Reaction Rates

Rate measures change in concentration per unit time: rate = -delta[reactant]/delta t, or +delta[product]/delta t. The negative sign for a reactant just keeps the rate a positive number, since its concentration is falling. Fast reactions, like gasoline combusting, move through huge concentration changes in tiny slivers of time. Slow ones, like iron rusting, inch along for years because the particles involved collide far less often or far less effectively. Temperature, concentration, surface area, and the natural tendency of a reaction all factor into that difference, but the core idea is just: how much stuff changes, how quickly.

Stoichiometry ties all the individual rates in a reaction together, but it does not make them equal. For N2 + 3H2 → 2NH3, three molecules of H2 vanish and two molecules of NH3 appear for every one N2 consumed. So H2 disappears three times faster than N2, and NH3 forms twice as fast as N2 disappears. To compare rates fairly, divide each rate of change by its coefficient: rate = -delta[N2]/delta t = -(1/3)delta[H2]/delta t = (1/2)delta[NH3]/delta t. That single shared value is the actual reaction rate.

Rate is not constant over time because it typically depends on concentration through a rate law, and concentration is always dropping as reactants get used up. Early on, with lots of reactant, collisions are frequent and the rate is high. As the limiting reactant runs low, collisions become rare and the rate crawls toward zero, which is why concentration-versus-time curves flatten out. This sets up the next idea in the unit: writing a rate law that states exactly how rate depends on each concentration, and defining the rate constant k that stays fixed while concentration and rate do not.