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7.2.2 The evidence from a comparison of enthalpies of hydrogenation

Below are five hydrogenation equations, together with the enthalpies of hydrogenation.

I've used the latest enthalpy values from https://webbook.nist.gov/chemistry/

The 5th equation assumes that benzene really is a 'triene', so we will show in sections 7.2.2 and 7.2.3 that this is not the real structure of benzene.

 

hydrogenation of cyclohexene to cyclohexane

ΔH_{hydrogenation}(cyclohexene) = -118 kJ mol^-1

This is a typical value for the hydrogenation of an alkene with one C=C bond.

 

Hydrogenation of a 1,3-cyclohexadiene to cyclohexane

ΔH_{hydrogenation}(cyclohexa-1,3-diene) = -229 kJ mol^-1

This is just below the expected value for hydrogenating two C=C bonds (2 x -118 = -236 kJ mol^-1).

 

Hydrogenation of a 1,4-cyclohexadiene to cyclohexane

ΔH_{hydrogenation}(cyclohexa-1,4-diene/) = -233 kJ mol^-1

This is just below the expected value for hydrogenating two C=C bonds (2 x -118 = -236 kJ mol^-1).

 

hydrogenation of the 'real' benzene to cyclohexane

ΔH_{hydrogenation}(benzene) = -208 kJ mol^-1

This is well below the expected value for hydrogenating three C=C bonds (3 x -118 = -354 kJ mol^-1).

 

hydrogenation of methylbenzene to methylcyclohexane

ΔH_{hydrogenation}(methylbenzene) = -205 kJ mol^-1

This is considerable below the expected value for hydrogenating three C=C bonds (3 x -118 = -354 kJ mol^-1).

It should be noticed that this 'addition' reaction requires a heated nickel catalyst because of the stability of the benzene ring.

The hydrogen molecules are adsorbed on the Ni surface and split into atoms.

These atoms are effectively very reactive hydrogen radicals (H•) which can break open the pi bond system and form new C-H bonds until the saturated cycloalkane is formed.

 

The enthalpy of hydrogenation for this is theoretically about -354 kJ mol^-1.

If we assume (incorrectly) that benzene has a cyclotriene structure with alternating single and double bonds.

 

Discussion of the above enthalpy of hydrogenation values:

Apart from the first and simplest alkene, ethene, most enthalpies of hydrogenation are about -120±6 kJ mol^-1 (for 1 mole H₂ per mole alkene) and in the case of dienes about double that for two moles of hydrogen per diene, which is what you might reasonably expect.

The special case of benzene - the aromatic ring structure of arenes - aromaticity

However, on the basis of these trends, the expected value for the complete hydrogenation of benzene and other aromatic compounds with a single benzene ring would be around 3 x -118 = ~-354 kJ mol-1, but not so!

In fact the energy released on hydrogenating benzene (208 kJ/mole) is even less than hydrogenating a diene!

So, something must be different about benzene but it can be explained with the enthalpy level diagram shown below and an examination of possible molecular structures.

Also note the comparison of equations 8. and 10. from above.

The 'top' molecule in the diagram shows the theoretical structure of a triene with the same molecular formula of benzene (C₆H₆) and, if it existed in this form it would be called cyclohexa-1,3,5-triene or 1,3,5-cyclohexatriene.

This molecular structure assumes there are simple alternate single (C-C) and double (C=C) carbon-carbon bonds.

BUT, according to the actual thermochemical data calculated and derived from e.g. enthalpies of hydrogenation, benzene is already more stable by ~149 kJ mol^-1, so, whatever its structure, it cannot have this 'triene' structure.

This lowering of the potential energy of the benzene molecules is referred to as the resonance stabilisation energy.

The concept of resonance structures is discussed in section 7.2.4

What happens in reality, is that the equivalent of 3 double bonds (C=C) and 3 single bonds (C-C) 'merge' to form a six equal bonds each involve an average of 3 shared electrons.

Two of the electrons for each bond are concentrated between the two carbon atoms OR between a carbon atom and hydrogen atom, both equivalent to a single bond known as a sigma (σ) bond.

The 4th electron per carbon atom is located in a ring orbital, of two sections, above and below the plane of the hexagonal ring of carbon atoms.

These are known as pi (π) orbitals and each one contains 3 pi (π) electrons which are delocalised around the ring.

This is indicated by the ring in the centre of the ring either in skeletal formula or structural formula and is a symmetrical planar hexagonal molecule.

Whenever charge is delocalised or 'spread out' the potential energy of the system is lowered and in the case of benzene about 149 kJ per mole compared to the triene structure of alternate single C-C and double C=C bonds.

 

Apart from the thermochemical evidence argued above, X-ray crystallography has shown:

The actual bond length measurements show that there are six carbon - carbon bonds all of equal length and intermediate between single and double bonds.

The electron density map shows a symmetrical distribution of the electron clouds that fitted a symmetrical hexagonal planar ring of six carbon atoms.

BUT, the electron density map also showed significant electron density in a ring above and below the plane of the ring of 6 carbon atoms.

