Reaction rates

For many reactions, raising the temperature by 10 °C roughly doubles the rate, which is why food keeps for days in a fridge but spoils on a warm counter.

An explosion is over in a fraction of a second. Rusting takes years. The rate of a reaction is how quickly reactants are used up or products are made, for example in moles per litre per second. Two reactions with the same equation can run at very different rates under different conditions, and a great deal of chemistry, from cooking to medicine to industry, is about controlling that speed.

Collision theory

Collision theory explains rates with three ideas. First, particles have to meet: the rate depends on how often reactant particles collide. Second, they have to meet the right way round, with the atoms that will bond facing each other. Third, they have to hit hard enough to reach the activation energy. Only collisions that pass all three tests lead to products. In the lab, watch the readout of collisions in the last five seconds and how many of them had enough energy. Most collisions are just bounces.

Five things that change the rate

The nature of the reactants matters first. Small pieces of sodium and iron left in air behave very differently: the sodium reacts completely overnight while the iron is barely touched. Calcium reacts with water at a moderate pace, sodium almost explosively. Surface area matters when a solid reacts with a liquid or gas, because the reaction happens only where they touch. Large logs smoulder, small sticks burn quickly and sawdust in the air can burn explosively. Iron powder fizzes in dilute acid much faster than an iron nail.

Temperature is the third factor. Hotter particles move faster, so they collide more often, and a much larger share of the collisions carry enough energy to react. For many processes the rate roughly doubles for every 10 °C rise. Concentration is the fourth: more particles in the same space means more collisions. Phosphorus burns far faster in pure oxygen than in air, which is only about 20% oxygen. The fifth is a catalyst, a substance that speeds up a reaction without being used up. Hydrogen peroxide decomposes very slowly in the bottle, but poured on a cut it foams at once, because substances in the tissue act as catalysts.

Putting a number on temperature

The Swedish chemist Svante Arrhenius described how the rate constant k depends on temperature: k = A·e^(−Eₐ/RT), where R is 8.314 J per mole per kelvin and T is the temperature in kelvins. The exponential term is the fraction of collisions with enough energy, and it grows quickly with temperature. For a reaction with an activation energy of 50 kJ per mole, going from 25 °C to 35 °C multiplies the rate by about 1.92, close to the doubling rule. Cooling the same reaction to 0 °C cuts the rate to about 0.16 of its value at 25 °C. The last line of the lab readout does this calculation for whatever temperature you choose.

The particle model in the lab uses a lowered energy barrier so that reactions appear within seconds instead of never. Its job is to show the pattern: more speed, more particles and a lower barrier all give a steeper green line. Notice also that the blue line flattens over time. As reactants are used up there are fewer A particles to collide, so the reaction slows down by itself even at a fixed temperature.