[Lt.Kill]

So for my chemistry experiment, I was to create aspirin. I followed all the steps (cooling solution until some solid formed, using vacuum to dry the solution, etc.). I took two separate trials, one with 0.10 g and another with 0.1468 g of aspirin, from what I was left with. I added 15 mL of 95% ethanol to both (separately). I stirred until fully dissolved. I added 2 drops of phenolphthalein indicator. It took 3 and 4 mL respectively of 0.2046 NaOH for the solution to turn light pink. How do I find the percent purity?

If it helps, I found and hit the melting point of the aspirin when performing a melting point experiment (started melting at 135.8 degrees C and fully melted at 138 degrees C).

My Work: So (for the .1 g case) first I multiplied .2046M x .003L=6.138 x 10-4 mols. Since NaOH is a 1:1 ratio with aspirin, then 6.138 x 10-4 mols * 180.157 g/mol= .110580366 g x 100/.1g=110.58%. Of course I used the same process for the other mass of aspirin.

Just wanted to check my work since that 110 feels weird.

[mjc123]

I'm a bit concerned that you report the molarity of NaOH to 4 significant figures, but the volume only to one. How accurately did you measure the volume of NaOH you added?

[Lt.Kill]

I preformed the actual experiment in lab and that was the concentration on the bottle. I was told to use the complete concentration as it says on the bottle.

[Lt.Kill]

Figured out the my possible sources of error may have stemmed from overshooting the endpoint, unreacted reactants etc. Consider this question solved.