DUNCANRIG SECONDARY ADVANCED HIGHER CHEMISTRY

UNIT TWO

BOOKLET 1 Molecular Orbitals

and

Hybridisation

In the inorganic unit we learned about atomic orbitals and how they could be

used to write the electron configuration of atoms and ions.

Molecular orbital theory is used to describe how atomic orbitals combine when

atoms combine to produce molecules. It explains many issues with regard to

chemical bonding such as molecular shape, the existence of double and triple

bonds and why some organic molecules are coloured while others are not.

Consider two hydrogen atoms approaching each other and attempting to bond.

Molecular orbital theory states that when any two atomic orbitals meet each

other and overlap they will form two new molecular orbitals (one called bonding

and one called anti-bonding) in which a maximum of two electrons can be found.

One of these molecular orbitals will be lower in energy than the two atomic

orbitals from which it was made; the other molecular orbital lies at higher

energy.

The lower-energy MO of H2 concentrates electron density between the two

hydrogen nuclei and is called the bonding molecular orbital. This sausage

shaped MO results from summing the two atomic orbitals so that the atomic

orbitals combine in the region between the two nuclei. Because an electron in

this MO is attracted to both nuclei, the electron is more stable (it has lower

energy) than it has in the 1s atomic orbital of an isolated hydrogen atom.

Further, because this bonding MO concentrates electron density between the

nuclei, it holds the atoms together in a covalent bond.

The higher-energy MO has very little electron density between the nuclei and

is called the antibonding molecular orbital.

Instead of combining in the region between the nuclei, the atomic orbitals

cancel each other in this region, leaving the greatest electron density on

opposite sides of the nuclei. Thus, this MO excludes electrons from the very

region in which a bond must be formed. An electron in this MO is repelled from

the bonding region and is therefore less stable (it has higher energy) than it is

in the 1s orbital of a hydrogen atom

In both these types of molecular orbital the electron density is concentrated

along the internuclear axis. This type of molecular orbital is called a sigma (σ)

orbital. To distinguish between them an antibonding oribital is designated as

sigma star (σ*).

We can draw energy diagrams showing how the atomic orbitals and molecular

orbitals are related.

The energy diagram for the molecular orbitals for hydrogen is shown below.

Similar to atomic orbitals, when placing electrons into molecular orbitals, the

Aufbau principle, Hund’s rule and the Pauli exclusion principle are obeyed.

Obviously the picture gets more complicated when dealing with larger

molecules with more and more electrons.

The second-row atoms of the Periodic Table have valence 2s and 2p orbitals,

and we need to consider how they interact to form MOs.

Due to their shape and orientation, when p-orbitals are involved in forming

bonds, they can overlap and therefore interact in two ways; END-ON and

SIDE- ON.

In general END-ON overlap affords a better contact between p-orbitals and

so this leads to a lower energy (a more stable) molecular orbital than results

from a side-on overlap. Orbital formed from side-on overlap are called pi(π)

orbitals.

In the diagram above the

green p- atomic orbitals will

form an end-on sigma

molecular orbital while the

others will form side-on pi

molecular orbitals.

Molecular orbital diagrams for N2, O2, F2 and Ne2

Only the valence shell electrons are shown as these will be the electrons which are

involved in bonding.

The diagram for Ne2

conforms that this

molecule is not allowed -

Ne2 does not exist.

N2 O2

F2 Ne2

Note there is one bonding

orbital and one anti-bonding

orbital for each atomic

orbital.

Bond Order – this tells us how many bonds exist between the atoms.

Bond order = ½(no of bonding electrons – number of anti- bonding electrons )

The picture on the right shows a

representation of a methane molecule

showing the familiar tetrahedral shape

with the central carbon atom having a

valency of four.

Consider the electronic structure and the shape of the

of the atomic orbitals in the valence shell of an isolated carbon atom.

This picture is problematic.

1. How does carbon form four bonds

with only 2 half- filled p-orbitals.

2. Why does methane have a

tetrahedral geometry when none of

the oribitals have this shape.

In methane and all other alkanes these problems are solved by orbital

hybridisation.

