Limiting Reactants and Percent Yield

The video tutorial covers stoichiometry, focusing on converting measurements between reactants and products using mole ratios. It explains the concept of limiting reactants with practical examples, demonstrating how to identify the limiting and excess reactants in a chemical reaction. The tutorial also discusses theoretical yield, actual yield, and percent yield, providing calculations to determine the efficiency of a reaction. Key concepts include using molar mass, Avogadro's number, and balanced chemical equations to perform stoichiometric calculations.

What is the primary focus of stoichiometry in chemical reactions?

• Balancing chemical equations

• Measuring the temperature of reactions

• Converting measurements of reactants or products into different units

• Determining the color of compounds

In the example problem, what is the first step to convert nitrogen gas to ammonia?

• Find the density of nitrogen gas

• Convert liters of nitrogen gas to moles

• Determine the temperature of the reaction

• Calculate the volume of ammonia

What is the mole-to-mole ratio used for converting nitrogen gas to ammonia in the example?

• 1:2

• 1:1

• 3:1

• 2:1

What is a limiting reactant in a chemical reaction?

• A reactant that speeds up the reaction

• A reactant that changes the color of the product

• A reactant that is left over after the reaction

• A reactant that is completely used up first

In the limiting reactant example, which reactant is the limiting one?

• Aluminum

• Ammonia

• Nitrogen

• Chlorine

What is the excess reactant in the aluminum and chlorine example?

• Nitrogen

• Ammonia

• Chlorine

• Aluminum

What does the theoretical yield represent in a chemical reaction?

• The speed of the reaction

• The calculated amount of product expected

• The actual amount of product obtained

• The amount of reactant used

How is percent yield calculated?

• Theoretical yield divided by actual yield, multiplied by 100%

• Actual yield divided by theoretical yield, multiplied by 100%

• Theoretical yield minus actual yield, multiplied by 100%

• Actual yield plus theoretical yield, divided by 2

In the percent yield example, what is the actual yield of iron?

• 105 grams

• 87.9 grams

• 150 grams

• 200 grams

Why should the percent yield never exceed 100%?

• Because it implies more product was obtained than theoretically possible

• Because it indicates a calculation error

• Because it means the reaction was too fast

• Because it suggests the reactants were not pure