Principles of Chemistry

What you'll learn

This section covers the fundamental building blocks of chemistry that underpin all other topics in your IGCSE course. You'll master atomic structure, the periodic table organisation, chemical bonding types, formulae writing, equation balancing and quantitative calculations including the mole concept. These principles form the foundation for understanding chemical reactions, patterns and properties tested throughout your examination.

Key terms and definitions

Atom — the smallest particle of an element that can exist, consisting of a nucleus containing protons and neutrons, surrounded by electrons in shells

Relative atomic mass (Ar) — the average mass of atoms of an element compared to 1/12th the mass of a carbon-12 atom

Relative formula mass (Mr) — the sum of the relative atomic masses of all atoms in a formula of a substance

Mole — the unit for amount of substance; one mole of any substance contains 6.02 × 10²³ particles (Avogadro's constant)

Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons

Ion — an atom or group of atoms that has lost or gained electrons, resulting in a positive or negative charge

Covalent bond — a shared pair of electrons between two non-metal atoms

Ionic bond — the electrostatic attraction between oppositely charged ions formed by electron transfer between metal and non-metal atoms

Core concepts

Atomic structure and the periodic table

The atom consists of a central nucleus containing protons (relative mass 1, charge +1) and neutrons (relative mass 1, charge 0), with electrons (relative mass 1/1840, charge -1) orbiting in shells or energy levels.

The atomic number (proton number) defines the element and equals the number of protons in the nucleus. In a neutral atom, the number of electrons equals the number of protons. The mass number equals protons plus neutrons.

Electron arrangement determines chemical properties:

• First shell holds maximum 2 electrons
• Second shell holds maximum 8 electrons
• Third shell holds maximum 8 electrons (for elements 1-20)
• Electronic configurations are written as 2,8,1 for sodium (11 electrons)

The modern periodic table arranges elements in order of atomic number. Elements in the same group (vertical column) have the same number of outer shell electrons, giving them similar chemical properties. The period (horizontal row) number indicates the number of electron shells.

Metals are found on the left and centre of the periodic table; non-metals on the right. The transition metals occupy the central block between Groups 2 and 3.

Isotopes have identical chemical properties because they have the same electron arrangement, but different physical properties due to different masses. Relative atomic mass accounts for the abundance of different isotopes: Ar = Σ(isotope mass × percentage abundance) ÷ 100

Chemical bonding and structure

Ionic bonding occurs between metals and non-metals. Metal atoms lose electrons to form positive ions (cations); non-metal atoms gain electrons to form negative ions (anions). Both achieve stable electron configurations, typically a full outer shell.

Properties of ionic compounds:

• High melting and boiling points due to strong electrostatic forces
• Conduct electricity when molten or dissolved (ions free to move)
• Often soluble in water
• Form giant lattice structures with regular arrangement of ions

Covalent bonding occurs between non-metal atoms sharing electron pairs. Each shared pair forms one covalent bond. Examples include H₂ (single bond), O₂ (double bond), N₂ (triple bond).

Simple molecular covalent substances:

• Low melting and boiling points (weak intermolecular forces)
• Do not conduct electricity (no free charged particles)
• Often gases or liquids at room temperature

Giant covalent structures (diamond, graphite, silicon dioxide):

• Very high melting points (many strong covalent bonds)
• Do not conduct electricity (except graphite)
• Very hard (except graphite which has layers)

Metallic bonding consists of positive metal ions in a regular arrangement surrounded by delocalised electrons that can move freely throughout the structure.

Properties of metals:

• Good conductors of electricity and heat (free electrons)
• Malleable and ductile (layers can slide)
• High melting points (strong metallic bonds)
• Lustrous appearance

Chemical formulae and nomenclature

Chemical formulae show the number and type of atoms in a substance. You must learn common ions:

Positive ions: Na⁺, K⁺, Mg²⁺, Ca²⁺, Al³⁺, NH₄⁺, Ag⁺, Fe²⁺, Fe³⁺, Cu²⁺, Zn²⁺, Pb²⁺

Negative ions: Cl⁻, Br⁻, I⁻, O²⁻, S²⁻, OH⁻, NO₃⁻, CO₃²⁻, SO₄²⁻

For ionic compounds, balance charges to give neutral formula:

• Sodium chloride: Na⁺ + Cl⁻ = NaCl
• Magnesium oxide: Mg²⁺ + O²⁻ = MgO
• Calcium chloride: Ca²⁺ + 2Cl⁻ = CaCl₂
• Aluminium oxide: 2Al³⁺ + 3O²⁻ = Al₂O₃

For covalent compounds, use prefixes or systematic naming based on valency.

Balancing chemical equations

Chemical equations must be balanced because atoms cannot be created or destroyed (law of conservation of mass).

Steps for balancing equations:

1. Write correct formulae for all reactants and products
2. Count atoms of each element on both sides
3. Add coefficients (numbers in front) to balance, starting with the most complex molecule
4. Never change subscript numbers within formulae
5. Check all elements are balanced

State symbols indicate physical state: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution (dissolved in water)

Example: Unbalanced: Fe + O₂ → Fe₂O₃ Balanced: 4Fe(s) + 3O₂(g) → 2Fe₂O₃(s)

Relative formula mass and percentage composition

Calculate Mr by adding Ar values for all atoms in the formula.