This theory assumes that the 2s orbital and the three p orbitals mix or

hybridise to form four new hybrid orbitals known as sp3orbitals.

This can also be shown as an energy diagram.

four sp3 hybridised orbitals

The hybrid orbitals are

degenerate and have

identical shapes.

As there is only a small energy gap between the 2s and 2p orbitals it is

energetically favourable for carbon to promote a 2s electron to the empty 2p

orbital. This Carbon will now have four unpaired electrons and so have a normal

valency of four.

It is energetically favourable because the energy required to promote one

electron is more than compensated for when bonds are formed. If carbon

forms four bonds instead of two then twice as much energy will be released.

When orbitals overlap bonds are formed. In a similar way naming molecular

orbitals, if the atomic orbitals overlap “end – on” the bond is called a

SIGMA(σ)BOND.

This diagram shows an sp3 hybrid

orbital forming a sigma bond with

the 1s orbital of a hydrogen atom

In methane, all four hybrid orbitals are

used to make 4 sigma bonds with

hydrogen atoms. This gives the familiar

tetrahedral arrangement for CH4

molecules.

All other alkanes bond in a similar way using sp3 hydridisation. Sigma bonds

form between carbon atoms as well as forming between hydrogen atoms and

carbon atoms.

The bonding in ethane is shown below.

It is useful to remember

that any carbon atom

bonded to FOUR other

groups is an sp3 hybridised

carbon atom.

sp3 hybridisation does not explain the bonding

and structure of ethene.

Ethene has bond angles of 120 degrees and all

the atoms lie in the same plane – it is

definitely not tetrahedral.

Furthermore it has a carbon to carbon double bond

which, unlike single bonds, does not allow free

rotation along the internuclear axis

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 time the 2s orbital mixes with only TWO of the p orbitals forming

THREE hybrid sp2 orbitals. One of the p orbitals remains unhybridised.

The carbon atom now has four electrons to use in bonding – 3 at 120 degrees

apart (on the same plane) in the sp2 hybrid orbitals and one at 90 degrees to

these in the unhybridised p orbital.

three sp2

hybrid orbitals

one unhybridised

p orbital

Have you ever thought it strange that

when ethene reacts with bromine one

of the carbon to carbon bonds breaks

while the other does not.

The reason for this is that the two bonds are NOT the same type of bond –

one of them is much stronger than the other.

The diagram shows that when two

sp2 hybridised carbon atoms form a

bond with each other the do so by end-on

overlap – a sigma bond – the four

hydrogen atoms also form sigma bonds

with carbon the same way they

did in ethane.

Now consider the remaining bonding electron in the

unhybridised p orbital. It is clear from the position

they approach each other that they will overlap in

a side-on fashion. When this happens a pi(π) bond

will form – pi bonds are weaker than sigma bonds – Why?

In simple terms, after forming a sigma-bond (a pre-requisite for pi-

bonds), the two atoms get locked along the internuclear axis. As a result,

the orbitals available for pi-bonding can only partially overlap, thus

forming a weaker bond.

The diagram below shows how sigma bonds form along the internuclear axis

while pi bonds form above and below this axis.

It is useful to

remember that

any double bond

consists of one

sigma and one pi

bond.

Benzene, C6H6, is an aromatic compound – a compound which contains a ring of

delocalised electrons.

We now have to know exactly what this really means – the answer lies in

hybridisation.

The benzene molecule is planar with carbon to carbon

bond angles of 120 degrees.

The carbon atoms are sp2 hybridised and joined by

sigma bonds. The hydrogen atoms are as usual joined to

the carbon atoms by sigma bonds

There are six unhybridised p orbitals (one per carbon atom) and these orbitals

meet side – on forming a doughnut shaped pi bonding system above and below

the plane of carbon atoms. The six electrons will be somewhere in this pi

system – they are delocalised.

This idea is shown in many different diagrams.

While many chemical compounds are coloured because they absorb visible light,

most organic molecules appear colourless.

Remember that colour in compounds

generally arises due to the fact

that electrons in the substance

absorb certain wavelengths of light

and move to higher energy levels.

The observed colour of the

compound is the complementary

colour to the colour that is

absorbed.

