Stoichiometry Calculation Mass Of MgO Produced From Burning Mg

Jul 13, 2025 by ADMIN 63 views

Stoichiometry Calculation The Burning of Magnesium with Oxygen to Form Magnesium Oxide

Introduction

In the realm of chemistry, stoichiometry serves as a fundamental tool for understanding and quantifying the relationships between reactants and products in chemical reactions. Stoichiometry, derived from the Greek words stoicheion (element) and metron (measure), essentially involves the calculation of relative quantities of reactants and products in chemical reactions. This intricate dance of atoms and molecules follows the law of conservation of mass, which dictates that matter cannot be created or destroyed in a chemical reaction. This law underscores the importance of balanced chemical equations, which accurately represent the number of atoms and molecules involved in a reaction. In this comprehensive article, we will delve into a specific stoichiometric calculation, focusing on the burning of magnesium ( $Mg$ ) with oxygen ( $O_2$ ) to form magnesium oxide ( $MgO$ ). Our aim is to determine the mass of $MgO$ produced when 28.0 g of $Mg$ undergoes complete combustion. This practical example will illuminate the step-by-step process of stoichiometric calculations, emphasizing the critical role of molar masses, mole ratios, and balanced chemical equations. By mastering these concepts, you will gain a deeper appreciation for the quantitative nature of chemistry and its ability to predict and explain the outcomes of chemical reactions. This knowledge is not only essential for academic pursuits but also finds wide applications in various industries, including pharmaceuticals, materials science, and environmental science. Therefore, let us embark on this journey of stoichiometric exploration and unravel the mysteries of chemical reactions.

The Balanced Chemical Equation

The cornerstone of any stoichiometric calculation lies in the balanced chemical equation. A balanced chemical equation is a symbolic representation of a chemical reaction that adheres to the law of conservation of mass. This law, a fundamental principle in chemistry, states that matter cannot be created or destroyed in a chemical reaction. In simpler terms, the number of atoms of each element must remain the same on both sides of the equation, ensuring that the number of atoms entering the reaction equals the number of atoms exiting the reaction. The balanced chemical equation for the burning of magnesium with oxygen to form magnesium oxide is:

$2Mg + O_2 ightarrow 2MgO$

This equation reveals that two moles of magnesium ( $Mg$ ) react with one mole of oxygen ( $O_2$ ) to produce two moles of magnesium oxide ( $MgO$ ). The coefficients in front of each chemical formula represent the stoichiometric coefficients, which indicate the relative number of moles of each reactant and product involved in the reaction. These coefficients are crucial for determining the mole ratios, which are the foundation of stoichiometric calculations. In this specific reaction, the mole ratio between $Mg$ and $MgO$ is 2:2, which simplifies to 1:1. This means that for every one mole of magnesium that reacts, one mole of magnesium oxide is produced. Understanding and interpreting balanced chemical equations is paramount for accurate stoichiometric calculations, as they provide the necessary quantitative information to predict the amounts of reactants and products involved in a chemical reaction. A correctly balanced chemical equation ensures that the calculation will accurately reflect the actual chemical process, leading to reliable results.

Determining Molar Masses

Central to stoichiometric calculations is the concept of molar mass. Molar mass is defined as the mass of one mole of a substance, expressed in grams per mole (g/mol). It serves as the bridge between mass, which is a measurable quantity in the laboratory, and moles, which represent the number of particles (atoms, molecules, or ions) involved in a chemical reaction. To determine the molar mass of a compound, we sum the atomic masses of all the atoms present in the chemical formula. The atomic masses of elements are readily available on the periodic table. For instance, the atomic mass of magnesium ( $Mg$ ) is approximately 24.31 g/mol, and the atomic mass of oxygen ( $O$ ) is approximately 16.00 g/mol.

Now, let's calculate the molar masses of the reactants and products in our reaction:

Magnesium ( $Mg$ ): The molar mass of $Mg$ is simply its atomic mass, which is 24.31 g/mol.
Oxygen ( $O_2$ ): Oxygen exists as a diatomic molecule, meaning it consists of two oxygen atoms. Therefore, the molar mass of $O_2$ is 2 times the atomic mass of oxygen, which is 2 * 16.00 g/mol = 32.00 g/mol.
Magnesium Oxide ( $MgO$ ): To find the molar mass of $MgO$ , we add the atomic masses of one magnesium atom and one oxygen atom: 24.31 g/mol + 16.00 g/mol = 40.31 g/mol.

The calculated molar masses are essential for converting between grams and moles, which is a crucial step in stoichiometric calculations. By knowing the molar mass of a substance, we can easily determine the number of moles present in a given mass, or vice versa. This conversion is vital for relating the amount of reactants used to the amount of products formed in a chemical reaction. Accurate determination of molar masses is thus a cornerstone of precise stoichiometric calculations.

