Molar Mass of Elements and Compounds

In the microscopic world, we measure mass in Atomic Mass Units ($u$). However, a single atom is too small to see or weigh. To bridge this gap, chemists use the Mole ($6.022 \times 10^{23}$ particles). The magic of the mole is that the mass of one atom in $u$ is numerically equal to the mass of one mole of those atoms in grams per mole ($g/mol$). This is called the Molar Mass ($M$). For example, a single Carbon atom weighs $12.01 u$, and one mole of Carbon atoms weighs exactly $12.01 g$.

To find the molar mass of a compound, we sum the molar masses of every atom in its chemical formula. For covalent compounds like $H_2O$, we add two Hydrogens and one Oxygen. For ionic compounds with polyatomic ions, like $Mg(OH)_2$, the subscript outside the parentheses acts as a multiplier for everything inside. This requires careful 'bookkeeping' to ensure no atoms are missed in the final tally.

You may notice atomic masses on the periodic table are rarely whole numbers. This is because elements exist as isotopes—atoms with the same number of protons but different numbers of neutrons. The mass shown is a weighted average based on the natural abundance of these isotopes. If an element has one heavy isotope and one light one, the average mass will lean toward the one that is more common in nature.