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Chemical reactions are transformations that involve the rearrangement of atoms to form new substances. During a chemical reaction, the bonds between atoms in reactant molecules are broken and new bonds are formed in the product molecules. This results in different physical and chemical properties from the original substances.
Stoichiometry is central to chemical reactions. It involves balancing chemical equations to ensure the law of conservation of mass is obeyed, meaning that the number of each type of atom is the same on both the reactant and product sides. By balancing these equations, we can predict the quantities of reactants and products involved in the reaction.
For example, when sulfur trioxide (\(\text{SO}_3\)) reacts with water (\(\text{H}_2\text{O}\)) to form sulfuric acid (\(\text{H}_2\text{SO}_4\)), all atoms must be accounted for to balance the equation correctly, which is:

• \(\text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4\)

Balancing chemical equations requires a systematic approach, often involving starting with more complex molecules and gradually balancing simpler atoms like hydrogen and oxygen.

Sulfuric acid formation is an exothermic process where sulfur trioxide (\(\text{SO}_3\)) reacts with water (\(\text{H}_2\text{O}\)) to produce sulfuric acid (\(\text{H}_2\text{SO}_4\)). This reaction is crucial in industrial processes, especially in the manufacture of fertilizers, dyes, and explosives.
The overall reaction is very straightforward:

• \(\text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4\)

Even though the equation is balanced as written, the actual reaction process involves a highly exothermic interaction, releasing a significant amount of heat.
This requires careful control to prevent hazards in industrial settings, where heat management is crucial to ensure safety and maximize efficiency.

Decomposition reactions are chemical changes where one compound breaks down into two or more simpler products. This type of reaction usually requires energy input, such as heat, light, or electricity, to break the chemical bonds in the reactant.
An example is the decomposition of mercury(II) nitrate (\(\text{Hg(NO}_3\text{)}_2\)) upon heating:

• \(2\text{Hg(NO}_3\text{)}_2 \rightarrow 2\text{HgO} + 4\text{NO}_2 + \text{O}_2\)

In this balanced decomposition reaction, mercury(II) nitrate breaks down into mercury(II) oxide (\(\text{HgO}\)), nitrogen dioxide (\(\text{NO}_2\)), and oxygen (\(\text{O}_2\)) gas.
The stoichiometry of decomposition reactions ensures the number of atoms is conserved during the transformation, illustrating the fundamental principle of conservation of mass.

Combustion reactions are exothermic processes where a substance rapidly reacts with oxygen, releasing energy in the form of light and heat. These reactions are typically characterized by the presence of oxygen as a reactant and often involve hydrocarbons or more complex compounds.
An example is the combustion of phosphine (\(\text{PH}_3\)):

• \(4\text{PH}_3 + 8\text{O}_2 \rightarrow 6\text{H}_2\text{O} + \text{P}_4\text{O}_{10}\)

In this reaction, phosphine burns in oxygen to form water vapor and solid tetraphosphorus decaoxide (\(\text{P}_4\text{O}_{10}\)).
By balancing the combustion equation, we can determine the proportionate amounts of reactants and products, ensuring a complete and accurate description of the chemical change. Combustion reactions are vital in energy production, powering engines and other industrial processes through the controlled release of energy.