General Chemistry/Limiting Reactants and Percent Yield

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Limiting Reactants

When chemical reactions occur, the reactants undergo change to create the products. The coefficients of the chemical equation show the relative amounts of substance needed for the reaction to occur. Consider the combustion of propane:

{\hbox{C}}_{3}{\hbox{H}}_{8}+5{\hbox{O}}_{2}\to 3{\hbox{CO}}_{2}+4{\hbox{H}}_{2}{\hbox{O}}

For every one mole of propane, there must be five moles of oxygen. For every one mole of propane combusted, there will be three moles of carbon dioxide and four moles of water produced (along with much heat). If a propane grill is burning, there will be a very large amount of oxygen available to react with the propane gas. In this case, oxygen is the excess reactant. There is so much oxygen that the exact amount doesn't matter—it will not run out.

On the other hand, there is not an unlimited amount of propane. It will run out far before the oxygen runs out, making it a limiting reactant. The amount of propane available will decide how far the reaction will go.

Example

2H₂ + O₂ → 2H₂O

If there are three moles of hydrogen, and one mole of oxygen, which is the limiting reactant? How much product is created?

Twice as much hydrogen than oxygen is required. However, there is more than twice as much hydrogen. Thus hydrogen is the excess reactant and oxygen is the limiting reactant. If the reaction proceeds to completion, all of the oxygen will be used up, and one mole of hydrogen will remain. You can imagine this situation like this:

3H₂ + O₂ → 2H₂O + H₂

The reactant that is left over after the reaction is complete is called the "excess reactant". Often, you will want to figure out how much of the excess reactant is left after the reaction is complete. to do this, first use mole ratios to determine how much excess reactant is used up in the reaction.

Here are the ratios that need to be used:

$\left({\frac {\text{moles of limiting reactant}}{1}}\right)*\left({\frac {\text{coefficient of product}}{\text{coefficient of limiting reactant}}}\right)={\text{moles of excess remaining}}$

Percent Yield

Usually, less product is made than theoretically possible. The actual yield is lower than the theoretical yield. To compare the two, one can calculate percent yield, which is ${\frac {actual~yield}{theoretical~yield}}\times 100$ .

The percent yield tells us how far the reaction actually went.