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Contents
1(a) Molecular orbitals
4
1(b) The bonding continuum
6
1(c) Hybridisation
7
3(m) Aromatic hydrocarbons
12
4(a) Absorption of visible light by organic molecules
15
4(b) Chromophores
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1(a) Molecular orbitals
Orbitals can be used to explain bonding between atoms. Atomic orbitals are
the volume of space in which the electrons are likely to be found. The atomic
orbitals containing the valence electrons (outer electrons) are the ones that
are important here.
When atomic orbitals overlap, they combine to form molecular orbitals.
In the case of two hydrogen atoms, the overlap of the two 1s atomic orbitals
results in the formation of a σ (sigma) molecular orbital. σ bonds are covalent
bonds formed between atoms when end-on overlap of orbitals occurs.
H
+
H
H
H
Figure 1 Two 1s orbitals overlapping to form a σ molecular orbital.
The resultant molecular orbital is more stable than each of the separate
atomic orbitals and contributes to the shape of the molecule. In the above
case this results in hydrogen molecules being linear, with the bonding
electrons most likely to be found between the atoms. Molecular orbitals
encompass the whole molecule and are not simply found between atoms
inside a molecule.
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The molecular orbital formed is a lower energy arrangement than the separate
atomic orbitals.
σ* anti-bonding
molecular orbital
Increasing
energy
1s
orbital
1s
orbital
σ bonding
molecular orbital
Figure 2 Comparison of the energies of atomic orbitals with molecular
orbitals.
The anti-bonding molecular orbital is a higher energy molecular orbital. It is
usually empty as higher energy orbitals are less stable and so electrons fill
the more stable lower energy state. The anti -bonding orbital is used when
electrons that are excited are promoted into it. These arrangements are often
unstable.
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1(b) The bonding continuum
The shape of the molecular orbital formed from overlapping atomic orbitals
will govern the type of intermolecular bonding that is observed.
Non-polar covalent bonds or pure covalent bonds are formed between two
atoms of the same element, or two atoms with a very low difference in
electronegativities. The molecular orbital formed from overlapping atomic
orbitals is symmetrical around a mid-point where the bonding electrons are
most likely to be found.
H
F
H
F
Figure 3 Symmetrical σ bonding orbitals of non-polar covalent bonds.
As shown above, in hydrogen and fluorine molecules, or any non-polar
covalent bond, the σ bonding orbital is symmetrical.
This is not always the case. When there is a large difference between the
electronegativities of the two elements involved in the bond, the bonding
molecular orbital will be asymmetrical.
Water molecules contain highly electronegative oxygen atoms. Because
oxygen has a greater attraction for the bonding electrons than hydrogen, the
molecular orbital formed will be asymmetrical. Figure 4 shows that because
the bonds in water molecules are polar, bonding electrons are more likely to
be found around the δ– oxygen atom.
O
H
H
Figure 4 Asymmetrical bonding orbitals exist in polar covalent bonds .
When ionic bonds form, there is extreme asymmetry and the bonding
molecular orbital is almost entirely around one atom.
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1(c) Hybridisation
In its ground state, an isolated atom of carbon has the electron arrangement
1s 2 2s 2 2p 2 .
1s
2p
2s
Figure 5 Orbital box notation for an atom of carbon in its ground state.
The orbital box notation for an unexcited atom of carb on suggests that only
two electrons are available for bonding. Clearly, carbon atoms are capable of
forming four covalent bonds thus the shapes of the atomic orbitals involved
cannot explain the bonding observed in hydrocarbon compounds such as
alkanes and alkenes.
H
H
H
C
C
H
H
H
H
C
C
H
C
H
H
H
H
Figure 6 Hydrocarbons show carbon atoms exhibiting four covalent bonds .
We can explain bonding in hydrocarbons by hybridisation. Hybrid theory
assumes that the 2s and 2p orbitals of carbon atoms can combine to form new
orbitals. In the case of alkanes such as ethane (Figure 6, above), the 2s orbital
and the three 2p orbitals of carbon combine or mix to form four degenerate
orbitals (i.e. orbitals of equal energy).
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z
z
y
y
x
x
2s orbital
2px orbital
z
z
y
y
x
x
2pz orbital
2py orbital
Figure 7 These four orbitals combine to form four new orbitals of equal
energy. We describe these hybridised orbitals as degenerate.
The hybrid orbitals found in the carbon atoms of alkane molecules are formed
from one s orbital and three p orbitals. These are known as sp 3 orbitals.
