Enthalpy tells you how much energy a reaction releases, but it says nothing about how fast that happens. Wood burning and iron rusting are both strongly exothermic, yet one takes minutes and the other takes years. Explosives react in a fraction of a second, while the reactions that spoil milk are slow enough that a refrigerator can hold them back for days. The branch of chemistry that studies how fast reactions go, and why, is called chemical kinetics. Controlling rate matters everywhere: speeding up industrial reactions saves energy, and slowing down food spoilage or corrosion saves money and materials.
What a Rate Is
A rate of reaction is the change in the amount or concentration of a reactant or product per unit of time. For substances in solution or in the gas phase, we usually track molar concentration, written with square brackets: means the concentration of A in mol/L. The average rate over a time interval is the change in concentration divided by the change in time. Reactant concentrations go down, so a minus sign keeps the rate positive. Rates almost always slow down as a reaction proceeds, because the reactants are being used up. On a graph of concentration versus time, the average rate between two times is the slope of the secant line joining them, and the instantaneous rate at one moment is the slope of the tangent line at that point.
- Gas volume: collect a gas product in an inverted graduated cylinder or gas syringe and record the volume over time.
- Mass loss: if a gas escapes an open flask on a balance, the mass drops as the reaction proceeds.
- Colour: a spectrophotometer measures how the intensity of a coloured reactant or product changes.
- pH or conductivity: useful when ions are used up or produced.
- Pressure: in a sealed container of gases, the pressure changes if the number of gas molecules changes.
Using Coefficients to Relate Rates
In , every time 1 mol of reacts, 3 mol of react and 2 mol of form. So hydrogen disappears three times as fast as nitrogen, and ammonia appears twice as fast as nitrogen disappears. Dividing each species' rate by its coefficient gives one single rate of reaction that is the same whichever species you measured.
Collision Theory
Collision theory explains rates at the particle level. For a reaction to happen, reactant particles must collide. But in a gas at room temperature, particles collide billions of times per second and only a tiny fraction of those collisions produce products. An effective collision needs two things at once: the particles must hit with the correct orientation, so the atoms that need to bond actually meet, and with at least a minimum amount of kinetic energy, the activation energy , so existing bonds can be stretched and broken. Anything that increases the number of effective collisions per second increases the rate.
| Factor | Change that speeds up the reaction | Collision-theory explanation |
|---|---|---|
| Concentration (or gas pressure) | Higher concentration; for gases, higher pressure from a smaller volume | More particles per litre means more collisions per second, so more effective ones |
| Temperature | Higher temperature | Particles move faster, so they collide more often and, far more importantly, a much larger fraction of collisions has at least the activation energy |
| Surface area | Smaller pieces, powder, or stirring | Only particles at the surface of a solid can collide with the other reactant; breaking it up exposes more of them |
| Nature of reactants | Reactants with fewer or weaker bonds to break, e.g. oppositely charged ions in solution | A lower activation energy means more collisions are energetic enough; ions attract each other and need no bonds broken to combine |
| Catalyst | Adding a suitable catalyst | Provides an alternative pathway with a lower activation energy, so more collisions are effective; the catalyst is regenerated |
The Maxwell-Boltzmann Distribution
Not all particles in a sample have the same kinetic energy. A Maxwell-Boltzmann distribution graphs the number of particles against kinetic energy: a few particles are slow, most are near the peak, and a long tail of particles has very high energy. Only the particles in the tail beyond can react. Raising the temperature flattens the curve and shifts it to the right. The total area under the curve stays the same, because the number of particles has not changed, but the area beyond grows dramatically. A catalyst does not change the curve at all; instead, it moves the line to the left, so more of the existing particles are past it.