Question d8778

Oct 26, 2015

Here's what's going on.

Explanation:

I can make an educated guess about what your experiment is, but the question can be answered without knowing the exact reaction.

Notice that the student added less acetic anhydride, ${\text{CH"_3"CO"_2"CH}}_{3}$, than what was required for the reaction.

The stoichiometric calculations determined that you need $\text{4.0 mL}$ if acetic anhydride, but the student only added $\text{0.4 mL}$.

This means that the acetic anhydride will most likely become a limiting reagent, since you have less than what you need.

As a result, the theoretical yield of the reaction, which tells you how much product you should expect to see for a 100% yield, will be affected, more specifically it will decrease.

The actual yield, which represents the amount of product you actually get if you take into account the reaction percent yield, will be affected as well.

This happens because you now have less reactants actually reacting. The acetic anhydride will limit the amount of (presumably) salicyclic acid that takes part in the reaction.

The percent yield will not change. If the reaction ahs an 80% yield, for example, then it will convert 80%# of the moles of reactants that take part in the reaction to products, regardless of how many moles you have.

I recommend using so actual values, which I assume you have, to test how the theoretical yield, the actual yield, and the percent yield are affected by this accident.

SIDE NOTE Are you making aspirin?