Course: CHEM 101, Michigan State University | Source: Compounds and Stoichiometry lecture notes
Difficulty: Introductory | Prerequisites: Basic atomic structure, periodic table familiarity
This topic sits at the foundation of general chemistry. Before you can do any meaningful work with chemical reactions, you need to know how chemists count atoms and measure substances. The mole concept bridges the gap between the atomic scale (individual atoms and molecules) and the lab scale (grams on a balance). If you are comfortable with atomic structure and the periodic table, you have everything you need to start here.
A molecule is atoms bonded covalently; ionic compounds use formula units instead. Molecular weight is the sum of all atomic weights in a formula (in amu). The mole (6.02 x 10²³ particles) lets you convert between particle counts and grams using molar mass.
Molecule
A combination of two or more atoms held together by covalent bonds. In simple terms, it is the smallest unit of a covalent compound that still behaves like that compound.
Formula unit
The simplest whole-number ratio of ions in an ionic compound. Think of it as the ionic equivalent of a molecule, since ionic compounds do not form discrete molecules.
Molecular weight (molecular mass)
The sum of the atomic weights of every atom in a molecule, measured in atomic mass units (amu). In simple terms, add up the masses from the periodic table for each atom in the formula.
Formula weight
The same calculation as molecular weight, but applied to ionic compounds. Used because ionic compounds are described by formula units rather than molecules.
Mole (mol)
The quantity of substance containing exactly 6.02 x 10²³ particles (atoms, molecules, ions, etc.). Think of it as the chemist's counting unit, the way a "dozen" means 12.
Avogadro's number
6.02 x 10²³. The number of particles in one mole of any substance.
Molar mass
The mass of one mole of a substance, expressed in grams per mole (g/mol). Numerically equal to the molecular or formula weight, but in grams instead of amu.
Gram equivalent weight
The mass of a compound that supplies one mole of a particle of interest. In acid-base chemistry, it is the mass of acid or base that yields one mole of H⁺ or OH⁻.
Normality (N)
A measure of concentration defined as equivalents of solute per litre of solution. Think of it as molarity adjusted for how many reactive particles each formula unit delivers.
Molecules are held together by covalent bonds. Water (H₂O) and methane (CH₄) are molecules.
Ionic compounds (e.g. NaCl) do not form discrete molecules. Oppositely charged ions arrange themselves in a repeating lattice in the solid state.
Because there is no single "molecule" of an ionic compound, we use the term formula unit to describe the simplest ratio of ions.
Molecular weight applies to covalent compounds. Add up the atomic masses (from the periodic table) of every atom in the molecular formula.
Example: H₂O = 2(1.008) + 16.00 = 18.02 amu
Formula weight is the same arithmetic, applied to ionic compounds.
Example: NaCl = 22.99 + 35.45 = 58.44 amu
Both are measured in amu at the particle level.
One mole is defined as the number of particles in exactly 12 grams of carbon-12.
That number is Avogadro's number: 6.02 x 10²³.
The mole connects the atomic scale to the laboratory scale. You cannot count atoms on a balance, but you can weigh out a mole of them.
Molar mass is numerically equal to the molecular weight or formula weight, but its unit is g/mol instead of amu.
Example: The molecular weight of H₂O is 18.02 amu, so the molar mass of H₂O is 18.02 g/mol.
One mole of a substance has a mass in grams equal to its molar mass.
Gram equivalent weight is the mass of a compound that delivers one mole of a "particle of interest."
In acid-base chemistry, the particle of interest is H⁺ (for acids) or OH⁻ (for bases).
HCl donates 1 H⁺ per formula unit, so its equivalent weight equals its molar mass.
H₂SO₄ donates 2 H⁺ per formula unit, so its equivalent weight is half its molar mass.
H₃PO₄ donates 3 H⁺ per formula unit, so its equivalent weight is one-third its molar mass.
Normality = equivalents of solute per litre of solution. For a monoprotic acid like HCl, normality equals molarity. For H₂SO₄, normality = 2 x molarity.
