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.

CF₄, H₂O and CO₂ 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.

CF₄, H₂O and CO₂ dot structures
.. : F : . . . . . . : F : C : F : ^{. . . . . .} : F : ' '	_.. ._. O / \ H H	_.. ._. O::C::O ^.. .^. or . . . . : O=C=O :

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 CH₄, NH₃, HF, OF₂, F₂, O₂, N₂, 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 SO₂

Solution

Put down number of valence electrons:

.. : O :	.. : S :	.. : O :

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 6-valence-electron 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 SO₂ 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 S-O 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: ' ' ' '	«	.. 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. CO₂ - :O::C::O: (plus two more dots for each of O)
NO₂ - ^.NO₂ (bent molecule due to the odd electron)
NO₂^- - :NO₂^- (same number of electron as SO₂)
HCO₂^- - H-CO₂
O₃ - (ozone, OO₂ same number of electron as SO₂)
SO₃ - (consider O-SO₂, and the resonance structures)
NO₃^- (see Example 2 below)
CO₃^2- (ditto) Notice that some of the resonance structures may not satisfy the octet rule. The NO₂ 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 NO₃^-

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 HNO₃
H₂SO₄
H₂CO₃
HClO₄
C₅H₅N
NO₃^-
SO₄^2-
CO₃^2-
ClO₄^-
Benzene C₆H₆
Cl₂CO 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 HNO₃, NO₃^-, H₂CO₃, CO₃^2-, C₅H₅N, C₆H₆, and Cl₂CO.

Exceptions to the octet rule

We can write Lewis dot structures that satisfy the octet rule for many molecules consisting of main-group elements, but the octet rule may not be satisfied for a number of compounds. For example, the dot structures for NO, NO₂, BF₃ (AlCl₃), and BeCl₂ 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 AlCl₃ and BeCL₂ 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, AlCl₃, 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, Al₂Cl₆, whose structure is shown above. It vaporizes as dimers, but further heating gives a monomer that has the same structure as the BF₃.

Arrange dots this way
::Ex

:::Ex = ..
:Ex
. .
:Ex:
In compounds PF₅, PCl₅, :SF₄, ::ClF₃, :::XeF₂ and :::I₃^-, the center atoms have more than 10 electrons instead of 8. In compounds SF₆, IOF₅, :IF₅, BrF₅, ::XeF₄, PF₆^- etc, the center atoms have 12 electrons.

Arrange dots this way
::Ex :::Ex	=	.. :Ex . . :Ex:

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 CO₂?

What about NO₂?
What is the formal charge for S in H₂SO₄ in the structure
O-H | (provide all dots yourself) O=S=O ? | O-H

The oxygen double bonded to S has formal charge of 0. What is the formal charge of C in O-C:::O?
What is the formal charge for S in H₂SO₄ in the structure
O-H | (provide all dots yourself) O-S-O ? | O-H

The oxygen double bonded to S has formal charge of -1. What is the formal charge of C in O-C:::O?
According to the formal charge rule, which structure in the two problems you have just work on is the best for H₂SO₄?
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 NO₂.
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 NO₂^-,
CO₂, NO₂, O₃, CO₃^-, or CO₂^-?

Elements of the 2nd period are: Li Be B C N O F Ne.
Draw Lewis dot structure for O₃ and NO₂^-.

.. .. .. : O : S : O : ' ' ' '	<- octet rule satisfied
-1 +2 -1	formal charge