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Chemistry: Unit F321: Atoms, Bonds & Groups
Covalent Bonding
o Covalent bonding occurs between non metal atoms. They overlap their outer
shells to give each a noble gas structure, with a stable octet of 8 outer
electrons. The resulting species is a molecule
o A covalent bond is a shared pair of electrons. There is a force of attraction
between the bonding pair of electrons and the nuclei of the atoms at each
side.
o Hydrogen atoms need only 2 electrons in their outer shell to reach the noble
gas structure of Helium.
The covalent bond holds the two atoms
together because the pair of shared
electrons is attracted to both nuclei. The
electrons will be held in the exact centre of
the bond because each atom is the same
A single covalent – or sigma (σ) – bond is
and therefore they both have the same
formed when one pair of electrons is
electronegativity.
shared.
o Some atoms make double covalent bonds: oxygen and
carbon share 2 pairs of electrons in this double
covalent bond in carbon dioxide.
o Some atoms make triple covalent bonds: nitrogen combines
with another nitrogen atom, sharing three pairs of electrons
to make a triple covalent bond.
o There are some exceptions to the octet rule: Boron for example
has a depleted octet, having only 6 electrons in its outer shell.
o Sulfur can have an expanded octet: in the compound sulphur
hexafluoride, the sulphur atom has 12 electrons in its outer shell
o A dative covalent bond (or coordinate bond) is a shared pair of electrons
which has been provided by one of the bonding atoms only. A dative covalent
bond can be formed between an ammonia molecule and an H+ ion (which is
electron deficient); both of the bonding electrons are donated by the
nitrogen atom.
o Boron triflouride-ammonia is a
second example of a dative
covalent bond being formed;
both of the bonding electrons
are donated by the nitrogen
atom.
o The shape of a molecule is determined by the repulsion of electron pairs
around a central atom. The electron pair repulsion theory states that;
 the four pairs of electrons around a central atom repel each other
 the pairs of electrons arrange themselves so that they are as far
apart as possible
 the repulsion between lone pairs is greater than the repulsion
between bonding pairs
o There are 6 shapes for molecules with up to 6 electron pairs surrounding a
central atom:
Shape: Tetrahedral
Bond Angle: 109.5
Fischer projection:
Shape: Trigonal Planar
Bond Angle: 120
Fischer projection:
Example: Methane, CH4 –
4 bonding pairs of electrons
Example: BCl3 – 3 bonding pairs,
no lone pairs
Shape: Non linear
Bond Angle: 104.5
Fischer projection:
Shape: Pyramidal
Bond Angle: 107
Fischer projection:
Example: Water, H2O – 2
bonding pairs, two lone pairs
Example: Ammonia, NH3 –
3 bonding pairs, 1 lone pair
Shape: Linear
Bond Angle: 180
Fischer projection:
Shape: Octahedral
Bond Angle: 90
Fischer projection:
Example: CO2 – 2 double bonding
pairs, no lone pairs
Example: SF6 – 6 bonding pairs
o A spectrum of bonding: As
the difference in
electronegativity
increases the bonds
become more polar, then
become ionic.
o Electronegativity is the ability of an atom to attract the bonding electrons in
a covalent bond.
o Pauling’s Scale of Electronegativity gives each element a value for
electronegativity.
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
As we move from left to right across a period, there is
an increased nuclear charge – more protons in the
nucleus. Electrons are added to the same shell, so
there is no increase in distance between the bonding
electrons and the nucleus, and no additional electron
shielding. The electronegativity therefore increases.
As we move down a group, there
are a greater number of
electron shells. This results in
increased electron shielding, so
there are more layers between
the bonding electrons and the
nucleus. There is a greater
distance between the bonding
electrons and the nucleus.
Although there are an increased
number of protons in the
nucleus of the atoms as we go
down the group, this is
outweighed by the increased
distance and electron shielding.
o In a non polar covalent bond, the bonding electrons are
shared equally between atoms which have the same
electronegativity values.
o In a polar covalent bond, a difference in electronegativity
results in unequal sharing of bonding electrons and a
permanent dipole. A permanent dipole is therefore a small
charge difference across a bond that results from the
difference in the electronegativities of the bonded atoms. The
greater the difference in electronegativity, the greater the permanent dipole.
o Overall Polarity – not all molecules containing polar bonds are polar overall.
Sometimes, the dipoles can cancel out, resulting in a net no overall dipole,
making the molecule as a whole non polar.
Net dipole
– polar
No Net dipole –
non polar
Net dipole
– polar
o Intermolecular Forces – Acting between molecules
1. Van der Waals’ forces – non polar molecules have small, temporary,
fluctuating dipoles due to random movement of electrons. This
instantaneous dipole induces a dipole on
another neighbouring molecule. The
larger the molecule, the greater the
number of electrons, the larger the
temporary dipoles and so the greater the
attractive force.
2. Permanent Dipole-Dipole Interaction – polar molecules have
permanent dipoles. The permanent dipole on one molecule attracts the
permanent dipole on another molecule to form the permanent dipoledipole interaction. Molecules with permanent dipoles have much
stronger intermolecular forces than molecules with induced dipoles.
3. Hydrogen Bonding – a hydrogen bond is a strong
dipole-dipole interaction between an electron
deficient hydrogen atom attached to an
electronegative atom and the lone pair of
electrons on a highly electronegative atom (O, N
or F) on a different molecule.
The atoms involved in hydrogen bonds are small; they therefore have a high charge
density, creating greater intermolecular forces, resulting in higher boiling points.
The peculiar nature of water…
1. Ice is less dense than water
Each water molecule is bonded to 4 others in an open lattice
structure. This means that the volume of the ice is larger than the
volume of the liquid, making it less dense. When ice melts, the
hydrogen bonds break and the structure breaks down, so water
molecules can then move closer together, filling all of the gaps.
2. Water has a higher boiling and freezing point than expected
Water has a much higher boiling point than
H2S, H2Se and H2Te. This is because of the
hydrogen bonding between water
molecules; the other compounds have only
weak Van der Waals forces, whereas
water, with hydrogen bonds, must be
heated to a higher temperature before the
hydrogen bonds can be overcome.