For H₂O: Mr = (2 × 1) + 16 = 18 For CaCO₃: Mr = 40 + 12 + (3 × 16) = 100 For Ca(OH)₂: Mr = 40 + 2(16 + 1) = 74

Percentage composition = (Ar × number of atoms of element ÷ Mr) × 100

Percentage of calcium in CaCO₃ = (40 ÷ 100) × 100 = 40%

Mole calculations and reacting masses

One mole of any substance contains 6.02 × 10²³ particles and has a mass in grams equal to its relative formula mass.

Key formula: number of moles = mass (g) ÷ Mr

For reactions, the balanced equation shows the molar ratio of reactants and products.

2Mg + O₂ → 2MgO indicates:

• 2 moles of Mg react with 1 mole of O₂
• To produce 2 moles of MgO
• The ratio is 2:1:2

To calculate reacting masses:

1. Write balanced equation
2. Identify known and unknown substances
3. Calculate moles of known substance
4. Use molar ratio from equation
5. Calculate mass of unknown substance

Empirical formula shows the simplest whole number ratio of atoms in a compound. To find it from masses or percentages:

1. Divide each mass by Ar to get moles
2. Divide all mole values by the smallest
3. If necessary, multiply to get whole numbers

Molecular formula shows the actual number of atoms in a molecule. It's a simple multiple of the empirical formula.

Worked examples

Example 1: Balancing equations (2 marks)

Question: Balance the equation for the combustion of propane: C₃H₈ + O₂ → CO₂ + H₂O

Solution: C₃H₈ + 5O₂ → 3CO₂ + 4H₂O ✓

(1 mark for correct coefficients for products; 1 mark for correct coefficient for oxygen)

Example 2: Mole calculation (3 marks)

Question: Calculate the mass of magnesium oxide produced when 6.0 g of magnesium burns completely in oxygen. (Ar: Mg = 24, O = 16)

Solution: Equation: 2Mg + O₂ → 2MgO ✓ Moles of Mg = 6.0 ÷ 24 = 0.25 mol ✓ Molar ratio Mg:MgO = 1:1, so moles MgO = 0.25 mol Mr of MgO = 24 + 16 = 40 Mass of MgO = 0.25 × 40 = 10 g ✓

(1 mark for correct equation or molar ratio; 1 mark for correct moles; 1 mark for correct final answer with working)

Example 3: Empirical formula (4 marks)

Question: A compound contains 52.2% carbon, 13.0% hydrogen and 34.8% oxygen by mass. Calculate its empirical formula. (Ar: C = 12, H = 1, O = 16)

Solution:

Element	C	H	O
Mass (g)	52.2	13.0	34.8
Moles (÷ Ar)	52.2 ÷ 12 = 4.35	13.0 ÷ 1 = 13.0	34.8 ÷ 16 = 2.175
Divide by smallest	4.35 ÷ 2.175 = 2	13.0 ÷ 2.175 = 6	2.175 ÷ 2.175 = 1

✓ for dividing by Ar ✓ for dividing by smallest value ✓ for whole number ratio

Empirical formula: C₂H₆O ✓

Common mistakes and how to avoid them

• Changing formulae when balancing equations — only add coefficients in front of formulae, never alter the subscript numbers within the formula itself (e.g., write 2H₂O not H₄O₂)

• Confusing mass number with Ar — mass number is for a specific isotope and is always a whole number; Ar is the weighted average and usually has decimal places

• Forgetting to multiply bracketed groups — in Ca(NO₃)₂, there are 2 nitrogen atoms and 6 oxygen atoms; always multiply everything inside brackets by the subscript number outside

• Using Mr instead of Ar in percentage composition — use Ar (single element) for the numerator and Mr (whole compound) for the denominator

• Incorrect molar ratios — always refer to the coefficients in the balanced equation, not the subscripts in formulae; in 2H₂ + O₂ → 2H₂O, the ratio is 2:1:2, not 4:2:2

• Not simplifying empirical formulae — after dividing by the smallest, if you get 1:1.5:2, multiply all by 2 to get whole numbers (2:3:4)

Exam technique for "Principles of Chemistry"

• Command words matter: "State" requires simple facts without explanation; "Explain" requires reasoning using scientific knowledge; "Calculate" requires numerical working shown step-by-step even if using a calculator

• Show all working for calculations — even if your final answer is incorrect, you can gain method marks for correct steps; write the formula you're using, then substitute values, then calculate

• Standard form and significant figures — give answers to the same number of significant figures as the data with the least sig figs in the question; for Avogadro's constant use 6.02 × 10²³ not 602000000000000000000000

• Check equation balancing — systematically count atoms of each element on both sides; examiners commonly give 2 marks for fully correct and 0 marks if any element is unbalanced

Quick revision summary

Atoms consist of protons, neutrons and electrons in shells. The periodic table arranges elements by atomic number; groups share outer electron numbers. Ionic bonding transfers electrons between metals and non-metals forming charged ions; covalent bonding shares electron pairs between non-metals. Balance equations using coefficients only. Calculate Mr by summing Ar values. The mole links mass to particles: moles = mass ÷ Mr. Use balanced equations to find molar ratios for reacting mass calculations. Empirical formulae show simplest atom ratios.