Edexcel IGCSE Chemistry Revision

Resources for Revision

• Free revision guide available on the website.
• Links to teaching videos for unclear topics.
• Free multiple-choice questions online.
• Predictive papers for exams with video walkthroughs.

Structure of Atoms

• Atoms consist of protons, neutrons, and electrons.
  • Protons: Mass = 1, Charge = +1
  • Neutrons: Mass = 1, Charge = 0
  • Electrons: Very small mass, Charge = -1
• Relative charge is used instead of actual charge in coulombs.
• Atomic mass based on carbon-12 standard.
• Atom's structure mostly empty space.
• Nucleus diameter vastly smaller than atom diameter.

Evolution of Atomic Structure

• Initial concept was that atoms were indivisible.
• Evolved to the solid sphere model with embedded electrons.
• Current understanding: nucleus with orbiting electrons.

Periodic Table and Atomic Numbers

• Mass number = protons + neutrons.
• Atomic number symbolized as Z.
• Mass number symbolized as A.
• Isotopes have same atomic number, different mass numbers.
• Isotopes: Same electron arrangement, different physical properties.

Definitions in Chemistry

• Relative molecular mass and relative atomic mass are compared to carbon-12.
• Mass spectrometry calculates atomic mass by isotope abundance.

Ionization Energy

• First ionization energy: energy needed to remove one electron from each atom in one mole of gaseous atoms to form +1 ions.
• Factors affecting ionization energy include atomic radius, electron shielding, and nuclear charge.
• Trend: Increased across periods, decreased at start of new periods.

Electron Configuration and Periodic Table Blocks

• Electrons fill orbitals based on increasing energy levels.
• Blocks: s, p, d, f based on orbitals.
• Periodic table divided into blocks aligning with electron configurations.

Bonding

• Ionic Bonding: transfer of electrons from metals to non-metals.
  • Example: Magnesium chloride (MgCl2).
  • Ionic compounds have high melting points, conduct electricity when molten.
• Covalent Bonding: sharing of electrons between non-metals.
  • Types: Single, Double, and Triple bonds.
  • Dative covalent bonding involves one atom providing both electrons for a bond.

Molecular Shapes and Theories

• Shapes: Linear, Trigonal Planar, Tetrahedral, Trigonal Pyramidal, etc.
• Bond angles vary by shape and lone pair electron repulsion.
• VSEPR theory explains shapes based on electron pair repulsion.

Electronegativity and Polarity

• Electronegativity: atom's ability to attract electrons.
• Trends: Increases across a period, decreases down a group.
• Bond types: Pure covalent to ionic based on electronegativity difference.

Intermolecular Forces

• London Forces: Weakest, present in all molecules.
• Permanent Dipole-Dipole Interactions: Occur in polar molecules.
• Hydrogen Bonding: Strongest, occurs in molecules with N, O, F bonded to H.

Solubility and Solvents

• Solubility affected by polarity of solvent and solute.
• Water is a polar solvent, dissolves many ionic compounds.
• Non-polar substances require non-polar solvents like oils or fats.

Metallic Bonding and Properties

• Metals consist of positive ions in a sea of delocalized electrons.
• Properties: Conduct electricity, malleable, ductile, high melting points.

Giant Covalent Structures

• Examples: Silicon Dioxide, Diamond, Graphite.
• High melting points, poor electrical conductivity (except graphite).

Simple Molecular Structures

• Low melting and boiling points due to weak intermolecular forces.
• Examples include water, ammonia, and carbon dioxide.

Oxidation States and Reactions

• Oxidation involves loss of electrons, reduction involves gain of electrons.
• Use of oxidation numbers to write and balance redox equations.
• Disproportionation reactions involve simultaneous oxidation and reduction of the same element.

Periodicity

• Periodic table arranged by atomic number.
• Groups have similar chemical properties.
• Trends: Atomic radius, ionization energy, and reactivity vary across periods and groups.

Group 2 and Group 7 Elements

• Group 2: Alkaline Earth Metals
  • Reactivity increases down the group.
  • Common compounds include oxides and hydroxides.
• Group 7: Halogens
  • Reactivity decreases down the group.
  • Displacement reactions show reactivity trends.

Chemical Analysis and Tests

• Flame tests identify metal ions by color.
• Precipitation reactions help identify halides and other ions.
• Use of acids and bases for testing carbonates and sulfates.

Calculations and Stoichiometry

• Use of moles, mass, and molecular mass in calculations.
• Ideal gas law and Avogadro's law in gas calculations.
• Concentration calculations involve converting between units and using balanced equations.

Organic Chemistry

• Alkanes: Saturated hydrocarbons, used as fuels.
• Alkenes: Unsaturated hydrocarbons, undergo addition reactions.
• Haloalkanes: Undergo nucleophilic substitution.
• Alcohols: Can be oxidized to aldehydes, ketones, or carboxylic acids.
• Esters: Formed by reaction of alcohols and acids.

Polymers

• Addition and condensation polymerization.
• Biodegradable polymers and their environmental impact.

Analytical Techniques

• Mass Spectrometry: Identifies compounds by fragmentation patterns.
• Infrared Spectroscopy: Identifies functional groups by absorption.
• NMR Spectroscopy: Identifies molecular structure and environment.

Energetics

• Enthalpy Changes: Exothermic and endothermic reactions.
• Hess's Law: Energy changes are independent of the pathway.
• Bond Enthalpies: Energy required to break/make bonds.

Kinetics

• Factors affecting rate: concentration, temperature, surface area, catalysts.
• Reaction mechanisms and rate equations.
• Arrhenius equation relates rate constant to temperature.

Equilibrium

• Reversible reactions and dynamic equilibrium.
• Le Chatelier's Principle predicts the effect of changes in conditions.
• Equilibrium constants (Kc, Kp) indicate reaction position.

Acids and Bases

• pH calculations for strong and weak acids/bases.
• Buffers maintain pH stability.

Transition Metals

• Form colored ions, variable oxidation states, act as catalysts.
• Complex ions and ligand exchange reactions.
• Redox reactions involving transition metals.

This guide encompasses key topics in Edexcel IGCSE Chemistry, designed to assist in revision and understanding of essential concepts. Each section provides a concise summary of fundamental principles and reactions encountered in the syllabus.