The Half Decent Pharmaceutical Chemistry Blog

Look at these two molecules: obviously they are two xanthines and rather similar too. Still, one is Dyphylline (aka diprophylline) and the other is etofylline and, if you're asked to distinguish them, I'd better not be in a hurry and have good music with you.

The IR spectra, for example, look exactly the same. Annoyingly, even the melting point can't help you very much: both melt between 160°C and 165°C.

And don't fool yourself: yes, go on trying to see a different absorbance (273 v. 270) with the UV spectroscopy...rumour has it that someone is still in the lab hopelessly trying to detect any difference

These molecules are so similar to each other to make the job in hand look impossible.
Unless you fiercely fight the enemy here: water.

Your only option is acetic anhydride (99%). By the way, once again both are soluble in this aprotic solvent, but this is just the beginning.

The mixture boils for 15 minutes at 160°C. Meanwhile, the "magic" happens: both substances are capable of reacting with the anhydride, yielding acetylated derivatives, whose melting points, however, are fortunately different.



The derivative of dyphylline melts between 142 and 148 degrees, while acetyl etofylline needs approximately 120.

But this is just the end of the work. Before you could reach a verdict, the flask has to cool down to room temperature, so that a mixture (20:80) of ether and petroleum could be added.

Then, there is the trickiest part: you place the flask on ice for about 20 minutes. Meanwhile, you regularly stir the solution, increasing the chances of yielding a nice crystalline precipitate of the acetate by rubbing vigorously the inner of the flask every so often.

So you can't go out of the lab to check your inbox...

Finally, the precipitate is purified (through filtration), washed with the same mixture of ether and benzine, recrystallized from alcohol and dryed much as possible (which means this procedure takes 30 minutes, at least).

To sum up, the message here is rather direct: water is the enemy, ether is your best friend.

And you'd better take your MP3 player with you in such a circumstance, because the lab is cold and depressing when everybody has left.

At the end of every lab course, I always take home some worthless souvenir. This time, after my last analytical chem lab, I chose one those small ignition tubes, used for the Lassaigne's test.

Although the second half of the course dealt with quantitative determinations (HPLC, GC, etc.), there's no doubt this small object epitomizes the first part, whose aim was to teach us how to identify unknown substances.

Obviously, the Lassaigne's test plays a key role in this type of analysis. It tells you whether your substance contains nitrogen, sulfur, both or neither. At least, it should.

Thing is a great part of my colleagues had hard times with it. That's no surprise: the test often gives you false positive or negative results, but, frankly, they are not that tough to detect.

I came across a false positive twice but, both times, I was pretty confident those results were misleading: the blue colour, for example, which should be the evidence of the presence of nitrogen, was too pail and/or not bright enough.

As I'm sure you know, this test is also known as the Sodium Fusion one: working by the book, in fact, metallic sodium ought to be used. However, since that would be too dangerous, according to my professor, we used a 1:2 mixture of magnesium dust and anhydrous potassium carbonate.

You put a bit of your substance in the tube, twice of the mixture and start heating.

So, in the end, the decomposition of the molecule yields an alkaline solution which may, eventually, contain cyanide and sulphide salts.

I actually saved a tube from a sad destiny: before you begin to look for the two elements, the test tube has to be quickly put in a beaker of cool water (well, room temperature). The ignition tube breaks due to the shock and its black, precious content is released and ready to be analyzed.


For what concerns the previous labs, I stole a Pasteur pipette, some potassium permanganate in a test tube (with a cap, of course) and some atropine (actually, a lot of atropine!).

I am sorry but I don't like analytical chemistry very much. In particular, I reckon working in an analytical chemistry lab with other 50 students (most of whom have no interest in what they are doing) to be a terrible waste of time.

That's partly due to the nature of Italian University, where they assume everyone must graduate, in the end, and no matter how long this will take. This means your lab-mates are, technically, idiots who shouldn't even think of graduating, if this was a responsible country.

Another waste of time is having three, MANDATORY courses (+lab) of analytical chemistry, all dealing (almost) with the same things.

I'm currently doing my third and, thank God, last lab course. Being the last, it's also the best one: we can at least learn useful techniques.

Now, the lab is split in two parts: surprisingly, the first half (that finished two weeks ago) is the most important. You work on your own and have to recognize substances.

In the latter, briefly, we work in pairs and concentrate on technical approaches to quantitative assays.

The first part ended with a sort of final test which, oddly, doesn't select who is suitable for the second part (where queues are, unsurprisingly, the major problem), but provides an entry test for the oral exam (which God only knows when will take place).

Students are guided through lists of substances (all those with a monograph on the European Pharmacopoeia), according to common properties: inorganic, organic, salt of organic substance, eventual presence of N and/or S, solubility in water or NaOH, etc.

Tables were created by my professor and published in her book (so, we all had to buy her crap).

Of course, the last identification has to be done following what is written on the E.P.

So, on the day of the "exam", my substance was organic, without nitrogen (Lassaigne Test gave easy to detect false positive) or sulfur, very poorly soluble in water (while the said book simply described it as insoluble), freely soluble in NaOH, ethanol and methanol.

Found the group, I queued for the IR spectrophotometer. Here is the very spectrum I did on that day.

Now, before showing you the silly table, I must explain the odd way they display IR spectra: instead of simply giving you the wavenumbers of the most characteristic peaks from the highest to lowest, they proudly report them according to their intensities.

Here it is.



I was a bit confused: the letter n, moreover, means that those data are not from literature, but where collected by people of the department. So, to quote my professor, they had been carefully checked.

Oh, that makes me feel really good!

My peaks matched those reported, but the order was different.

However, once checked the melting point, I had no doubt I had found the right one.

As it turned out, my IR skills are much better than those of the people who registered the spectra for my professor.

Mind you, can you recognise the substance?