Some organic molecules are coloured. Indeed, these molecules are often

present in many of the coloured substances in nature.

This molecule is responsible for the

orange/ yellow colour of some fruits and vegetables

Carrots, melons and peppers all contain carotene.

Vitamin A is a yellow coloured compound.

This vitamin is required for good eyesight.

Vitamin A deficiency can lead to an inability to see in the dark . Carotene is

converted to vitamin A by the body and this is why there is some truth to the

fact that carrots help you see in the dark.

CH3 CH3 CH3 H3C

CH3

H3C

H3C CH3

CH3 CH3

OH

CH3 CH3 CH3 H3C

CH3

Lycopene provides the red colour in fruit and vegetables.

There is some evidence that lycopene can help prevent

some forms of cancer.

Absorption of visible light by organic molecules

Why are these molecules coloured?

Energy from photons (light) 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* anti-bonding 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.

Consider the transitions shown in the diagram (The diagram is not to scale).

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. Therefore no visible

light is absorbed and the compound is colourless.

σ*

π *

Excitations of electrons in compounds containing simple π bonds, like

those shown below, 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.

All these molecules are colourless.

Consider the coloured molecules we looked at previously.

All these molecules have a chain of alternating double(π) and single (σ) bonds.

This is called CONJUGATION. The molecules shown above are conjugated.

When this happens the electrons in the conjugated part of the molecules are

DELOCALISED along the length of the conjugated chain. Delocalised electrons

LOWER the energy gap between the HOMO and the LUMO to such an extent

that compounds with enough conjugation will absorb light in the visible region

of the spectrum as electrons are promoted from HOMO to LUMO.

OH

CH3 CH3 CH3 H3C

CH3

CH3 CH3 CH3 H3C

CH3

H3C

H3C CH3

CH3 CH3

The greater the number of atoms spanned by the delocalised electrons,

the smaller the energy gap will be and the compound will absorb light of

lower energy (towards the red end of the visible part of the spectrum)

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.

The part of the molecule responsible for causing the colour is termed the

chromophore. Coloured compounds arise because visible light is absorbed by

electrons in the chromophore, which are then promoted to a higher energy

molecular orbital.

The chromophore of vitamin A is highlighted in red.

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

Many pH indicators are coloured due to conjugated electron systems.

Phenolphthalein is colourless in acid but pinky/purple in alkali. Why?

+ 2H+

No conjugated system

Big energy gap between

HOMO LUMO

Absorbs light in UV spectrum

Colourless

Conjugated system

Lower energy gap between

HOMO LUMO

Absorbs light in Visible spectrum

See complementary colour

1. A student was asked to draw a diagram to illustrate the bonding in ethene, C2H4. The

diagram drawn by the student is shown below.

Use your knowledge of chemistry

to comment on this diagram.

2. How many sigma and how many pi bonds do the following molecules have?

a. b. c. d.

3. The electronic spectra of molecules can be described in terms of the wavelength of

maximum absorbance, max.

The table below shows a number of compounds with their corresponding max values.

a. Compound 1 is buta-1,3-diene. Name compound 2.

b. Draw the most likely structure for the compound with max = 291nm.

c. The compounds shown have a system of alternation double and single bonds. What word

is used to describe this type of system?

d. Explain why compound 4 has the highest max value.

e. - carotene, max = 452 nm gives the orange colour to carrots and has the structure

whereas - carotene max = 434 nm is found in oranges and has the structure:

Explain why there is a difference in the max values for these two structures.

f. The pink colour of cooked salmon and lobster is due to astaxanthin which has the

structure

(i) Circle the chromophore in astaxanthin.

(ii) The molecule is optically active.

Circle any asymmetric carbon atom responsible for this optical activity.

4. The carbon atoms in benzene are sp2 hybridised.

a. Describe what is meant by sp2 hybridised.

b. Describe how a pi bond is formed.

c. How many electrons are found in the delocalised pi system in benzene?

5. Why is there a conflict between the electronic configuration of carbon and the

formula of carbon tetrachloride?

6. Ethanal and ethene both have sp2 hybrid orbitals.

a. Draw both molecules extended structural formulae and circle the sp2 carbon atom in

ethanal.