Increasing
energy
2p
sp3 hybridised orbitals
2s
Figure 8 Energy diagram showing the formation of the new orbitals.
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Figure 9 An sp 3 hybridised orbital.
The sp 3 orbitals formed are all half-filled, with the electron far more likely to
be found in the larger lobe. When drawn, the smaller lobe is often omitted.
Since electrons repel each other, the four sp 3 hybridised orbitals surrounding
a central carbon atom result in a familiar tetrahedral shape, with a maximum
possible angle between each orbital of 109.5°.
Figure 10 Four sp 3 hybridised orbitals arranged around a central carbon
atom.
In methane, all four hybrid orbitals are used to form σ bonds between the
central carbon atom and hydrogen atoms. Carbon-to-carbon single bonds in
alkanes result from overlapping sp 3 orbitals forming σ bonds. σ bonds are
covalent bonds formed by end-on overlap of two atomic orbitals and since σ
bonds must lie along the line joining both atoms, there will be free rotation
around these orbitals.
In alkenes the bonding observed is also due to hybridisation. As with alkanes,
an electron from the 2s shell is promoted to the empty 2p orbital. This results
in the formation of three hybrid orbitals, with one remaining unhybridised 2p
orbital.
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unhybridised
2p orbital
Increasing
energy
2p
sp2 hybridised
orbitals
2s
Figure 11 Energy diagram showing the formation of hybrid orbitals in
alkenes.
The hybrid orbitals formed in alkenes from one s orbital and two p orbitals
are called sp 2 orbitals. The three sp 2 orbitals repel each other, resulting in a
bond angle of 120° between them. The hybrid orbitals overlap to form σ
bonds joining their central carbon atoms to both carbon and hydrogen.
2p orbital
Three sp2 orbitals
Figure 12 An sp 2 hybridised carbon atom showing the unhybridised p orbital.
+
σ bond
Figure 13 sp 2 orbitals overlap to form σ bonds.
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The unhybridised p orbitals are perpendicula r to the plane of the molecule.
The p orbitals of the carbon atoms are parallel and close enough to overlap
sideways. This sideways overlap between the 2p orbitals produces a new
molecular orbital between the two carbon atoms. This new orbital is called a
pi (π) orbital or more commonly a π bond. A π bond is a covalent bond
formed by the sideways overlap of two parallel atomic orbitals.
Figure 14 Sideways overlap of p orbitals results in formation of a π bond.
Looking at information comparing σ and π bonds, we can see that double
bonds are stronger than single bonds, but not twice as strong. This is because
the sideways overlap (π bond) is weaker than the end-on overlap (σ bond).
Bond type
Bonding
orbitals
present
Bond length
Mean bond
enthalpy
C
C
1σ
154 pm
370 kJ mol –1
C
C
1σ+1π
134 pm
602 kJ mol –1
C
C
1σ+2π
121 pm
835 kJ mol –1
Table 1 A comparison of the length and strength of various carbon -to-carbon
covalent bonds.
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3(m) Aromatic hydrocarbons
Aromatic compounds differ to other hydrocarbons as they contain de localised
electrons. The simplest aromatic compound is benzene (C 6 H 6 ). Chemists
initially represented a molecule of benzene as shown in Figure 15 but this
structure does not accurately depict the nature of the bonds within the
molecule.
H
H
H
C
C
C
C
C
H
C
H
H
Figure 15 An early representation of benzene.
Benzene is a very stable, saturated structure that does not undergo addition
reactions. The structure in Figure 15 would suggest addition reactions were
possible and it does not show the delocalisation of electrons within the
molecule. This model also does not explain why all the bonds in benzene are
observed to be the same length. From the structure above, it could be
(wrongly) assumed that analysis of benzene would show three shorter double
bonds and three longer single bonds.
In benzene, each carbon atom has used three of its four valence electrons to
form σ bonds. The fourth electron of each carbon atom is delocalised over the
entire ring, and is not involved in forming double covalent bonds.
The σ bonding can be described as existing between six sp 2 hybridised
orbitals.
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There are six C–C σ bonds in the molecule and so each carbon atom has two
σ bonds to adjacent carbon atoms and every carbon atom also has a σ bond to
an atom of hydrogen. This results in a planar molecule with the unfilled 2p
orbital of each carbon atom above and below the plane of the mo lecule.
Figure 16 Benzene showing the unused 2p orbitals.
These unused 2p orbitals all combine to form a set of delocalised π molecular
orbitals above and below the plane of the mole cule, each containing three
electrons that are spread over the whole system.
