# Delocalised benzene model

Watch
Announcements

Page 1 of 1

Go to first unread

Skip to page:

Right so carbon has 4 outer electrons of which it uses 3 to bond to 2 carbon atoms and 1 hydrogen. That leaves one electron in an orbital for the sideways overlap of orbitals to form a ring of electron density. I get all that . But I don’t understand why that electron that’s left is in a p orbital . Surely it should be s orbital because the bonding takes 2 electrons from 2p and 1 from 2s? Any help appreciated😁

0

reply

Report

#2

I can't think of a good way to explain this in A-level terms, but basically there's a thing called the

We can model bonding by saying that atomic orbitals of a particular atom can interact with each other to form hybrid orbitals. n atomic orbitals overlap to form n hybrid atomic orbitals. E.g. the 2s and three 2p orbitals can interact to form four sp3 hybrid atomic orbitals. These each point at 109.47 degrees to each other and lead to tetrahedral geometries.

You can also get one 2s and two 2p overlapping to form 3 sp2 hybrid atomic orbitals, or one 2s and one 2p to form two sp hybrid atomic orbitals.

The carbons in benzene are sp2 hybridised, so they have three sp2 orbitals (at 120 degrees, in the plane of the ring) and one p atomic orbital (above/below plane of ring).

The sp2 orbitals are lower in energy than the p-orbital and have the correct symmetry to overlap with each other and with H 1s, so they form the bonds C-C and C-H sigma bonds (we call this the sigma framework). That leaves one electron in each 2p orbital. The 2p orbitals all overlap to form a set of delocalised molecular orbitals which are occupied by six electrons in total.

N.B. hybridisation model is just a model to help us rationalise bonding. For a more theoretically "correct" description, you need to invoke molecular orbital theory, which uses detailed quantum mechanical calculations. This is beyond the scope of this answer.

I hope I helped, and possibly increased your appetite for more advanced chemistry study!

**hybridisation model**.We can model bonding by saying that atomic orbitals of a particular atom can interact with each other to form hybrid orbitals. n atomic orbitals overlap to form n hybrid atomic orbitals. E.g. the 2s and three 2p orbitals can interact to form four sp3 hybrid atomic orbitals. These each point at 109.47 degrees to each other and lead to tetrahedral geometries.

You can also get one 2s and two 2p overlapping to form 3 sp2 hybrid atomic orbitals, or one 2s and one 2p to form two sp hybrid atomic orbitals.

The carbons in benzene are sp2 hybridised, so they have three sp2 orbitals (at 120 degrees, in the plane of the ring) and one p atomic orbital (above/below plane of ring).

*(e.g. "less repulsion").***There's one electron in each orbital since this is the lowest energy state**The sp2 orbitals are lower in energy than the p-orbital and have the correct symmetry to overlap with each other and with H 1s, so they form the bonds C-C and C-H sigma bonds (we call this the sigma framework). That leaves one electron in each 2p orbital. The 2p orbitals all overlap to form a set of delocalised molecular orbitals which are occupied by six electrons in total.

N.B. hybridisation model is just a model to help us rationalise bonding. For a more theoretically "correct" description, you need to invoke molecular orbital theory, which uses detailed quantum mechanical calculations. This is beyond the scope of this answer.

I hope I helped, and possibly increased your appetite for more advanced chemistry study!

(Original post by

Right so carbon has 4 outer electrons of which it uses 3 to bond to 2 carbon atoms and 1 hydrogen. That leaves one electron in an orbital for the sideways overlap of orbitals to form a ring of electron density. I get all that . But I don’t understand why that electron that’s left is in a p orbital . Surely it should be s orbital because the bonding takes 2 electrons from 2p and 1 from 2s? Any help appreciated😁

**avacados1234**)Right so carbon has 4 outer electrons of which it uses 3 to bond to 2 carbon atoms and 1 hydrogen. That leaves one electron in an orbital for the sideways overlap of orbitals to form a ring of electron density. I get all that . But I don’t understand why that electron that’s left is in a p orbital . Surely it should be s orbital because the bonding takes 2 electrons from 2p and 1 from 2s? Any help appreciated😁

Last edited by K-Man_PhysCheM; 1 year ago

2

reply

(Original post by

I can't think of a good way to explain this in A-level terms, but basically there's a thing called the

We can model bonding by saying that atomic orbitals of a particular atom can interact with each other to form hybrid orbitals. n atomic orbitals overlap to form n hybrid atomic orbitals. E.g. the 2s and three 2p orbitals can interact to form four sp3 hybrid atomic orbitals. These each point at 109.47 degrees to each other and lead to tetrahedral bonding.

