6 - shapes of molecules and intermolecular forces

• Electron pair repulsion theory

  An electron has a negative charge, so electron pairs repel one another:

  
  

  - The electron pairs surrounding a central atom determine the shape of the molecule or ion.

  - The electron pairs repel one another as far apart as possible.

  - The arrangement of electron pairs minimises repulsion and thus holds the bonded atoms in a definite shape.

  - Different numbers of electron pairs result in different shapes.
• What does a solid line represent?


  A bond in the plane of the paper.
• What does a solid wedge represent?


  A bond coming out of the plane of the paper (towards you).
• What does a dotted wedge represent?


  A bond going into the plane of the the paper (away from you).
• Why does a lone pair of electron repel more strongly than a bonding pair of electrons?


  A lone pair of electrons is slightly closer to the central atom, and occupies more space than a bonded pair.
• Types of shapes molecules can adapt to

  - Linear

  - Trigonal planar

  - Non-linear (bent)

  - (Trigonal) pyramidal

  - Tetrahedral

  - Trigonal bipyramidal

  - Octahedral.
• Bonding in linear shaped molecules


  Bonding regions: 2

  Lone pairs: 0
• Bond angle in linear shaped molecules


  180 degrees
• Bonding in trigonal planar shaped molecules


  Bonding regions: 3

  Lone pairs: 0
• Bond angle in trigonal planar shaped molecules


  120 degrees
• Bonding in non-linear (bent) shaped molecules


  Bonding regions: 3 or 4

  Lone pairs: 1 or 2
• Bond angle in non-linear (bent) shaped molecules

  <120 degrees with 3 bonding regions 1 lone pair.

  < 109 degrees with 4 bonding regions and 2 lone pairs.
• Bonding in tetrahedral shaped molecules


  Bonding regions: 4

  Lone pairs: 0
• Bond angle in tetrahedral shaped molecules


  109.5 degrees
• Bonding in trigonal pyramidal shaped molecules

  Bonding regions: 4

  Lone pairs: 1
• Bond angle in trigonal pyramidal shaped molecules


  107 degrees
• Bonding in trigonal bipyramidal shaped molecules


  Bonding regions: 5

  Lone pairs: 0
• Bonding in octahedral shaped molecules


  Bonding regions: 6

  Lone pairs: 0
• Bond angle in octahedral shaped molecules


  90 degrees
• Example of an octahedral shaped molecule


  SF₆
• Electronegativity


  The relative tendency of an atom to attract the pair of electrons in a covalent bond towards itself.
• How does nuclear charge effect electronegativity?


  As the number of protons increases, the positive charge of the nucleus increases so the attraction between the nucleus and the bonding pair of electrons increases.
• Three factors effecting electronegativity


  1. Nuclear charge

  2. Atomic radius

  3. Shielding
• How does atomic radius effect electronegativity?


  The smaller the atomic radius the closer the bonding electrons will be to the nucleus of the atom, so electronegatvity increases.
• How does shielding effect electronegativity?


  Electrons in the inner shells screen electrons in the outer shell from the positive charge of the nucleus, because of this, the greater number of inner shells the lower the electronegativity.
• How does electronegtivity change across a period?


  Increases
• How does electronegativity change down a group?


  Decreases
• Elelctronegativity difference in a covalent bond


  0
• Electronegativity difference in a polar covalent bond


  0 to 1.8
• Electronegativity difference in an ionic bond


  greater than 1.8
• Non-polar bond


  When two atoms in a covalent bond have the same electronegativity the covalent bond is non-polar.
• When is a bond non-polar?


  When the bonded atoms:

  - are the same

  - have same or similar electronegativity
• Pure covalent bond


  Two identical non-metals sharing electrons equally.
• Polar bond


  When two atoms in a covalent bond have different electronegativities the covalent bond is polar and the electrons will be drawn towards the more electronegative atom.
• What is the result of a bond being polar?


  - Electron distribution is asymmetric.

  - Less electronegative atom gets a δ+ charge.

  - More electronegtive atom gets a δ- charge.
• How does polarity change with increasing difference in electronegativities?


  The greater the difference in electronegativity the more polar the bond becomes.
• Dipole


  Opposite charges on atoms in a bond.
• Dipole moment


  A measure of how polar a bond is.
• How is the direction of a dipole moment shown?


  An arrow pointing towards the partially negtive charged end of the dipole.
• What two things need to be taken into consideration when determining if a molecule with more than two different atoms is polar?

  - The polarity of each bond.

  - How the bonds are arranged in the molecule.