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Question
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).
Answer
Here is a point by point response
to match your question type.
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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.
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Not reliably, acetylsalicylic
acid is a weak acid, most any other acid is going to throw off your
results.
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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.
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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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