Calculating Moles of Magnesium

The problem states that 28.0 g of magnesium ( $Mg$ ) is burned. To determine the amount of magnesium oxide ( $MgO$ ) produced, we must first convert the mass of magnesium to moles. This conversion is achieved using the molar mass of magnesium, which we previously calculated to be 24.31 g/mol. The formula for converting mass to moles is:

Moles = Mass / Molar Mass

Plugging in the values for magnesium, we get:

Moles of $Mg$ = 28.0 g / 24.31 g/mol

Moles of $Mg$ ≈ 1.15 moles

Therefore, 28.0 g of magnesium is equivalent to approximately 1.15 moles. This conversion is a pivotal step in stoichiometric calculations because it allows us to relate the amount of magnesium in grams, a measurable quantity in the laboratory, to the number of moles, which is directly related to the number of atoms involved in the reaction. The mole concept is central to stoichiometry, as it provides a common unit for comparing the amounts of different substances in a chemical reaction. By converting the mass of magnesium to moles, we can now use the mole ratio from the balanced chemical equation to determine the number of moles of magnesium oxide produced. This step highlights the importance of molar mass as a bridge between the macroscopic world of grams and the microscopic world of atoms and molecules.

Applying the Mole Ratio

Now that we have determined the number of moles of magnesium ( $Mg$ ) that reacted (1.15 moles), we can use the mole ratio from the balanced chemical equation to calculate the number of moles of magnesium oxide ( $MgO$ ) produced. The balanced chemical equation:

$2Mg + O_2 ightarrow 2MgO$

reveals that the mole ratio between $Mg$ and $MgO$ is 2:2, which simplifies to 1:1. This signifies that for every one mole of magnesium that reacts, one mole of magnesium oxide is formed. Therefore, the number of moles of $MgO$ produced is equal to the number of moles of $Mg$ reacted.

Moles of $MgO$ = Moles of $Mg$

Moles of $MgO$ = 1.15 moles

This simple yet crucial step demonstrates the power of stoichiometry in predicting the quantitative outcome of a chemical reaction. The mole ratio acts as a conversion factor, allowing us to translate the amount of one substance (in moles) to the equivalent amount of another substance in the reaction. In this case, the 1:1 mole ratio directly links the moles of magnesium consumed to the moles of magnesium oxide produced. This direct relationship stems from the fundamental principles of the law of conservation of mass, which dictates that atoms are neither created nor destroyed in a chemical reaction. The mole ratio, derived from the balanced chemical equation, ensures that we accurately account for the atoms involved in the reaction, leading to a precise prediction of the amount of product formed. This step underscores the importance of the balanced chemical equation as the foundation for all stoichiometric calculations.

Calculating the Mass of Magnesium Oxide

Having determined the number of moles of magnesium oxide ( $MgO$ ) produced (1.15 moles), we can now calculate the mass of $MgO$ using its molar mass. We previously calculated the molar mass of $MgO$ to be 40.31 g/mol. The formula for converting moles to mass is:

Mass = Moles × Molar Mass

Plugging in the values for magnesium oxide, we get:

Mass of $MgO$ = 1.15 moles × 40.31 g/mol

Mass of $MgO$ ≈ 46.36 g

Therefore, the burning of 28.0 g of magnesium ( $Mg$ ) will produce approximately 46.36 g of magnesium oxide ( $MgO$ ). This final calculation provides the answer to the initial problem, demonstrating the complete stoichiometric process from start to finish. By converting the initial mass of magnesium to moles, applying the mole ratio from the balanced chemical equation, and then converting the moles of magnesium oxide back to mass, we have successfully predicted the amount of product formed in the reaction. This calculation showcases the practical application of stoichiometry in chemistry, allowing us to quantify chemical reactions and predict the yield of products based on the amount of reactants used. The ability to accurately calculate the mass of products formed is essential in various fields, including chemical synthesis, industrial production, and environmental monitoring. This final step underscores the importance of a thorough understanding of molar masses, mole ratios, and the mole concept in stoichiometric calculations.

Conclusion

In conclusion, we have successfully calculated the mass of magnesium oxide ( $MgO$ ) produced by the burning of 28.0 g of magnesium ( $Mg$ ) using stoichiometric principles. The process involved several key steps, each crucial for accurate determination of the result. First, we established the balanced chemical equation for the reaction: $2Mg + O_2 ightarrow 2MgO$ . This equation served as the foundation for our calculations, providing the necessary mole ratios between reactants and products. Next, we determined the molar masses of the substances involved, which are essential for converting between mass and moles. We then converted the given mass of magnesium (28.0 g) to moles using its molar mass. This step allowed us to work with the mole concept, which is central to stoichiometric calculations. By applying the mole ratio from the balanced chemical equation, we determined the number of moles of magnesium oxide produced. Finally, we converted the moles of magnesium oxide back to mass using its molar mass, arriving at the answer of approximately 46.36 g of $MgO$ . This comprehensive example illustrates the power and utility of stoichiometry in chemistry. By mastering these principles, we can predict the quantitative outcomes of chemical reactions, which is vital in various fields, including chemical research, industrial chemistry, and environmental science. The ability to accurately calculate the amounts of reactants and products involved in a chemical reaction is fundamental to understanding and controlling chemical processes. Therefore, a thorough grasp of stoichiometric principles is essential for anyone pursuing a career in chemistry or related fields. This exercise not only provides a numerical answer but also reinforces the importance of a systematic approach to problem-solving in chemistry.