# How can you qualitatively tell when the hydrolysis of aspirin is over?

I'm doing an experiment on the effect of temperature on the time taken to produce acetic acid and salicylic acid. Are there any qualitative observations I can make about when the reaction has finished?

I am using a 0.01 M aspirin in pure water.

• You might want to check this answer and the spectrum therein: chemistry.stackexchange.com/a/150095/79678. You might have to run a series of experiments to get a feel for about how long it takes for the hydrolysis to effectively reach completion.
– Ed V
Apr 27 at 12:27

First I was thinking that maybe a chemical indicator could be used, however the pKa values of all three components are relatively close to each other (especially aspyrin and salicylic acid). It might still be a bit of a help to stick a pH electrode in the system and observe the change - however, this is not exactly something you can see easily.

Another thing that came to mind is that salicylic acid (as all phenols) will form a colorful complex with Fe(III). However, the appearance of the red-ish color will only tell you when the reaction has started, since it's borderline impossbile to observe that the solution is not getting any darker anymore. Again, if you have some machinery, you could stick your solution into a cuvette and put it in a photometer.

Apart from the two ideas above, I don't really see an easy way of observing anything.

Why be qualitative when with just a little more effort you can be quantitative?

Acetylsalicylic acid has a pK$$_a$$ of 2.97, i.e., it is fairly acidic (Ref 1). Salicylic acid, one of the hydrolysis products, is comparably acidic (pK$$_a$$ = 2.79 Ref 2) and its sodium salt is fairly neutral (pH = 6.0 - 8.5 for a commercial material Ref 3; calculations suggest ~7.4 at 0.01 M).

Acetic acid, the other hydrolysis product, is not so acidic: pK$$_a$$ = 4.75, and its sodium salt, in solution, has a pH of 8.375 (Ref 4). This is how the calculation plays out:

pH = 7 + 0.5(pK$$_a$$ + logconcentration)

= 7 + 0.5(4.75 - 2) = 8.375

Approximately. The other ions in solution will probably have some effect.

One more fact: phenolphthalein is colorless below pH = 8.5 and pink/red above pH = 9 (Ref 5).

So, your 0.01 M acetylsalicylic acid will dissolve (soluble in water 3g/L = 0.0166 M). As it hydrolyzes, it will produce salicylic acid with essentially the same strength (pK$$_a$$ ~same) and concentration (1:1). But at the same time, it will generate acetic acid, somewhat weaker, but still titratable.

If you take 10.0 mL of your experimental solution, add a few drops of phenolphthalein indicator solution and then titrate with 0.01 N NaOH (or KOH), at time = 0, it will require 10.01 mL to turn the phenolphthalein red. As the hydrolysis proceeds, the acetic acid produced requires additional alkali to get the pH high enough to get a red color. In fact, at the end of the hydrolysis, it will require 20.01 mL of 0.01 NaOH to turn red. (The decimal places are for instructive purposes, not to indicate accuracy - you know what I mean.)

In fact, you don't need to use exactly 0.01 M alkali - just realize that whatever you need to turn red at time zero will double when the hydrolysis is complete. And you can weigh the amount of alkali added rather than use a burette. And you can do a reaction rate study by checking every 15 minutes or so.