The complete lecture — read on the left while the live panel on the right shows each idea through a real, everyday picture: cars in a car park, a cyclist climbing a hill, a crowded room, a hot afternoon, a stopwatch over a fizzing flask. Scroll down; the animation keeps pace.
The rate of a reaction is how fast it happens — formally, the change in concentration of a reactant or product per unit time. Its unit is mol dm⁻³ s⁻¹. An explosion is over in an instant; the rusting of an iron gate takes years.
Why are some reactions fast and others slow? Collision theory answers it. For particles to react they must collide — but not every collision works. Think of a busy car park: two cars only end up neatly parked together if they meet at the correct angle and the right speed. A fast, head-on, well-lined-up meeting succeeds; a slow or badly-angled one just glances off.
Picture a cyclist facing a hill. To reach the valley on the far side they must first pedal hard up and over the top. That climb is the activation energy: a reaction cannot start until the particles have enough energy to get over the hill. Once over the crest, they roll down to the products.
On an energy profile the reactants climb the Ea barrier to the transition state, then fall to products. If the products sit lower than the reactants (downhill landing), the reaction is exothermic — it releases heat. If they sit higher, it is endothermic.
Back to the hill. A catalyst is like boring a tunnel straight through it. The cyclist starts and finishes at exactly the same two points — so the overall energy change is unchanged — but the barrier in between is much lower, so far more riders make it across. Crucially, the tunnel is still there afterwards: the catalyst is regenerated and not consumed.
| Type | Description | Example |
|---|---|---|
| Homogeneous | same phase as reactants | acid in ester hydrolysis |
| Heterogeneous | different phase (usually solid) | Fe in Haber; Pt in catalytic converters |
For a solid reactant, only the particles on the surface can be hit. Drop a whole sugar cube into tea and it dissolves slowly; crush it to powder first and it vanishes almost at once. Crushing exposes a far larger surface, so many more collisions happen each second and the rate shoots up. This is also why fine flour or sugar dust can explode while a solid lump only smoulders.
| Factor | Effect (increase →) |
|---|---|
| Concentration (of reactant) | faster — more collisions |
| Temperature | much faster — more energetic & frequent collisions |
| Surface area (solid) | faster — more contact, more collisions |
| Catalyst | faster — lowers activation energy |
| Pressure (gases) | faster — molecules pushed closer |
Imagine a few people wandering a large hall — they rarely bump into one another. Now squeeze many more people into the same room: collisions happen constantly. Raising the concentration of a reactant does exactly this — more particles in the same space means more collisions each second, so a faster rate.
We capture this in the rate law, found by experiment:
| Order | Rate | Units of k |
|---|---|---|
| zero | rate = k | mol dm⁻³ s⁻¹ |
| first | rate = k[A] | s⁻¹ |
| second | rate = k[A]² | mol⁻¹ dm³ s⁻¹ |
Picture the same crowd on a cool morning versus a scorching afternoon. In the heat people dash about, bump harder and more often. Molecules behave the same way: raise the temperature and they move faster, so collisions are both more frequent and more energetic — and far more of them now clear the activation-energy hill.
A rise of only ~10 °C roughly doubles the rate. The Maxwell–Boltzmann distribution of molecular energies shifts to the right, so a much larger fraction of molecules have energy ≥ Ea.
How do we actually measure a rate in the lab? The simplest way for a reaction that gives off gas is a stopwatch and a bubble counter: start timing, then count how quickly the gas bubbles come off or how fast a syringe fills. Fast bubbling early on, slowing as the reactants run out — exactly the shape kinetics predicts.