\(\mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) is emprical formula of a crystalline compound of iron. It is used in water and sewage treatment to aid in the removal of suspended impurities. Calculate the mass percentage of iron, sulphur and oxygen in this compound.

Understanding molar mass is fundamental in chemistry when you’re dealing with the quantitative aspects of substances.

Molar mass is defined as the mass of one mole of a substance, usually expressed in grams per mole (g/mol). It's the sum of the atomic masses of the elements that make up a compound, each multiplied by the number of times the element appears in the formula. The atomic masses are found on the periodic table of elements.

For instance, in our exercise, we calculate the molar mass of the compound \( \mathrm{Fe}_2\left(\mathrm{SO}_4\right)_3 \) as follows. Iron (\(\mathrm{Fe}\)) has an atomic mass of 55.845 g/mol, sulfur (\(\mathrm{S}\)) has an atomic mass of 32.065 g/mol, and oxygen (\(\mathrm{O}\)) has an atomic mass of 15.999 g/mol.

The crucial step is to multiply each atomic mass by the number of times each element appears in the chemical formula. For the iron in the compound, we multiply the atomic mass of iron by 2 since iron is present twice in the formula. This approach is repeated for sulfur and oxygen, adjusting the atomic mass with the quantity present. With the molar mass calculated, you can then tackle many stoichiometric problems, including finding mass percentages.

The empirical formula of a compound offers a straightforward representation of the elemental proportions within it, using the smallest whole number ratio.

In chemistry, the empirical formula is the simplest expression of a compound's composition. It differs from the molecular formula, which may represent the actual number of atoms in a molecule. Sometimes, as with ionic compounds, the empirical formula is the same as the molecular formula, but not always. For covalent compounds that form molecules, the empirical formula can be a simplification.

To illustrate, the compound used in the exercise, \(\mathrm{Fe}_2\left(\mathrm{SO}_4\right)_3\), is already in the form of an empirical formula. It shows the smallest whole number ratio of iron to sulfur to oxygen atoms. The '2' next to iron indicates there are twice as many iron atoms as sulfur atoms (which you can understand from the subscript '1' that is not written but implied), and the oxygen atoms are in a fixed ratio to sulfur of 4:1, with each sulfur atom within the sulfate ion \(\mathrm{SO}_4^{2-}\) being surrounded by four oxygen atoms.

Stoichiometry is a branch of chemistry that deals with the measurement and quantitative relationships of the reactants and products in a chemical reaction. It's rooted in the law of conservation of mass, where matter is neither created nor destroyed.

In practical terms, stoichiometry allows us to predict how much product will form from a certain amount of reactants, or conversely, how much of a reactant is needed to create a desired quantity of product. To apply stoichiometry, we often use a balanced chemical equation, which contains the molar ratios of the substances involved.

However, stoichiometry doesn't only apply to reactions. It's also essential for calculating the composition of compounds, as we see in our exercise. You use stoichiometry to determine the mass percentage of each element in a compound. Starting from the empirical formula and incorporating the molar masses of the involved elements, we can calculate the total mass of each constituent element and then their proportion in the entire compound. In this way, stoichiometry bridges the gap between the microscale of atoms and molecules and the macroscale quantities that we can measure.