Electron- Pair Repulsion Theory: Electron pairs arranged as far apart as possible to minimise
repulsion. Number of bond pairs and lone pairs of electrons surrounding central atom determine
shape. This holds the bonded atoms in a definite shape.
– Name shape.
– State number of bonded pairs of electrons and lone pairs of electrons.
– If a lone pair state electron pairs repel but lone pair repels more than bonded pairs.
– If no lone pairs state bonding pairs of electrons repel one another equally- so equal bond
angles.
• 3D Diagrams: A solid wedge comes out of plane of paper. A dashed wedge goes into plane of
paper. Narrowest part of wedge closest to central atom.
• Lone Pair: Lone pair occupies more space than a bonded pair as closer to central atom. Lone pair
repels more strongly than a bond pair. Therefore lone pairs repel bonded pairs slightly closer
together, decreasing the bond angle between the bonded pair by 2.50
for each lone pair. Imagine
if lone pair is normal bond and then take 2.5 from angle.
• Multiple Bonds: All the multiple bonds in one overlap is treated as one bonding region.
• Octahedral: SF6 is given octahedral name as it describes the shape of the molecule.
• Non Linear: 2 bonding pairs and 2 lone pairs around central atom. 104.5 bond angle.
• Use dot and cross diagrams to work out shape and bond angles.
• Know the example below for each shape.
• Make sure there are brackets and charge for a displayed formula of an ion.
• When there is a charge- an ion- means more or less electrons in outershell, like in this azide ion
here.