\text{Molecular weight} = \sum (\text{atomic mass of each atom in the formula})\text{Molar mass (g/mol)} = \text{Molecular weight (amu), same number}\text{Number of moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}\text{Number of particles} = n \times 6.02 \times 10^{23}\text{Equivalent weight} = \frac{\text{molar mass}}{\text{number of equivalents per formula unit}}\text{Normality (N)} = \text{Molarity} \times \text{equivalents per formula unit}Pharmacists use molar mass every time they calculate drug dosages: the active ingredient's mass in a tablet is converted to moles to ensure the correct number of molecules reaches the bloodstream. Industrial chemists rely on equivalent weight and normality when preparing titration solutions for quality control in manufacturing.
Students often confuse molecular weight (amu, single-particle level) with molar mass (g/mol, bulk level). They are numerically the same, but the units and contexts differ.
Students sometimes try to write a "molecular formula" for an ionic compound like NaCl. Ionic compounds only have empirical (formula unit) representations.
The mole is not a mass. It is a count. One mole of helium and one mole of lead contain the same number of atoms; they differ in mass.
Equivalent weight depends on the reaction type. The same compound can have different equivalent weights in acid-base vs. redox contexts.
⚠️ Converting between grams, moles, and particles is the single most tested skill in introductory stoichiometry. Know the formulas cold.
⚠️ Be ready to calculate the equivalent weight of polyprotic acids (H₂SO₄, H₃PO₄). A common exam question asks you to find normality from molarity.
⚠️ Know the distinction between molecular weight (covalent compounds) and formula weight (ionic compounds). The terms are not interchangeable.
⚠️ Expect a question requiring you to convert amu to g/mol or vice versa. The numerical value is the same; the context is what changes.
True or False: One mole of any substance contains 6.02 x 10²³ particles. (True)
Fill in the blank: The molar mass of a compound is numerically equal to its ______ but uses units of g/mol. (molecular weight / formula weight)
True or False: NaCl has a molecular weight. (False, it has a formula weight)
Fill in the blank: H₂SO₄ provides ______ molar equivalents of H⁺. (2)
True or False: Equivalent weight is always the same as molar mass. (False, it depends on the number of equivalents per formula unit)
Q: What is the molar mass of calcium chloride (CaCl₂)?
A: Ca = 40.08, Cl = 35.45 x 2 = 70.90. Molar mass = 40.08 + 70.90 = 110.98 g/mol.
Q: How many moles are in 36.04 g of water (H₂O)?
A: Molar mass of H₂O = 18.02 g/mol. Moles = 36.04 / 18.02 = 2.00 mol.
Q: What is the equivalent weight of H₃PO₄ in an acid-base reaction?
A: Molar mass of H₃PO₄ = 98.00 g/mol. It donates 3 H⁺, so equivalent weight = 98.00 / 3 = 32.67 g/equivalent.
Q: A 0.5 M solution of H₂SO₄ has what normality?
A: H₂SO₄ donates 2 H⁺, so N = 0.5 x 2 = 1.0 N.
Q: How many molecules are in 3 moles of CO₂?
A: 3 x 6.02 x 10²³ = 1.806 x 10²⁴ molecules.
The mole concept feeds directly into stoichiometry (covered in the next set of notes), where you use mole ratios from balanced equations to predict how much product a reaction yields. Molar mass also connects to solution chemistry: molarity (moles per litre) is how you describe concentrations in virtually every later chapter. If you are heading toward gas laws, remember that 1 mole of any ideal gas occupies 22.4 L at STP, a fact that ties moles to measurable volume.
molecule, formula unit, ionic compound, covalent bond, molecular weight, molecular mass, formula weight, amu, atomic mass unit, mole, mol, Avogadro's number, 6.02 x 10²³, molar mass, grams per mole, g/mol, equivalent weight, gram equivalent weight, normality, concentration, monoprotic acid, diprotic acid, triprotic acid, polyprotic acid, HCl, H₂SO₄, H₃PO₄, carbon-12