IB syllabus > bonding (hl) > 14.3 

14.3 - Delocalisation


14.3.1:- Describe the delocalization of p electrons and explain how this can account for the structures of some species. Examples shoud include NO3-, NO2-, CO32-, O3, RCOO- and benzene. TOK: Kekulé claimed that the inspiration for the cclic structure of benzene came from a dream. What role do the less rational ways of knowing play in the acquisition of scientific knowledge? What distinguishes a scientific from a non-scientific hypothesis in its origins or how it is tested?


Delocalisation in organic systems

When singly occupied 'p' orbitals laterally overlap (after the formation of a sigma bond of course) they create a molecular 'pi' orbital as described in section 14.2.

If a series of p orbitals on neighbouring atoms can all overlap (they must have the same orientation) then they can make a delocalised molecular orbital over all of the involved atoms. This delocalisation of the electrons effectively spreads their charge over a larger region of space lowering the energy of the system making the molecule more stable.

This situation can easily be recognised in conjugated systems i.e. systems that have alternating double and single bonds such as CH2=CH-CH=CH2 (butadiene)

Each of the carbon atoms is sp2 hybridised and has one p orbital at 90º to the plane of the sp2 orbitals. These p orbitals can overlap to make a delocalised orbital that extends over the whole molecule lowering the internal energy associated with the molecule. This cannot be measured directly but its effect can be shown in reactions when the enthalpy (energy) change due to the reaction will be affected by the molecules greater stability.


Example 1: In the hydrogenation of ethene the enthalpy change can be measured:

CH2=CH2 + H2 --> CH3CH3 ΔH = -138 kJ/mol

If butadiene is hydrogenated the enthalpy change would be expected to be equal to exactly twice this value = -272 kJ/mol

However, when measured experimentally the value is:

CH2=CH-CH=CH2 + 2H2 --> CH3CH2CH2CH3 ΔH = -240 kJ/mol

The 'missing' energy is explained by saying that the delocalisation if the electrons over the molecule in the butadiene has stabilised the molecule and consequently the energy produced on hydrogenation is less by 32 kJ/mol


Benzene

The molecule benzene is a ring structure with six carbon atoms arranged in a circle. Each of the carbons is sp2 hybridised and all of the p orbitals are oriented at 90º to the plane of the ring. They all overlap to produce a circular molecular orbital above and beneath the ring of carbon atoms.

The 'p' orbitals overlap to form the delocalised molecular 'pi' system. The six electrons are now free to rotate over the whole ring system.

Once again the effect of the delocalisation is reflected in enthalpies lower than those expected. It also results in the bonds being all of equal length, somewhere between the length of a single bond and the length of a double bond.

  • single bond length: 1,54 nm
  • double bond length: 1,34 nm
  • benzene bond length: 1,39 nm

Example 2: In the hydrogenation of cyclohexene the enthalpy change is:

C6H10 + H2 --> C6H12 ΔH = -119 kJ/mol

If benzene is hydrogenated the enthalpy change would be expected to be equal to exactly three times this value = -357 kJ/mol

However, when measured experimentally the value is:

C6H6 + 3H2 --> C6H12 ΔH = -207 kJ/mol

The 'missing' energy is again explained by the delocalisation of the electrons over the whole molecule, consequently the energy produced on hydrogenation is less by 150 kJ/mol

+ H2 --> ΔH = -119 kJ/mol

+ 3H2 --> ΔH = -207 kJ/mol



 
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