Chemistry Molecular Shapes and Bonds

CAIE AS Level Chemistry Summary Notes

•  Updated to 2023-2025 syllabus
• Summarized notes on the theory syllabus

1. Atoms, Molecules, and Stoichiometry

• 1.1. Relative Mass
  • Atomic mass (Ar): weighted average mass of an atom
  • Molecular mass (Mr): mass of a molecule
  • Formula mass: mass of one formula unit of a compound
  • Isotopic mass: mass of a particular isotope of an element
  • Unified atomic mass unit: u = 1.66 x 10⁻²⁷ kg
• 1.2. The Mole
  • Mole: amount of substance with the same number of particles as 12g of carbon-12 isotope
  • Avogadro's constant: number of particles in a mole (6.02 × 10²³)
• 1.3. Mass Spectra
  • Abundance of isotopes can be visualized on a mass spectra diagram
  • Relative Abundance = (Peak Height / Total Height) x 100%
• 1.4. Empirical and Molecular Formulae
  • Empirical Formula: simplest whole number ratio of atoms in a compound
  • Molecular Formula: actual number of atoms of each element in a molecule
  • Molecular Formula = (Empirical Formula)n where n = (Molecular Mass/Mass of Empirical Formula)
  • % Composition: percentage of each element within the compound

2. Chemical Bonding

• 2.1. Electronegativity and Bonding
  • Electronegativity: power of an atom to attract electrons
  • Increases across periods, decreases down groups
  • Dipole Moment: slight charges on atoms in a bond due to differences in electronegativity
• 2.2. Ionic (Electrovalent) Bonding
  • Electrostatic attraction between oppositely charged ions (ions)
  • Structure: giant ionic lattice / crystalline solids; High m.p./b.p.
  • Coordination Number: number of ions surrounding a given ion in an ionic solid
• 2.3. Dot and Cross Diagrams
• 2.4. Metallic Bonding
  • Strong electrostatic forces of attraction between metal cations and delocalized mobile electrons
  • Structure: lattice of +ve ions surrounded by mobile electrons
• 2.5. Covalent Bonding
  • Sharing of electron pairs between nuclei of two atoms.
  • Bonding electrons: electrons involved in bond formation
  • Non-bonding electrons: electron pairs not involved in bonding (lone pairs).
• 2.6. Coordinate (Dative Covalent) Bonding
  • Both electrons in the bond come from the same atom
• 2.7. Orbital Overlap
  • Atomic orbitals containing unpaired valence electrons need to overlap to form a covalent bond.
• 2.8. Shapes of Molecules
  • Molecular shape determined by electron pairs surrounding central atom (lone pairs and bonded).
  • VSEPR Theory used to determine the molecular geometry and bond angles.
  • VSEPR (valence shell electron pair repulsion) Theory used to describe the 3D shape of molecules.

3. States of Matter

• 3.1. Basic Assumptions of Kinetic Theory
  • Real gases deviate from ideal behaviour at high pressure and low temperature as volume of molecules become considerable and intermolecular forces attract each other significantly.
  • Idea gases have rapidly moving and randomly moving particles. The distance between gases is greater than the size of molecules. Forces of attraction do not exist between the molecules. No forces of attraction/repulsion between molecules.

4. Chemical Energetics

• 4.1. Energy Change In Reactions
  • Exothermic reactions: release energy (heat)
  • Surroundings get warmer
  • Bond making.
  • ΔH (enthalpy change) negative
  • Endothermic reactions: absorb energy (heat)
  • Surroundings get colder
  • Bond breaking
  • ΔH positive
• 4.2. Reaction Pathway Diagrams
• 4.3. Activation Energy
• 4.4. Enthalpy Change Definitions

5. Equilibria

• 5.1. Introduction
  • Reversible reaction: can proceed forward and backward.
  • Dynamic Equilibrium: forward and backward reactions occur at equal rates.

6. The Periodic Table: Chemical Periodicity

7. Group 2

8. Group 17

9. Nitrogen and Sulfur

10. An Introduction to AS Level Organic Chemistry

11. Hydrocarbons

12. Halogen Compounds

13. Hydroxy Compounds

14. Carbonyl Compounds

15. Carboxylic Acids and Derivatives

16. Nitrogen Compounds

17. Polymerisation

18. Organic Synthesis

19. Analytical Techniques