Figure 17 Delocalised electron clouds above and below the plane of the
molecule in benzene.
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The structure of benzene is drawn as shown below to represent the
delocalised electron clouds.
Figure 18 A modern representation of benzene.
Benzene is an important feedstock for the chemical industry. While it does
not undergo addition reactions, substituted benzene rings (called phenyl
groups) are found in many medicines, antiseptics, drugs and other useful
products.
R
Figure 19 A phenyl group (C 6 H 5 ).
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4(a) Absorption of visible light by organic
molecules
While many chemical compounds are coloured because they absorb visible
light, most organic molecules appear colourless. Energy from photons is used
to promote electrons from bonding or non -bonding orbitals into higher energy
anti-bonding orbitals. Several transitions are possible. The σ* and π* antibonding orbitals are normally empty. When absorptions occur, electrons are
excited and promoted from a filled orbital (an electron in a σ or π bonding
orbital or from a lone pair in a non-bonding orbital) into a higher energy anti bonding orbital.
π* anti-bonding orbital
σ* anti-bonding orbital
Increasing
energy
non-bonding orbital
π bonding orbital
σ bonding orbital
Figure 20 The two possible transitions in a compound containing only σ
bonds.
Organic compounds that contain only σ bonds are colourless. The σ bonding
orbital is the highest occupied molecular orbital (HOMO), and the lowest
unoccupied molecular orbital (LUMO) is the σ* anti-bonding orbital. The
transition between these orbitals (as shown above) is quite large (high
energy) and corresponds to the UV part of the spectrum.
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π* anti-bonding orbital
σ* anti-bonding orbital
Increasing
energy
non-bonding orbital
π bonding orbital
σ bonding orbital
Figure 21 Four transitions are possible in a compound containing both σ
bonds and simple π bonds.
Excitations of electrons in compounds containing simple π bonds still involve
a large transition to promote an electron from HOMO ( π bonding orbital) to
LUMO (σ* anti-bonding orbital) , and thus these compounds also absorb in
the UV region of the spectrum.
Organic molecules that are coloured contain delocalised electrons spread over
a number of atoms. These molecules are known as conjugated systems.
Previously we have considered conj ugation in small molecules such as
benzene.
H
H
H
H
H
H
Figure 22 Benzene exhibits a small degree of conjugation.
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Coloured organic compounds contain a larger degree of conjugation. For
bonds to be conjugated in long carbon chains, alt ernating double and single
bonds must be present, as shown in the structure of vitamin A (Figure 23).
H3C
CH3
CH3
CH3
OH
CH3
Figure 23 Vitamin A.
Vitamin A contains a long chain of alternating σ and π bonds. The molecular
orbitals contain delocalised electrons which stretch along the length of the
conjugated system. The greater the number of atoms spanned by the
delocalised electrons, the smaller the energy gap will be between the
delocalised orbital and the next unoccupied orbital .
Exciting the delocalised electrons will therefore require less energy so light
of a lower frequency is absorbed. If this falls within the visible part of the
electromagnetic spectrum, the compounds appear coloured.
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4(b) Chromophores
A chromophore is a group of atoms within a molecule that is responsible for
the colour observed. Coloured compounds arise because visible light is
absorbed by electrons in the chromophore, which are then promoted to a
higher energy molecular orbital.
By comparing chromophores, we can find out about the energy of light that is
being absorbed.
We can explain different colours being observed in different organic
molecules by considering the length of the conjugated system. The colours we
observe are not absorbed by the molecule. If the chromophore absorbs light
of one colour, then the complementary colour is observed.
Compound
Number of C=C in
conjugated system
Main
colour
absorbed
Colour
compound
appears
Vitamin A
5
Violet
Yellow
β-carotene
11
Blue
Orange
Lycopene
11
Green
Red
Table 2 A selection of organic compounds exhibiting colour and the length of
their conjugated system.
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H3C
CH3
CH3
H3C
CH3
CH3
CH3
H3C
CH3
Figure 24 β- carotene.
Figure 25 Lycopene.
If a compound absorbs any portion of the spectrum in the visible light region,
it will exhibit an observable colour. Since violet light has higher energy than
blue or green, when it is absorbed we observe the yellow light that is
transmitted.
As molecules with greater conjugation absorb lower energy light, the greater
the degree of conjugation, the more likely the compound is to have a red
colour. Similarly, less conjugation would result in compounds appearing
yellow.
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CH3