Overview
This lecture covers the complete content required for A Level Chemistry, summarizing modules 2 to 6. It includes atomic structure, bonding, periodicity, organic chemistry, analytical techniques, and key practical skills.
Atomic Structure & The Periodic Table
- Atoms consist of protons (+1 charge, mass 1), neutrons (0 charge, mass 1), and electrons (-1 charge, negligible mass).
- The nucleus is tiny compared to the atom; most of the atom is empty space.
- Atomic number (Z) = number of protons; mass number (A) = protons + neutrons.
- Isotopes have the same atomic number but different mass numbers (different neutrons).
- Relative atomic mass is the weighted average of isotopes.
Chemical Bonding & Formulas
- Ions are formed by loss/gain of electrons for noble gas configuration; group 1: +1, group 2: +2, group 6: -2, group 7: -1.
- Transition metals have variable oxidation states.
- Balancing equations requires counting elements and adding state symbols (s, l, g, aq).
- Ionic compounds must have balanced charges; know key polyatomic ions.
Moles & Calculations
- Mole = amount containing Avogadro’s number (6.02 × 10²³) of particles.
- n = m / Mr (moles = mass/Mr), particles = moles × Avogadro’s number.
- 1 mol of gas at RTP = 24 dm³; ideal gas equation: PV = nRT.
- Empirical formula = simplest ratio, molecular formula = actual numbers.
- Hydrated salts: water of crystallization can be calculated by mass loss.
Practical Skills & Titrations
- Practicals include finding composition using gas collection, determining relative atomic mass, and empirical formula by heating.
- Making standard solutions requires accurate transfer and dilution.
- Acid-base titrations: calculate moles, use balanced equations, and determine unknown concentrations.
- Always include state symbols and accuracy with significant figures in calculations.
Chemical Reactions & Redox
- Oxidation state rules: uncombined elements = 0, group 1 = +1, O = -2 (except in peroxides), H = +1 (except in hydrides).
- Redox: oxidation is loss of electrons, reduction is gain; identified by changes in oxidation state.
- Disproportionation: same species is both oxidized and reduced.
Structure, Bonding & Properties
- Ionic: between metals and non-metals, high melting/boiling points, conduct when molten/aqueous.
- Covalent: sharing electrons between non-metals, forms molecules or giant lattices.
- Metallic: delocalized electrons, conduct, malleable, ductile.
- Intermolecular forces: permanent dipole, induced/London, hydrogen bonding.
Shapes of Molecules & Isomerism
- Shape determined by electron pair repulsion (VSEPR): linear (180°), trigonal planar (120°), tetrahedral (109.5°), etc.
- E/Z isomerism arises with restricted rotation around C=C and different groups.
- Optical isomerism: chiral carbons produce non-superimposable mirror images.
The Periodic Table & Trends
- Arranged by increasing atomic number; groups have similar properties.
- Atomic radius decreases across a period and increases down a group.
- Ionization energy generally increases across a period, decreases down a group; drops at new subshells or electron pairing.
Group Chemistry
- Group 2: reactivity increases down the group, oxides and hydroxides more soluble, important uses in industry and medicine.
- Group 7 (halogens): reactivity decreases down the group, displacement reactions, silver nitrate tests for halides.
Energetics & Equilibria
- Enthalpy change (ΔH): exothermic (-), endothermic (+), measured in kJ/mol.
- Standard conditions: 298 K, 100 kPa.
- Hess’s law: enthalpy change is independent of pathway.
- Bond enthalpy: energy to break 1 mol of bonds.
- Equilibrium: rate forward = rate reverse; Le Chatelier’s principle explains shifts.
Rates of Reaction & Kinetics
- Rate affected by concentration, pressure, temperature, and catalysts.
- Rate = change in concentration/time; catalysts lower activation energy.
- Rate equations: rate = k[A]^x[B]^y; units depend on overall order.
- Arrhenius equation relates k to temperature and activation energy.
Organic Chemistry & Mechanisms
- Naming uses IUPAC rules: identify longest carbon chain, branches, and functional groups.
- Alkanes (single bonds), alkenes (double bonds), alcohols, halogenoalkanes, amines, carboxylic acids, esters, etc.
- Key mechanisms: free radical substitution (alkanes), electrophilic addition (alkenes), nucleophilic substitution (halogenoalkanes, amines), nucleophilic addition (carbonyls).
- Isomerism: chain, position, functional group, optical, and E/Z.
Analysis & Practical Techniques
- Mass spectrometry: molecular ion peak gives Mr, fragmentation for structure.
- IR spectroscopy: characteristic absorptions for O-H, C=O, etc.
- NMR: number of peaks (environments), chemical shift, splitting (n+1 rule), integration (proton number).
- Chromatography (TLC, column, gas): separates based on solubility/retention, use Rf values.
Transition Metals & Complexes
- Form complex ions with variable oxidation states, colored compounds, catalysis.
- Ligands: monodentate, bidentate, polydentate; shapes include octahedral, tetrahedral, square planar.
- Precipitation, ligand exchange, redox, and color changes are key reactions.
Key Terms & Definitions
- Isotope — Atoms with same protons but different neutrons.
- Empirical formula — Simplest whole-number ratio of elements.
- Mole (mol) — Avogadro’s number of particles.
- Enthalpy change (ΔH) — Heat change at constant pressure.
- Ionization energy — Energy to remove an electron from an atom/gas.
- Oxidation state — Charge an atom would have if electrons assigned by rules.
- Electronegativity — Ability of an atom to attract electrons in a bond.
Action Items / Next Steps
- Review free revision guide and practice questions via provided links.
- Practice balancing equations, titration calculations, and drawing mechanisms.
- Complete assigned homework, readings, and revisit topics with supplementary videos as needed.