A gaseous system at 298\,K is established by the reversible reaction
N2O4(g)⇌2NO2(g)
At equilibrium the brown color of NO2 gives the mixture a light amber appearance. A student observes that a freshly prepared 2.0\,L sealed flask containing only these two gases at the same temperature appears *much* darker than the reference equilibrium mixture even though the total pressure is identical.
Answer the following questions about the system, citing particulate–level reasoning where appropriate.
(a) Describe qualitatively how the instantaneous value of the reaction quotient, Qp, for the freshly prepared flask compares with the equilibrium constant, Kp. Justify your answer using macroscopic evidence given in the stimulus.
(b) Predict the direction in which the net reaction will proceed as the system moves toward equilibrium. Explain your reasoning in terms of the relative rates of the forward and reverse reactions.
(c) After sufficient time has passed, the color of the flask is observed to become lighter but remains darker than the reference equilibrium mixture. Explain what this observation indicates about the magnitude of Qp relative to Kp at that later time.
(d) Using a molecular–level description, explain how changes in the relative number of NO2 and N2O4 molecules account for both the color change and the shift you predicted in part (b).