0.19 Reaction kinetics  (Page 5/7)

The answer must be that the O-O bond does not break to create a free O atom during this reaction. That must cost too much energy. Instead, the O atom is handed off in a single step between the O ₂ and the Cl atom. The Cl-O bond is formed at the same instant that the O-O bond is broken. The bonding electrons rearrange from one bond to the other, so the electron energy is never raised by very much energy.

This is a lesson worth remembering. Activation energies and bond energies are not the same thing. In general, activation energies are much, much lower than bond energies, so reactions do not typically occur by consecutively breaking the reactant bonds and then forming the product bonds.

As a final note on Equation (7), the constant A must have some physical significance. We have accounted for the probability of collision between two molecules and we have accounted for the energetic requirement for a successful reactive collision. We have not accounted for the probability that a collision will have the appropriate orientation of reactant molecules during the collision. Moreover, not every collision that occurs with proper orientation and sufficient energy will actually result in a reaction. There are other random factors relating to the internal structure of each molecule at the instant of collision. The factor A takes accounts for all of these factors and is essentially the probability that a collision with sufficient energy for reaction will indeed lead to a reaction. A is commonly called the “frequency factor.”

Observation 4: rate laws for more complicated reaction processes

Our collision model in the previous section accounts for the concentration and temperature dependence of the reaction rate, as expressed by the rate law. The concentration dependence arises from calculating the probability of the reactant molecules being in the same vicinity at the same instant. Therefore, we should be able to predict the rate law for any reaction by simply multiplying together the concentrations of all reactant molecules in the balanced stoichiometric equation. The order of the reaction should therefore be simply related to the stoichiometric coefficients in the reaction. However, the data in Table 4 of the previous study shows that this is incorrect for many reactions.

Consider, for example, the apparently simple reaction

2ICl(g) + H ₂ (g) → 2HCl(g) + I ₂ (g)

Based on the collision model, we would assume that the reaction occurs by 2 ICl molecules colliding with a single H ₂ molecule. The probability for such a collision should be proportional to [ICl] ² [H ₂ ]. However, we experimentally observe that the rate law for this reaction is

Rate = k[ICl][H ₂ ]

As a second example, consider the reaction

NO ₂ (g) + CO(g) → NO(g) + CO ₂ (g)

It would seem reasonable to assume that this reaction occurs as a single collision in which an oxygen atom is exchanged between the two molecules. However, the experimentally observed rate law for this reaction is

Rate = k[NO ₂ ] ²

In this case, [CO] does not affect the rate of the reaction at all, and [NO ₂ ] is squared. These examples demonstrate that the rate law for a reaction cannot always be predicted from the stoichiometric coefficients and therefore that the collision model often does not account for the rate of the reaction. There must be something seriously incomplete with the collision model.