Representing
Molecules
Bonding
Chemical bonds are forces that cause a
group of atoms to behave as a unit.
Bonds result from the tendency of a system
to seek its lowest possible energy.
Bond breaking always requires energy, and
bond formation always releases energy.
Types of Bonds
The type of bonding depends upon the
nature of the atoms that are combined.
A metal and a non-metal will form ionic bonds
when electrons are transferred from the metal to
the non-metal. The resulting attraction between
oppositely charged ions forms a stable crystal.
Types of Bonds
When metals bond with each other, the
valence electrons are shared by the atoms in the
entire crystal. The electrons are no longer
associated with a specific nucleus, and are free to
move throughout the sample.
Lewis Structures
Lewis Structures, also known as Lewis dot
diagrams, show how the valence electrons are
arranged among the atoms in the molecule.
For ionic compounds, it shows the end
result when the metal loses its electrons to the
non-metal.
Covalent compounds
exist as discrete
molecules, whereas
ionic compounds
consist of an
aggregate of cations
and anions.
Covalent Bonds
When two (or more) non-metals form
bonds, electrons are shared. The result is a
covalent bond.
Covalent bonds form because the attraction
of electrons for the nuclei in the atoms is greater
than the electron-electron repulsion or the
nucleus-nucleus repulsion.
Types of Bonds
There is usually
an optimum bond
length or internuclear
distance where
attractions between
electrons and the
nuclei are optimized
and repulsions are
minimized.
Covalent Bonding
Bond Energy
Bond formation releases energy, and bond
breaking requires energy.
Types of Covalent Bonds
Atoms bonded together may share one, two
or three pairs of electrons to make single, double
or triple bonds.
Double and triple bonds are stronger and
shorter than single bonds between the same
atoms.
Covalent Bonding
When atoms of the same element form a
covalent bond, the electrons are shared equally.
Such a bond is called non-polar, or pure covalent.
All homonuclear diatomic molecules contain
non-polar bonds.
Electronegativity
An electronegativity scale, developed by Linus
Pauling (1901-1995), is used to predict the
direction of the polarity of bonds.
Electronegativity is a relative scale that
reflects the ability of an atom to attract the
electrons in a bond.
Electronegativity
The small atoms in the upper right corner of
the table, having high values of Zeff, also have
high electronegativity values. Fluorine has the
highest value at 4.0.
The noble gases generally do not form
compounds, and are not given electronegativity
values.
Electronegativity
The large metal atoms in the lower left hand
corner of the periodic table have the lowest
electronegativity values.
Electronegativity
Electronegativity
Note that hydrogen has an electro-negativity
value of 2.1, consistent with non-metals.
Covalent Bonding
When atoms of different elements form a
covalent bond, the electrons often are not
shared equally. The electrons in the bond may
spend, on average, more time on one of the
atoms. This atom will have a slightly negative
charge, indicated by the symbol δ–.
The other atom will be slightly positive,
indicated as δ+.
Covalent Bonding
Covalent bonds with unequal sharing of the
electrons in the bond are called polar bonds.
An example is the molecule HF. The
electrons spend more time on fluorine than on
hydrogen. As a result, HF is a polar molecule.
δ+
H F
δ-
Dipole Moment
Polar molecules have a dipole moment. This is a
measure of the tendency of a molecule to line up
in an electric field.
Dipole Moment
Dipole moment, μ, depends upon the size of
the partial charges and the distance between the
charges. It is measured in Debye (D). A
positive charge and a negative charge separated
by 100 pm has a dipole moment of 4.80 D.
Polarity of Molecules
Molecules with polar bonds may be polar,
having a permanent dipole moment. Both the
polarity of the bonds and the shape of the
molecule must be considered.
Once the Lewis structure (dot diagram) for
the molecule has been determined, it is possible
to predict the shape of the molecule and its
polarity.
Polar Molecules
Lewis Structures – Covalent
Molecules
Once the Lewis structure has been obtained,
the Valence Shell Electron Pair Repulsion
approach can be used to predict the shape of the
molecule and its polarity.
CO2 has polar carbon to oxygen bonds.
Even though the bonds are polar, the molecule
is non-polar because of its linear shape.
Polarity of Molecules
Since CO2 is a linear molecule, the dipoles
cancel out, and the molecule is non-polar.
