Simple rate equations, orders of reaction and rate constants — common mistakes

A common mistake is to try and deduce the order of reaction from the stoichiometry of the overall chemical equation. Remember, orders can only be found from experimental rate data. Also, be very careful with calculating the units of k. Always show your working by rearranging the rate equation and cancelling units – this can often earn you a mark even if your final calculation is incorrect.

If a reactant has an order of zero, its concentration does not affect the rate of the reaction. This might seem strange, but it can happen for several reasons. For example, the reactant might be involved in a step of the reaction mechanism that is much faster than the rate-determining step, so its concentration is never the limiting factor. It could also mean the reactant is in such large excess (like water as a solvent) that its concentration doesn't change noticeably during the reaction.

Yes, although it's less common at A-Level. Fractional orders often indicate a complex reaction mechanism involving radicals. Negative orders can occur when a substance, often a product, inhibits the reaction, meaning an increase in its concentration causes the reaction rate to decrease.

The 'rate of reaction' is the speed at which reactants are used up or products are formed (units: mol dm⁻³ s⁻¹). It changes as the reaction proceeds because reactant concentrations decrease. The 'rate constant' (k) is a constant of proportionality that links the rate to the concentrations. For a specific reaction at a constant temperature, k does not change, whereas the rate does.