All matter — solids, liquids, gases, living and non-living — is made of atoms. Understanding how atoms are organised, how they combine, and how matter can be classified gives us the foundation of chemistry and material science.
Elements, Compounds, and Mixtures
- Element – A pure substance made of only one kind of atom. Cannot be broken down by chemical means. Examples: Oxygen (O), Iron (Fe), Gold (Au). There are 118 known elements.
- Compound – A pure substance made of two or more different elements chemically combined in a fixed ratio. Properties differ from constituent elements. Example: Water (H2O), Salt (NaCl), Carbon dioxide (CO2).
- Mixture – Two or more substances physically combined; components retain their own properties and can be separated by physical means. Examples: Air, soil, saltwater.
Laws of Chemical Combination
Law of Conservation of Mass (Lavoisier): In a chemical reaction, the total mass of reactants equals the total mass of products. Matter is neither created nor destroyed.
Law of Definite Proportions (Proust): A pure chemical compound always contains the same elements in the same mass ratio, regardless of source or amount.
Mole Concept
The mole is the SI unit for amount of substance. One mole of any substance contains 6.022 x 1023 particles (atoms, molecules, or ions). This number is called Avogadro's number (NA).
Key formulas
Valency
Valency is the combining capacity of an element. It equals the number of electrons an atom gains, loses, or shares to achieve a stable electronic configuration (usually 8 electrons in the outermost shell — octet rule).
- Examples:
- Hydrogen: valency 1 (1 electron in outermost shell)
- Oxygen: valency 2 (needs 2 electrons to complete octet)
- Nitrogen: valency 3
- Carbon: valency 4
Writing Chemical Formulae
The formula of a compound is written by crossing the valencies of the combining elements.
Example: Calcium (valency 2) + Chlorine (valency 1) → CaCl2 (cross valencies: Ca2 and Cl1, giving CaCl2).
Molecular Mass and Formula Unit Mass
Key formulas
States of Matter and Kinetic Theory
- Matter exists in three main states depending on the arrangement and energy of particles:
- Solid – Particles tightly packed, fixed positions, definite shape and volume.
- Liquid – Particles close but free to move, definite volume, no fixed shape.
- Gas – Particles far apart, moving rapidly, no fixed shape or volume.
The kinetic theory states that all particles of matter are in continuous random motion and that temperature is a measure of the average kinetic energy of particles.
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Calculate the molecular mass of sulphuric acid (H2SO4).
Atoms: 2 H + 1 S + 4 O.
Mass = 2(1) + 32 + 4(16) = 2 + 32 + 64 = 98 u (unified atomic mass units).
How many moles are present in 44 g of carbon dioxide (CO2)? Molar mass of CO2 = 12 + 2(16) = 44 g/mol. Moles = 44 / 44 = 1 mole.
Calculate the number of molecules in 2 moles of water.
Number of molecules = 2 x 6.022 x 1023 = 1.2044 x 1024 molecules.
Write the chemical formula for aluminium oxide. Aluminium valency = 3; Oxygen valency = 2. Cross valencies: Al2O3. Formula: Al2O3.
Verify the Law of Conservation of Mass for: 2H2 + O2 → 2H2O.
Mass of reactants: 2(2) + 32 = 4 + 32 = 36 u. Mass of products: 2(18) = 36 u. Mass is conserved. Verified.
What is the percentage of oxygen in water (H2O)?
Molecular mass of H2O = 18 u. Mass of O = 16 u. % O = (16/18) x 100 = 88.9%.
Iron (Fe, molar mass 56 g/mol) and oxygen (O2, molar mass 32 g/mol) react to form iron oxide. If 56 g of Fe is used, how many moles of Fe react? Moles = 56 / 56 = 1 mole of Fe.
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Key Formulas
Key formulas
Common mistakes
- Confusing atomic mass (mass of a single atom in u) with molar mass (mass of one mole in grams) — numerically the same value but different units and meaning.
- Using the wrong valency when writing formulae — always double-check by referring to standard valency tables.
- Forgetting that in ionic compounds we calculate formula unit mass, not molecular mass.
Summary
Matter is composed of elements, compounds, and mixtures. Chemical combination follows the Laws of Conservation of Mass and Definite Proportions. The mole concept links macroscopic masses to numbers of atoms and molecules via Avogadro's number. Valency determines how elements combine, and molecular/formula unit mass is the sum of constituent atomic masses. These atomic foundations underpin all of chemistry.