Practical investigations
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Metal-acid reactions: a second tutorial
More on this popular investigation. Here is the first metal-acid reactions tutorial:
Investigating the rate of the reaction of metals with acids
This apparently simple reaction continues to produce surprises and to reveal procedural difficulties. It seems that student investigation, which, unlike genuine research, does not have to be totally original work, has nevertheless strayed into the research arena. It seems that students, to their very considerable credit, are eager to know what the mechanism actually is.
Student investigations and research
It might therefore be helpful to review a few fundamentals about research.
The first thing to get clear is that there is no ultimately authoritative document anywhere which states with certainty what the mechanism of a reaction is. We are not in the business of trying to second-guess something which is known but which is hidden from us. There is no ‘right answer’.
However there are many wrong answers. Investigators ask a question, propose an answer on the basis of available evidence, and then try to demonstrate that this answer cannot be correct. If no-one can demolish the answer, it is accepted as correct until, eventually, someone does find fault with it or proposes an answer which is even more appealing than the first. When this happens the process begins all over again. If this view of things seems to owe much to the philosopher Karl Popper, I am not apologising!
An essential element in this process is publication. Researchers publish their conclusions and the evidence on which they are based. Other workers then join in, sometimes in support of what is proposed, sometimes in apparent contradiction. It is this step which is missing in the case we are now considering.
Students write up their work and it is assessed by their teachers, and possibly by Coursework Moderators. These assessors may spot flaws in students’ arguments and take these into account when rewarding the quality of the work. This is exactly what they are supposed to do, but they are not part of the research process because there is no obligation to join the discussion.
Consequently, we are neither confirming evidence nor advancing the debate.
What can we take as correct?
1 Is it true that in the magnesium-acid reaction, the order is 2 with respect to hydrochloric acid? In the 1970 article by Williams and Hacker Download the article from here, they considered this to be true over all the concentrations they studied.
2 Is it true that in the same reaction, the order is 1 with respect to sulphuric acid? Williams and Hacker did not think this to be the case. They thought the reaction was second order except that there were deviations from second order kinetics at high concentrations. Recent questions to Re-act have suggested first order kinetics.
3 Is there any evidence that either of these orders of reaction is different if a different experimental method is used? (See below for comments about experimental methods.)
4 Is it true that if a weak acid such as ethanoic acid is used, the reaction is slower than it is when hydrochloric acid of the same concentration is used?
5 Is it true that, if the answer to the previous question is ‘yes’, the reaction is not as much slower as is suggested by the Ka of the acid?
6 Has anyone tried to find out if the rate of reaction is affected by agitation of the reaction mixture?
What steps are there in a metal-acid reaction?
What, if anything, have I missed in this list?
• Diffusion of some particle such as a hydrated proton towards the metal surface
• Collision between this particle and the metal surface
• Loss of electron(s) from a metal atom (or metal surface - see the model of a metal structure)
• Acceptance of an electron by the attacking particle
• Breakaway of a hydrogen atom from the attacking particle
• Breakaway of the metal ion from its lattice
• Hydration of the metal ion
• Diffusion of the metal ion away from the surface
• Combination of hydrogen atoms into molecules
What is the nature of the ‘particle’ mentioned here? If it is a hydrated proton, H3O+(aq), the answers to questions 4 and 5 above could be important.
Williams and Hacker proposed two mechanisms for the rate-determining stage. These remain unchallenged if their experimental conclusions are supported by more recent work, though they tend to combine steps which I have shown as separate in the above list.
If the diffusion in the first step is rate-determining, the answer to 6 above could be important, since hydrogen bubbles would impede diffusion and agitation would tend to remove bubbles. Williams and Hacker have some results here tending to confirm that agitation does affect rate.
In future, we suggest that mechanistic discussion should start from the Williams and Hacker article.
All references to the Williams and Hacker article and to this summary, as well as the first Metal-acid reactions tutorial, must be acknowledged in the write-up of your investigation.
