Limiting Reactants and Percent Yield

The video tutorial covers the concepts of limiting reactants and percent yields in stoichiometry. It begins with a simple analogy of making sandwiches to explain limiting reactants, followed by a detailed example of a chemical reaction between hydrogen and oxygen to form water. The tutorial then explains theoretical and actual yields, and how to calculate percent yield. A step-by-step calculation is provided to determine the limiting reactant and percent yield for a given chemical reaction.

What is the limiting reactant in the sandwich analogy if you have 6 slices of bread, 12 pieces of lunch meat, and 5 slices of cheese?

• Cheese

• Bread

• Lunch meat

• None of the above

In the reaction of hydrogen and oxygen to form water, which reactant is in excess if you have 10 moles of hydrogen and 7 moles of oxygen?

• Hydrogen

• Oxygen

• Water

• Both are limiting

What does the percent yield represent in a chemical reaction?

• The speed of a chemical reaction

• The ratio of actual yield to theoretical yield, expressed as a percentage

• The amount of product formed in a reaction

• The amount of reactants used in a reaction

If you have 68.3 grams of hydrogen and 85.4 grams of oxygen, which is the limiting reactant?

• Oxygen

• Water

• Hydrogen

• Both are limiting

What is the theoretical yield of water if all the hydrogen reacts?

• 68.3 grams

• 609 grams

• 96.2 grams

• 86.4 grams

How do you calculate the percent yield of a reaction?

• Subtract actual yield from theoretical yield

• Add actual yield and theoretical yield

• Divide actual yield by theoretical yield and multiply by 100

• Divide theoretical yield by actual yield and multiply by 100

What is the actual yield of water collected in the problem-solving example?

• 86.4 grams

• 96.2 grams

• 609 grams

• 68.3 grams

What is the percent yield of water in the example problem?

• 89.8%

• 96.2%

• 100%

• 68.3%

Which of the following is NOT a step in calculating the theoretical yield?

• Using mole ratios from the balanced equation

• Multiplying by the molar mass of the product

• Converting grams to moles

• Subtracting the mass of reactants from the mass of products

Why is the actual yield often less than the theoretical yield?

• Due to experimental errors and incomplete reactions

• Because the theoretical yield is always overestimated

• Because reactants are always impure

• Due to incorrect stoichiometric calculations