4
$\begingroup$

A few months back I designed and carried out an experiment to measure the rate of the reaction between ferric chloride and copper.The dependent variable was the rate of the reaction; the independent variable was the the concentration of ferric chloride.

I searched up the reaction and got it from chemiday. There, it stated it as the following:

$$\ce{2FeCl3 + Cu -> 2FeCl2 + CuCl2}$$

So the method I used to measure the rate of this reaction was to use a spectrometer to find a rising absorbance peak for one of the reactants, which I selected to be: $\ce{CuCl2}$.

Then I got some anhydrous $\ce{CuCl2}$ and made a standard curve (Absorbance vs Concentration) for it, so that I could use that to convert the absorbance at it's $\lambda_\mathrm{max}$ to a concentration value.

I carried out the experiment taking the readings at that $\lambda_\mathrm{max}$ at set time intervals, and I reached the conclusion that the relationship between the concentration of ferric chloride and the rate of the reaction between it and copper was a linear one, at least for the concentrations I tested ($\pu{0.1 .. 0.6 M}$).

But now I'm doing more reading to write my report and I found out that this was a complex reaction, meaning that there were several steps to the reaction. After searching up some papers, the consensus seems to be:

\begin{align} \ce{FeCl3 + Cu &-> FeCl2 + CuCl}\tag{1}\\ \ce{FeCl3 + CuCl &-> FeCl2 + CuCl2}\tag{2}\\ \ce{CuCl2 + Cu &-> 2CuCl}\tag{3} \end{align}

As seen from above, it isn't as simple as the original single reaction that I mentioned earlier. First, $\ce{CuCl}$ is produced, then it becomes a reactant in the second reaction to produce $\ce{CuCl2}$, and lastly $\ce{CuCl2}$ becomes a reactant to produce $\ce{CuCl}$.

I'm seriously questioning my method after finding out that a series of reactions takes place, not a just single one. Although the data that I got and the conclusion that I reached matched nicely with other papers that I've read, I'm not sure if the explanation behing my method is correct.

Any input would be greatly appreciated. Thank you.

$\endgroup$
  • 1
    $\begingroup$ All this mechanism doesn't matter that much as you're reacting solid with liquid. Your experiment is probably okayish, at least linear dependence seems OK. $\endgroup$ – Mithoron Aug 31 '17 at 21:17
  • $\begingroup$ This is new for me; how does reacting solid with liquid make the steps irrelevant? Thank you for your input. $\endgroup$ – Hamza Qayyum Aug 31 '17 at 21:39
  • $\begingroup$ Redox itself is quite fast but only surface of metal can react; that's like you were slowly feeding your solution with metal, and pace of this "feeding" determines change in conc. in solution. $\endgroup$ – Mithoron Aug 31 '17 at 21:57
  • $\begingroup$ The setup I had was to throw all the copper, which was very little, by the way (0.15g), into the solution. So, in terms of 'feeding' the metal, it was all fed at once. $\endgroup$ – Hamza Qayyum Aug 31 '17 at 22:03
  • $\begingroup$ The solid in liquid part doesn't matter. Having a solid in there means that the solution should be saturated in that reactant (unless the reaction is very fast) which is exactly what is desired in most circumstances, as the rate-dependence on copper concentration is being measured. $\endgroup$ – jheindel Aug 31 '17 at 22:41
3
$\begingroup$

This is a perfectly valid method of determining rates of chemical reactions. It is obviously limited to those reactions where either the product or the initial reactant can be monitored easily using spectroscopy as well as reactions which are slow enough to get an absorbance measurement at set intervals. I have used this method multiple times in undergraduate chemistry labs and it works, and in principle it does not matter if there are intermediates.

The intermediates should not matter because the rate of a reaction is defined as the rate at which product is formed. That is, the rate of a first-order reaction is, $$ \text{rate}=k_{tot}[A] $$ If the reaction involves only one elementary step then the prefactor $k_{tot}=k_1$, the rate constant of the one step (assuming no backwards reaction). Note that $[A]$ can be the concentration of either a product or reactant, but you might need a stoichiometric coefficient in there depending on the reaction. On the other hand, if there are intermediate steps, the prefactor can become arbitrarily messy. Maybe there's an equilibrium between two intermediates so you have a term in there $k_2/k_{-2}$ and whatever else you can imagine. This doesn't matter though because your experiment measures the change in concentration at either end, and the slope is then the effective rate constant for the reaction which has all of the intermediate stuff lumped together.

In fact, this method you use is probably the most direct means of measuring a rate in the sense that a rate is defined as a change in concentration over time. You are measuring a change in absorbance over time, so there is only step (and one assumption) in going from your data straight to the reaction rate. That assumption is simply the validity of Beer's law, which is usually not a problem for a large range of concentrations.


As a sanity check, look at this website which has a description of the same experimental setup and what appears to be real data. Note that they are measuring change in absorbance of a reactant while you say you are measuring the change in absorbance of a product, $\ce{CuCl2}$

$\endgroup$
  • $\begingroup$ Thank you very, very much for your input. I haven't studied rate equations as part of my studies yet, but I will certainly study them now so that I can understand that part of your answer. As for Beer's Law, I didn't think of that, but I see a clear pathway now as to how I can write about this. Thanks again. $\endgroup$ – Hamza Qayyum Aug 31 '17 at 22:49

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.