949640

Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Publisher:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Chapter15: Principles Of Chemical Reactivity: Equilibria

Section: Chapter Questions

Problem 40GQ: Consider the following equilibrium: COBr2(g) CO(g) + Br2(g)Kc = 0.190 at 73 C (a) A 0.50 mol sample...

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At 2300 K the equilibrium constant for the formation of NO(g) is 1.7 103. N2(g) + O2(g) 2 NO(g) (a) Analysis shows that the concentrations of N2 and O2 are both 0.25 M, and that of NO is 0.0042 M under certain conditions. Is the system at equilibrium? (b) If the system is not at equilibrium, in which direction does the reaction proceed? (c) When the system is at equilibrium, what are the equilibrium concentrations?
Two molecules of A react to form one molecule of B, as in the reaction 2 A(g) B(g) Three experiments are done at different temperatures and equilibrium concentrations are measured. For each experiment, calculate the equilibrium constant, Kc. (a) [A] = 0.74 mol/L, [B] = 0.74 mol/L (b) [A] = 2.0 mol/L, [B] = 2.0 mol/L (c) [A] = 0.01 mol/L, [B] = 0.01 mol/L What can you conclude about this statement: If the concentrations of reactants and products are equal, then the equilibrium constant is always 1.0.
1’he reaction in Exercise 12.33 was repeated. This time, the reaction began when only NO was injected into the reaction container. 110.200 mol L_l NO was injected, what were the equilibrium concentrations of all species? The following reaction establishes equilibrium at 2000 K: N2(g) + O2(g) ^2 NO K = 4.1 X 10~4 If the reaction began with 0.100 mol L-1 of N2 and 0.100 mol L"' ofO2, what were the equilibrium concentrations of all species?
A mixture of N2, H2, and NH3 is at equilibrium [according to the equationN2(g)+3H2(g)2NH3(g)] as depicted below: The volume is suddenly decreased (by increasing the external pressure) and a new equilibrium is established as depicted below: a. If the volume of the final equilibrium mixture is 1.00 L, determine the value of the equilibrium constant, K. for the reaction. Assume temperature is constant. b. Determine the volume of the initial equilibrium mixture assuming a final equilibrium volume of 1.00 L and assuming a constant temperature.
A solution is prepared by dissolving 0.050 mol of diiodocyclohexane, C5H10I2, in the solvent CCl4.The total solution volume is 1.00 L When the reaction C6H10I2 C6H10 + I2 has come to equilibrium at 35 C, the concentration of I2 is 0.035 mol/L. (a) What are the concentrations of C6H10I2 and C6H10 at equilibrium? (b) Calculate Kc, the equilibrium constant.
A mixture of SO2, O2, and SO3 at 1000 K contains the gases at the following concentrations: [SO2] = 5.0 103 mol/L, [O2] = 1.9 103 mol/L, and [SO3] = 6.9 103 mol/L. Is the reaction at equilibrium? If not, which way will the reaction proceed to reach equilibrium? 2 SO2(g) + O2(g) 2 SO3(g) Kc = 279
For the equilibrium 2 SO2(g) + O2(g) 2 SO3(g) Kc = 245 (at 1000 K) the equilibrium concentrations are [SO2] = 0.102, [O2] = 0.0132, and [SO3] = 0.184. The concentration of SO2 is suddenly doubled. Show that the forward reaction takes place to reach a new equilibrium.
What is Le Chteliers principle? Consider the reaction 2NOCI(g)2NO(g)+Cl2(g) If this reaction is at equilibrium. what happens when the following changes occur? a. NOCI(g) is added. b. NO(g) is added. c. NOCI(g) is removed. d. Cl2(g) is removed. e. The container volume is decreased. For each of these changes, what happens to the value of K for the reaction as equilibrium is reached again? Give an example of a reaction for which the addition or removal of one of the reactants or products has no effect on the equilibrium position. In general, how will the equilibrium position of a gas-phase reaction be affected if the volume of the reaction vessel changes? Are there reactions that will not have their equilibria shifted by a change in volume? Explain. Why does changing the pressure in a rigid container by adding an inert gas not shift the equilibrium position for a gas-phase reaction?
The experiment in Exercise 12.33 was redesigned so that the reaction started with 0.15 mol each of N2 and O2 being injected into a 1.0-L container at 2500 K. The equilibrium constant at 2500 K is 3.6 X 10“’. What was the composition of the reaction mixture after equilibrium was attained? The following reaction establishes equilibrium at 2000 K: N2(g) + O2(g) *2 2 NO K = 4.1 X IO-4 If the reaction began with 0.100 mol L-1 of N2 and 0.100 mol L-’ ofO2, what were the equilibrium concentrations of all species?
Consider the following reactions at some temperature: 2NOCl(g)2NO(g)+Cl2(g)K=1.6105 2NO(g)N2(g)+O2(g)K=11031 For each reaction, assume some quantities of the reactants were placed in separate containers and allowed to come to equilibrium. Describe the relative amounts of reactants and products that would be present at equilibrium. At equilibrium, which is faster, the forward or reverse reaction in each case?
At room temperature, the equilibrium constant Kc for the reaction 2 NO(g) ⇌ N2(g) + O2(g) is 1.4 × 1030. Is this reaction product-favored or reactant-favored? Explain your answer. In the atmosphere at room temperature the concentration of N2 is 0.33 mol/L, and the concentration of O2 is about 25% of that value. Calculate the equilibrium concentration of NO in the atmosphere produced by the reaction of N2 and O2. How does this affect your answer to Question 11?
12.101 An engineer working on a design to extract petroleum from a deep thermal reservoir wishes to capture toxic hydrogen sulfide gases present by reaction with aqueous iron(II) nitrate to form solid iron(II) sulfide. (a) Write the chemical equation for this process, assuming that it reaches equilibrium. (b) What is the equilibrium constant expression for this system? (c) How can the process be manipulated so that it does not reach equilibrium, allowing the continuous removal of hydrogen sulfide?