8.1 Rate of reaction

Rate of reaction is the change in concentration of a reactant or product per unit time. It can be expressed as a positive product-formation rate or a negative reactant-consumption change with the sign interpreted.

Collision theory explains rate: particles must collide with enough energy and suitable orientation. Increasing concentration, pressure, surface area or temperature changes the frequency or success of collisions.

A gas reaction speeds up when pressure rises because particles collide more often in the smaller volume. A powdered solid reacts faster than the same mass in lumps because more surface is exposed.

A faster rate does not mean a larger final yield. Rate describes time dependence; equilibrium or limiting reactants determine how far a reaction proceeds.

For a collision to produce products, particles must collide with energy at least equal to the activation energy and with an orientation that allows bonds to rearrange.

Many collisions are unsuccessful. Changing conditions can alter the fraction that is energetic enough or the number of collisions per second.

Heating a reacting mixture increases the fraction of particles with energy above Eₐ, so more collisions become successful even though the activation-energy threshold itself is unchanged.

Particles do not all have the same kinetic energy, and simply colliding is not enough. Both energy and orientation matter.

Average rate = change in concentration divided by time. For a reactant, use the decrease in concentration as a positive rate magnitude; for a product, use the increase.

Read the correct interval and units from the data. Tangent gradients give an instantaneous rate at one point, whereas a secant over two readings gives an average rate.

If a product rises from 0.10 to 0.34 mol dm⁻³ in 60 s, average rate = (0.34−0.10)/60 = 4.0×10⁻³ mol dm⁻³ s⁻¹.

Do not divide by the final time when the interval begins later, and do not confuse concentration change with total amount when volume changes.