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chemistry Unit 14 learning aim B
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chemistry Unit 14 learning aim B
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Unit 14B: Understand the reactions and properties of aromatic compounds.Aromatic ring chemistry for designer chemicals
P2: Explain the structure of benzene using sigma and pi bonding,
providing evidence for the structure.
At room temperature, benzene is a colourless or light-yellow liquid
compound that is also very combustible. The molecular formula of benzene is
C6H6. It has an empirical formula of CH, a molecular mass of 78, and a
molecular formula of CH. When a little amount of energy is applied to
carbon, the s orbital can be lifted to a p orbital, resulting in 1s22s12p3 in the
sp3 hybridisation. In the sp3 hybridisation, the four orbitals hybridise with
each other to produce four new orbitals that are all equivalent. The benzene
structure is exceedingly stable and consistently structured, as evidenced by
infrared, x-ray, and thermochemical analyses. When the x-ray evidence was
utilised to measure the bond lengths of the benzene, which was 0.139nm, it
was proven. This is an example of Kekule being inaccurate since the bonds
are all the same length. The thermochemical data showed that benzene has a
higher stability than Kekule's model. Kekule indicated that the enthalpy of
formation was +252kjmol-1, but when measured, it was found to be
+82kjmol-1, a significant discrepancy. Because benzene lacks a core of
symmetry, the infrared data indicated that Kekule's regular hexagonal
structure is incorrect. Benzene is planar, cyclic, and possesses alternating
double and single bonds, according to Kekule. However, because benzene's
structure does not undergo electrophilic addition, there is no c=c bond. This
disproves Kekule's theory that the carbon-carbon bond lengths are
alternating, as the benzene has equal lengths of bonds due to only having
single c-c bonds in the structure. The c=c bonds are shorter and would show
Kekule accurate if they were there, but they aren't. The molecular formula of
benzene is C6H6. It has an empirical formula of CH, a molecular mass of 78,
and a molecular formula of CH. When a little amount of energy is applied to
carbon, the s orbital can be lifted to a p orbital, resulting in 1s22s12p3 in the
sp3 hybridisation. In the sp3 hybridisation, the four orbitals hybridise with
each other to produce four new orbitals that are all equivalent.
Evidence 1, X-ray crystallography studies of bond lengths:
The method of crystallography is used to determine the length of
covalent connections. The bond length of carbon - carbon single bonds
is 0.154 micrometres. Double bonds between carbon atoms have a
bond length of 0.134 micrometres, which is somewhat less than single
, bonds. The bond lengths in benzene were measured via
crystallography. The bond was 0.14 micrometres long. This indicated
that benzene lacked both single and double carbon bonds.
Evidence 2, Infrared spectroscopy:
Another technique used to study the bonds inside a molecule is
infrared spectroscopy. Each link in a molecule stores energy, and each
degree of energy is distinct. The sample is surrounded by
electromagnetic radiation. The level of each sort of electromagnetic
wave is likewise distinct. When waves flow through a sample without
interruption and collide with a bond of the same energy level, the
bond absorbs the energy. The sample will vibrate because its energy
levels have grown, and the wave will have less energy because the bulk
of it has been absorbed. The infrared spectrums of benzene did not
show the typical single carbon bond energy absorption, but neither did
they show the standard double carbon bond energy absorption.
Evidence 3, Thermochemical:
The computed enthalpy of production from its component elements of
carbon and hydrogen in their standard states for one mole of gaseous
benzene with Kekule's structure is +252KJmol-1. When the actual
enthalpy of formation was measured, it was found to be far lower, at
just +82 kjmol-1. This suggests that benzene's structure is substantially
more stable than Kekule's model.
Sigma and pi bonds:
Covalent bonds are created when orbitals overlap. For example, when
the sp2 from each carbon overlaps to create a single c-c bond, a sigma
bond is formed. When the two 2p orbitals overlap to form a second
bond, it is called a pi bond. This causes an increase in the planar
arrangement around the c=c bonds. Every carbon atom in the benzene
structure will have one electron in the atomic p orbitals. This is due to
the fact that the carbon atoms employ three of their electros to make
three sigma orbitals bonds with the atoms nearby. Negative charges
will build above and below the plane, as seen in the diagram below. As
a result, a delocalized pi electron system will be formed.
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