Overview
This lecture covers the fundamental concepts of Cambridge IGCSE Chemistry, including atomic structure, bonding, the periodic table, chemical reactions, separation methods, calculations, and practical applications like analysis and organic chemistry.
Atoms, Elements, Compounds & Mixtures
- Atoms are the smallest units of matter; elements are types of atoms in the periodic table.
- Compounds contain two or more different atoms chemically bonded, e.g., H₂O.
- Mixtures are combinations of substances not chemically bonded, e.g., air or salt water.
- Word and chemical equations represent reactions; atoms are conserved, so equations must be balanced.
Separation Techniques
- Filtration separates insoluble solids from liquids.
- Crystallization isolates a solute by evaporating the solvent.
- Distillation and fractional distillation separate mixtures based on different boiling points.
- Chromatography separates mixtures using stationary and mobile phases; Rf value = distance moved by substance/distance moved by solvent.
States of Matter & Changes
- Solids, liquids, and gases differ in particle arrangement and energy.
- Melting, evaporation, and condensation are physical changes; no new substances form.
- State symbols: (s), (l), (g), (aq) in equations.
Atomic Structure & Isotopes
- Atoms consist of a nucleus (protons, neutrons) and electrons in shells.
- Atomic number = protons; mass number = protons + neutrons.
- Isotopes are atoms of the same element with different neutrons.
- Electron configuration fills shells: 2, 8, 8, 2 (up to Ca).
Periodic Table & Groups
- Metals (left) lose electrons, form positive ions; non-metals (right) gain electrons.
- Group number = outer shell electrons.
- Group 1: Alkali metals, more reactive down the group.
- Group 7: Halogens, less reactive down the group.
- Group 0: Noble gases, unreactive.
Bonding & Structure
- Ionic bonding: metal + non-metal, transfer of electrons, form lattice, conduct electricity when molten/aqueous.
- Covalent bonding: non-metals share electrons, form molecules.
- Simple covalent molecules: low melting points, do not conduct electricity.
- Giant covalent: e.g., diamond (hard, high melting point), graphite (conducts electricity, layers slide).
- Metallic bonding: lattice of ions with delocalized electrons; metals conduct electricity.
- Alloys are stronger due to disrupted layers.
Chemical Calculations
- Relative atomic/formula mass: sum of atomic masses in a formula.
- Mole = mass ÷ Mr (relative formula mass).
- Conservation of mass: total mass of reactants = products.
- Limiting reactant determines amount of product.
- Concentration: moles/volume (mol/dm³).
Reactivity & Extraction
- Reactivity series ranks metals by reactivity; more reactive metals displace less reactive from compounds.
- Extraction by carbon possible if metal is less reactive than carbon; smelting and reduction.
- Oxidation: loss of electrons; reduction: gain of electrons (OIL RIG).
Acids, Bases & Salts
- Acid + alkali = salt + water (neutralization).
- Strong acids fully dissociate; weak acids partially dissociate.
- pH scale is logarithmic; lower pH = higher H⁺ concentration.
- Tests for gases: hydrogen (squeaky pop), oxygen (relights splint), CO₂ (limewater), chlorine (bleaches litmus).
Electrolysis
- Electrolysis splits compounds using electricity.
- Anions oxidized at anode; cations reduced at cathode.
- In solution, less reactive ions are discharged.
- Used in metal extraction (e.g., aluminum).
Energetics
- Exothermic reactions release energy (temperature rises), endothermic absorb (temperature falls).
- Activation energy is energy needed to start a reaction.
- Energy profile diagrams show energy changes.
Rates of Reaction
- Rate = change in quantity ÷ time.
- Increased concentration, pressure, temperature, surface area, or catalyst increases rate.
- Catalysts lower activation energy but are not used up.
- Tangents on graphs give instantaneous rate.
Reversible Reactions & Equilibrium
- Reversible reactions can go both ways; equilibrium reached in closed systems.
- Le Chatelier's principle: system counteracts changes in conditions.
- Higher pressure or concentration favors side with fewer molecules; higher temperature favors endothermic reaction.
Organic Chemistry (Triple)
- Alkanes: CₙH₂ₙ₊₂; saturated hydrocarbons.
- Alkenes: CₙH₂ₙ; unsaturated, contain C=C double bond, tested with bromine water.
- Alcohols: contain –OH group; oxidize to carboxylic acids.
- Cracking produces alkenes from alkanes.
- Addition and condensation polymerization produce polymers.
Analysis and Testing
- Flame tests distinguish metal ions by color.
- Precipitate tests for metal ions and halides.
- Sulfate ions form white precipitate with BaCl₂ and HCl.
Environmental Chemistry
- Greenhouse gases trap heat; CO₂ and methane increase global warming.
- Water purification uses filtration and sterilization or desalination.
- Corrosion (rusting) can be prevented by coating or sacrificial metals.
Key Terms & Definitions
- Isotope — Atoms of the same element with different neutron numbers.
- Ionic bond — Electrostatic attraction between oppositely charged ions.
- Covalent bond — Shared pair of electrons between non-metals.
- Mole — Amount containing Avogadro’s number of particles.
- Electrolysis — Decomposition of substances using electricity.
- Activation energy — Minimum energy required to start a reaction.
- Equilibrium — State where forward and reverse reaction rates are equal.
- Catalyst — Substance that increases rate of reaction without being used up.
Action Items / Next Steps
- Review electron configuration and practice drawing dot and cross diagrams.
- Memorize the flame test colors and common ion precipitate results.
- Practice balancing equations and chemical calculations for mole and concentration.
- Read textbook/revision guide on organic chemistry if taking triple science.