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Lewis Dot Structures
Skills to develop
 Draw the Lewis dot structure of a given molecule or ion.
 Draw resonance structures of some molecules.
 Assign formal charge to an atom in a dot structure.
 Assess the stability of a structure by considering formal charges of atoms.
 Give examples for molecules and ions that do not follow the octet rule.
Lewis Dot Structures
Lewis symbols of the main group elements


H•
       He:


Li •
 • Be •
 _{•} • B •
 _{•} • C • ^{•}
 _{•} : N • ^{•}
 _{•} : O • ''
 _{. .} : F • ''
 _{. .} : Ne : ''

Na K Rb Cs
 Mg Ca Sr Ba
 Al Ga In Tl
 Si Ge Sn Pb
 P As Sb Bi
 S Se Te Po
 Cl Br I At
 Ar Kr At Rn

G.N. Lewis used dots to represent the valence electrons in his teaching
of chemical bonding. He eventually published his theory of chemical
bonding in 1916. He put dots around the symbols so that we see
the valence electrons for the main group elements. Formation of chemical
bonds to complete the requirement of eight electrons for the atom becomes
a natural tendency. Lewis dot symbols of the first two periods are given
here to illustrate this point. In fact, the entire group (column) of
elements have the same Lewis dot symbols, because they have the same
number of valence electrons.
CF_{4}, H_{2}O and CO_{2} dot structures


..
: F :
. . . . . .
: F : C : F :
^{. . . . . . }
: F :
' '
 _{.}. ._{.}
O
/ \
H H
 _{.}. ._{.}
O::C::O
^{.}. .^{.}
or
. . . . : O=C=O :


Lewis dot structures are useful in explaining the chemical bonding
in molecules or ions. When several dot structures are reasonable for a
molecule or ion, they all contribute to the molecular or ionic structure
making it more stable.
The representation of a molecular or ionic structure by several
structures is called resonance. The more stable the dot structure is,
the more it contributes to the electronic structure of the molecule or ion.
You need to know what dot structures represent, how to draw them,
and what the formal charges for the atoms in the structure are.
When several dot structures are possible, consider the
resonance structures to interpret the real structure. Apply
some simple rules to explain which of the resonance structures
are major contributors to the electronic structure.
Drawing Lewis Dot Structures and Resonance Structures
Follow these simple steps to draw Lewis dot structures:

Draw the atoms on paper and put dots around them to represent
valence electrons of the atom. Be sure to have the correct number
of electrons.

If the species is an ion, add or subtract electrons corresponding to the
charge of the ion. Add an electron for every negative () charge,
and subtract an electrons for every positive (+) charge.

Consider bonding between atoms by sharing electrons, some may
come from one atom.

If possible, apply the octet rule to your structure.
Some structures don't obey the octet rule, but explain why.

Assign formal charges to atoms in the structure.
Exercise
Draw Lewis dot structures for CH_{4}, NH_{3}, HF, OF_{2},
F_{2}, O_{2}, N_{2}, Cl^{} and some compounds
you know.
Formal Charge
The formal charge on any atom in a Lewis structure is a number
assigned to it according to the number of valence electrons of the atom
and the number of electrons around it.
The formal charge of an atom is equal to the number of valence electrons,
N_{v.e.} subtract the number of unshared electrons,
N_{us.e.} and subtract half of the bonding electrons,
½ N_{b.e.}.
Formal charge = N_{v.e.}  N_{us.e.}  ½ N_{b.e.}
Some practice of assigning formal charge is necessary before you master
this technique. Some examples of drawing Lewis structure and assigning
formal charge are given below.
The formal charge is a hypothetical charge from the dot structure.
The formal charges in a structure tell us the quality of the dot structure.
Formal charge rules
Often, many Lewis dot structures are possible. These are possible resonance
structures, but often we should write a reasonable one, which is stable.
The formal charge guides us about the stability of the dot structure.
The guidance are called formal charge rules
 Formulas with the lowest magnitude of formal charges are more stable.
 More electonegative atoms should have negative formal charges.
 Adjacent atoms should have opposite formal charges.
Example 1.
Draw Lewis dot structure for SO_{2}
Solution
Put down number of valence electrons:
Put all atoms together to make a molecule and check to see if it satisfy
the octet rule.
.. .. .. : O : : S : : O :  <= octet rule not satisfied


