Experiments to Find the Order of Reaction

The order of reaction with respect to a reactant cannot be read off the balanced equation. It has to come from experiment. Two practical approaches dominate A level courses: the initial rates method (often using a clock reaction) and the continuous monitoring method. Both give you sets of concentration and rate data that you can compare to decide whether a reactant is zero, first or second order. Before working through this page it helps to be comfortable with the basic ideas in introduction to kinetics, since the definitions of rate and order are assumed here.

In the initial rates method, you run a series of separate experiments. In each one you change the starting concentration of a single reactant while keeping every other concentration (and the temperature) fixed, then measure the initial rate, the rate right at the start before any concentration has changed much. Comparing pairs of experiments where only one concentration has changed tells you the order with respect to that reactant.

If doubling a reactant's concentration:

• leaves the rate unchanged, the reaction is zero order in that reactant;
• doubles the rate, the reaction is first order in that reactant;
• quadruples the rate, the reaction is second order in that reactant.

This works because the rate equation has the form \( \)rate\( = k[A]^m[B]^n \), and changing only \([A]\) by a factor while holding \([B]\) constant isolates the effect of the exponent \(m\).

Measuring an initial rate directly is hard, so many experiments use a clock reaction instead. The classic example is the iodine clock reaction. A small, fixed amount of sodium thiosulfate and starch indicator is added to a mixture that slowly produces iodine (for example, from potassium iodide and hydrogen peroxide, or from a persulfate). The thiosulfate reacts with iodine as fast as it forms, keeping the solution colourless, until the thiosulfate runs out. At that point, free iodine builds up and the starch indicator turns the solution blue–black almost instantly.

Because the same small amount of iodine has to be produced each time before the colour appears, the time \(t\) taken is inversely proportional to the initial rate: \( \)rate\( \propto \dfrac{1}{t} \). By timing how long the colour change takes for different starting concentrations of a reactant, you get a set of relative rates without ever measuring a rate directly, which is exactly the data the initial rates method needs.

Instead of running many separate experiments, continuous monitoring follows a single reaction mixture all the way through. A property that changes smoothly with concentration is tracked over time, for example colour intensity by colorimetry, volume of gas produced, pressure, or conductivity. The result is a concentration–time graph like the one below.

To turn this single curve into order data, you draw tangents at several points along the curve and measure their gradients. Each gradient is the rate at that instant, paired with the concentration at that same instant. This gives several (rate, concentration) pairs from one experiment, which can then be plotted and analysed exactly as if they had come from separate initial rate experiments. The full technique for extracting gradients from these curves is covered in factors affecting rate of reaction and its companion topics on rate graphs.

Once you have several (concentration, rate) pairs, whether from clock experiments or from tangents on a continuous monitoring curve, plotting rate against concentration shows the order directly from the shape of the graph.

A first order reactant gives a straight line through the origin, since rate is directly proportional to concentration:

A second order reactant gives a curve, because rate is proportional to the square of concentration:

A zero order reactant would give a horizontal line, since changing its concentration has no effect on rate at all. Recognising these three shapes, straight through the origin, curved upward, or flat, is usually enough to assign an order without any calculation.

Three experiments are carried out, changing the concentrations of reactants \(A\) and \(B\) and measuring the initial rate each time.

Comparing experiments 1 and 2, only \([A]\) has doubled and the rate has also doubled, so the reaction is first order in \(A\). Comparing experiments 2 and 3, only \([B]\) has doubled but the rate has quadrupled, so the reaction is second order in \(B\). The rate equation is therefore \( \)rate\( = k[A][B]^2 \), and the overall order is \(1 + 2 = 3\).

This pairwise comparison is the core skill this topic is building toward: once you can find each individual order this way, you can go on to calculate the rate constant \(k\) and write the complete rate equation, which is covered in finding the order of reaction and rate equation.

The initial rates (clock) method needs several repeat experiments but each one is quick, making it well suited to reactions where a distinctive colour change (like the iodine clock) can be triggered. Continuous monitoring needs only one run but requires a property that can be measured accurately throughout, and it produces far more data points from a single trial. Both methods rest on the same underlying idea: change a concentration, measure how the rate responds, and use that response to assign an order.

Related lessons

• Uk A Level Chemistry