Kinetics Experiments - The Nuts and Bolts

Last week in lab we performed a kinetics experiment in which we observed the iodination of acetone. For those of you who may be curious, there were a number of things that went into the planning of that experiment that are required to make it work as smoothly as it did. There is a certain art to experimental design, but it is all firmly rooted in the science that is being observed. For the kinetics of the iodination of acetone, the experimental design considerations are (in my opinion) fascinating, largely because they all make very good sense and can be understood with a Gen-Chem-level of knowledge.

1. Relative concentrations of the reactants. Some of you may have noticed that the concentration of the iodine was very much less than the concentration of the other reagents. This was not an accident. First, since we were observing the color of the solution and that color was due to the iodine, the concentration of iodine had to be high enough to be easily observable but low enough to react in a reasonable amount of time. In addition to this very practical consideration, there is a chemical reason for the concentrations used. The rate of a chemical reaction is dependent upon the concentration of the reactants. As “good” scientists, we always want to design our experiments in such a way that only one variable is changing. If, for example, the concentration of iodine was 1M and the concentration of acetone was also 1M, then as the iodine concentration changed (and affected the rate of the reaction), the acetone concentration would also change and would also affect the rate of the reaction. Although this data could be mathematically interpreted, it would be much more convenient if only one of the concentrations was changing during the reaction. Is this possible? Of course not. Iodine cannot disappear unless it reacts with acetone (in this reaction). BUT, we can make it seem like the acetone concentration is not changing by making the initial concentration of acetone SO high, that the change will be very small, indeed “negligible”, compared to the change in the iodine concentration. To put some numbers to it, if the initial concentration of iodine is 5mM and the initial concentration of acetone is 1.000M and the reaction is allowed to proceed, when all of the iodine has reacted, the final concentration of acetone would be 0.995M. Yes, the concentration has changed, but it has changed by only 0.5%. This means that the change in rate caused by the change in concentration of acetone can be ignored. The reaction (and reaction rate) will observe the rate law order with respect to iodine only.
2. Limits of the Spec-20s. When we did the experiment, we said that the sample did not have to be put in the spectrometer until most of the color had faded. If anyone put their samples in the Spec-20 immediately, you might have noticed that the Spec-20 was unable to read the absorbance of the solution until most of the iodine color faded. This can be explained by thinking about the nature of “absorbance”. Absorbance is a logarithmic scale related to percent transmittance. An absorbance of “0” is equivalent to a percent transmittance of 100%. If the percent transmittance drops to 10%, the absorbance is “1”, which means 90% of the light that is shining on the sample is being absorbed. If the absorbance climbs to “2”, it means that only 1% of the incident light is getting through the sample (99% is absorbed). An absorbance of “3” means only 0.1% of the light is getting through. As we can see, increasing absorbance ver quickly decreases the amount of light getting through the sample. This brings us to the limits of the detector we are using. It can pretty easily discern the difference between 5% of the light being absorbed and 85% of the light being absorbed, but it's less reliable when trying to distinguish 97.8% absorbed and 98.5% absorbed. It is usually best practice to design experiments so the absorbances used are ~1.5 or less. If the absorbance is too high, the data will either get very “noisy”, or will just be abnormally low, making trends that should be linear appear curved.
3. Aggregation of solutes. This probably wouldn't have been a big problem in our iodination of acetone experiment, but Beer's Law assumes that the solute particles are separate and independent in solution. If the concentration of colored solute is relatively high, it's possible that the colored solute particles will start to interact in a way that will make them absorb differently than if they were truly separate and independent in solution. This is often solved by addressing the instrument limits mentioned in #2; if the sample is diluted sufficiently bring to the absorbance down below ~1.5, it is usually dilute enough to avoid interactions between solute particles.

There are other considerations, but these are the big ones that often come up in experimental design for kinetics experiments and any experiment where we are observing color using a spectrometer. For most of our experiments, we set things up behind the scenes to minimize or eliminate these problems, but I hope that some of you were curious enough to wonder why the experiment was set up the way it was.

See you in the morning.