0 0 0  formal charge

Adjust bonding electrons so that octet rules apply to all the atoms
.. .. .. : O : S : : O : ''  < octet rule satisfied

1 +1 0  formal charge

Since the left O has 6 unshared plus 2 shared electrons, it effectively has
7 electrons for a 6valenceelectron O, and thus its formal charge is 1.
Formal charge for O = 6  6  (2/2) = 1.
Formal charge for S = 6  2 (6/2) = +1.
There is yet another structure that does not satisfy the octet rule,
but it's a reasonable structure:
.. .. .. : O : S : O :
' ' ' '
 < octet rule satisfied

1 +2 1
 formal charge

Resonance Structures
When several structures with different electron distributions among
the bonds are possible, all structures contribute to the electronic
structure of the molecule. These structures are called
resonance structures. A combination of all these resonance structures
represents the real or observed structure. The Lewis structures of some
molecules do not agree with the observed structures. For such a molecule,
several dot structures may be drawn. All the dot structures contribute
to the real structure. The more stable structures contribute
more than less stable ones.
For resonance structures, the skeleton of the molecule (or ion) stays in
the same relative position, and only distributions of electrons in the
resonance structures are different.
Let us return to the SO_{2} molecule.
The molecule has a bent structure due to the lone pair of electrons on S.
In the last structure that has a formal charge, there is a single SO bond
and a double S=O bond. These two bonds can switch over giving two
resonance structures as shown below.
..
S
/ \
:O: :O:
' ' ' '
 «

..
S
// \
:O: :O:
' '
 «

..
S
/ \\
:O: :O:
' '
 «

..
S
// \\
:O: :O:


1   2   3   4


..
S
/.' ^{''} '.\
:O: :O:


In structure 1, the formal charges are +2 for S, and 1 for both O atoms.
In structures 2 and 3, the formal charges are +1 for S, and 1 for the
oxygen atom with a single bond to S. The low formal charges of S
make structures 2 and 3 more stable or more important
contributors.
The formal charges for all atoms are zero for structure 4, given earlier.
This is also a possible resonance structure, although the octet rule is not
satisfied. Combining resonance structures 2 and 3 results in
the following structure:
Exercise
Draw the Lewis dot structures and resonance structures for the following.
Some hints are given.
CO_{2}  :O::C::O: (plus two more dots for each of O)
NO_{2}  ^{.}NO_{2} (bent molecule due to the odd electron)
NO_{2}^{}  :NO_{2}^{} (same number of electron as SO_{2})
HCO_{2}^{}  HCO_{2}
O_{3}  (ozone, OO_{2} same number of electron as SO_{2})
SO_{3}  (consider OSO_{2}, and the resonance structures)
NO_{3}^{} (see Example 2 below)
CO_{3}^{2} (ditto)
Notice that some of the resonance structures may not satisfy the octet rule.
The NO_{2} molecule has an odd number of electrons,
and the octet rule cannot be satisfied for the nitrogen atom.
Example 2.
Draw the resonance structures of NO_{3}^{}
Solution
1
:O:

N
/ \
:O: :O:
' ' ' '
 «

. . 1
:O:

N
// \
:O: :O:
' '
 «

. . 1
:O:

N
/ \\
:O: :O:
' '


