Question

Calcium oxide or quicklime \((\mathrm{CaO})\) is used in steelmaking, cement manufacture, and pollution control. It is prepared by the thermal decomposition of calcium carbonate: $$ \mathrm{CaCO}_{3}(s) \longrightarrow \mathrm{CaO}(s)+\mathrm{CO}_{2}(g) $$ Calculate the yearly release of \(\mathrm{CO}_{2}\) (in \(\mathrm{kg}\) ) to the atmosphere if the annual production of \(\mathrm{CaO}\) in the United States is \(1.7 \times 10^{10} \mathrm{~kg}\).

Short answer

The given chemical reaction is already balanced as:\[ \mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaO}(s) + \mathrm{CO}_{2}(g) \]

Step-by-step solution

Write the Balanced Equation

The given chemical reaction is already balanced as:\[ \mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaO}(s) + \mathrm{CO}_{2}(g) \]

Key concepts

Thermal Decomposition

Thermal decomposition is a chemical reaction where a compound breaks down into two or more simpler substances when heated. For calcium carbonate \(\mathrm{CaCO}_3\), this process involves heating it to a high temperature, causing it to decompose into calcium oxide \(\mathrm{CaO}\) and carbon dioxide \(\mathrm{CO}_2\). This is an endothermic reaction, meaning it absorbs heat from its surroundings.

In industrial applications, such as producing quicklime (calcium oxide), large limestone kilns are used to achieve the necessary temperatures for this decomposition to occur efficiently. The reaction can be summarized as:

• The starting material (reactant) is calcium carbonate (\(\mathrm{CaCO}_3\))
• The products are calcium oxide (\(\mathrm{CaO}\)) and carbon dioxide (\(\mathrm{CO}_2\))
• The process requires heat to drive the reaction

Understanding thermal decomposition is essential for many industries, especially those involved in construction materials and environmental management.

CO2 Emission

Carbon dioxide (\(\mathrm{CO}_2\)) emission refers to the release of carbon dioxide gas into the atmosphere as a result of chemical reactions or combustion processes. During the decomposition of calcium carbonate to form calcium oxide, carbon dioxide is a byproduct of the reaction.

In the context of the production of calcium oxide, the emission of \(\mathrm{CO}_2\) is significant as it contributes to greenhouse gases, which impact climate change. For every mole of calcium carbonate decomposed, one mole of carbon dioxide gas is released.

Given that the yearly production of calcium oxide in the United States is significant, calculations can show the substantial levels of carbon dioxide emitted annually:

• By totaling the molar masses of the substances involved, one can deduce the mass of \(\mathrm{CO}_2\) released per kilogram of \(\mathrm{CaO}\) produced.
• This involves precision in understanding the stoichiometry of the reaction and the molar relationships between the reactants and products.

Managing CO2 emissions is a critical part of reducing the industrial carbon footprint and addressing environmental concerns.

Chemical Reaction Balancing

Balancing chemical equations is a fundamental skill in chemistry, ensuring that the law of conservation of mass is satisfied in a chemical reaction. This law states that matter cannot be created or destroyed in an isolated system.

For the decomposition of calcium carbonate, the chemical equation is:\[ \mathrm{CaCO}_{3}(s) \rightarrow \mathrm{CaO}(s) + \mathrm{CO}_{2}(g) \]This equation is balanced as written, with the same number of each type of atom on both sides:

• There is one calcium atom on the reactant side matched by one calcium atom in the product \(\mathrm{CaO}\)
• Three oxygen atoms in \(\mathrm{CaCO}_{3}\) are split between \(\mathrm{CaO}\) and \(\mathrm{CO}_2\)
• One carbon atom in the reactant becomes the carbon in the product \(\mathrm{CO}_2\)

Balancing ensures the chemical equation obeys the conservation laws, making predictions about the amounts of products formed or reactants needed feasible.

CaCO3 Decomposition

The decomposition of calcium carbonate (\(\mathrm{CaCO}_{3}\)) into calcium oxide (\(\mathrm{CaO}\)) and carbon dioxide (\(\mathrm{CO}_2\)) is not only crucial in industrial processes but also naturally occurring. Known as calcination, this process typically occurs at temperatures above 840°C.

In industry, the preparation of calcium oxide from \(\mathrm{CaCO}_3\) is vital for several applications, such as:

• Manufacturing cement by providing the main ingredient (quicklime)
• In steelmaking, where it acts as a flux
• Environmental applications, where it helps neutralize acidic emissions

From a chemical perspective, understanding this breakdown process helps in designing equipment and processes that optimize efficiency, minimize emissions, and improve yield. Moreover, knowledge of decompositions like these is integral to developments in green chemistry, where reducing energy usage and emissions are paramount.