Skip to main content
Pearson+ LogoPearson+ Logo
Ch.15 - Chemical Equilibrium
All textbooksMcMurry 8th EditionCh.15 - Chemical EquilibriumProblem 113
Previous problem
Next problem
Chapter 15, Problem 113

When 9.25 g of ClF3 was introduced into an empty 2.00-L container at 700.0 K, 19.8% of the ClF3 decomposed to give an equilibrium mixture of ClF3, ClF, and F2. ClF3 (g) ⇌ ClF (g) + F2 (g). (a) What is the value of the equilibrium constant Kc at 700.0 K? (b) What is the value of the equilibrium constant Kp at 700.0 K? (c) In a separate experiment, 39.4 g of ClF3 was introduced into an empty 2.00-L container at 700.0 K. What are the concentrations of ClF3, ClF, and F2 when the mixture reaches equilibrium?

Verified step by step guidance
1
Step 1: Calculate the initial concentration of ClF3. Use the formula: \( \text{Concentration} = \frac{\text{moles}}{\text{volume}} \). First, convert the mass of ClF3 to moles using its molar mass.
Step 2: Determine the change in concentration due to decomposition. Since 19.8% of ClF3 decomposes, calculate the moles of ClF3 that decomposed and the moles of ClF and F2 formed.
Step 3: Write the expression for the equilibrium constant \( K_c \) using the equilibrium concentrations of ClF3, ClF, and F2. \( K_c = \frac{[\text{ClF}][\text{F}_2]}{[\text{ClF}_3]} \). Substitute the equilibrium concentrations into this expression.
Step 4: Relate \( K_c \) to \( K_p \) using the equation \( K_p = K_c(RT)^{\Delta n} \), where \( \Delta n \) is the change in moles of gas, \( R \) is the ideal gas constant, and \( T \) is the temperature in Kelvin.
Step 5: For the separate experiment, calculate the initial concentration of ClF3 with the new mass. Use the same decomposition percentage to find the equilibrium concentrations of ClF3, ClF, and F2, and apply the equilibrium constant to verify the concentrations.
Previous problem
Next problem
Related Practice
Open Question
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?
Textbook Question
The equilibrium constant Kc for the reaction N21g2 + 3 H21g2 ∆ 2 NH31g2 is 4.20 at 600 K. When a quantity of gaseous NH3 was placed in a 1.00-L reaction vessel at 600 K and the reaction was allowed to reach equilibrium, the vessel was found to contain 0.200 mol of N2. How many moles of NH3 were placed in the vessel?
432
views
Open Question
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?
Textbook Question
The reaction of fumarate with water to form L-malate is catalyzed by the enzyme fumarase; Kc = 3.3 at 37°C. When a reaction mixture with [fumarate] = 1.56 * 10-3 M and [l -malate] = 2.27 * 10-3 M comes to equilibrium in the presence of fumarase at 37 °C, what are the equilibrium concentrations of fumarate and L-malate? (Water can be omit- ted from the equilibrium equation because its concentration in dilute solutions is essentially the same as that in pure water.)
834
views
Open Question
Calculate the equilibrium concentrations of SO2, Cl2, and SO2Cl2 at 298 K if the initial concentrations are [SO2] = 1.50 M and [Cl2] = 0.85 M. The equilibrium constant Kc for the reaction SO2(g) + Cl2(g) ⇌ SO2Cl2(g) is 8.40 × 10^-3 at 298 K.
Open Question
Calculate the equilibrium concentrations of H2O(g), Cl2(g), HCl(g), and O2(g) at 298 K if the initial concentrations are [H2O] = 0.050 M and [Cl2] = 0.100 M. The equilibrium constant Kc for the reaction H2O(g) + Cl2(g) ⇌ 2 HCl(g) + O2(g) is 8.96 × 10^-9 at 298 K.