4.2 Covalent Bonding
4.2.1 Describe the covalent bond as the result of electron sharing.
4.2.2 Draw the electron distribution of single and multiple bonds in molecules
4.2.3 Deduce the Lewis structures of molecules and ions for up to 4 electron pairs on
each atom.
4.2.4 State and explain the relationship between the number of bonds, bond length and
bond strength.
4.2.5 Predict whether a compound of two or more elements would be covalent from the
position of the elements in their periodic table or from their electronegativity values.
4.2.6 Predict the relative polarity of bonds based on electronegativity values
4.2.7 Predict the shape and bond angles for molecules with four charge centres on the
central atom.
4.2.8 Predict molecular polarity based on bond polarity and molecular shape.
4.2.9 Describe and compare the structure and bonding in the 3 allotropes of carbon
(diamond, graphite and C60 fullerene)
4.2.10 Describe the structure of and bonding in silicon and silicon dioxide
Pure covalent bonds
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Sharing of electrons between two or
more of the same type of non-metal
atoms.
HOBrFINCl elements are all covalently
bonded.
H2, O2, Br2, F2, I2, N2, Cl2
Pure covalent bonds
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Equal sharing of electrons when
forming the bond
H2(g) forms a single bond (shared pair)
Polar covalent bond
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Unequal sharing of electrons.
One atom will have a higher electronegativity than
the other, so it will “pull” the shared electrons closer
to itself making that atom slightly more negative
than the other.
The Cl (3.00) is more negative than the H (2.20)
Naming simple molecules
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Must memorize the
prefixes
RULES: if there is only
one of the first atom than
don’t use a prefix,
otherwise use a prefix.
Ex: CO = carbon
monoxide
Ex: P2O4 =
diphosphorous tetroxide
Prefix
Mono
Number
1
Di
Tri
Tetra
2
3
4
Penta
Hexa
Hepta
Octa
5
6
7
8
Nona
Deca
9
10
Chemical structures
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Need to show the structure of a
molecule.
Will use Lewis structures (electron dot
diagrams) to show where there are
lone pairs (filled orbitals) and bonding
pairs (places where bonds most likely
occur)
Drawing Lewis Structures
1.
2.
3.
Look at valence electrons of all atoms
Pick a central atom (least
electronegative usually, has most
bonding sites)
Align all atoms so that each have their
ideal amount of valence electrons
achieved through sharing.
Carbon tetrachloride
Carbon is the central
atom.
 It has 4 bonding pairs.
 Chlorine wants to share
one bonding site each.
 Need 4 chlorines for
every one carbon
(Cl has 3 lone pairs and 1
bonding pair)

Some examples
Practice drawing and naming
Lewis Structures
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H 2O
CH2O
Tricky ones!

Try ozone O3
What about ions?
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Count up all valence electrons that you
are allowed to place.
Still pick the central atom.
Still have the correct number of
electrons around each atom (usually 8,
except for H and He)
Add extra electrons if an anion and
take away electrons if a cation
Practice with a cation
Practice with an anion
Oxygen has an unshared pair of electrons, but since this is
an anion it receives an extra electron which will fill up the
outer orbital.
Coordinate covalent bonds
(dative)
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A covalent bond that occurs between two atoms in
which both electrons shared in the bond come from
the same atom.
Both electrons from the nitrogen are shared with the
upper hydrogen
Ammonium has 3 polar covalent bonds and 1
coordinate (dative) covalent bond.
Examples
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Hydronium (H3O+)
Carbon monoxide
(CO)
Free Radicals
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A molecule with an odd amount of electrons, or a
broken bond causing a particle with an uneven
amount of electrons
Free radicals are very unstable and react quickly
with other compounds, trying to capture the needed
electron to gain stability, but causing a new free
radical to form in the process.
It’s a chain reaction which usually involves the
destruction of living cells
Vitamin E (fat soluble) and C (water soluble)are
antioxidants which are able to neutralize the
damage by ‘donating’ an electron causing the chain
to stop
Free Radicals
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NO is usually a slow reaction with nitrogen and oxygen gases,
but can occur more quickly in the presence of a catalyst or
high temperatures
NO is a common free radical that is primarily found due to
internal combustion engines (car exhaust).
Cars have catalytic converters to reverse the reaction
(decompose NO)
It reacts to form nitric acid, causing more problems with acid
rain, and reacts with ozone to produce NO2
VSEPR
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Valence shell electron pair repulsion
theory
Bonding pairs and lone pairs around an
atom in a molecule adopt positions
where their mutual interactions are
minimized.
Electron pairs are negatively charged
and will get as far apart from each other
as possible. (Same charge = repulsion)
Bond angles
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Lone pairs occupy more space than
bonding electron pairs.
Double bonds occupy more space
than single bonds.
LP-LP > LP-BP > BP-BP
Lone pairs are more repulsive than
bonding pairs
Chemistry SL Shapes
Sets
(group of
Lone
Pairs
Shape
0
2
0
1
0
Linear 180 o
Bent 104.5 o
Triagonal Planar 120 o
Pyramidal 107.3 o
Tetrahedral 109.5 o
bonding pairs)
2
2
3
3
4
Examples
Arrangement of
electron pairs on
central atom
Number of
bonding
electron pairs
Example
Linear
2
BeCl2
Planar triangular
3
BCl3
Tetrahedral
4
CH4
Practice Lewis structure and
state the shape
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SO2
SO3
[SO4 ] -2
AsCl3
SI2
CH3F
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CH2F2
NH4+
NO2NO2+
H3O+
Advanced structural drawings
(3 D)
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The dashed wedge = bond going back
Solid wedge = bond going forward
Unbroken line = plane of the paper
Polarity and shape
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The shape of the molecule directly
influences the overall polarity of the
molecule.
If there is symmetry the charges
cancel each other out, making the
molecule non-polar
If there is no symmetry, then its polar
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Polar bonds do not guarantee a polar
molecule
Ex: CCl4 and CO2 both have polar bonds,
but both are NON-POLAR molecules. They
have a dipole moment of zero
The greater the dipole moment, the more
polar the molecule
The symetry of the molecule
Cancels out the “charges”
Making this NON-POLAR
No overall DIPOLE
The bent shape creates an
overall positive end and negative end
of the molecule = POLAR
Summary of Polarity of
Molecules
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Linear:
When two atoms attached to central atom are the
same, the molecule will be Non-Polar (CO2)
 When the two atoms are different the dipoles will
not cancel, and the molecule will be Polar (HCN)
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Bent:

