A student mistakenly measured out 5.0 g of salicylic acid but correctly measured

Question

A student mistakenly measured out 5.0 g of salicylic acid but correctly measured out 5.0 mL of acetic anhydride. Is this amount of anhydride sufficient to react all of the salicylic acid? How many milliliters of acetic anhydride are needed to completely react all 5.0 g of salicylic acid? Bused on 5.0 g of salicylic acid being the limiting reactant, what will the theoretical yield of aspirin be for this reaction? After the reaction was carried out, the student reported a yield of 3.9 g of aspirin What is the percentage yield?

Explanation / Answer

(b)

Moles of salicylic acid taken = Mass/MW = 5/138.12 = 0.0362

SInce 1 mole of salicylic acid reacts with 1 mole anhydride to produce 1 mole each of aspirin and acetic acid, so moles of aspirin produced theoretically = 0.0362

Mass of aspirin produced = Moles*MW = 0.0362*180.15 = 6.521 g

(a)

Assuming the molarity of anhydride solution as 'x',

Moles of anhydride taken = x*0.005

Moles of anhydride required = Moles of acid taken = 0.0362

So, if x*0.005 >= 0.0362, then the amount oh anhydride taken will be sufficient, else not.

Hope this helps!

(c)

Mass of aspirin produced = 3.9 g

So, % yield = (Mass produced/Theoretical mass) * 100 = (3.9/6.521) * 100 = 59.8%

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