Summary

This chapter has built, step by step, the quantitative toolkit that underlies the whole of chemistry:

• Elements, atoms and molecules (1.1) set out the basic vocabulary of composition.
• The laws of chemical combination — conservation of mass, definite proportions, multiple proportions, Gay-Lussac's law of gaseous volumes, and Avogadro's law — are experimental facts about how masses and gas volumes combine, and Dalton's atomic theory (1.2) explains all of them at once by proposing that matter is built from indivisible atoms that combine in fixed, simple, whole-number ratios.
• Atomic mass, molecular mass and formula mass (1.3) put atoms and molecules on a common relative scale, anchored to $\dfrac{1}{12}$th the mass of a carbon-12 atom, and account for the fact that most elements are natural mixtures of isotopes.
• The mole, and Avogadro's number $N_A = 6.022 \times 10^{23}\ \text{mol}^{-1}$ (1.4), is the counting unit that finally lets chemists convert between a number of atoms/molecules and a mass that can be weighed — via molar mass — and, for gases, a volume at STP via the molar volume, $22.4\ \text{L mol}^{-1}$.
• Percentage composition and empirical/molecular formulas (1.5) show how the composition of an unknown compound is determined purely from mass data, and how the true molecular formula is recovered once the molecular mass is separately known.
• Stoichiometry and the limiting reagent (1.6) use the mole ratios written into a balanced equation to predict exactly how much product a reaction will give — and identify which reactant actually controls that amount when reactants are not mixed in the exact ratio the equation demands.
• Concentration terms (1.7) — molarity, molality, normality, mole fraction and percentage concentration — give precise, standardised ways to state how much solute is dissolved in a solution, each suited to a different kind of calculation. …