The dipoles created from this molecule will not
cancel creating a net dipole moment and the
molecule will be Polar (H2O)
Summary of Polarity of
Molecules
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Pyramidal:

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The dipoles created from this molecule will not
cancel creating a net dipole and the molecule will
be Polar (NH3)
Trigonal Planar:
When the three atoms attached to central atom
are the same, the molecule will be Non-Polar
(BF3)
 When the three atoms are different the dipoles
will not cancel, resulting in a net dipole, and the
molecule will be Polar (CH2O)

Tetrahedral
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When the four atoms
attached to the
central atom are the
same the molecule
will be Non-Polar
When three atoms
are the same, and
one is different, the
dipoles will not
cancel, and the
molecule will be Polar
Summary of Polarity of Molecules
Examples to Try
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Determine whether the following
molecules will be polar or non-polar
 SI2
 CH3F
 AsI3
 H2O2
Angular = bent
triangular pyramid = pyramidal
Testing a liquid’s polarity
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As the liquid is flowing bring a
magnetically charged object close.
If the stream of liquid is attracted to the
rod, it is polar
If the stream is unaffected, it is nonpolar.

Can we explain why this would happen?
Why is molecular polarity
important?
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Polar molecules have higher melting
and boiling points (for example the BP
of HF is 19.5° C, and the BP of F2 is –
188° C).
Polar solvents dissolve ionic and polar
molecules more efficiently than nonpolar solvents
Covalent bond strength
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Two forces operating:
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increased overlap of atomic orbitals
(better sharing) brings atoms together
closer distance between nuclei increases
positive-positive charge repulsion
balance of these forces = its bond
length
Measured in pm (10-12 m) or Ǻ(10-10 m)
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In a molecule as you increase the number of
electrons shared between two atoms (from
single to double to triple bond), you increase
the bond order, increase the strength of the
bond, and decrease the distance between
nuclei.
Bond strength is measured by how much
energy it takes to break the bond (kJ/mol)
Bond Length and Bond strength
Bond enthalpy (energy needed
to break the bond as a gas)
Properties of molecules
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The forces between discrete
molecules are relatively weak
(Intermolecular forces) so
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Low boiling points and melting points
Quite soft if solid
Do not conduct electricity
Tend to be more soluble in non-polar
solvents than polar solvents.
Allotropes of carbon
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elements can exist in two or more
different forms because the element's
atoms are bonded together in a
different manner
Carbon has 3 allotrophes
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Diamond
Graphite
Fullerenes (C60)
Diamonds
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carbon atoms are bonded
together in a tetrahedral lattice
arrangement (3D framework)
Giant covalent structure
Very strong, so they require a
lot of energy to break them
M.P is 3820 K
Does NOT conduct electricity
4x harder than any other
natural mineral
Graphite
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has a sheet like structure
where the atoms all lie in a
plane and are only weakly
bonded to the sheets
above and below. (2D
framework)
Much softer, conducts
electricity.
The C-C bonds are still
quite strong.
Fullerene C60
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consists of 60 carbon atoms
bonded in the nearly spherical
configuration
C60 is highly electronegative,
meaning that it readily forms
compounds
it is a yellow powder which
turns pink when dissolved in
certain solvents such as
toluene.
Also includes nanotubes
(cylindrical)
Silicon
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Has almost
identical crystal
structure to
diamond
Silicon dioxide
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Sometimes called
silica
Occurs as quartz
and sand
Oxygen atoms
bridge the silicon
atoms
Bibliography and good sites
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http://www.chemguide.co.uk/atoms/bo
nding/dative.html
http://en.wikipedia.org/wiki/Coordinate
_covalent_bond
http://en.wikipedia.org/wiki/Diamond
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Use links to find out about fullerenes