Experimental methods
Some students are finding it difficult to match experimental methods with processing methods. In particular, some seem uncertain whether a given method is ‘continuous’ or ‘initial rate’ in design. In answer to one such query, we offered the following note.
Continuous method
This method only works if the concentration of only one reactant is varying to a significant extent - obviously the case if there is only one reactant.
Let us take a hypothetical example to illustrate this. Suppose you have a reaction involving a substance X reacting with an alkaline solution which, of course, contains OH- ions. The reaction has the equation:
X + OH- -> products
You set up one batch of reaction mixture with the concentration of X in large enough excess for it not to vary significantly. You start the clock when the mixture is made. At measured time intervals you withdraw a sample of the mixture and titrate it with an acid. The more acid you need, the greater the OH- ion concentration. As the reaction goes on, this titration value will decrease because the OH- ions are being used up. You now have a list of times and a list of titration values which are proportional to the concentrations of OH- ions. You plot titration value against time.
If your graph is a straight line, the reaction is zero order with respect to OH-.
If your graph is a curve with constant half-life, the reaction is first order with respect to OH-
If your graph is a deeper curve with a half-life which increases markedly, the reaction is probably second order. Orders higher than this are rare, and outside the scope of the half-life method.
Notice the characteristic feature: you only have one reaction mixture, and you make continuous measurements on it.
Initial rate method
Again, let’s take an example. You have a 1-cm long piece of magnesium, and you drop it into an excess of hydrochloric acid with concentration 1 mol dm-3. You time how long it takes for the piece of magnesium to react completely. In this time, the acid concentration will not have changed much so you assume that it has not changed at all.
You repeat the experiment with another 1 cm of magnesium but this time you use acid of concentration 0.9 mol dm-3. It takes a bit longer this time.
You repeat again several times with a lower concentration each time.
You now have a list of concentrations and a corresponding list of times.
Here’s the tricky bit: for each of these times, t, the same amount of magnesium has reacted so the rate of the reaction (the initial rate) is given by:
moles Mg / t
Since the amount of magnesium is the same for each separate experiment, the initial rate is proportional to 1/t.
Notice the characteristic feature of the method: you do lots of experiments, each one giving you one piece of information.
You now compare with each rate equation in turn so see which fits the data best.
You now have two ways to proceed, a clumsy way and an elegant way!
The clumsy way is to plot a graph of rate against concentration. If n= 1 the graph will be a straight line. If it isn’t, try plotting rate against concentration2. A straight line resulting from this indicates an order of 2. The weakness of this method is that experimental error may result in the line not being quite straight in either case. The elegant method gets round this.
Taking logarithms of both sides of the rate equation gives:
Lg(rate) = lg(constant) + n lg(concentration)
If you plot lg(rate) against lg(concentration) you should get a straight line whatever the order of the reaction. Draw the best straight line you can and find its gradient, which is n, the order of the reaction.
Something new
Finally, there was a sudden rash of questions about the feasibility of quite a novel procedure for following the metal-acid reaction.
The suggestion involved taking a decently large volume of quite dilute acid (the vagueness here is deliberate) and adding magnesium ribbon to it, starting timing. At measured time intervals, pipette volumes of acid would be removed and analysed by titration with sodium hydroxide solution of known concentration. The reaction mixture would need stirring, but quenching of the samples before titration should be unnecessary.
Has anyone tried this? Does it work? Please use the React button below.
Always carry out a risk assessment and check with your teacher before starting any practical work.
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Vendrus Sci
Sulphuric acid: VERY clear change from first order to second order between 0.8M and 1.2M. As in very near perfect half-lives.
11 February 2007
Robyn
A really useful site, only sorry I didn't stumble on it last week before I handed by c/w in! But just thought I'd say it put my mind at rest - I found that both hydrochloric and sulphuric acid are 2nd order and was a bit worried that no one else had!
05 April 2006
Dante Ierubino
Agitation increases the activation energy!!! arg!!!
This is because the local temperatures and concentrations caused by the exothermic reation and non-agitation causes massive errors! be warned! I made this mistake already!
30 March 2006
updated: 12 January 2007
