Benzene is a marvellous molecule.
C6H6 is a bonding and structural puzzle.
Try building a structural or display formula for C6H6.
You will have to produce a structure in which there are multiple carbon carbon bonds.
Your molecule will be unsaturated.
And therein lies our problem.
Because benzene's chemical properties do not suggest it is unsaturated.
Simply, it will not decolorize red bromine water or acidified purple potassium manganate (VII)!!
Then again take its molecular structure.
If it is a linear unsaturated molecule then some carbon-carbon bonds will be longer than others since
double bonds are shorter than single bonds.
Augustus Kekule struggled with the structure of benzene in the mid 19th century.
He could not see how the molecule could be saturated C6H6.
This is the quotation where Kekule describes how he came to see for the first time what benzene's structure could be:
I was sitting writing at my textbook but the work did not progress; my thoughts were elsewhere. I turned my chair to the fire and dozed. Again the atoms were gambolling before my eyes. This time the smaller groups kept modestly in the background. My mental eye, rendered more acute by the repeated visions of the kind, could now distinguish larger structures of manifold confirmation: long rows, sometimes more closely fitted together all twining and twisting in snake like motion. But look! What was that? One of the snakes had seized hold of its own tail, and the form whirled mockingly before my eyes. As if by a flash of lightning I awoke; and this time also I spent the rest of the night in working out the rest of the hypothesis. Let us learn to dream, gentlemen, then perhaps we shall find the truth... But let us beware of publishing our dreams till they have been tested by waking understanding.
What Kekule had realised in his dozing dream was that the molecule of C6H6 was not linear but circular.
"One of the snakes seized its own tail."
This is one of those cases of serendipity in scientific discovery.
This flash of brilliance did not fully resolve the problem of benzene's structure but it went a good deal of the way to doing so.
Now it was possible to suggest that the structure of benzene could be this:
without the snake you understand!!
But as you can see, this does not adequately solve the problem of benzene because we still have double and single bonds in an unsaturated structure for a molecule that does not behave as if it were unsaturated.
Kekule's suggestion is not much of an improvement, hailed at the time for its brilliance though it was.
Now, if we look at hydrogenation of benzene to cyclohexane and compare its value with that of a related molecule cyclohexene we see these results:
:
If we were to think of a molecule of benzene with three distinct carbon-carbon double bonds then what would the enthalpy of hydrogenation be?
We would expect a value of around three times that for cyclohexene: say -360 kJ/mol.
Why do we not see that value but measure the enthalpy of hydrogenation to be only -208kJ/mol?
If less energy is released then the structure of benzene requires more energy to break apart its carbon carbon bonding.
That difference is 152kJ/mol, the difference between the hydrogenation energy for a hypothetical Kekule structure and the hydrogenation energy of the actual benzene molecule.
Consider the diagram below: (there is a mistake in this diagram, the end product of hydrogenation is always cyclohexAne not cyclohexene)
with this diagram which includes the hypothetical Kekule structure:
You can see in the middle the hypothetical Kekule structure called here 1, 3, 5-cyclohexatriene.
(The hydrogenation energies are given in kJ/mol and in kcal/mol for our American friends!!)
So how do we now account for the fact that the bonding between carbon atoms in the benzene ring structure is much stronger than expected, requiring 152kJ/mol more energy to disrupt than expected?
We need to use the Molecular Orbital Theory for ethene as a template here.
Remember in ethene the model for the arrangement of three bonds around the carbon atoms a double and two single bonds is described as being sp2 hybridised.
Well a similar model is applied to the situation involving benzene.
Energy promotes an electron from the ground state in each of the six carbon atoms in benzene.
The resultant atomic orbitals in each carbon atom then hybridise to form three sp2 hybrid orbitals and leave one remaining p orbital
This model allows then for the formation of three sigma bonds between the six carbon atoms and the six hydrogen atoms.
But what about the p orbital?
Here is the genius of the model for benzene.
The six p orbitals overlap side-on to form a ring molecular orbital called a π bond above and below the ring of six carbon atoms.
In this ring there are six electrons and they are free to move: they are said to be delocalised.
The π bond is sometime referred to as a delocalised system. The diagram below helps to visualise this new and original situation.
This diagram just shows the π bond, the σ bonds have been rendered as lines for clarity.
Now the genius of the model is to suggest that in real benzene molecules something like this actually happens and a π bond of this kind requires that extra 152 kJ/mol to disrupt it before benzene will react with hydrogen.
The 152kJ/mol is called the delocalisation energy.
Here are the representations of the benzene structure:
The structure of benzene has since been confirmed to be a perfect hexagonal ring.
Note that all six Carbon Carbon bonds are of equal length as the model predicts: each is 140pm.
The Carbon Carbon Carbon internal bond angle is 120o as the model predicts.
The molecule is planar in shape and is represented by the circle in the hexagon although most courses and exam authorities accept the Kekule structures as well in equations and mechanisms.
Lastly, and beautifully, if the π bond has six delocalised moving electrons then these six electrons moving ought to set up a magnetic field.
The ring current should do this and this is actually what is observed and has real consequences for measurements of benzene type molecules in nuclear magnetic resonance studies.
We have Kekule to thank for starting the ball rolling on benzene and he is rightly celebrated as one more famous Belgian.
