Chemistry- Aspirin lab question

Question

In lab today, we measured the actual yield and melting point of aspirin that we had made in lab last week.  Our actual yield was 2.87 grams.  The theoretical yield was 2.9 g, so the percent yield was 98.96%.  When we tested the melting point, the range was 106-118.  Since the melting point of pure aspirin is 138-140, this means that our aspirin was impure.  My question is what would have caused us to have a high percent yield, but low purity?  What went wrong?  The aspirin was completely dry, so it wasn't that there was water making the aspirin heavier, giving us a higher actual yield and therefore a higher percent yield, but obviously there was something in our aspirin that was making it heavier right? Or could we have measured the melting point wrong?

Answer 1

Hello Dana,

 

It looks like the Aspirin still had insoluble impurities in it, causing a melting point depression and a high percent yield. This may have occurred during the early stages of the crystallization process.

 

During the crystallization process, if the Aspirin and solvent are not filtered (Gravity filtration) before starting, then impurities such as dirt and other contaminants could bond to the crystallizing Aspirin. This error would translate all the way to your melting point and percent yield because the Aspirin is no longer pure.

 

 

Answer 2

That's a whopping melting point depression -- a ridiculously high level of impurities, or substantial water. If your synthesis was terminated (quenched) by adding water, you may have trapped some in the crude product. Did you have a purification step of recrystallizing from anhydrous ethanol (which would tend to remove water, acetic acid, and phosphoric acid impurities)?

Also, when you say that the product was completely dry, as a chemist you'd better be ready to provide some evidence. Merely drying e.g. at 110C to a constant weight is not evidence! Water may still be physically trapped inside crystals. Karl Fischer titration is what you need, with perhaps some modification of chamber solutions if a substantial amount of free salicylic acid is present.

Though (depending on your equipment and experience) substantial error in mp measurement itself is unlikely, there may be a few contributing factors. When you loaded your melting point tubes, did you push crystals in by ungloved fingers (I sure hope not!) Also, if you have an automated mp detector, things can interfere (melted sample can be viscous and trap air bubbles, or even decompose -- these are unlikely with aspirin, however)

Answer 3

two things point to the fact that there were impurities in your sample: a broad mp range, and a large deviance from the expected value.

another red flag is that you got nearly 100% yield in the chemical reaction. most organic reactions don't go all the way to completion since they are equilibrium driven.

recrystallization of the product should sharpen up the melting point range and increase it as well.

Answer 4

Hi Dana,

 

You are comparing the purity based on published melting point of commercially available Aspirin and the one that you isolated in the lab. Assuming that you followed the same protocol and ingredients, yours is likely to contain some water vapor since commercially available aspirin undergoes additional treatment to remove every trace of water vapor. Lab equipment is not always equally efficient in that aspect.   That is the most likely scientific and logical explanation of lower melting point. I work in the lab and our vacuum is not  working at optimum efficiency, therefore, it takes longer to dry any sample.

 

I hope this helps.

 

 

Answer 5

My thinking is that weight even if 100% yield is not a guarantee of purity. They are different values. when the impurity is found it should add up to your sample weight. Is there a given magin of error. And the option to retest is available.