A 0.025-g sample of a compound composed of boron and hydrogen, with a molecular mass of 28 amu, bums spontaneously when exposed to air, producing \(0.063 \mathrm{g}\) of \(\mathrm{B}_{2} \mathrm{O}_{3} .\) What are the empirical and molecular formulas of the compound?

Stoichiometry is a branch of chemistry that focuses on the quantitative relationships between reactants and products in a chemical reaction. It serves as a mathematical approach to predicting how much product will form from a given amount of reactants, or conversely, how much reactants are needed to produce a certain amount of product.

For instance, in the given problem, stoichiometry is used when calculating the moles of boric oxide formed during the combustion process. By understanding that reactions proceed according to their stoichiometry, you can derive the starting amounts of the elements involved in the compound by looking at the results of the reaction.

Practical Application of Stoichiometry

Using the molar mass of a compound (in this case, boric oxide), we convert the mass of the substance to its corresponding number of moles. This process is crucial because chemical reactions depend on the number of moles, not the mass, which provides a universal way to describe quantities in chemistry.

The mole concept is central to understanding stoichiometry and managing the quantitative aspects of chemistry. One mole is defined as the amount of substance that contains as many entities (atoms, molecules, or other particles) as there are atoms in exactly 12 g of pure carbon-12. This number is known as Avogadro's number, approximately equal to \(6.022 \times 10^{23}\) entities.

In our exercise, the mole concept is employed to relate the mass of boric oxide produced to the amount of boron atoms it contains. By knowing the molar mass of the substance, you can convert between mass and moles, allowing the calculation of moles of hydrogen and boron in the original compound based on the product's composition.

Calculating Moles

For instance, to calculate moles of a substance, you use the formula:
\[ \text{moles} = \frac{\text{mass}}{\text{molar mass}} \]
It allows you to establish the ratio of elements in a compound and find the empirical and molecular formulas, which gives insight into the substance's makeup and proportions.

Chemical formulas represent the composition of substances using symbols and numerical subscripts. They provide a wealth of information about the elements within a compound and their respective ratios. There are two types of chemical formulas commonly discussed: empirical and molecular formulas.

The empirical formula of a compound provides the simplest whole-number ratio of the elements in that compound. In contrast, the molecular formula gives the actual number of atoms of each element in a molecule of the compound. Often, the molecular formula is a whole-number multiple of the empirical formula.

In our case, the empirical formula is determined by dividing the moles of each element by the smallest number of moles calculated. The molecular formula then considers the actual molecular mass of the compound and uses the empirical formula as a foundation for determining the actual number and proportion of atoms in the molecule.

• Empirical formula: Simplified ratio of elements (e.g., CH for benzene).
• Molecular formula: Actual number of atoms in the molecule (e.g., C6H6 for benzene).

Understanding these formulas is crucial for tasks such as naming compounds, writing reaction equations, and even designing chemical syntheses.