Atoms and Molecules

What you'll learn

• Laws of chemical combination: conservation of mass, definite proportions.
• Dalton's Atomic Theory — postulates and limitations.
• What atoms and molecules are; atomic mass, molecular mass.
• How to write chemical formulae using valency.
• The mole concept and Avogadro's number.

Key concepts

Laws of chemical combination

1. Law of Conservation of Mass (Lavoisier, 1774)

• Mass is neither created nor destroyed in a chemical reaction.
• Total mass of reactants = Total mass of products.
• Example: 3 g of carbon + 8 g of oxygen → 11 g of carbon dioxide.

2. Law of Definite Proportions (Proust, 1799)

• A pure compound always contains the same elements in the same mass ratio, regardless of source or method of preparation.
• Example: Water (H₂O) always has H:O = 1:8 by mass, whether from a river or a lab.

Dalton's Atomic Theory (1808)

Limitations of Dalton's theory:

• Atoms are NOT indivisible — they have subatomic particles (electrons, protons, neutrons).
• Isotopes: atoms of the same element can have different masses.
• Isobars: atoms of different elements can have the same mass.
• Does not explain allotropy (diamond vs graphite — both carbon).

Atoms

• Atom: the smallest particle of an element that retains its chemical identity.
• Atoms are extremely small: radius ~0.1 nm (1 nm = 10⁻⁹ m).
• Atoms of most elements do not exist independently — they form molecules or ions.
• Noble gases (He, Ne, Ar) exist as single atoms (monoatomic).

Atomic mass

• Atomic mass unit (amu or u): 1/12th the mass of one carbon-12 atom.
• 1 u = 1.66 × 10⁻²⁴ g.

Molecules

• Molecule: the smallest particle of a substance that can exist independently and retains all properties of that substance.
• Atomicity: number of atoms in one molecule.

• Molecular mass: sum of atomic masses of all atoms in one molecule.
• Example: H₂O = 2(1) + 16 = 18 u.
• CO₂ = 12 + 2(16) = 44 u.

Chemical formulae and valency

• Valency: combining capacity of an atom; number of electrons it can share/donate/accept.
• Chemical formula shows the types and numbers of atoms in one molecule/formula unit.

Common valencies

Writing formulae using criss-cross rule

• Write the symbols with their valencies, then swap (criss-cross) the valencies as subscripts.
• If subscripts have a common factor, simplify.

Examples:

• MgO: Mg(2), O(2) → Mg₂O₂ → simplify → MgO
• Al₂O₃: Al(3), O(2) → Al₂O₃
• CaCl₂: Ca(2), Cl(1) → CaCl₂
• NH₃: N(3), H(1) → NH₃
• H₂SO₄: H(1), SO₄(2) → H₂SO₄

Common polyatomic ions (radicals)

Mole concept

• Mole: the amount of substance that contains the same number of particles as there are atoms in exactly 12 g of carbon-12.
• Avogadro's number (Nₐ) = 6.022 × 10²³ particles/mol.
• Named after Amedeo Avogadro (Italian scientist).

Key relationships

Examples

• 1 mole of H₂O = 18 g; contains 6.022 × 10²³ molecules.
• 1 mole of O₂ = 32 g; contains 6.022 × 10²³ molecules.
• 1 mole of Ca = 40 g; contains 6.022 × 10²³ atoms.
• 2 moles of CO₂ = 2 × 44 = 88 g; contains 2 × 6.022 × 10²³ = 1.2044 × 10²⁴ molecules.

Standard calculation:

How many moles in 9 g of water? n = 9 ÷ 18 = 0.5 mol

How many molecules in 9 g of water? N = 0.5 × 6.022 × 10²³ = 3.011 × 10²³ molecules

Formula unit mass (for ionic compounds)

• Ionic compounds (like NaCl, MgO) do not form molecules — they form lattices.
• Formula unit mass = sum of atomic masses in the formula.
• NaCl: 23 + 35.5 = 58.5 u
• MgO: 24 + 16 = 40 u
• CaCO₃: 40 + 12 + 3(16) = 100 u

Quick check

• State the Law of Conservation of Mass. Give one numerical example.
• What are the postulates of Dalton's Atomic Theory? Name one limitation.
• What is the molecular mass of H₂SO₄?
• What is the mole? State Avogadro's number.
• Calculate: How many moles in 46 g of Na? How many atoms does this contain?

Open the Practice tab for graded questions on Atoms and Molecules.