Learning Objectives
- Understand the laws of chemical combination
- Learn about atoms, molecules, and ions
- Write chemical formulae of compounds
- Understand the concept of mole and molar mass
- Calculate molecular mass and formula unit mass
Key Concepts
Laws of Chemical Combination
Law of Conservation of Mass (Lavoisier): Mass can neither be created nor destroyed in a chemical reaction. Total mass of reactants = Total mass of products.
Law of Constant Proportions (Proust): In a chemical substance, the elements are always present in definite proportions by mass. Example: Water always has hydrogen and oxygen in 1:8 ratio by mass.
Dalton's Atomic Theory
- All matter is made of tiny, indivisible particles called atoms.
- Atoms of a given element are identical in mass and properties.
- Atoms of different elements have different masses and properties.
- Atoms combine in simple whole-number ratios to form compounds.
- Atoms cannot be created or destroyed in a chemical reaction.
Atoms and Their Symbols
An atom is the smallest particle of an element that retains the chemical properties of that element. Atomic symbols are one or two letter abbreviations. Examples: H (Hydrogen), O (Oxygen), Na (Sodium - from Latin Natrium), Fe (Iron - from Latin Ferrum).
Atomic mass: Relative mass of an atom compared to 1/12 the mass of a carbon-12 atom. Unit: u (unified atomic mass unit).
Molecules
A molecule is a group of two or more atoms chemically bonded together. It is the smallest particle of a substance that can exist independently.
Atomicity: Number of atoms in a molecule. Monoatomic (1 atom: He, Ar), Diatomic (2 atoms: H₂, O₂, N₂), Triatomic (3 atoms: O₃, H₂O), Polyatomic (more than 3: P₄, S₈).
Ions
Ions are charged particles formed when atoms gain or lose electrons.
Cation: Positively charged ion (e.g., Na⁺, Ca²⁺). Anion: Negatively charged ion (e.g., Cl⁻, O²⁻).
Polyatomic ions: Group of atoms carrying a charge. Examples: NH₄⁺ (ammonium), SO₄²⁻ (sulphate), CO₃²⁻ (carbonate), NO₃⁻ (nitrate).
Writing Chemical Formulae
Use the criss-cross method: write the valencies of the two ions and cross them over to get the subscripts. Example: Aluminium oxide → Al³⁺ and O²⁻ → Al₂O₃.
Mole Concept
Mole: The amount of substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.
Avogadro's Number: 6.022 × 10²³ particles per mole.
Molar Mass: Mass of 1 mole of a substance in grams. Numerically equal to its atomic/molecular mass in u.
Number of moles = Given mass / Molar mass
Number of particles = Number of moles × 6.022 × 10²³
Summary
Matter is made of atoms, which combine to form molecules and ions. Chemical combination follows the laws of conservation of mass and constant proportions. Chemical formulae represent the composition of compounds. The mole concept connects the microscopic world of atoms to measurable quantities through Avogadro's number.
Important Terms
- Atom: Smallest particle of an element that retains chemical identity
- Molecule: A group of two or more atoms chemically bonded
- Ion: A charged atom or group of atoms
- Mole: Amount of substance containing 6.022 × 10²³ particles
- Avogadro's Number: 6.022 × 10²³
- Atomic Mass Unit (u): 1/12 the mass of a carbon-12 atom
Quick Revision
- Law of Conservation of Mass: total mass of reactants = total mass of products
- Law of Constant Proportions: elements combine in fixed mass ratios
- 1 mole = 6.022 × 10²³ particles
- Molar mass (g) = Atomic/Molecular mass (u) numerically
- Number of moles = Given mass / Molar mass