6 - Shapes & IM forces

• electron pair repulsion theory

  • pairs of electrons around a central atom repel each other
  • they move as far apart as possible to minimise this repulsion
  • this means molecules containing 3 or more atoms have shapes
• wedges are used to help visualise the structure of molecules in 3-dimensions:

  • a solid line represents a bond in the plane of the paper
  • a solid wedge comes out of the plane of the paper
  • a dotted wedge goes into the plane of the paper

  
• if there are 2 pairs of electrons (or 2 regions of electron density) repelling around a central atom:

  • the greatest bond angle possible around this atom is 180 degrees
  • the molecule is linear

  
• double and triple bonds count as one region of electron density

  • CO2 and HCN are counted as linear
• 3 bonded pairs / 3 regions of electron density

  • shape = trigonal planar
  • bond angle = 120
  • all on one plane
• 2 bonded pairs / 2 regions of electron density

  • shape = linear
  • bond angle = 180
  • all on one plane
• 4 bonded pairs / 4 regions of electron density

  • shape = tetrahedral
  • bond angle = 109.5
  • multiple planes
• 5 bonded pairs / 5 regions of electron density

  • shape = trigonal bipyramidal
  • bond angle = 120 and 90
  • multiple planes
• 6 bonded pairs / 6 regions of electron density

  • shape = octahedral
  • bond angle = 90
  • multiple planes
• shape = linear

  bond angle = 180

  
• shape = linear

  bond angle = 180

  
• shape = trigonal planar

  bond angle = 120

  
• shape = tetrahedral

  bond angle = 109.5

  
• shape = trigonal bipyramidal

  bond angle = 90 and 120

  
• shape = octahedral

  bond angle = 90

  
• a lone pair of electrons is

  • slightly closer to the central atom and
  • occupies more space than a bonded pair so
  • it repels more than a bonded pair
• a bonded pair and another bonded pair have the lowest repulsion
• a lone pair and another lone pair have the greatest repulsion
• • lone pairs repel more strongly than bonded pairs.
  • therefore, lone pairs push bonded pairs slightly closer together, decreasing the bond angle
  • the bond angle is reduced by about 2.5 degrees for each lone pair.
• complete the table

  A) tetrahedral

  B) pyramidal

  C) non-linear

  D) 109.5

  E) 107

  F) 104.5

  G) 1

  H) 2

  I) 2

  J) 3

  
• ammonium ion:

  • 1 dative bond
  • 4 areas of electron density
  • that repel each other as far apart as possible
  • shape = tetrahedral
  • bond angle = 109.5

  
• predict the bond angle of this molecule (below) and explain your reasoning

  • bond angle = 117.5
  • 3 regions of electron density
  • 1 double bond
  • 1 single bond
  • 1 lone pair
  • the pairs of electrons will move as far apart as possible to minimise repulsion
  • 120 - 2.5 = 117.5

  
• what makes the molecule OCl2 polar?

  the two lone pairs of electrons on the oxygen
• electronegativity - a measure of the ability of an atom in a molecule to attract a pair of electrons in a covalent bond to itself
• electronegativity

  • the higher the value of electronegativity, the more strongly the bonding electrons are attracted to that atom
  • the most electronegative atom is fluorine, followed by oxygen and nitrogen
  • electronegativity increases up a group
  • electronegativity increases across a period (up to group 7)
  • group 1 metals have the least electronegative atoms