Calculating Theoretical And Percent Yields In Combustion Reactions

This video tutorial explains how to calculate the theoretical and percent yield of a compound. It begins with balancing a chemical equation for the combustion of propane, identifying limiting and excess reactants, and then calculating the theoretical yield of CO2. The tutorial concludes with determining the percent yield based on the actual yield obtained from an experiment.

What is the first step in calculating the theoretical and percent yield of a compound?

• Identify the limiting reactant

• Write a balanced chemical equation

• Measure the actual yield

• Calculate the molar mass

In a combustion reaction, what are the typical products formed?

• Carbon dioxide and water

• Carbon dioxide and hydrogen

• Carbon monoxide and water

• Oxygen and water

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

• To balance the chemical equation

• To measure the actual yield

• To calculate the theoretical yield

• To determine the excess reactant

What is the theoretical yield in a chemical reaction?

• The amount of reactant used

• The maximum amount of product possible

• The amount of product actually obtained

• The efficiency of the reaction

How is percent yield calculated?

• Actual yield plus theoretical yield

• Theoretical yield minus actual yield

• Theoretical yield divided by actual yield times 100

• Actual yield divided by theoretical yield times 100

What is the molar mass of propane (C3H8)?

• 50 grams per mole

• 30 grams per mole

• 44 grams per mole

• 36 grams per mole

What is the molar ratio of C3H8 to CO2 in the combustion reaction?

• 1:1

• 1:2

• 1:3

• 2:3

How many grams of CO2 are produced from 30 grams of C3H8 if the reaction is 100% efficient?

• 100 grams

• 80 grams

• 70 grams

• 90 grams

If the actual yield of CO2 is 70 grams, what is the percent yield?

• 80%

• 85%

• 90%

• 77.7%

What does a percent yield of less than 100% indicate?

• The reaction used more reactants than needed

• The reaction was more efficient than expected

• The reaction produced more product than possible

• The reaction was not perfectly efficient