Calculating the Yield of Iron

Oct 13, 2025 by pynkocean.com 107 views

Introduction

Iron oxide, also known as rust, is a common compound found in nature. It can be reduced to its metal form through a chemical reaction with carbon monoxide (CO). The reaction is as follows:

Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)

In this article, we will explore how to calculate the yield of iron obtained from 1.00 kg of iron oxide and determine the percent yield if 654 g of iron is obtained from the reaction.

Calculating the Yield of Iron

To calculate the yield of iron, we need to know the molar mass of iron oxide (Fe2O3) and iron (Fe). The molar mass of Fe2O3 is 159.69 g/mol, and the molar mass of Fe is 55.85 g/mol.

First, let's calculate the number of moles of iron oxide:

1.00 kg Fe2O3 = 1000 g Fe2O3 Number of moles Fe2O3 = mass Fe2O3 / molar mass Fe2O3 = 1000 g / 159.69 g/mol = 6.26 mol

According to the balanced equation, 1 mole of Fe2O3 produces 2 moles of Fe. Therefore, the number of moles of Fe produced is:

Number of moles Fe = 2 x number of moles Fe2O3 = 2 x 6.26 mol = 12.52 mol

Now, let's calculate the mass of Fe produced:

Mass of Fe = number of moles Fe x molar mass Fe = 12.52 mol x 55.85 g/mol = 699.41 g

Calculating the Percent Yield

If 654 g of iron is obtained from the reaction, we can calculate the percent yield as follows:

Percent yield = (actual yield / theoretical yield) x 100 = (654 g / 699.41 g) x 100 = 93.4%

Therefore, the percent yield of iron obtained from the reaction is 93.4%.

Discussion

The percent yield of a reaction is a measure of how much of the theoretical yield is actually obtained. In this case, the percent yield is 93.4%, which means that 93.4% of the theoretical yield of iron is obtained from the reaction.

There are several reasons why the actual yield may be less than the theoretical yield. These include:

Incomplete reaction: The reaction may not be complete, resulting in some of the reactants not being converted to products.
Impurities: The reactants or products may contain impurities that affect the yield of the reaction.
Equipment limitations: The equipment used to perform the reaction may have limitations that affect the yield of the reaction.

Conclusion

In conclusion, the yield of iron obtained from 1.00 kg of iron oxide can be calculated using the molar masses of iron oxide and iron. The percent yield of iron obtained from the reaction can also be calculated using the actual and theoretical yields. The percent yield is a measure of how much of the theoretical yield is actually obtained, and it can be affected by several factors, including incomplete reaction, impurities, and equipment limitations.

References

CRC Handbook of Chemistry and Physics, 97th ed., CRC Press, 2016.
Chemistry: An Atoms First Approach, 2nd ed., Steven S. Zumdahl, Cengage Learning, 2014.

Glossary

Molar mass: The mass of one mole of a substance.
Percent yield: The ratio of the actual yield to the theoretical yield, expressed as a percentage.
Theoretical yield: The maximum amount of product that can be obtained from a reaction, based on the stoichiometry of the reaction.
Actual yield: The amount of product that is actually obtained from a reaction.
Iron Oxide Reduction: Q&A ==========================

Introduction

In our previous article, we explored how to calculate the yield of iron obtained from 1.00 kg of iron oxide and determine the percent yield if 654 g of iron is obtained from the reaction. In this article, we will answer some frequently asked questions (FAQs) related to the reduction of iron oxide.

Q&A

Q: What is the molar mass of iron oxide (Fe2O3)?

A: The molar mass of iron oxide (Fe2O3) is 159.69 g/mol.

Q: How many moles of iron oxide are in 1.00 kg of iron oxide?

A: To calculate the number of moles of iron oxide, we need to divide the mass of iron oxide (1000 g) by the molar mass of iron oxide (159.69 g/mol).

Number of moles Fe2O3 = mass Fe2O3 / molar mass Fe2O3 = 1000 g / 159.69 g/mol = 6.26 mol

Q: How many moles of iron are produced from 1 mole of iron oxide?

A: According to the balanced equation, 1 mole of Fe2O3 produces 2 moles of Fe.

Q: What is the mass of iron produced from 1 mole of iron oxide?

A: To calculate the mass of iron produced, we need to multiply the number of moles of iron by the molar mass of iron (55.85 g/mol).

Mass of Fe = number of moles Fe x molar mass Fe = 2 mol x 55.85 g/mol = 111.70 g

Q: How many grams of iron are produced from 1.00 kg of iron oxide?

A: To calculate the mass of iron produced, we need to multiply the number of moles of iron oxide by the mass of iron produced per mole of iron oxide.

Number of moles Fe2O3 = 6.26 mol Mass of Fe per mole Fe2O3 = 111.70 g Mass of Fe = number of moles Fe2O3 x mass of Fe per mole Fe2O3 = 6.26 mol x 111.70 g/mol = 699.41 g

Q: What is the percent yield of iron obtained from the reaction?

A: To calculate the percent yield, we need to divide the actual yield by the theoretical yield and multiply by 100.

Percent yield = (actual yield / theoretical yield) x 100 = (654 g / 699.41 g) x 100 = 93.4%

Q: What are some common reasons for a low percent yield?

A: Some common reasons for a low percent yield include:

Incomplete reaction: The reaction may not be complete, resulting in some of the reactants not being converted to products.
Impurities: The reactants or products may contain impurities that affect the yield of the reaction.
Equipment limitations: The equipment used to perform the reaction may have limitations that affect the yield of the reaction.

Conclusion

In conclusion, the reduction of iron oxide is a complex process that requires careful consideration of the molar masses of the reactants and products. By understanding the stoichiometry of the reaction and the factors that affect the yield, we can optimize the reaction conditions to produce the desired product.

References

CRC Handbook of Chemistry and Physics, 97th ed., CRC Press, 2016.
Chemistry: An Atoms First Approach, 2nd ed., Steven S. Zumdahl, Cengage Learning, 2014.

Glossary

Molar mass: The mass of one mole of a substance.
Percent yield: The ratio of the actual yield to the theoretical yield, expressed as a percentage.
Theoretical yield: The maximum amount of product that can be obtained from a reaction, based on the stoichiometry of the reaction.
Actual yield: The amount of product that is actually obtained from a reaction.