In this lesson, we will learn:
- To recall the types of experiments done to find the order of reaction.
- To understand the iodine clock reaction as an example of an initial rates test.
- To use initial rates experimental data to find the order of reaction.
- In chemistry, to investigate the order of a reaction and reaction mechanisms, you start by doing an experiment. Usually, the experiment is measuring the concentration of one of the reactants; how quickly is it being used up in the reaction?
Depending on what reaction you are studying and the properties of the reactants, there are several ways to do this. These include:
This lesson will primarily focus on how we use these experiments to get measurements of the rate.
- Measuring the volume of gas produced by a reaction. This is common for decomposition reactions and quite a few reactions between two aqueous solutions.
- Measuring the time taken for a precipitate to form, which blocks out visibility through the solution. For example, looking through your reaction flask at a marked surface underneath it and recording when it is is no longer visible.
- “Clock” reactions, such as iodine clocks.
- The iodine clock is a well-known type of reaction that exploits one of its key properties; it reacts with starch to make a deep blue solution.
If we investigate a reaction that produces iodine (for example, where I- is oxidised by hydrogen peroxide) then we know a blue-black colour will form when iodine is produced:
2I- (aq) + 2H+ (aq) + H2O2 (aq) → I2 (aq) + 2H2O (l)
There is a problem in that this colour will appear almost immediately. To get around this, we add sodium thiosulfate which reacts with I2:
S2O32- (aq) + I2 (aq) → S4O62- (aq) + 2I- (aq)
If we add starch and a consistent, small amount of thiosulfate to our solution, then we know the first instance of a blue-black colour in our solution shows the iodine reacting with starch, which will only happen when there is no more thiosulfate remaining to react the I2 away.
This could be repeated with different concentrations of iodide, hydrogen peroxide, different temperatures and so on, to see how the reaction rate changes. The important point is that only one variable is changed per experiment!
- Investigating the rate with the iodine clock reaction is an example of an initial rates test. In these tests, you’re using the reaction’s initial rate, its highest and most distinct, to compare different reaction conditions.
- For experiments that produce gas products, find the slope of their volume of gas over time graphs. Plot volume of gas over time; it will usually look something like this:
Take a section of the beginning of the graph, and the slope is just volume over time. If you compare it to an experiment where the concentration was half, or double, or the temperature was higher you could get a different slope.
This is going to reveal how the rate depends on these factors.
- With the iodine clock reaction, you are taking one measurement of time. If you just want to compare reaction conditions, this is all you need per experiment. Run another experiment(s) with different initial concentration, and then find 1/time for each time recorded for these experiments. Plot a graph of initial concentration on the x-axis against 1/time on the y-axis.
Because 1/time is a measure of reaction rate, a graph of reactant concentration against 1/time reveals how rate depends on concentration.
- A zero-order reaction will have a gradient of zero because concentration is independent of the rate – the flat line is telling you rate does not change when you increase or decrease reactant concentration.
- A straight line with a gradient means you have a first-order dependency. The positive constant slope tells you there is a fixed relationship between rate and reactant concentration.
- A curve could mean a second order dependency, but it is not clear just from this.
Remember that rate is calculated by:
Rate = k [A]n, where n is order of reaction with respect to [A].
Taking the logarithm of both sides, we can find:
Log[rate] = log[k.An]
A property of logarithms is that log[x.y] = log[x] + log[y]
Log[rate] = log[k] + log[An]
Another property of logarithms is that log[xy] = ylog[x]
Log[rate] = log[k] + nlog[A]
This fits the form y = mx + c, where n = m, the slope of the straight line. Plot log[A] on the x-axis and log[1/t] on the y-axis, and the slope of this straight line will be n, the order of the reaction.