# The theoretical yield of a reaction is 82.5 grams, but the reaction actually yields 12.3 grams less than expected. What is the percent yield for this reaction?

Apr 5, 2016

85.1%

#### Explanation:

Percent yield is defined as the ratio between the actual yield and the theoretical yield of the reaction, multiplied by $100$

$\textcolor{b l u e}{| \overline{\underline{\textcolor{w h i t e}{\frac{a}{a}} \text{% yield" = "what you actually get"/"what you should theoretically get} \times 100 \textcolor{w h i t e}{\frac{a}{a}} |}}}$

So, you know that your reaction has a theoretical yield of $\text{82.5 g}$. This means that if all the moles of reactants that take part in the reaction end up producing moles of product according to the reaction's stoichiometric coefficients, you will get $\text{82.5 g}$ of product.

In other words, the theoretical yield tells you how much product is produced for a 100% yield.

Now, the reaction is said to produce $\text{12.3 g}$ less than expected. This means that you only collected

$\text{actual yield" = "82.5 g" - "12.3 g" = "70.2 g}$

The reaction's percent yield will thus be

"% yield" = (70.2 color(red)(cancel(color(black)("g"))))/(82.5color(red)(cancel(color(black)("g")))) xx 100 = color(green)(|bar(ul(color(white)(a/a)"85.1%"color(white)(a/a)|)))

The answer is rounded to three sig figs.