$$ \ce{NO(g) + NO3(g) -> 2NO2(g)} $$ $$ \text{rate} = k[\ce{NO}][\ce{NO3}] $$

The reaction represented above occurs in a single step that involves the collision between a particle of \(\ce{NO}\) and a particle of \(\ce{NO3}\). A scientist correctly calculates the rate of collisions between \(\ce{NO}\) and \(\ce{NO3}\) that have sufficient energy to overcome the activation energy. The observed reaction rate is only a small fraction of the calculated collision rate. Which of the following best explains the discrepancy?

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Only a fraction of collisions occur because of orientation. \(\ce{NO}\) is not symmetrical, so it if collides with \(\ce{NO3}\) on the nitrogen side, the rate of reaction may be different than if it collides on the oxygen side.

There is no third particle present in the reaction.

We can consider the last two statements. The concentration of the reactant particles and the temperature do not change in the question, so it would not affect the activation energy.