Lesson 2 bonding and reactions of alkanes

Advanced Higher Chemistry
Unit 3
Bonding and reactions of alkanes
Alkanes
 General Formula CnH2n+2
 Saturated hydrocarbons


Each carbon atom forms 4 single bonds.
Each carbon atom has 4 bonding pairs of
electrons.
 The 4 bonding pairs of electrons repel one
another resulting in a tetrahedral
arrangement.
Bonding in Alkanes
Definitions :Covalent Bond
Two half filled atomic orbitals from two different
atoms overlap and attract the nuclei of each
atom.
This gives rise to a MOLECULAR ORBITAL.
Molecular Orbital
Orbital of 2 bonding electrons moving under the
influence of 2 nuclei.
 Electron arrangement of a ground state carbon
atom
1s2 2s2 2p2
 How many unpaired electrons are there in a
carbon atom (Remember Hund’s Rule)?
Answer - There are 2 unpaired electrons in a
carbon atom (the two electrons in the p
orbital). Carbon should therefore only form 2
covalent bonds.
 How many covalent bonds does carbon
usually form?
Answer - carbon usually forms 4 covalent bonds.
WHY?
Hund’s Rule of
Maximum Multiplicity
“When electrons occupy degenerate orbitals, the
electrons fill each orbital singly, keeping their
spins parallel before spin pairing occurs”
Hybrid Orbitals
A 2s electron is
promoted to the
third 2p orbital
Carbon atom now has 4
singly occupied orbitals that
‘mix’ to form 4 hybrid
orbitals of equal energy.
2s Orbital
Three 2p orbitals (px, py and pz)
Hybridisation
Four sp3 hybrid orbitals
 Hybridisation is possible as the 2s and 2p
sublevels are close in energy.
 Hybridised orbitals are more directional in
shape than unhybridised orbitals, this
provides better overlap when bonds form.
 In diagrams the small lobe of the hybridized
orbital is usually omitted.
 The sp3 hybridised carbon atom is often
referred to as a tetrahedral carbon atom.
Bonding in Alkanes
 sp3 hybridised carbons form bonds when
each of the four sp3 orbitals (each with one
electron) overlap with an orbital of another
atom that contains only one electron.
e.g. methane
Ethane
• When orbitals from one atom overlap with the
orbital from another atom a MOLECULAR
ORBITAL is formed.
• The four molecular orbitals in methane (the four CH bonds) can be described as
sp3-s
molecular orbitals
 A molecular orbital between two carbon
atoms (a C-C bond) could be described as a
sp3-sp3 molecular orbital
 All bonds in alkanes are formed by this end-
on overlap of orbitals.
 All the molecular orbitals formed lie along the
axis between the atoms (i.e. if you drew a line
between the two atoms the molecular orbital
would lie on that line).
 When covalent bonds form along the axis
between the atoms they are called sigma
bonds ().
 All covalent bonds in alkanes are sigma
bonds.
Reactions of Alkanes
 Alkanes are relatively unreactive.
e.g. paraffin (from Latin - unreactive) is a
mixture of alkanes used to store alkali metals.
 Only two types of reaction to consider
1) Combustion (alkanes are important fuels)
e.g.
natural gas
petrol
kerosene
diesel
2) Reaction with chlorine and bromine.
Halogenation
 Reaction of bromine with alkanes is relatively
slow, but it does happen.
 Reaction will only start in the presence of
sunlight.
 Reaction is described as a substitution reaction
or a free radical chain reaction.
 A chain reaction consists of 3 steps.
Initiation
 Propagation
 Termination

 Free Radical - An atom or molecule with an
unpaired electron (e.g. Br•, CH3•,
Cl•), always very reactive.
Step 1- Initiation : Homolytic fission of bromine
molecule.(energy supplied by UV light from sunlight)
Br—Br  Br• +
•Br
Step 2 - Propagation 1 : Reaction of Br radical with alkane
R—H
•Br  R• +
H-Br
Propagation 2 : Reaction of alkane radical with Br2
Br—Br
•R  Br• +
R-Br
• The new Br radical can now react with another alkane
molecule, and so on and so on and so on.
• Propagation will continue until one of the reactants is
used up.
Step 3 - Termination : When radicals meet.
Br•
•Br  Br—Br
OR
R•
•Br  R—Br
OR
R•
•R  R—R
Notes on Chain Reactions
 Chain reactions are very difficult to control and
will produce a mixture of products.
e.g.
CH4 + Br2 CH3 Br + CH2 Br2 + CH Br3 + CBr4
 The more complicated the alkane the more
complicated the mixture.
 Controlling the proportions of reactants can
give some control
e.g. CH4 + excess Br2 Mainly CBr4
excess CH4 + Br2 Mainly CH3Br
 The ease of formation for halogen free
radicals is related to bond strength.
I-I
Br-Br
Cl-Cl

Stronger bond

Harder to form radicals

More reactive radicals
 Bromination is slow, but chlorination can be
very fast
e.g. chlorination of methane when initiated
using a flash gun .