• Consider the reaction N2O4 (g) ⇌ 2 NO2 (g). At equilibrium, a 2.00-L reaction vessel contains NO2 at a pressure of 0.269 atm and N2O4 at a pressure of 0.500 atm. The reaction vessel is then compressed to 1.00 L. What will be the pressures of NO2 and N2O4 once equilibrium is re-established?

Explanation:

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At first, it is convenient to compute Kp, since it remains constant during the proposed situation:

Now, since the volume is halved, the pressures of N₂O₄ and NO₂ are doubled, 1.00 atm and 0.538 atm respectively, but now they are considered as the initial pressures, so a change is now present in order to take the reaction again to the equilibrium. Besides, the reaction changes as the products have more moles as:

Thus, the law of mass action goes (considering the reverse reaction):

Hence, solving for by using solver or the quadratic equation, one obtains:

Nonetheless, the feasible solution is 0.072 atm as 0.50 atm leads to a negative pressure of NO₂, thus, the new equilibrium pressures are:

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the molecules move in straight line until they collide with each other or with the walls of the container. collisions are perfectly elastic; when two molecules collide, they change their directions and kinetic energies, but the total kinetic energy is conserved.

explanation:

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