Limiting Reactants and Percent Yield

Mr. Anderson explains the concept of limiting reactants and percent yield in chemical reactions. Using analogies like cookie recipes, he illustrates how limiting reactants determine the extent of a reaction. He provides examples of hydrogen and methane combustion to show how to identify limiting reactants. The video also covers the calculation of limiting reactants using moles and mass, and explains percent yield, highlighting the difference between actual and predicted yields.

What is a limiting reactant in a chemical reaction?

• The reactant that does not participate in the reaction

• The reactant that is completely used up first

• The reactant that is left over after the reaction

• The reactant that is in excess

In the hydrogen combustion example, what was the limiting reactant?

• Carbon dioxide

• Water

• Hydrogen gas

• Oxygen gas

During the combustion of methane, what happens to the excess oxygen?

• It forms additional water molecules

• It remains unreacted

• It forms carbon monoxide

• It reacts with nitrogen

What is the first step in determining the limiting reactant in a reaction?

• Identify the products

• Calculate the percent yield

• Balance the chemical equation

• Convert masses to moles

Why is it important to identify the limiting reactant in a chemical reaction?

• To balance the chemical equation

• To predict the amount of product formed

• To calculate the reaction rate

• To determine the amount of excess reactant

What is the predicted yield in a chemical reaction?

• The amount of reactant used

• The amount of reactant left over

• The actual amount of product obtained

• The amount of product expected if all reactants react completely

How is percent yield calculated?

• Predicted yield minus actual yield, divided by predicted yield

• Actual yield divided by predicted yield, multiplied by 100

• Actual yield plus predicted yield, divided by 2

• Predicted yield divided by actual yield, multiplied by 100

What does a percent yield of 90% indicate about a reaction?

• The reaction is highly efficient

• The reaction is inefficient

• The reaction produced no product

• The reaction used all reactants

Why is the percent yield never 100% in real reactions?

• The reaction is always exothermic

• All reactants are used up completely

• Some reactants are always lost or do not react

• The products are always pure

In a chemical reaction, why might a chemist choose to make the most expensive reactant the limiting reactant?

• To ensure all reactants are used

• To maximize the reaction rate

• To increase the temperature of the reaction

• To minimize waste of costly materials