Chemistry 545
Inorganic Chemistry
Lecture 1.
Lewis Dot Structures
VSEPR
G. N. Lewis
was probably
the best chemist
who never won
the Nobel Prize
Gilbert Newton Lewis (1875-1946)
Lewis Dot Structures (revision)
Lewis dot structures present a simple approach to bonding that
allows us to rationalize much molecular structure. The idea is that
atoms share electrons in the valence shell to form the chemical
bond, with one pair of electrons per bond. Note that each H-atom
has two electrons, which is the structure of He, the next inert gas.
Electron pair = single bond
Valence electrons
H-atom
H-atom
(Each H-atom has one valence electron)
H2 molecule
Lewis Dot Structures (contd.):
Two shared pairs of electrons
= double bond
O-atom
O-atom
O2 molecule
Periodic table
1
2
3
Oxygen has six valence electrons
4
5
6
7
8
The octet rule
Electrons are shared in forming bonds such that atoms have the
same number of electrons in their valence shells as the nearest
noble gas, including the electrons shared with the atom to which
they are bonded.
O-atom
O-atom
O2 molecule
Each oxygen atom in the O2 molecule now has eight
valence electrons, including those it shares with the
other oxygen atom = number of electrons (8 = octet)
in the nearest inert gas = neon.
8.1 Chemical Bonds, Lewis Symbols, and
the Octet rule.
Chemical bonding involves mainly the attempt to
achieve the rare gas number of valence electrons,
i.e. an octet. This can be achieved in several
ways.
Ionic bond: Electrons are mainly the property of
one of the two atoms forming the bond.
Covalent bond: Electrons are shared so that each
atom has a noble gas electronic configuration.
Metallic bonds. Electrons are lost into the
conduction band.
8.2 Ionic Bonding.
This occurs between metallic elements from
the left-hand side of the periodic table and
non-metallic elements from the right hand
side of the periodic table.
Note that Na gives up its lone valence
electron to Cl, so that they both end up with
an octet of electrons.
8.3 Covalent bonding.
Here the two atoms share the electrons
to achieve a covalent bond.
two pairs of electrons equally shared
between the two oxygen atoms
Multiple Bonds and bond order:
The sharing of a single pair of electrons
consititutes a single bond. Sharing of two
pairs of electrons constitutes a double bond,
and sharing three pairs of electrons
constitutes a triple bond.
H:H
Single bond
.. ..
:O::O:
double bond
:N:::N:
triple bond
Bond order: a single bond has bond order =
1, a double bond has bond order = 2, and a
triple bond has a bond order = 3. Fractional
bond orders such as 1½ or 1⅓ are also
possible, as discussed below.
Some more examples of Lewis dot structures:
The N2 molecule:
N-atom
N-atom
triple bond
N2 molecule
Periodic table
1
2
3
4
5
6
7
8
Examples of Lewis dot diagrams:
Methane, CH4:
One shared
pair of electrons
= single bond
Carbon has four
valence electrons (red)
H
H
C
H
H
Hydrogens
achieve two
electrons like He
Carbon achieves
octet of electrons
single line
= single bond
Examples of Lewis dot diagrams:
Carbon dioxide: (CO2)
Carbon has four
valence electrons (red)
oxygens have six
valence electrons (black)
O=C=O
double line =
double bond
Carbon and both oxygens
achieve an octet of electrons
two shared
pairs of electrons
= double bond
Examples of Lewis dot diagrams:
Sulfur dioxide: (SO2)
double
bond?
single
bond?
O=S-O
(or O-S=O ?)
SO2 is an example where a
actual structure is average
molecule can be written in
of the two (bond order = 1½) :
two ways and actual structure
is the average of the two. This
is called RESONANCE (see later)
O
S
O
Slightly different Lewis dot
representations:
One can also represent molecules/ions with a
combination of dots and lines for bonds,
remembering that each line represents a shared
pair of electrons, e.g. the phosphate anion:
8.6 Resonance structures: Ozone (O3)
bond order = 1½
..
:
:
The ozone molecule can
be written with two
equivalent Lewis dot
structures. In such a
situation the actual
structure is the average
of these two structures,
with the two O-O bond
lengths equal.
O
=
O
O
O
:
O
O
:
:
O
::
: :
:
O
O
..
double arrow
= resonance
O-O bonds = 2.78 Å
O
O
O
The ozone molecule
Resonance structures – the nitrite anion: (NO2-)
In drawing up a Lewis dot diagram, if we are dealing with
an anion, we must put in an extra electron for each
negative charge on the anion:
negative charge
on anion
One extra electron
in Lewis dot
N
diagram because
O
Bond order
O
of single negative
= 1½
charge on anion
:
:
:
::
..
:
:
:
-
Two resonance structures
-
:
:
:
O
..
N
O
:
:
O
::
: :
:
O
..
N
=
-
N
O
O
average structure
The nitrate anion:
O
:
Number of canonical structures
:
O
: O:
:
.. N O
O.
.
:
:
O
O
N
B.O. = 1
..
:
O
N
..
:
:
:
:
average bond
order (B.O.)=
2 + 1 + 1 = 1⅓
3
O
..
O
B.O. = 1
..
:
O
N
:
B.O. = 2
-
..
..
O
:
..
to work out bond order,
pick the same bond in
each structure and
average the bond order
for that bond
Resonance in benzene.
H
C
H
C
C
H
C
C
C
H
H
H
H
H
C
C
C
C
H
H
or
C
C
H
H
There are two
canonical structures
for benzene, which
means that the C to C
bonds have a bond
order of (2+1)/2 = 1.5.
The benzene ring has
a very high stability
due to this resonance,
which is called
aromaticity.
Short-hand versions for the benzene ring
8.7. Exceptions to the octet rule.
BF3. This can be written as F2B=F with three
resonance structures. To complete its octet, BF3
readily reacts with e.g. H2O to form BF3.H2O. The
actual structure of BF3 appears not to involve a
double bond and does not obey the octet rule:
Possible resonance
structure for BF3,
but is not important
as this would
involve the very
electronegative
F donating e’s to B
Best representation of
BF3 with B
having only
6 electrons
in its valence
shell
Exceptions to the octet rule: free radicals
There are some molecules that do not obey the octet
rule because they have an odd number of electrons.
Such molecules are very reactive, because they do
not achieve an inert gas structure, and are known as
free radicals. Examples of free radicals are chlorine
dioxide, nitric oxide, nitrogen dioxide, and the
superoxide radical:
odd electrons
nitric oxide
chlorine dioxide
Exceptions to the Octet rule: Heavier atoms (P, As,
S, Se, Cl, Br, I) may attain more than an octet of
electrons:
Example: PF5.
In PF5, the P atom has ten electrons in its valence
shell, which occurs commonly for heavier non-metal
atoms:
leave off F
F
electrons not
shared with P
P
F
F
F
P has
10 valence
electrons
F
PF5
Many phosphorus compounds do obey the
octet rule:
PF3 and [PO4]3- :
three blue electrons are
from charge on anion
Some compounds greatly exceed an
octet of electrons:
IF7
XeF6
(both I and Xe have 14 valence e’s)
(Think about [XeF8]2-)