The following notes were written for the previous IB syllabus (2009). The new IB syllabus for first examinations 2016 can be accessed by clicking the link below.

IB syllabus for first examinations 2016

Bonding (hl)

14.1 - Shapes of Molecules

The actual arrangement of the atoms around the central atom in a molecule.

In order to find this we need to consider the number of regions of electron density (electron pairs or unpaired single electrons in some cases) around the central atom. The electrons will repel as far apart as possible as they have the same charge.

The valence shell electron pair repulsion theory (VSEPRT) gives us a means of working out the shapes of molecules and ions.

  1. Draw the Lewis structure to show all of the valence shell electrons.

  2. Then count the number of regions of electron density - this can then be translated into the electronic shape.

  3. Now consider the number of attached atoms and their orientation keeping the lone (non-bonding) pairs as far apart as possible without changing the electronic shape.

  4. The resultant molecular arrangement gives the molecular shape. Note that it may be slightly distorted by repulsions between regions of electron density.


For a fuller account see VSEPRT

Molecules in three dimensions

Click here and when the viewing screen opens right click on the image to see the options.

note: the MDL Chime plug-in must previously be installed (download here)

BF3 CH4 PCl5 SF6
NH3 XeF4

Predicting molecular shapes


Summary of the possible geometries

No. of electron pairs around the central atom electronic shape no of attached atoms molecular shape
2 linear 2 linear
3 trigonal planar 3 trigonal planar
4 tetrahedral 4 tetrahedral
4 tetrahedral 3 pyramidal
4 tetrahedral 2 angular or bent
5 trigonal bipyramidal 5 trigonal bipyramidal
5 trigonal bipyramidal 4 saw-horse (seesaw)
5 trigonal bipyramidal 3 T-shaped
5 trigonal bipyramidal 2 linear
6 octahedral 6 octahedral
6 octahedral 5 umbrella
6 octahedral 4 square planar
More information about the shapes of molecules  

14.2 - Hybridisation

This model explains the tetrahedral geometry of carbon and other atoms.

The electron structure of carbon is 1s2 2s2 2p2 suggesting that it should only be able to form two bonds (using the two singly occupied orbitals). However it is known to make four single bonds in many compounds and indeed never forms just two bonds. This can be explained by hybridisation - the mixing of atomic orbitals producing degenerate orbitals used for bonding.

  • sp3 hybridisation occurs when the 2s and 2p orbitals merge to become sp3 orbitals (all of equal energy, length etc.).

  • sp2 is the same except only two of the p orbitals are hybridised, leaving one p orbital unchanged

  • sp is the same except only one of the p orbitals is hybridised and two p orbitals are left unchanged

hybridisation geometry example carbon in
sp3 tetrahedral methane CH4
sp2 trigonal planar ethene CH2=CH2
sp linear ethyne CH=CH
More information about hybridisation  

14.3 - Delocalisation of electrons

When a particular molecule can be represented as several different Lewis structures is is generally not actually any of these, but a hybrid (mixture) of all of them. This can be represented either by using delocalised electrons, or through resonance (where each possible structure is drawn and the actual state 'resonates' between them. The delocalisation of these pi electrons (which is effectively what happens) makes the molecule more stable (as evidenced by lower energy) and gives the bonds a shorter length than would be expected.

Examples:

  • benzene

  • O3

  • SO42-

More information about delocalisation of electrons  



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