Here is how he has sometime been seen:
Pages on the "Mole" and "Using the Mole" in chemical calculations are here
C6H6 is a bonding and structural puzzle.
Try building a structural or display formula for C6H6.
You will have to produce a structure in which there are multiple carbon carbon bonds.
Your molecule will be unsaturated.
And therein lies our problem.
Because benzene's chemical properties do not suggest it is unsaturated.
Simply, it will not decolorize red bromine water or acidified purple potassium manganate (VII)!!
Then again take its molecular structure.
If it is a linear unsaturated molecule then some carbon-carbon bonds will be longer than others since
double bonds are shorter than single bonds.
Augustus Kekule struggled with the structure of benzene in the mid 19th century.
He could not see how the molecule could be saturated C6H6.
This is the quotation where Kekule describes how he came to see for the first time what benzene's structure could be:
I was sitting writing at my textbook but the work did not progress; my thoughts were elsewhere. I turned my chair to the fire and dozed. Again the atoms were gambolling before my eyes. This time the smaller groups kept modestly in the background. My mental eye, rendered more acute by the repeated visions of the kind, could now distinguish larger structures of manifold confirmation: long rows, sometimes more closely fitted together all twining and twisting in snake like motion. But look! What was that? One of the snakes had seized hold of its own tail, and the form whirled mockingly before my eyes. As if by a flash of lightning I awoke; and this time also I spent the rest of the night in working out the rest of the hypothesis. Let us learn to dream, gentlemen, then perhaps we shall find the truth... But let us beware of publishing our dreams till they have been tested by waking understanding.
What Kekule had realised in his dozing dream was that the molecule of C6H6 was not linear but circular.
"One of the snakes seized its own tail."
This is one of those cases of serendipity in scientific discovery.
This flash of brilliance did not fully resolve the problem of benzene's structure but it went a good deal of the way to doing so.
Now it was possible to suggest that the structure of benzene could be this:
But as you can see, this does not adequately solve the problem of benzene because we still have double and single bonds in an unsaturated structure for a molecule that does not behave as if it were unsaturated.
Kekule's suggestion is not much of an improvement, hailed at the time for its brilliance though it was.
Now, if we look at hydrogenation of benzene to cyclohexane and compare its value with that of a related molecule cyclohexene we see these results:
If we were to think of a molecule of benzene with three distinct carbon-carbon double bonds then what would the enthalpy of hydrogenation be?
We would expect a value of around three times that for cyclohexene: say -360 kJ/mol.
Why do we not see that value but measure the enthalpy of hydrogenation to be only -208kJ/mol?
If less energy is released then the structure of benzene requires more energy to break apart its carbon carbon bonding.
That difference is 152kJ/mol, the difference between the hydrogenation energy for a hypothetical Kekule structure and the hydrogenation energy of the actual benzene molecule.
Consider the diagram below: (there is a mistake in this diagram, the end product of hydrogenation is always cyclohexAne not cyclohexene)
with this diagram which includes the hypothetical Kekule structure:
(The hydrogenation energies are given in kJ/mol and in kcal/mol for our American friends!!)
So how do we now account for the fact that the bonding between carbon atoms in the benzene ring structure is much stronger than expected, requiring 152kJ/mol more energy to disrupt than expected?
We need to use the Molecular Orbital Theory for ethene as a template here.
Remember in ethene the model for the arrangement of three bonds around the carbon atoms a double and two single bonds is described as being sp2 hybridised.
Well a similar model is applied to the situation involving benzene.
Energy promotes an electron from the ground state in each of the six carbon atoms in benzene.
The resultant atomic orbitals in each carbon atom then hybridise to form three sp2 hybrid orbitals and leave one remaining p orbital
This model allows then for the formation of three sigma bonds between the six carbon atoms and the six hydrogen atoms.
But what about the p orbital?
Here is the genius of the model for benzene.
The six p orbitals overlap side-on to form a ring molecular orbital called a π bond above and below the ring of six carbon atoms.
In this ring there are six electrons and they are free to move: they are said to be delocalised.
The π bond is sometime referred to as a delocalised system. The diagram below helps to visualise this new and original situation.
This diagram just shows the π bond, the σ bonds have been rendered as lines for clarity.
Now the genius of the model is to suggest that in real benzene molecules something like this actually happens and a π bond of this kind requires that extra 152 kJ/mol to disrupt it before benzene will react with hydrogen.
The 152kJ/mol is called the delocalisation energy.
Here are the representations of the benzene structure:
The structure of benzene has since been confirmed to be a perfect hexagonal ring.
Note that all six Carbon Carbon bonds are of equal length as the model predicts: each is 140pm.
The Carbon Carbon Carbon internal bond angle is 120o as the model predicts.
The molecule is planar in shape and is represented by the circle in the hexagon although most courses and exam authorities accept the Kekule structures as well in equations and mechanisms.
Lastly, and beautifully, if the π bond has six delocalised moving electrons then these six electrons moving ought to set up a magnetic field.
The ring current should do this and this is actually what is observed and has real consequences for measurements of benzene type molecules in nuclear magnetic resonance studies.
Here is how he has sometime been seen:
Pages on the "Mole" and "Using the Mole" in chemical calculations are here
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