Energy of activation vs Bond energy In the second chapter, Methane Energy of activation and Transition state in the textbook Organic Chemistry by Morrison and Boyd, Energy of activation is given the following definition:


  The minimum amount of energy that must be provided by a collision for a reaction to occur is called the energy of activation, Eact.


The Eact for the reaction:

Cl(free radical) + CH3-H ---------> H-Cl + CH3(free radical)

is 4 kcal/mol. Therefore, on an average, it is 4/NA kcal per individual collision (where NA is the Avogadro number).

However, the bond energy of the carbon-hydrogen bond in methane is 104 kcal/mol. Therefore, on an average, it would be 104/ NA kcal for a single molecule of methane. In other words, 104/ NA kcal should be the energy provided by a collision that involves a methane molecule in order to break the carbon-hydrogen bond. Why does this not agree with the statement above which says 4/NA kcal per individual collision is what is required?
The bond dissociation energy is the energy needed for the reaction leading to the rupture the $\ce{C-H}$ bond. I.e. it is the enthalpy of the reaction:

$$\ce{CH_3-H \rightarrow CH_3^{\bullet} + H^{\bullet}}$$

In the case of the reaction in the question, formation of the $\ce{H-Cl}$ bond will counterbalance the energy cost for breaking the $\ce{C-H}$ bond.