Molar mass

A mole is a counting unit: one mole is 6.022 × 10²³ particles, Avogadro's number. Molar mass is what makes that count usable in the laboratory, because balances measure grams, not particles. By definition, the molar mass of an element in grams per mole is numerically equal to its atomic mass in atomic mass units, which is why the periodic table doubles as a molar mass table.

Calculating a compound's molar mass is straightforward addition. Take each element in the formula, multiply its atomic mass by the number of atoms present, and total the results. Water, H₂O, gives 2(1.008) + 16.00 = 18.02 g/mol. Watch subscripts outside parentheses — in Ca(OH)₂ the subscript applies to both the oxygen and the hydrogen. The single relationship that carries the most weight is n = m / M: moles equal mass divided by molar mass, rearranged as needed to find any of the three.

In stoichiometry you convert a measured mass to moles, apply the mole ratio from the balanced equation, and convert back to mass. It also connects to gas behavior through the ideal gas law: since n = m / M, substituting into PV = nRT gives M = mRT / (PV), which yields an unknown gas's molar mass from measured pressure, volume, temperature, and mass. Graham's law extends the idea to effusion, where rate varies inversely with the square root of molar mass, so lighter gases escape faster. Molar mass also distinguishes empirical from molecular formulas: dividing the compound's actual molar mass by its empirical formula mass gives the whole-number multiplier.

MCAT chemistry tests molar mass throughout the stoichiometry material and again in the gas phase content, where it appears inside ideal gas and effusion calculations rather than as a standalone question. The ACT science section assumes it as prior knowledge, so computing a molar mass quickly from a formula is worth the memorization.