A container has a mixture of NO2 gas and N2O4 gas in equilibrium. The chemical reaction between the two gases is described by the equation The partial pressure of NO2 is 61.2 MPa and the partial pressure of N2O4 is 38.8 MPa. Several weights are added to the lid of the container, increasing the total pressure on the container to 200 MPa. What are the most likely partial pressures when equilibrium is established again?

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indigognat193Lv1

1 Apr 2020

A container has a mixture of NO₂ gas and N₂O₄ gas in equilibrium. The chemical reaction between the two gases is described by the equation

The partial pressure of NO₂ is 61.2 MPa and the partial pressure of N₂O₄ is 38.8 MPa. Several weights are added to the lid of the container, increasing the total pressure on the container to 200 MPa. What are the most likely partial pressures when equilibrium is established again?

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Suppose that nitrogen gas and hydrogen gas are mixed in a container of constant volume at 700 K and in the presence of a catalyst such that the initial partial pressures of nitrogen gas is 2.0 atm and the initial partial pressure of hydrogen gas is 6.0 atm. Under these conditions equilibrium between the reactants and the product ammonia is established. The equilibrium constant, K_P, for this reaction at 700 K is 1.0 x 10^-4. Compute partial pressures of the three gases when equilibrium is established

2. Solve the following related problems involving the equilibrium between dinitrogen tetroxide and nitrogen dioxide.

N₂O₄ <==> 2 NO₂

The equilibrium constant for this reaction is K_p = 0.60 at 350 degrees Kelvin. A 1.0 L vessel contains the above two gases in equilibrium at 350 K. The total pressure of the two gases, P_(NO₂) + P_(N₂O₄), in the vessel is 0.50 atm. From this information, calculate the equilibrium partial pressure for each gas. This is the initial equilibrium position for the reaction, which you will use to consider the effects of the following perturbations.

This equilibrium mixture is subjected to the following four perturbations:

a. addition of N₂O₄ gas to the vessel (which had been at the above equilibrium) to give a total concentration (before readjustment of the equilibrium position) of 0.375 atm in the reactant gas;

b. decreasing the total volume of the vessel to 0.50 L (in answering b-d, go back to the initial equilibrium mixture before then applying the perturbation);

c. increasing the total pressure inside the vessel to 1.5 atm by adding 1.0 atm of N₂ gas;

d. increasing the temperature of the gas mixture to 700 degrees K, a temperature at which the equilibrium constant is 2.40.

3. Hints: Note that the equilibrium constant is a K_p value and that reactant and product concentrations are given in terms of partial pressures.

For each of these perturbations, you can solve for the equilibrium partial pressures of the two gases. To do this you may have to set up an ICE table. Note that in ï¿½bï¿½, the inverse relationship between P and V in the ideal gas law allows you to easily determine the partial pressures of the reactant and product gases immediately after the volume increase. Then you can calculate the final equilibrium concentrations following this perturbation.

Finally, the exact solution to part ï¿½dï¿½ is challenging. You should first consider the effect, on the partial pressures of the reactant and product gases, of an increase in temperature from 350 to 700 degrees K. (Remember that the ideal gas law indicates a direct proportionality of P and T.) Then you should consider a shift in the equilibrium concentrations, along with the new equilibrium constant for the higher temperature.

Even if you are not able to determine numerical values for the equilibrium positions following the perturbations, you still should understand the concepts and be able to prepare the following written response.

4. Now, prepare an explanation: Attempt to solve for the extent and direction of the shift in the equilibrium as a result of the above four perturbations. If you are unable to solve for numerical values of the final equilibrium partial pressures, you should still know the predicted directions of the shifts.

After studying and solving the above problem, write an explanation of the Le Chatelier Principle. This explanation should give a general statement of the principle and then should explain how the above four types of perturbations are rationalized in terms of this principle.

Consider the apparatus shown in the drawing.

(a) When the stopcock between the two containers is opened and the gases are allowed to mix, how does the volume occupied by the N₂ gas change? What is the partial pressure of N₂ after mixing?

(b) How does the volume of the O2 gas change when the gases mix? What is the partial pressure of O₂ in the mixture?

(c) What is the total pressure in the container after the gases mix?