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March 29, 2017

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A student prepared aspirin (C9H8O4) in a laboratory experiment using the following reaction.


C7H6O3 + C4H6O3 C9H8O4 + HC2H3O2

The student reacted 2.00 g salicylic acid (C7H6O3) with 2.67 g acetic anhydride (C4H6O3). The yield was 2.00 g aspirin. Calculate the theoretical yield and the percent yield for this experiment.

the first one:

2.00g C7H6O3 x (mol C7H6O3/138.12g C7H6O3) x (molC9H8O4/mol C7H6O3) = .014

is this right??
and for the other one i got .026

i changed .014 mol to C9H8O4g as the theoretical

where did i go wrong??

  • chemistry - ,

    First you didn't add an arrow to separate the products from the reactants. It's easy to get mixed up when you don't know the difference.
    Second you didn't carry out the 0.014 far enough. You were given 2.00 g which is three significant figures so you should have at least 3 s.f. for the moles salicylic acid. That is 2.00/138.123 = 0.01448 moles salicylic acid which will produce, as the limiting reagent, 0.01448 moles aspirin. That IS the theoretical yield but its in MOLES. You were given actual yield in grams; therefore, convert 0.01448 moles aspirin to grams (I get 2.609 g and percent yield = (2.00/2.609)*100 = ?? and round to 3 s.f. You should get something like 77% or so.

  • chemistry - ,

    YOu went wrong by not calculating what .014moles is in grams.

    Yield=massproduct/massreactant.

  • chemistry - ,

    thanks soo much

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