transition state
Imagine pushing a heavy cart over the crest of a steep hill. There is one exact instant at the very top when the cart is balanced on the peak — not rolling back, not yet rolling forward, poised between the two valleys. It cannot rest there; the slightest nudge sends it down one side. The transition state is the chemical equivalent of that knife-edge moment at the summit of a reaction.
More precisely, the transition state is the arrangement of atoms at the highest point of energy along the path from reactants to products — the top of the energy barrier. At this point old bonds are partly broken and new bonds are partly formed; the molecule is neither reactant nor product but a fleeting halfway configuration. Its energy above the reactants is the activation energy, and only collisions energetic enough to reach this summit can cross over into products.
Why it matters: the height of the transition state sets how fast a reaction goes, so anything that lowers it — like a catalyst — speeds the reaction up. The crucial subtlety is that a transition state is not a real molecule you can bottle: it exists for only about the time of a single bond vibration and sits at a maximum of energy, so it can never be isolated, only inferred and modelled. This is exactly what separates it from a true intermediate.
When a hydroxide ion attacks bromomethane, there is a split-second arrangement in which the carbon atom is loosely bonded to both the incoming oxygen and the departing bromine at once — a flat, five-connected carbon. That strained, half-made, half-broken arrangement is the transition state; a heartbeat later it has collapsed into product.
Old bonds half-broken, new bonds half-formed — the strained summit of the energy hill.
A transition state sits at an energy maximum and cannot be isolated; an intermediate sits in an energy minimum (a small dip) and has a real, if brief, lifetime. Mixing the two up is the most common error here.