The resonance structure is shown on the right here.
Note that only the locations of double and single bonds change here.
What are the formal charges for the N atoms?
What are the formal charges for the oxygen atoms that are single bonded
and double bonded to N respectively? Please work these numbers out.
 Formal charges: N, +1; =O, 0; O, 1
 The most stable structure has the least formal charge.
 In a stable structure, adjacent atoms should have formal charges
of opposite signs.
The more stable the structure, the more it contributes to the
resonance structure of the molecule or ion. All three structures
above are the same, only the double bond rotates.
Exercise
Draw the Lewis dot structures and resonance structures for
HNO_{3}
H_{2}SO_{4}
H_{2}CO_{3}
HClO_{4}
C_{5}H_{5}N
NO_{3}^{}
SO_{4}^{2}
CO_{3}^{2}
ClO_{4}^{}
Benzene C_{6}H_{6}
Cl_{2}CO
You have to do these on paper, because putting dots around the symbols
is very difficult using a any word processor. The Octet rule should be
applied to HNO_{3}, NO_{3}^{},
H_{2}CO_{3}, CO_{3}^{2},
C_{5}H_{5}N, C_{6}H_{6},
and Cl_{2}CO.
Exceptions to the octet rule
We can write Lewis dot structures that satisfy the octet rule for many
molecules consisting of maingroup elements, but the octet rule may not
be satisfied for a number of compounds. For example, the dot structures
for NO, NO_{2}, BF_{3} (AlCl_{3}), and
BeCl_{2} do not satisfy the octet rule.
^{.}N:::O:
compare with
:C:::O:

.
N
// \
:O: :O:
' '

..
:F:

B
/ \
:F: :F:
' ' ' '

. . . .
:Cl : Be : Cl:
' ' ' '


The above are structures for the gas molecules. The solids of AlCl_{3}
and BeCL_{2} are polymeric with bridged chlorides.
. . . .
:Cl: :Cl: :Cl:
\ / \ /
Al Al
/ \ / \
:Cl: :Cl: :Cl:
. . . .

Polymeric solid structures

:Cl: :Cl: :Cl:
/ \ / \ / \
Be Be
\ / \ / \ /
:Cl: :Cl: :Cl:


Alumunum chloride, AlCl_{3}, is a white, crystalline solid, and an
ionic compound. However, it has a low melting point of 465 K (192°C),
and the liquid consists of dimers, Al_{2}Cl_{6}, whose
structure is shown above. It vaporizes as dimers, but further heating
gives a monomer that has the same structure as the BF_{3}.
Arrange dots this way


::Ex :::Ex  =

..
:Ex
. .
:Ex:

In compounds PF_{5}, PCl_{5}, :SF_{4},
::ClF_{3}, :::XeF_{2} and :::I_{3}^{},
the center atoms have more than 10 electrons instead of 8.
In compounds SF_{6}, IOF_{5}, :IF_{5},
BrF_{5}, ::XeF_{4}, PF_{6}^{} etc,
the center atoms have 12 electrons.
The formulas given above follows a systematic patter according to
the positions of the elements on the periodic table. As the number of
atoms bonded to it decreases, the number of unshared electrons increase.
Confidence Building Questions

What is the total number of valence electrons in CO_{2}?
What about NO_{2}?

What is the formal charge for S in H_{2}SO_{4} in the
structure
OH
 (provide all dots yourself)
O=S=O ?

OH
The oxygen double bonded to S has formal charge of 0.
What is the formal charge of C in OC:::O?

What is the formal charge for S in H_{2}SO_{4} in the
structure
OH
 (provide all dots yourself)
OSO ?

OH
The oxygen double bonded to S has formal charge of 1.
What is the formal charge of C in OC:::O?

According to the formal charge rule, which structure in the two problems
you have just work on is the best for H_{2}SO_{4}?
 What is the formal charge for N in the structure
.
N = O

:O:
.. ?
What is the formal charge if the structure is
.
:O:
 ..
N = O : ?
..
and which one you think is the best? You can write two resonance structures
for each of the two to give 4 resonance structure for NO_{2}.

What is the formal charge for B in the structure
F
\
B = F ? (Supply your dots for unshared electrons)
/
F
What about B(F)_{3}, all single bond?

What is the formal charge of I in I(Cl)_{3} (Supply your dots)
What about Cl in the same structure?

Which one of the following compounds have the same number of
valence electrons as NO_{2}^{},
CO_{2}, NO_{2}, O_{3}, CO_{3}^{}, or CO_{2}^{}?
Elements of the 2nd period are: Li Be B C N O F Ne.
Draw Lewis dot structure for O_{3} and NO_{2}^{}.
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