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Ch.15 - Chemical Equilibrium
All textbooksMcMurry 8th EditionCh.15 - Chemical EquilibriumProblem 112
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Chapter 15, Problem 112

At 45 °C, Kc = 0.619 for the reaction N2O4(g) ⇌ 2 NO2(g). If 46.0 g of N2O4 is introduced into an empty 2.00-L container, what are the partial pressures of NO2 and N2O4 after equilibrium has been achieved at 45 °C?

Verified step by step guidance
1
Calculate the initial concentration of N2O4 by dividing the mass of N2O4 by its molar mass to find moles, then divide by the volume of the container to find molarity.
Set up an ICE (Initial, Change, Equilibrium) table to track the concentrations of N2O4 and NO2. Initially, the concentration of NO2 is zero, and the concentration of N2O4 is the initial concentration calculated.
Define the change in concentration for the reaction as x, where N2O4 decreases by x and NO2 increases by 2x, according to the stoichiometry of the reaction.
Write the expression for the equilibrium constant Kc in terms of the equilibrium concentrations: Kc = [NO2]^2 / [N2O4]. Substitute the expressions from the ICE table into this equation.
Solve the quadratic equation for x to find the equilibrium concentrations of N2O4 and NO2. Use these concentrations to calculate the partial pressures using the ideal gas law: P = (n/V)RT, where n is the number of moles, V is the volume, R is the ideal gas constant, and T is the temperature in Kelvin.
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Related Practice
Open Question
At 500 °C, F2 gas is stable and does not dissociate, but at 840 °C, some dissociation occurs: F2(g) ⇌ 2 F(g). A flask filled with 0.600 atm of F2 at 500 °C was heated to 840 °C, and the pressure at equilibrium was measured to be 0.984 atm. What is the equilibrium constant Kp for the dissociation of F2 gas at 840 °C?
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The reaction NO(g) + NO2(g) ⇌ N2O3(g) takes place in the atmosphere with Kc = 13 at 298 K. A gas mixture is prepared with 2.0 mol NO and 3.0 mol NO2 and an initial total pressure of 1.65 atm. (a) What are the equilibrium partial pressures of NO, NO2, and N2O3 at 298 K?
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