The laws of chemical combination
Two laws, established by careful experiments in the late 1700s, led to the idea of atoms.
Law of conservation of mass (Antoine Lavoisier): mass can neither be created nor destroyed in a chemical reaction. The total mass of the reactants equals the total mass of the products.
Example. 12 g of carbon burns completely in 32 g of oxygen. The carbon dioxide formed has a mass of 12 + 32 = 44 g.
Law of constant proportions (Joseph Proust): a chemical substance always contains the same elements combined in the same proportion by mass. In water, hydrogen and oxygen are always in the mass ratio 1 : 8, whatever the source of the water.
Example. If 9 g of water is decomposed, it gives 1 g of hydrogen and 8 g of oxygen.
Dalton's atomic theory
John Dalton explained these laws with his atomic theory (1808):
All matter is made of very tiny particles called atoms.
Atoms cannot be created or destroyed in a chemical reaction.
Atoms of a given element are identical in mass and properties; atoms of different elements differ.
Atoms combine in small whole-number ratios to form compounds.
The relative number and kinds of atoms in a compound are constant.
Postulate 2 explains conservation of mass, and postulates 4 and 5 explain constant proportions. (Later discoveries showed that atoms are made of smaller particles and that atoms of one element can differ in mass; you will meet these ideas in the next chapter.)
Atoms: size and symbols
Atoms are extremely small. The radius of a hydrogen atom is about 10⁻¹⁰ m. Millions placed side by side would barely span the thickness of a sheet of paper.
Each element has a symbol of one or two letters, set by IUPAC. The first letter is always capital and the second small: H (hydrogen), O (oxygen), Ca (calcium), Cl (chlorine), Co (cobalt). Some come from Latin names: Na (sodium, natrium), K (potassium, kalium), Fe (iron, ferrum), Cu (copper, cuprum), Ag (silver, argentum), Au (gold, aurum).
Atomic mass
Atoms are too light to weigh in grams conveniently, so we compare them with a standard. One atomic mass unit (u) is 1/12 of the mass of one carbon-12 atom.
Element | Atomic mass (u) |
|---|---|
Hydrogen | 1 |
Carbon | 12 |
Nitrogen | 14 |
Oxygen | 16 |
Sodium | 23 |
Magnesium | 24 |
Sulphur | 32 |
Chlorine | 35.5 |
Calcium | 40 |
Molecules
A molecule is the smallest particle of an element or compound that can exist independently and shows its properties.
Molecules of elements contain one kind of atom. The number of atoms is the atomicity: argon (Ar) is monatomic, oxygen (O₂) diatomic, ozone (O₃) triatomic, phosphorus (P₄) tetra-atomic, and sulphur (S₈) polyatomic.
Molecules of compounds contain different atoms: H₂O, NH₃, CO₂.
Ions and valency
An ion is a charged particle. A positive ion is a cation (Na⁺, Mg²⁺); a negative ion is an anion (Cl⁻, O²⁻). A group of atoms carrying a charge is a polyatomic ion: ammonium NH₄⁺, hydroxide OH⁻, nitrate NO₃⁻, carbonate CO₃²⁻, sulphate SO₄²⁻.
The valency of an element is its combining capacity. For ions, it equals the size of the charge.
Writing chemical formulae
Write the symbols, cation first.
Write the valencies under them.
Criss-cross the valencies and simplify if needed. Put polyatomic ions in brackets when more than one is needed.
Compound | Ions | Formula |
|---|---|---|
Aluminium oxide | Al³⁺, O²⁻ | Al₂O₃ |
Calcium chloride | Ca²⁺, Cl⁻ | CaCl₂ |
Magnesium oxide | Mg²⁺, O²⁻ | MgO (2 : 2 simplifies to 1 : 1) |
Calcium hydroxide | Ca²⁺, OH⁻ | Ca(OH)₂ |
Ammonium sulphate | NH₄⁺, SO₄²⁻ | (NH₄)₂SO₄ |
Molecular mass and formula unit mass
Molecular mass is the sum of the atomic masses of all atoms in a molecule.
H₂O = 2 × 1 + 16 = 18 u. CO₂ = 12 + 2 × 16 = 44 u. H₂SO₄ = 2 + 32 + 64 = 98 u.
For ionic compounds such as NaCl we use formula unit mass: 23 + 35.5 = 58.5 u.
The mole concept
Chemists count particles in moles, just as we count eggs in dozens. One mole of any substance contains 6.022 × 10²³ particles. This number is Avogadro's number (N₀).
The mass of one mole of a substance, in grams, equals its atomic or molecular mass in u. This is its molar mass.
1 mole of oxygen atoms = 16 g = 6.022 × 10²³ atoms.
1 mole of water molecules = 18 g = 6.022 × 10²³ molecules.
Number of moles = given mass ÷ molar mass = number of particles ÷ N₀
Example. How many moles are in 36 g of water? 36 ÷ 18 = 2 mol, which is 2 × 6.022 × 10²³ = 1.2044 × 10²⁴ molecules.
Summary
Mass is conserved in reactions; compounds have fixed proportions by mass.
Dalton's atomic theory explains both laws.
1 u = 1/12 the mass of a carbon-12 atom.
Formulae are written by criss-crossing valencies.
1 mole = 6.022 × 10²³ particles and has a mass equal to the molar mass in grams.