A closed flask contains a mixture of nitrogen dioxide (NO2\text{NO}_2NO2, a brown gas) and dinitrogen tetraoxide (N2O4\text{N}_2\text{O}_4N2O4, a colourless gas) at a constant temperature. The system reaches dynamic equilibrium according to the following equation:
2NO2(g)⇌N2O4(g)ΔH<0 2\text{NO}_2\text{(g)} \rightleftharpoons \text{N}_2\text{O}_4\text{(g)} \quad \Delta H < 0 2NO2(g)⇌N2O4(g)ΔH<0The graph below shows how the rates of the forward and reverse reactions change over time for this system:
Which of the following statements is correct for this system once equilibrium has been achieved?
The rate of consumption of NO2\text{NO}_2NO2 in the forward reaction is exactly equal to its rate of production in the reverse reaction.
The concentration of NO2\text{NO}_2NO2 must be exactly twice the concentration of N2O4\text{N}_2\text{O}_4N2O4 due to the 2:12:12:1 stoichiometric ratio.
Both the forward and reverse reactions have completely stopped, preventing any further change in the intensity of the brown colour.
The rate of the forward reaction is twice the rate of the reverse reaction because two moles of NO2\text{NO}_2NO2 react to form one mole of N2O4\text{N}_2\text{O}_4N2O4.