Percent Yield
Percent Yield
The percent yield of a reaction is very important as it tells us how efficient a reaction is. A reaction that has a low percent yield is not very useful in industry. If you are making a new medicine or pesticide and your reaction has a low percent yield then you would search for a different way of doing the reaction. This reduces the amount of (often very expensive) chemicals that you use and reduces waste.
The percent yield can be calculated using: \begin{align*} \% \text{yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100 \end{align*} where the actual yield is the amount of product that is produced when you carry out the reaction and the theoretical yield is the amount of product that you calculate for the reaction using stoichiometric methods.
Example: Percent Yield
Question
Sulfuric acid (\(\text{H}_{2}\text{SO}_{4}\)) reacts with ammonia (\(\text{NH}_{3}\)) to produce the fertiliser ammonium sulfate (\((\text{NH}_{4})_{2}\text{SO}_{4}\)) according to the following equation:
\(\text{H}_{2}\text{SO}_{4}\text{(aq)} + 2\text{NH}_{3}\text{(g)} → \text{(NH}_{4}\text{)}_{2}\text{SO}_{4}\text{(aq)}\)
A factory worker carries out the above reaction (using \(\text{2.0}\) \(\text{kg}\) of sulfuric acid and \(\text{1.0}\) \(\text{kg}\) of ammonia) and gets \(\text{2.5}\) \(\text{kg}\) of ammonium sulfate. What is the percentage yield of the reaction?
Step 1: Determine which is the limiting reagent
We determined the limiting reagent for this reaction with the same amounts of reactants in the previous worked example, so we will just use the result from there.
Sulfuric acid is the limiting reagent. The number of moles of ammonium sulfate that can be produced is \(\text{20.4}\) \(\text{mol}\).
Step 2: Calculate the theoretical yield of ammonium sulphate
From the previous worked example we found the maximum mass of ammonium sulfate that could be produced.
The theoretical yield (or maximum mass) of ammonium sulfate that can be produced is \(\text{2.69}\) \(\text{kg}\).
Step 3: Calculate the percentage yield
\begin{align*} \% \text{yield} &= \frac{\text{actual yield}}{\text{theoretical yield}} \times \text{100}\\ & = \frac{\text{2.5}}{\text{2.694}}(\text{100})\\ & = \text{92.8} \% \end{align*}
This reaction has a high percent yield and so would therefore be a useful reaction to use in industry.
This lesson is part of:
Quantitative Aspects of Chemical Change