Polarity of Molecules
Water also has polar bonds, and a bent
shape. As a result, water is a polar molecule.
Polarity of Molecules
The polarity of molecules has a profound
effect on the properties and behavior of a
substance. It will affect solubility, melting and
boiling points, and other important aspects of
molecular behavior.
The Continuum of Bond Types
Between non-metals
Metal and a non-metal
Lewis Structures – Covalent
Molecules
In many covalent molecules, the non-metals
share valence electrons. Electrons are shared so
that each atom in the molecule has a full valence
shell.
H. . H
H:H or H..H or H-H
Each hydrogen has access to two electrons,
as does the noble gas helium.
Lewis Structures – Covalent
Molecules
For elements in period 2, the non-metals
generally share enough valence electrons so that
each atom obtains the same number of valence
electrons as neon (a total of eight electrons).
This may involve making multiple (double or
triple) bonds between atoms.
Lewis Structures – Covalent
Molecules
Provide Lewis structures for elemental
nitrogen, oxygen and fluorine.
Exceptions to the Octet “Rule”
The elements B and Be sometimes form
compounds with less than four electrons pairs
on them. The are called electron deficient, and
are often highly reactive.
The elements in period 3 and below may
accommodate more than four electron pairs.
Exceptions to the Octet “Rule”
In addition, some molecules, such as NO or
NO2, have an odd number of electrons and do
not obey the octet rule.
Practice

Write Lewis dot diagrams for CO, NH3 and
CO2
Resonance
Some molecules may have more than one
valid Lewis structure. These structures differ in
the placement of multiple bonds.
In molecules with resonance, none of the
Lewis structures accurately represents the true
bonding in the molecule.
Resonance
The molecule SO2 has two resonance
structures:
: :
:
: :
: :
:
: :
[ O=S-O: ] ↔ [ :O-S=O]
The molecule has two equivalent bonds
between sulfur and oxygen.
Resonance
The sulfur-oxygen bonds are identicallonger than double bonds, and shorter than
single bonds.
: :
:
: :
: :
:
: :
[ O=S-O: ] ↔ [ :O-S=O]
The true structure of the molecule is in
between the two Lewis structures drawn.
Formal Charges
Formal charge is a way to keep track of the
electrons in a covalent molecule. The formal
charges can also be used to determine if one
Lewis structure is more valid than another.
Formal Charges
The formal charge on an atom is a
comparison between the number of valence
electrons on each atom and the number of
electrons it has in the Lewis structure.
Formal Charges
Consider the ion SCN-1. There are three
valid Lewis structures for the ion.
-1
: :
: :
:
:
-1
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
Formal charges can be used to determine the
major contributor(s) to the actual structure of
the ion.
Formal Charges
Divide the bonds in half and determine the
number of electrons on each atom.
-1
: :
: :
:
:
-1
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
6e 4e 6e
7e 4e 5e
5e 4e 7e
Formal Charges
Compare the number of electrons in the
structure to the number of valence electrons.
-1
: :
: :
:
:
-1
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
6e 4e 6e
6e 4e 5e
7e 4e 5e
6e 4e 5e
5e 4e 7e
6e 4e 5e
Formal Charges
The net charge is the formal charge on each
atom.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
6e 4e 6e
6e 4e 5e
7e 4e 5e
6e 4e 5e
5e 4e 7e
6e 4e 5e
Formal Charges
The net charge is the formal charge on each
atom.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
The sum of the formal charges must equal
the charge on the ion.
Formal Charges
There are two rules used to determine the
most likely Lewis structure(s).
1. Atoms try to achieve formal charges as
close to zero as possible.
2. Any negative formal charges should
reside on the most electronegative atoms.
Formal Charges
The third Lewis structure is unlikely, due to
the high formal charge on nitrogen.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
Since nitrogen is more electronegative than
sulfur, the first structure should be the major
contributor.
Formal Charges
The actual molecule will be somewhere in
between the first and second structures.
0
0
-1
0
-1
+1
0 -2
: :
: :
:
:
-1
-1 0
-1
[:S=C=N:] ↔ [:S-C N:] ↔ [:S C-N:]
The sulfur-carbon bond should be slightly
longer than a double bond, and the carbonnitrogen bond should be slightly shorter than a
double bond.