Covalent Bonding 3
Molecular Shape
Molecular Shape
• Lewis diagrams are of use when we are learning
what a covalent bond is, but the most useful
representation of molecules in the structural
formula.
• In this representation each pair of
electrons, bonding and non-bonding
pairs alike, is shown as a simple line.
▫ Non-bonding pairs can also be shown as two dots.
Molecular Shape
• The actual shape of the molecule is also
shown.
• This shape has an important part to play in
determining the chemical and physical
properties of a molecule.
Molecular Shape
• Each pair of electrons will be repelled
from the others as far as possible in
three-dimensional space because
electrons all carry a negative charge.
• This is known as the valence electron pair
repulsion or VSEPR theory.
▫ This theory states that the electron pairs
around an atom repel each other.
Molecular Shape
▫ The electrostatic repulsion of pairs of electrons
determines the geometry of the atoms in the
molecule.
▫ The non-bonding pairs and bonding pairs
of electrons are arranged around the
central atom so as to minimize this
electrostatic repulsion between the nonbonding and bonding pairs of electrons.
Shape of Molecule
• The relative magnitude of the electron
pair repulsion is:
Non-bonding pair-non-bonding pair > bonding
pair-non-bonding pair > bonding pair-bonding
pair
• A non-bonding pair will spread out more than a
bonding pair and, therefore, the repulsion will
be greatest between non-bonding pairs.
Molecular Shape
• While non-bonding pairs of electrons can be
important in determining the overall shape of a
molecule, they are no actually considered part of
the shape.
▫ Shape describes the position of the atoms only;
however, non-bonding electrons repel other pairs
of electrons and so influence the final shape of a
molecule.
Molecular Shape
• Negative charge centre or region refers to pairs
of electrons on the central atom.
• This includes both the non-bonding pairs and
bonding pairs of electrons in single, double or
triple bonds.
▫ Each double or triple bond is counted as one
negative charge centre.
Molecular Shape
• For central atoms that have an octet of electrons
in the valence level, you can classify most shapes
into five categories.
Linear
• A combination of two double bonds or a single
bond and a triple bond results in two negative
charge centres.
• This results in a linear shape.
• Examples
▫ Carbon dioxide (CO2)
▫ Hydrogen cyanide (HCN)
• The bond angle, the angle formed by two
bonds, in linear molecules is 180o.
Carbon dioxide (CO2)
Trigonal Planar
• Three negative charge centres around a
central atom can be achieved by a combination
of two single bonds and one double bond
resulting in a trigonal planar shape.
• Example
▫ Methanal (CH2O)
▫ Boron trifluoride (BF3)
• All three bond angles are 120o.
Boron trifluoride (BF3)
• There are some exceptions to the octet rule.
• Boron is one of these exceptions. It is a
metalloid in group 3 and it has 3 valence
electrons.
• It forms three covalent bonds and ends up with
6 electrons.
Tetrahedral
• The most common situation encountered in
molecules is the existence of four pairs of
electrons, either bonding or non-bonding,
surrounding each atom.
• The most widely spaced arrangement of four
pairs of electrons in three-dimensional space is
known as the tetrahedral arrangement, in which
each atom can be imagined at the vertex of a
regular triangular-based pyramid.
• The bond angle in this arrangement is
109.5o.
Methane
Molecular Shape
• Shapes involving only bonding pairs are very
symmetrical because repulsive forces between all
electron groupings are the same.
• When non-bonding pairs are introduced into a
shape, the forces between electron groupings are
no longer the same resulting in a distortion of
the symmetry.
Pyramidal
• One non-bonding pair and three single
bonds result in a pyramidal shape.
• The non-bonding pair exerts a stronger force on
the bonded pairs than they do on each other,
reducing the bond angle to less than
109.5o found in the tetrahedral shape.
• It is not possible to predict exactly how much the
angle will be changed because the result depends
on the properties of the elements and the bonds
between the atoms.
Ammonia (NH3)
Bent
• When the central atom has two non-bonding
pairs and two single bonds, the resulting
shape is bent.
• The two non-bonding pairs exert greater force
on each other than on the bonding pairs and
reduce the angle between the two bonds even
more than in the pyramidal shape.
Water (H2O)
Predicting Molecular Shape
1. Draw a preliminary Lewis structure of the
molecule based on the formula given.
2. Determine the total number of negative charge
centres around the central atom. Remember that
a double or triple bond is counted as one charge
centre.
3. Determine the types of charge centres (bonding or
non-bonding pairs).
4. Determine which one of the five shapes will
accommodate this combination of charge centres.
The Ammonium Ion (NH4+)
• The ammonium ion is formed when a hydrogen
ion bonds to the non-bonding pair of an
ammonia molecule.
• Although the bond between the H+ and the
nitrogen is a covalent bond, it is formed in a
slightly different way from the other covalent
bonds in this molecule.
• The H+ ion has no electrons to share in the bond,
so both bonding electrons come from the
nitrogen atom.
The Ammonium Ion (NH4+)
• The H+ ion has no electrons to share in the bond,
so both bonding electrons come from the
nitrogen atom.
• This is shown by the arrow in the structural
formula.
Dative Covalent Bonds
• A covalent bond which is formed in this
way is called a dative covalent bond.
▫ You might also see this called co-ordinate covalent
bond.
• A dative covalent bond behaves in the same way
as any other single covalent bond.
The Ammonium Ion (NH4+)
• In the ammonium ion, the nitrogen is
surrounded by four bonding pairs of electrons,
rather than three bonding and one non-bonding,
so the ammonium ion has the same shape as a
methane molecule—it is tetrahedral, with bond
angles of 109.5o.
Sulfur dioxide (SO2)
• The structure of sulfur dioxide is more unusual
than those we have seen so far.
• The central atom, sulfur, has 6 valence electrons.
• Each oxygen atom has 6 electrons and needs to
make two bonds to fill its valence shell.
• This Lewis structure suggests that SO2 contains
a single bond and a double bond.
Sulfur dioxide (SO2)
• However, experimental measurements of bond
lengths indicate that the bonds between the S
and each O are identical.
• The two bonds have properties that are
somewhere between a single and a double bond.
• In effect, the SO2 molecule contains two “oneand-a-half” bonds.
Sulfur dioxide (SO2)
• To communicate the bonding in SO2 more
accurately, chemists draw two different
resonance Lewis structures side by side
separated by a double-headed arrow.
• Resonance structures are models that
give the same relative position of atoms
as in Lewis structures, but show different
placing of their bonding and non-bonding
pairs.
Resonance Structures
• Many molecules and ions require resonance
structures to represent their bonding.
• It is important to keep in mind that resonance
structures do not exist in reality.
• An actual SO2 molecule is a combination—a
hybrid—of its two resonance structures.
Practice Problems
• Work on practice problems 6 & 7 on page 193 of
the textbook.