Chemistry Glossary for Molar Mass Work

Amount of substance is measured in moles (mol). One mole contains Avogadro's number (6.022 × 10²³) of particles. Example: one mole of glucose C₆H₁₂O₆ is 180.16 g and contains 6.022 × 10²³ molecules. The mole links the atomic scale to laboratory mass.

Molar mass is mass per mole of a substance, in g/mol. For sodium chloride NaCl, molar mass ≈ 58.44 g/mol. Formula mass is the same sum of atomic masses for one formula unit; the term is often used for ionic solids. Molecular mass applies specifically to discrete molecules like water H₂O (18.02 g/mol).

Stoichiometry uses balanced equation coefficients as mole ratios. The limiting reagent is the reactant consumed first, capping product amount. Theoretical yield is the maximum product mass from that limit; percent yield compares actual isolated mass to theoretical. Example: precipitating AgCl from AgNO₃ and NaCl requires identifying which nitrate or chloride is limiting.

Molarity (M) = mol solute / L solution. Molality (m) = mol solute / kg solvent. Mass percent = (part mass / total mass) × 100%. For sulfuric acid H₂SO₄ solutions, molarity describes acid strength in titrations; molality appears in colligative property problems with water H₂O as solvent.

Empirical formula gives the simplest atom ratio (CH₂O for glucose). Molecular formula gives the true count (C₆H₁₂O₆). Mass percent composition lists each element's share of total mass. Polyatomic ions like SO₄²⁻ in sulfuric acid H₂SO₄ count as grouped atoms when parsing formulas for molar mass.

Atomic mass (relative atomic mass) is a weighted average of an element's naturally occurring isotopes, dimensionless on the periodic table but numerically equal to molar mass in g/mol. Isotopes are atoms of the same element with different neutron counts, hence different masses. Periodic trends describe how atomic mass, atomic radius, and electronegativity change predictably across periods and down groups — useful for sanity-checking calculated molar masses, as with potassium hydroxide KOH (56.11 g/mol) versus sodium hydroxide NaOH (40.00 g/mol).

Ionic bonding transfers electrons between metal and nonmetal, forming charged ions held together electrostatically, as in sodium chloride NaCl. Covalent bonding shares electrons between nonmetals, forming discrete molecules like water H₂O or methane CH₄. A formula unit is the smallest repeating ratio of ions in an ionic solid; a polyatomic ion (like sulfate, SO₄²⁻, or nitrate, NO₃⁻) is a bonded group of atoms carrying an overall charge, treated as one unit when parsing a formula's parentheses.

The ideal gas law, PV = nRT, relates pressure, volume, moles, and temperature, and can be rearranged to find molar mass from a gas's measured density. Standard temperature and pressure (STP) is commonly defined as 0 °C and either 1 bar (giving 22.7 L/mol) or 1 atm (giving 22.4 L/mol) depending on the convention used. A balanced chemical equation has equal numbers of each element's atoms on both the reactant and product sides, and its coefficients define the mole ratios used in every stoichiometry calculation — as in ammonia synthesis, N₂ + 3 H₂ → 2 NH₃.