X-ray crystallography also showed all the bond angles are 120^o, that is C-C-C  or  C-C-H due to 3 groups of bonding electrons around each carbon atom of the benzene ring

It was also shown that benzene is a planar molecule giving the 'benzene or aromatic ring' a symmetrical hexagonal shape.

Summary of types of carbon - carbon bonds.

σ = single sigma bond; π = delocalised pi bond system

There term bond order refers to the 'electronsworth' of the bond i.e. the average number of shared electrons involved in that particular bond.

Typical bond lengths: single bond C-C is 0.154 nm (bond order 1) e.g. in typical alkanes.

An aromatic ring bond is 0.139 nm in length (bond order 1.5) e.g. in arenes like benzene and methylbenzene. X-ray diffraction has shown this bond is shorter than a single C-C bond, but not as short as an alkene C=C bond.

A double bond C=C is 0.134 nm in length (bond order 2)  e.g. in typical alkenes.

Incidentally the triple bond in alkynes has a typical length of 0.120 nm in length (C≡C, bond order 3).

There is a clear trend of decreasing bond length with increasing in bond order and increasing bond strength shown by the increasing bond enthalpy.

How do not always think a high bond strength is always associated with low reactivity e.g. compare the reactivity of alkanes and alkenes.

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7.2.5 Influence of the stability of the benzene ring on the chemistry of aromatic compounds

In chemistry, aromaticity is a property of cyclic, planar structures with pi bonds in resonance that gives increased stability compared to other geometric or connective arrangements with the same set of atoms, bonding with the same valencies.

Aromatic rings e.g. the benzene ring are very stable and do not readily change in a completed reaction.

The hexagonal aromatic ring of carbon atoms e.g. in arenes like benzene and methyl benzene tends to undergo electrophilic substitution reactions, rather than electrophilic addition reactions like alkenes.

It is the very great stability of the benzene ring in aromatic compounds means that, unlike alkenes, they are reluctant to undergo addition reactions, and this is best understood by considering the basics of electrophilic substitution reactions undergone by aromatic compounds..

Addition of a pair of atoms to a benzene ring means the stable aromatic ring is destroyed - this can be done, but not easily e.g. uv light and chlorine or hydrogen, catalyst and a high temperature.

Therefore aromatic compounds tend to undergo electrophilic substitution reactions involving the generation of a powerful electrophile (electron pair acceptor) which subsequently attacks the electron rich π (pi) electron system of the double bond.

In other words, arenes like benzene and methylbenzene tend to undergo substitution, rather than addition, because substitution in the ring allows the very stable benzene ring to remain intact.

With the model portrayed above we can now consider the attack of an electrophile on the pi bond electrons.

The general mechanism for electrophilic substitution is shown above, though the source and/or generation of the electrophile is not shown, so it is often, but not always, a 3-step mechanism.

The electrophile E⁺ is an electron pair acceptor.

The benzene ring is capable of donating a pair of electrons from the pi orbitals of the aromatic ring to the incoming electrophile.

The electrophile E⁺ attacks the pi orbitals (the electron rich clouds - high electron density) above and below the plane of the benzene ring.

A C-E bond is formed and yielding a carbocation - but notice that the delocalised system only extends across five of the six carbon atoms of the benzene ring - you must draw this accurately in exams.

It is effectively an electrophilic addition, just like with alkenes, but here the similarity ends.

Instead of a further addition of a negative ion, as with electrophilic addition to alkenes, the mechanistic pathway follows a course that reforms the very stable delocalised electron system of the benzene ring, which is at a lower energy than a saturated system - think of the hydrogenation data and argument.

Since, in the carbocation, the stable delocalised system of the benzene ring is broken, a proton (H⁺) is expelled and the associated C-H bond pair of electrons rejoin the delocalised system, so, the complete benzene ring of pi orbitals is re-formed (aromaticity restored), yielding the stable substituted aromatic molecule.

e.g. a substituted benzene molecule C₆H₅E  or  a substituted methylbenzene molecule EC₆H₄CH₃.

In the above diagram for methylbenzene I've just assumed that the substituent is on carbon atom 2 in aromatic nomenclature, but other positions are available!

Overall the mechanism  =  an electrophilic addition  +  an elimination (expulsion)  ===>  substitution

Technically, the final step of this electrophilic substitution mechanism involves the expelled proton adding to a base (not shown in the above diagram)

The base is an electron donor, so one of two things can happen.

H⁺  +  :B  ===>  HB^- (anion formed)  OR  H⁺  +  :B^-  ==> HB (neutral molecule formed).

The diagram above shows a general reaction progress profile for the electrophilic substitution mechanism for aromatic compounds e.g. arenes like benzene and methyl benzene.

I've omitted the source and formation of the electrophile and just used benzene for simplicity.

The dip is decreases in potential energy for the formation of the carbocation.

E_a1 = the much larger initial activation energy for the attacking electrophile to form a C-E bond from benzene's pi electrons.

E_a2 = the much smaller activation energy for the expulsion of a proton (that will combine with a base) and the re-formation of the stable aromatic ring to complete the substitution reaction.

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