The carbons in benzene are sp2 hybridised, so they have three sp2 orbitals (at 120 degrees, in the plane of the ring) and one p atomic orbital (above/below plane of ring). There's one electron in each orbital since this is the lowest energy state (e.g. "less repulsion").

The sp2 orbitals are lower in energy than the p-orbital and have the correct symmetry to overlap with each other and with H 1s, so they form the bonds C-C and C-H sigma bonds (we call this the sigma framework). That leaves one electron in each 2p orbital. The 2p orbitals all overlap to form a set of delocalised molecular orbitals which are occupied by six electrons in total.

N.B. hybridisation model is just a model to help us rationalise bonding. For a more theoretically "correct" description, you need to invoke molecular orbital theory, which uses detailed quantum mechanical calculations. This is beyond the scope of this answer.

I hope I helped, and possibly increased your appetite for more advanced chemistry study!

**K-Man_PhysCheM**)I can't think of a good way to explain this in A-level terms, but basically there's a thing called the

**hybridisation model**.We can model bonding by saying that atomic orbitals of a particular atom can interact with each other to form hybrid orbitals. n atomic orbitals overlap to form n hybrid atomic orbitals. E.g. the 2s and three 2p orbitals can interact to form four sp3 hybrid atomic orbitals. These each point at 109.47 degrees to each other and lead to tetrahedral bonding.

The carbons in benzene are sp2 hybridised, so they have three sp2 orbitals (at 120 degrees, in the plane of the ring) and one p atomic orbital (above/below plane of ring). There's one electron in each orbital since this is the lowest energy state (e.g. "less repulsion").

The sp2 orbitals are lower in energy than the p-orbital and have the correct symmetry to overlap with each other and with H 1s, so they form the bonds C-C and C-H sigma bonds (we call this the sigma framework). That leaves one electron in each 2p orbital. The 2p orbitals all overlap to form a set of delocalised molecular orbitals which are occupied by six electrons in total.

N.B. hybridisation model is just a model to help us rationalise bonding. For a more theoretically "correct" description, you need to invoke molecular orbital theory, which uses detailed quantum mechanical calculations. This is beyond the scope of this answer.

I hope I helped, and possibly increased your appetite for more advanced chemistry study!

0

reply

Report

#4

(Original post by

Thanks for alll that😁😁. Took a while to get into my head but I sort of understand it so that’s good. I’d been thinking of this question for months now but just never asked anyone cos I thought I was being dumb somehow but clearly not. Thanks for all ur help

**avacados1234**)Thanks for alll that😁😁. Took a while to get into my head but I sort of understand it so that’s good. I’d been thinking of this question for months now but just never asked anyone cos I thought I was being dumb somehow but clearly not. Thanks for all ur help

*A lot of what you learn in chemistry GCSE/A-level are quite highly simplified models of what we now think might be going on, but the models you learn at each stage of education work well to rationalise most/all of the bonding that you'll meet at that level. If you study chemistry further, you'll find some more unusual types of bonding.

0

reply

(Original post by

Ahaha no worries! Maybe someone else will come along and give a simple A-level explanation lol, I can't think of one now because I've forgotten what you learn at A-level*. But yeah, the basic idea is that you have four orbitals and there's one electron in each because this is the minimum energy configuration for them in this case. Electrons really like minimum energy / minimum repulsion configurations!!

*A lot of what you learn in chemistry GCSE/A-level are quite highly simplified models of what we now think might be going on, but the models you learn at each stage of education work well to rationalise most/all of the bonding that you'll meet at that level. If you study chemistry further, you'll find some more unusual types of bonding.

**K-Man_PhysCheM**)Ahaha no worries! Maybe someone else will come along and give a simple A-level explanation lol, I can't think of one now because I've forgotten what you learn at A-level*. But yeah, the basic idea is that you have four orbitals and there's one electron in each because this is the minimum energy configuration for them in this case. Electrons really like minimum energy / minimum repulsion configurations!!

*A lot of what you learn in chemistry GCSE/A-level are quite highly simplified models of what we now think might be going on, but the models you learn at each stage of education work well to rationalise most/all of the bonding that you'll meet at that level. If you study chemistry further, you'll find some more unusual types of bonding.

1

reply

X

Page 1 of 1

Go to first unread

Skip to page:

### Quick Reply

Back

to top

to top