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7.2.2 The evidence from a comparison of enthalpies of hydrogenation

7.2.2 to 7.2.6 describe evidence of the ring structure of benzene and its presence in other aromatic compounds

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 in a trigonal planar arrangement around each carbon atom of the benzene ring.

So, it has been shown conclusively, that benzene is a planar molecule giving the 'benzene or aromatic ring' a symmetrical hexagonal shape.

In 1920s the chemist Kathleen Lonsdale proved from X-ray diffraction that all the six internal bond angles of hexachlorobenzene C₆Cl₆ and hexamethylbenzene C₆(CH₃)₆ were precisely 120^o, since the molecule was derived from benzene, it seemed illogical not to assume that benzene had the same perfect hexagonal ring of carbon atoms.

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 aliphatic alkenes.

Incidentally the triple bond in aliphatic 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.

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

The final electronic picture of benzene

The two pi orbital rings above and below the hexagonal plane of the carbon atoms is considered to be formed by the overlap of the six 2nd 2p orbitals of the carbon atoms (outer electrons 2s²2p²) i.e. from the 4th outer valency electron of the carbon atom.

For each carbon atom, three outer electrons (s + s + p) contribute to three sigma bonds, two C--C and one C-H.

The fourth electron (origin 2p orbital) is delocalised in the circular pi orbital, so in total there are three electrons in each pi orbital, six all together and non associated with a particular carbon atom.

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

An introduction to the electrophilic substitution reactions of aromatic compounds

7.2.2 to 7.2.6 describe evidence of the ring structure of benzene and its presence in other aromatic compounds

The lack of aromatic ring reactivity compared to the C=C double bond in aliphatic alkenes

Consider a series of addition reaction with solutions of the electrophiles bromine (Br₂, Br^δ+Br^δ-, weak electrophile on collision) and hydrogen bromide (strong acid, H^δ+Br^δ-, and stronger electrophile).

All of the alkene electrophilic addition reactions readily occur at room temperature and no catalyst involved either.

+  Br₂ ===>  cyclohexene gives 1,2-dibromocyclohexane

+  HBr ===> cyclohexene gives bromocyclohexane

+  2Br₂ ===> cyclohex-1,3-dene gives 1,2,3,4-tetrabromocyclohexane

or    +  Br₂  or  HBr ===> no reaction at all

+  Br₂  or  HBr ===> expect an addition reaction

Alkenes readily undergo addition reactions with electrophilic reagents like bromine, chlorine and hydrogen halides. The more localised pi electron orbitals of alkenes are much more susceptible to electrophilic attack than benzene ring of aromatic compounds (diagram on the right).

Though only part of the electrophile adds in the first step.

See 2.3 Bonding in alkenes, comparing alkane/alkene reactivity, electrophilic addition with hydrogen halides (HBr, HCl)

However, with aromatic compounds like benzene or methylbenzene, no such addition reactions occur under the same conditions i.e. room temperature, no catalyst and no raised temperature.

If benzene has a cyclo triene structure (last equation above) you would expect the rapid addition of three moles of bromine per benzene molecule as cyclohexene adds one molecule of bromine and cyclohexa-1,3-diene adds two molecules of bromine - so there must be something very different about the reactivity of 'aromatic' benzene compounds compared to 'aliphatic' alkenes.

These observations suggests that benzene compounds have some particular stability built into their structure.

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.

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 of pi electrons is destroyed - this can be done, but not easily e.g. uv light and chlorine or hydrogen using a 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 initial electrophilic reagent attack looks the same as the mechanism for addition of an electrophile to an alkene, but ...

(ii) For benzene compounds, all or part of the electrophile adds on to the benzene ring to give a saturated carbon atom (diagram above), a highly unstable intermediate due to loss of pi electron ring.

This intermediate carbocation expels a proton to give the very stable benzene ring and the substitution product - the retention of aromatic character of the delocalised pi electron ring of benzene/benzene compound.

The above diagram shows substitution in the 2 position of the benzene ring of a monosubstituted benzene compound, but you can also get substitution in the 3 and 4 positions.

