Definitions Glossary

Term	Definition	Key Context
Mole (mol)	Amount of substance containing $6.022 \times 10^{23}$ entities	SI base unit for amount of substance
Avogadro's Number (Nₐ)	$6.022 \times 10^{23}$ $mol^{-1}$ ; number of entities per mole	Named after Amedeo Avogadro; CODATA value
Atomic Mass Unit (amu)	1/12 the mass of one C-12 atom = $1.66054 \times 10^{-24}$ g	Also called Dalton (Da)
Molar Mass	Mass of one mole of a substance in g/mol; numerically equals molecular/atomic mass in amu	Key bridge between mass (g) and moles
Molar Volume (STP)	22.4 L/mol; volume occupied by 1 mol of ideal gas at 0 °C and 1 atm	Valid at STP only (not room temperature)
Empirical Formula	Simplest whole-number ratio of atoms in a compound	$CH_{2}O$ for glucose (actual: $C_{6}H_{12}O_{6}$ )
Molecular Formula	Actual number of atoms of each element in one molecule	$C_{6}H_{12}O_{6}$ for glucose; = n × empirical formula
Stoichiometry	Quantitative relationship between reactants and products using balanced equation	Mole ratios from balanced equation
Limiting Reagent	Reactant completely consumed first; limits the amount of product formed	Identified by dividing moles by stoichiometric coefficient
Theoretical Yield	Maximum mass of product obtainable if limiting reagent fully reacts	Actual yield ≤ theoretical yield
Molarity (M)	Moles of solute per litre of solution (mol/L)	Temperature-dependent
Molality (m)	Moles of solute per kilogram of solvent (mol/kg)	Temperature-independent; used for colligative properties
Mole Fraction (x)	Moles of component / total moles of all components	Dimensionless; sum = 1 for all components
Normality (N)	Equivalents of solute per litre of solution; N = M × n-factor	Reaction-dependent; n-factor varies
n-factor	Number of $H^{+}$ / $OH^{-}$ exchanged (acid-base) or electrons transferred (redox) per formula unit	$H_{2}SO_{4}$ n-factor = 2 in neutralisation
Equivalent Weight	Molar mass ÷ n-factor	$H_{2}SO_{4}$ eq. wt. = 49 g/eq in neutralisation
ppm	Parts per million = mg of solute per kg of solution	Used for trace pollutant concentrations
Law of Definite Proportions	A pure compound always contains the same elements in the same mass ratio	Proust (1799); basis for chemical formulas
Law of Multiple Proportions	When two elements form multiple compounds, mass ratios (for fixed mass of one) are simple whole numbers	Dalton (1803); e.g., CO and $CO_{2}$
Gay-Lussac's Law of Gaseous Volumes	Gases combine in simple whole-number volume ratios at same T and P	Precursor to Avogadro's Law