Chapter 6 - Shapes of Molecules

• • The solid lines are in the plane of the paper
  • The solid wedge shows a bond coming towards you
  • The dotted wedge shows a bond going away from you

  
• Electron pairs repel and try to get as far apart as possible to minimise repulsion
• Areas of electron density:

  Either a bond (single, double or triple) or a lone pair around the central atom
• Name: Tetrahedral

  Lone pairs: 0

  Bonded pairs: 4

  Bond angle: 109.5
• Name: Pyramidal

  Lone pairs: 1

  Bonded pairs: 3

  Bond angle: 107
• Name: non-linear

  Lone pairs: 2

  Bonded pairs: 2

  Bond angle: 104.5
• Name: Linear

  Electron pairs/regions: 2

  Bond angle: 180
• Name: Trigonal Planar

  Electron pairs/regions: 3

  Bond angle: 120
• Name: Tetrahedral

  Electron pairs/regions: 4

  Bond angle: 109.5
• Name: Octahedral

  Electron pairs/regions: 6

  Bond angle: 90
• Factors that could cause an unequal share of the electrons

  • Different sized atoms (a)
  • Different nuclear charges (b)
  • Leading to the shared pair being closer to one atom

  
• Electronegativity:

  A measure of the attraction of a bonded atom for the pair of electrons in a covalent bond

  
• In a non-polar bond, the bonded pair of electrons is shared equally between the bonded atoms
• Non-polar covalent bond:

  • The bonded atoms are the same (eg H2)
  • The bonded atoms have the same or similar electronegativities (eg C-H)
• In a polar bond, the bonded pair of electrons is shared unequally between the bonded atoms. The bonded atoms must have different electronegativities.
• Criteria for polar molecules

  • Polar bonds
  • A lone pair on the central atom
  • Different terminal atoms
• • Hydrogen has an electronegativity of 2.1 and chlorine 3.0
  • Chlorine is more electronegative than hydrogen
  • Chlorine has a partial negative charge and hydrogen a partial positive charge
  • The HCl molecule has a permanent dipole
  • The partial negative charge of the chlorine attracts the partial positive charge of the hydrogen of another HCl molecule

  
• London Forces

  • Electrons move around
  • Creates an instantaneous dipole
  • This induces a dipole on a neighbouring atom/molecule
  • This induced dipole induces further dipoles on neighbouring molecules, which then attract one another
• The more electrons in each molecule:

  • the larger the instantaneous and induced dipoles
  • The greater the induced dipole-dipole interaction
  • The stronger the attractive forces between molecule
  • So, the more energy is then needed to overcome the intermolecular forces - so higher the boiling point
• Simple Molecular Substances

  • Form a regular structure called a simple molecular lattice
  • Molecules are held together by weak intermolecular forces
  • The atoms within each molecule are bonded together strongly by covalent bonds
• Simple molecular substances have a low melting and boiling point due to weak intermolecular forces between molecules
• Solubility of non-polar simple molecular substances

  Non-Polar Solvent

  • Intermolecular forces weaken between substance in simple molecular lattice
  • Compound dissolves and mixes in with non polar solvent
• Solubility of non-polar simple molecular substances

  Polar Solvent

  • Solvent and simple molecular do not mix
  • Intermolecular bonding within solvent is too strong to be broken
• Solubility of polar simple molecular substances

  Polar Solvent

  • May dissolve
  • Depends on the strength of the interactions formed between polar solvent and polar simple molecule
• Electrical Conductivity

  • No mobile charged particles in simple molecular structures so simple molecules are insulators
• Criteria for Hydrogen bonding:

  • An electronegative atom with a lone pair of electrons (NOF atoms)
  • A hydrogen atom directly attached to the NOF atom
• Hydrogen bonding is a special type of permanent dipole-dipole interaction.