In the sections describing the electrophilic substitutions of arenes, I've fully described the formation of the electrophile via the catalyst and its regeneration.

Most electrophilic substitution of benzene and methylbenzene require a catalyst, which is also a reflection of the stability of the benzene ring.

So alkenes react with electrophiles by addition, usually without a catalyst.

Aromatic benzene compounds react with electrophiles by substitution and usually need a catalyst.

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 sigma covalent 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 - also note the loss of the extra stability of the original aromatic benzene ring.

At this point 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 aromatic pi electron system of the benzene ring, which is at a lower energy than a saturated system - think of the hydrogenation data 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 stable 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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7.2.6 Can spectroscopy tell us anything about benzene and the structure of aromatic compounds?

7.2.2 to 7.2.6 describe evidence of the ring structure of benzene and its presence in other aromatic compounds

Mass spectra

The mass spectrum of benzene doesn't give that much information about its structure.

There is a characteristic peak for the m/z ion of 77 (also the most abundant  base ion peak), but this occurs in the mass spectra of many aromatic benzene compounds, particularly mono-substituted benzene compounds with the C₆H₅- group (mass = 77).

The mass spectrum of benzene

The m/z 77 ion also appears as a small peak in the mass spectrum of cyclohexene and would probably appear in the theoretical mass spectrum of a symmetrical C₆H₆ cyclotriene structure, but perhaps not in some isomeric aliphatic di-ynes mentioned below!

The mass spectrum of cyclohexene

Infrared spectra

For benzene the principal absorption band for C=C stretching vibrations is ~1480 cm^-1 and is significantly different from that observed for non-conjugated C=C bonds in aliphatic alkenes (cyclic r open chain)

e.g. for cyclohexene, the principal absorption band for C=C stretching vibrations peaks at 1640 cm^-1 and is typical for aliphatic mono-alkenes.

The peak at ~1440 cm^-1 is attributed to CH₂ vibrations, not C=C stretching vibrations with cyclohexene.

The infrared spectrum of benzene

The infrared spectrum of cyclohexene

1H NMR spectra

All protons in benzene are equivalent to each other, so you get a single and unsplit 1H NMR resonance line, implying all the hydrogen atoms are in the same chemical environment.

(But, you might 'theoretically' observe a single ¹H NMR line for a symmetrical C₆H₆ cyclotriene structure too!)

Isomeric C₆H₆ hexacyclodienes will give more than one ¹H NMR resonance line e.g. hexa-1,3-diene  has three ¹H chemical environments (1:1:2 or 2:2:4 ratio)  and hexa-1,4-diene two ¹H chemical environments (1:1 or 4: 4 ratio)

The H-1 NMR spectrum of benzene

The H-1 NMR spectrum of cyclohexene

Some aliphatic C₆H₆ isomeric structures which do exist (note an ...yne instead of an ...ane or ...ene!)

CH₃-C≡C-C≡C-CH₃ hexa-2,4-diyne has only one proton chemical environment

HC≡C-CH₂-CH₂-C≡CH hexa-1,5-diyne has two proton chemical environments

13C NMR spectra

All carbon atoms in benzene are equivalent to each other, so you get a single ¹³C NMR resonance line, implying all the hydrogen atoms are in the same chemical environment.

(But, you might 'theoretically' observe a single ¹³C NMR line for a symmetrical C₆H₆ cyclotriene structure too!)

Isomeric hexacyclodienes (C₆H₆) would give more than one ¹³C NMR resonance line.

  three ¹³C chemical environments  and two ¹³C chemical environments

The C-13 NMR spectrum of benzene

The C-13 NMR spectrum of cyclohexene

Some aliphatic C₆H₆ isomeric structures which do exist

CH₃-C≡C-C≡C-CH₃ hexa-2,4-diyne has three ¹³C NMR chemical environments

HC≡C-CH₂-CH₂-C≡CH hexa-1,5-diyne also has three ¹³C NMR chemical environments

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