Q: If 21.5 g of $Ag_2C_6H_4N_2SO_7$ is obtained from a reaction where the theoretical yield is 71.4 g, what is the percent yield?

$\frac{21.5}{71.4} \times 100 = 30.1\%$

Q: Suppose a reaction has a theoretical yield of 10.0 g and an actual yield of 7.8 g. What is the percent yield?

$\frac{7.8}{10.0} \times 100 = 78\%$

Q: A student mixes 50.0 g of $C_6H_4N_2SO_7$ with 50.0 g of $Ag_2SO_4$ to produce $Ag_2C_6H_4N_2SO_7$. If the theoretical yield is 71.4 g, but only 21.5 g is obtained, what is the percent yield?

$\frac{21.5}{71.4} \times 100 = 30.1\%$

Question 1

Which of the following best describes the concept of a limiting reactant in a chemical reaction?

Accepted Answer

The reactant that will be completely consumed first, thus limiting the amount of product formed.

Question 2

In the Haber Process, nitrogen gas reacts with hydrogen gas to form ammonia. If you need to produce $1.00 	imes 10^3$ g of ammonia ($NH_3$), which calculation step must you perform first?

Accepted Answer

Convert the mass of ammonia to moles using its molar mass.

Question 3

Given the reaction $KClO_3 (s) + P_4 (s) \rightarrow P_2O_5 (s) + KCl (s)$, if you start with 52.9 g of $KClO_3$ and excess $P_4$, what is the first step to determine the mass of $P_2O_5$ produced?

Accepted Answer

Convert grams of $KClO_3$ to moles using its molar mass.

Question 4

A reaction between $2Sb (s) + 3I_2 (s) \rightarrow 2SbI_3 (s)$ is performed with 1.20 g of Sb and 2.40 g of $I_2$. Which of the following correctly identifies the limiting reactant?

Accepted Answer

Antimony ($Sb$) is the limiting reactant.

Question 5

In a real-world chemical reaction, why is the actual yield often less than the theoretical yield?

Accepted Answer

Because some reactants are lost or side reactions occur.

Question 6

If 21.5 g of $Ag_2C_6H_4N_2SO_7$ is obtained from a reaction where the theoretical yield is 71.4 g, what is the percent yield?

Accepted Answer

$\frac{21.5}{71.4} 	imes 100 = 30.1\%$

Question 7

Suppose a reaction has a theoretical yield of 10.0 g and an actual yield of 7.8 g. What is the percent yield?

Accepted Answer

$\frac{7.8}{10.0} 	imes 100 = 78\%$

Question 8

In the reaction $2Sb (s) + 3I_2 (s) \rightarrow 2SbI_3 (s)$, if the percent yield is 78.2% and the theoretical yield of $SbI_3$ is 3.60 g, how many grams of $SbI_3$ are actually formed?

Accepted Answer

$3.60 	imes 0.782 = 2.82$ g

Question 9

Which of the following steps is NOT part of the process for determining the limiting reactant in a chemical reaction?

Accepted Answer

Calculate the percent yield before identifying the limiting reactant.

Question 10

A student mixes 50.0 g of $C_6H_4N_2SO_7$ with 50.0 g of $Ag_2SO_4$ to produce $Ag_2C_6H_4N_2SO_7$. If the theoretical yield is 71.4 g, but only 21.5 g is obtained, what is the percent yield?

Accepted Answer

$\frac{21.5}{71.4} 	imes 100 = 30.1\%$

Question 11

Why is it important to use a balanced chemical equation when performing stoichiometric calculations?

Accepted Answer

Because it ensures the correct ratio of reactants and products is used.

Question 12

In a reaction, if you have 10.0 g of reactant A and 15.0 g of reactant B, and reactant A is the limiting reactant, what does this mean?

Accepted Answer

Reactant A will be completely consumed before reactant B.

Question 13

Which formula is used to calculate percent yield in a chemical reaction?

Accepted Answer

$	ext{Percent Yield} = \frac{	ext{Actual Yield}}{	ext{Theoretical Yield}} 	imes 100$

Question 14

A chemist reacts 52.9 g of $KClO_3$ with excess $P_4$ to produce $P_2O_5$. If the molar mass of $KClO_3$ is 122.6 g/mol, how many moles of $KClO_3$ are used?

Accepted Answer

$\frac{52.9}{122.6} = 0.432$ mol

Question 15

In the reaction $2Sb (s) + 3I_2 (s) \rightarrow 2SbI_3 (s)$, if you start with 1.20 g of Sb and 2.40 g of $I_2$, and the molar masses are 121.8 g/mol for Sb and 253.8 g/mol for $I_2$, which reactant will be left in excess after the reaction?

Accepted Answer

$I_2$ will be left in excess.

Question 16

A student calculates the theoretical yield of a product to be 50.0 g, but only 40.0 g is obtained. What is the percent yield?

Accepted Answer

$\frac{40.0}{50.0} 	imes 100 = 80\%$

Question 17

Which of the following is a correct sequence for solving a stoichiometry problem involving mass-to-mass conversion?

Accepted Answer

Convert mass of reactant to moles, use mole ratio from balanced equation, convert moles of product to mass.

Question 18

If a reaction is performed and the actual yield is greater than the theoretical yield, what is the most likely explanation?

Accepted Answer

There was a measurement error or contamination.

Stoichiometry and Limiting Reactant: Study Notes for General Chemistry

Study Guide - Practice Questions

Study Guide - Flashcards