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HL4-5

Delocalized electrons in pi bonds 14.3.1. HL4-5.ppt. What is a delocalized electron?. 14.3.1 – Describe the delocalization of pi electrons and explain how this can account for the structures of some species. Examples should include NO 3 – , NO 2 – , CO 3 2– , O 3 , RCOO – and benzene.

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HL4-5

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  1. Delocalized electrons in pi bonds 14.3.1 HL4-5.ppt

  2. What is a delocalized electron? • 14.3.1 – Describe the delocalization of pi electrons and explain how this can account for the structures of some species. Examples should include NO3–, NO2–, CO32–, O3, RCOO– and benzene. • The examples given above are all species with resonance structures. • Delocalized electrons show a tendency to be shared between more than one bonding position.

  3. Resonance Structures • Resonance occurs when more than one Lewis structure exists for a molecule • We would expect one leg of the molecule to be shorter than the other two because of its double bond. • When we measure them in that lab, that is not the case. • So what do we do when theory does not match observation?

  4. Resonance Structures • We theorize that the pi electrons have delocalized and spread themselves equally between all 3 possible bonding positions.

  5. Resonance Structures • Here you can see what is happening with the hybrid orbitals.

  6. Benzene • A classic example of π bond delocalization is found in the cyclic molecule benzene (C6H6) which consists of six carbon atoms bound together in a hexagonal ring. Each carbon has a single hydrogen atom attached to it.

  7. Benzene • The lines in this figure represent the σ bonds in benzene. • The basic ring structure is composed of σ bonds formed from overlap of sp2 hybrid orbitals on adjacent carbon atoms. • The unhybridized carbon pz orbitals project above and below the plane of the ring. They are shown here as they might appear if they did not interact with one another.

  8. Benzene • But what happens, of course, is that the lobes of these atomic orbitals meld together to form circular rings of electron density above and below the plane of the molecule. • The two of these together constitute the "second half" of the carbon-carbon double bonds in benzene.

  9. Benzene • This computer-generated plot of electron density in the benzene molecule is derived from a more rigorous theory that does not involve hybrid orbitals; the highest electron density (blue) appears around the periphery of the ring, while the lowest (red) is in the "doughnut hole" in the center.

  10. So what does it all mean? • Resonance structures with delocalized electrons show equal intermediate bond lengths

  11. So what does it all mean? • Resonance structures with delocalized electrons confer great stability to the molecule making ones like benzene quite unreactive. • This is because delocalization spreads the electrons as far apart as possible and so minimizes the repulsions between them.

  12. So what does it all mean? • When you have giant covalent molecules like the carbon allotropes with delocalized electrons spread out throughout the entire structure, they get to act like metals and conduct electricity in the solid state!

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