Half-life and experimental techniquesEdexcel International A Level Chemistry: Revision notes

The half-life (t½) is the time taken for the concentration of a reactant to fall to half its value. Read it from a concentration–time graph by choosing a concentration, finding half of it, and reading the time difference.

Repeat this from at least two or three different starting concentrations. If the half-life is constant, the reaction is first order with respect to that reactant, because the time to halve does not depend on the concentration. If the half-lives double each time, the reaction is second order. For zero order the graph is a straight line and there is no constant half-life.

Choose a method that measures a quantity that changes measurably during the reaction, and justify the choice by naming that change.

• Gas produced: collect the gas in a gas syringe and record the volume against time.
• Gas lost: record the loss in mass of the mixture on a balance (works best for dense gases such as CO₂).
• Coloured species: use colorimetry. The absorbance is proportional to the concentration of the coloured species, and no samples are needed.
• Reaction that can be stopped: remove samples at intervals, quench them and titrate.
• Change in the number or type of ions: measure the conductivity or pH of the mixture.

In continuous monitoring, one reaction mixture is followed over time. A concentration–time or volume–time graph is plotted. The rate at any time is the gradient of a tangent to the curve at that time. Rates at different concentrations give a rate–concentration graph, which shows the order.

The advantage is that a single experiment gives a whole graph. The orders with respect to other reactants need further experiments.

In the initial-rate method, separate experiments are done with different initial concentrations of one reagent, keeping everything else (other concentrations, total volume, temperature) constant. The initial rate in each experiment is compared.

A clock reaction is an acceptable approximation. A small fixed amount of a second reagent is included, and the time t for a sudden visible change is measured. Because the same amount of reaction happens in every experiment, the rate is proportional to 1/t. A graph of 1/t against concentration shows the order. A straight line through the origin means first order.

Acid, propanone and iodine solution are mixed and a timer is started. At regular intervals a 10.0 cm³ aliquot is removed with a pipette and added to excess sodium hydrogencarbonate solution. This neutralises the acid catalyst and quenches the reaction.

The iodine left is titrated with sodium thiosulfate solution, with starch added near the end-point (blue-black to colourless): I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻. The titre is proportional to the concentration of iodine. A plot of titre (or [I₂]) against time is a straight line, so the reaction is zero order with respect to iodine.

The Harcourt–Esson (iodine clock) reaction is H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O. A fixed small amount of sodium thiosulfate and starch is added to each mixture. The thiosulfate reacts with the iodine as soon as it forms, so the blue-black colour appears only when all the thiosulfate has been used.

The time for the colour to appear is measured for different concentrations of one reagent, with everything else constant. Plot 1/t against concentration. A straight line through the origin shows first order with respect to that reagent.