In an experiment using acid-base (NaOH solution) titration with an indicator to determine the amount of acetylsalicylic acid in an aspirin tablet (placed in cold ethanol):

1. Why do we not worry if all the ground aspirin tablet dissolves during the titration?

2. If there were other ingredients in the aspirin tablet that were acidic, say vitamin C, could this titration be used to determine the amount of ASA in the tablet?

3. Caffeine, a weak base, is sometimes added to cold medications. How would this compound affect the results for the titration of aspirin?

4. If titrated properly, the neutralized aspirin solutions should fade in color over time. Why might this be? (relates to CO2 in the air).


Here is a point by point response to match your question type.

  1. Aspirin can have a lot of crap in it, coatings, fillers, etc.  Some of these things are expected to be insoluble not only in your cold ethanol but most anything.  So the solids are not to be unexpected.  Since the aspirin is ground you should get enough surface area to extract any of the remaining aspirin from something that remains undissolved.

  2. Not reliably, acetylsalicylic acid is a weak acid, most any other acid is going to throw off your results.

  3. Again, the weak base will end up telling you there is less acid there than expected because it is essentially able to neutralize a small part of the acetylsalicylic acid already there so you won't see it with the titration.  Also it might help to buffer the solution making the endpoint of the titration less distinct.

  4. If you are using something like phenolphthalein which is going to go from colorless in acid to pink in base, the pink color would fade as it becomes more acidic.  And in your example the carbon dioxide from the air could dissolve in the solution yielding carbonic acid (a weak acid) which could back titrate your solution and make it loose its color as it again becomes acidic.


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