Elementary Reactions

An elementary reaction describes one molecular event: one or two particles (rarely three) colliding and turning directly into products, with nothing else happening in between. Because the equation for an elementary step IS the actual collision, you can read its rate law straight off the stoichiometry. If a step is A → products, rate = k[A]. If it's A + B → products, rate = k[A][B]. If it's 2A → products, rate = k[A]^2. This only works for elementary steps, never for an overall balanced equation.

Elementary steps are classified by how many species collide. Unimolecular steps involve one molecule falling apart or rearranging. Bimolecular steps involve two molecules colliding, the most common and most probable case. Termolecular steps need three particles colliding at the same instant with the right energy and orientation, and are rare but possible. Four or more colliding simultaneously is essentially never seen, since the odds of that many particles meeting at once, correctly aligned and energetic enough, shrink toward zero as the number of participants grows.

This is why a balanced overall equation is not automatically elementary. The overall equation only tells you the net stoichiometry, the starting materials and final products after every step is added together. Most reactions actually proceed through a mechanism, a sequence of elementary steps that often create and consume intermediates along the way. You cannot read the overall rate law from the overall equation; you have to find it from the mechanism, usually from whichever step is slowest. That idea, the rate-determining step, is where this topic connects to the rest of kinetics.