Rates of Reaction Explained
Why does hydrogen peroxide decompose slowly at room temperature but almost explosively in the presence of manganese dioxide? Why does bread go stale faster on a warm day than in the fridge? Both questions are answered by the same concept: the rate of a chemical reaction and what controls it.
Defining Reaction Rate
The rate of a reaction measures how quickly reactants are converted into products. It is defined as the change in concentration (or amount) of a reactant or product per unit of time. Experimentally, you can track rate by measuring: the decrease in mass as a gas escapes, the volume of gas collected over time, a colour change using a colorimeter, or the change in electrical conductivity. Plotting concentration against time gives a curve whose gradient at any point equals the instantaneous rate at that moment.
Collision Theory
Reactions occur when reactant particles collide. But not every collision leads to a reaction. For a collision to be successful — that is, to produce new bonds and new products — two conditions must both be met:
- The colliding particles must have energy equal to or greater than the activation energy (Ea).
- The particles must collide with the correct orientation so that the reactive parts of each molecule meet.
The activation energy is the minimum kinetic energy that colliding particles must possess in order to break existing bonds and begin forming new ones. Think of it as the energy needed to push the system over the top of a hill before it can roll down to products. Reactions with a high activation energy proceed slowly at room temperature because very few molecules in the distribution of energies happen to have enough energy at any given moment.
Factors That Affect Reaction Rate
| Factor | Change | Effect on rate | Reason (collision theory) |
|---|---|---|---|
| Temperature | Increase | Increases | Particles move faster; more collisions per second; a much larger proportion of particles exceed Ea |
| Concentration (solutions) | Increase | Increases | More particles per unit volume; collision frequency rises |
| Pressure (gases) | Increase | Increases | Same as concentration: particles are closer together |
| Surface area (solids) | Increase | Increases | More surface exposed to reactant; more possible collision sites |
| Catalyst | Added | Increases | Provides alternative pathway with lower Ea |
Temperature and the Maxwell-Boltzmann Distribution
At any given temperature, particles in a gas or solution do not all move at the same speed. Their energies are spread across a range described by the Maxwell-Boltzmann distribution. The curve is bell-shaped but skewed to the right: most particles have intermediate energies, a few have very low energies, and a tail of particles have very high energies. The area under the curve to the right of the activation energy line represents the fraction of particles that can react.
When temperature rises, the whole distribution shifts to higher energies and the peak flattens. Crucially, the high-energy tail extends further to the right, so the area beyond Ea increases significantly. A temperature rise of just 10°C roughly doubles the rate of many reactions at room temperature, because even a modest shift in the distribution moves a substantially larger fraction of particles above Ea.
Surface Area in Practice
When a solid reactant is involved, only the surface particles are exposed to the other reactant. Grinding a solid into a fine powder dramatically increases the total surface area without changing the number of moles present. This is why powdered calcium carbonate reacts with hydrochloric acid much faster than an equivalent-mass marble chip. The same principle explains why grain-dust explosions in silos are a serious industrial hazard: finely divided solid combustibles ignite far more readily than bulk material.
Catalysts and Activation Energy
A catalyst speeds up a reaction by providing an alternative reaction pathway that has a lower activation energy. On a Maxwell-Boltzmann diagram, lowering Ea moves the threshold line to the left, so a much larger fraction of the existing particle distribution now exceeds it — rate rises substantially even though the temperature and particle energies are unchanged. The catalyst is not consumed: it participates in an intermediate step, then reforms and is available to catalyse the next cycle.
Homogeneous catalysts are in the same phase as the reactants (for example, aqueous iron(III) ions catalysing the reaction between aqueous persulfate and iodide). Heterogeneous catalysts are in a different phase, typically a solid surface over which gaseous or liquid reactants flow; the platinum-rhodium catalytic converter in a car exhaust is a classic example, converting toxic CO and NOx into CO2 and N2.
Measuring Rate: A Quick-Start Guide
For the marble chips / hydrochloric acid experiment, place the flask on a balance and record mass every 30 seconds as CO2 escapes. Plot mass remaining against time. The gradient at any point (change in mass / change in time) gives the rate at that moment. As reactants are used up, the gradient decreases and the curve levels off. Repeating with the same mass of powdered chalk shows a steeper initial gradient — confirming the surface area effect — but the same final mass loss, because the same number of moles of reactant is used in both runs.
Summary
Reaction rate is controlled by collision frequency and the fraction of collisions that are successful. Any factor that increases the number of collisions per second — higher temperature, greater concentration, increased pressure for gases, larger surface area for solids — speeds up a reaction. Catalysts provide a shortcut with a lower activation energy. The Maxwell-Boltzmann distribution explains why temperature has such a powerful effect: even a small rise shifts a disproportionately large fraction of molecules above the activation energy threshold.