Drawing Lewis Structure Models

Essential Standard 2.2Understand chemical bonding and

chemical interactions.

Learning Objective 2.2.3

Predict chemical formulas and names

for simple compounds based on

knowledge of bond formation and

naming conventions.

I Can StatementsAt the end of this lesson, you should be

able to say, with confidence:

• I can determine the total number of valence

electrons for any molecule and how many

covalent bonds each nonmetal can form.

• I can determine the position of each atom in

a covalent molecule and where single,

double, or triple bonds will be located.

• I can use traditional steps to draw the Lewis

Structure for various covalent molecules.

Covalent Molecules

Recall that covalent

molecules consist of

only non-metals that

share electrons by

forming covalent bonds. Non-Metals

Ammonia, NH3, is an

example of a covalent

molecule.

Molecular Structure The two models, below,

can be used to show the

relative locations of atoms

and electrons in an

Ammonia molecule.

Lewis Structure Ball and Stick Model

Molecular Formula

Molecular Structures

The ball and stick model

lets you see a 3-D

model of the molecule.

But before you can

create a ball and stick

model, you have to

create a Lewis

Structure model.

Drawing Lewis Structures

1. Predict the location of each atom

a. Hydrogen is always a terminal atom.

b. The central atom will usually be the atom

for which there is only one.

N

H

HHAmmonia

Molecular Formula

2. Find the Total Number of Valence Electrons

1 5

1 N (5 VE) + 3 H (1 VE)

Total VE = 5 + 3 = 8

Drawing Lewis Structures

3. Draw single bonds from each terminal atom to

the central atom.

N

H

HH

Drawing Lewis Structures

N

H

HH

4. Subtract the number of valence electrons used

from the total number of valence electrons.

Since each covalent bond

represents a pair of

electrons and there are 3

covalent bonds, then 6

valence electrons have

already been used.

8 – 6 = 2 valence electrons not used

Drawing Lewis Structures

N

H

HH

Because each hydrogen

atom can only form one

bond, place the remaining

2 electrons on the nitrogen

atom.

5. Place the remaining valence electrons as lone

pairs of electrons on each of the atoms, except

hydrogen.

Drawing Lewis Structures

N

H

HH

Each line (bond)

represents 2 electrons

6. Make sure all atoms meet the octet rule,

unless its hydrogen which only needs 2

valence electrons.

Hydrogen(s)

1 bond = 2 electrons

Yes

Nitrogen

3 bonds = 6 electrons

1 lone pair = 2 electrons

6 + 2 = 8

Yes

Drawing Lewis Structures

Double Bond ExampleCarbon Dioxide or CO2

There is only one carbon atom, so carbon is

the central atom.

1. Predict the arrangement of the atoms

CO O

Double Bond Example2. Determine the total number of valence

electrons

4 6

1 C (4 VE) + 2 O (6 VE)

Total VE = 4 + 12 = 16

Double Bond Example3. Draw single bonds (lines) from each terminal

atom to the central atom.

O C O

4. Subtract the number of valence electrons used

from the total number of valence electrons.

Double Bond Example

O C O2 2

4 valence

electrons used.

Total of 16 valence electrons

– 4 used electrons

12 unused valence electrons

Double Bond Example

CO O

12 unused valence electrons

5. Place the remaining valence electrons as lone

pairs of electrons on each of the atoms, except

hydrogen.

Double Bond Example

CO O

6. Make sure all atoms meet the octet rule.

Each line (bond)

represents

2 electrons

Oxygen (s)

1 bond = 2 electrons

2 lone pairs = 4 electrons

2 + 4 = 6

No

Carbon

2 bonds = 4 electrons

2 lone pairs = 4 electrons

4 + 4 = 8

Yes

Double Bond Example

CO O

O C O

O C O

7. If any atom does not meet the octet rule,

convert lone pairs into bonded pairs.

Step Two

Step One

Step Three

Double Bond Example

CO O

7. If any atom does not meet the octet rule,

convert lone pairs into bonded pairs. (Continued)

Check Octet rule

again

Oxygen

2 bond = 4 electrons

2 lone pairs = 4 electrons

2 + 4 = 8

Yes

Carbon

4 bonds = 8 electrons

Yes

Triple Bond ExampleNitrogen Molecule or N2

N

1. Predict the arrangement of the atoms

N

2. Determine the total number of valence

electrons

2 N (5 VE) = 10 VE

5

Triple Bond Example3. Draw a single bond (lines) between the two

atoms.

N N

4. Subtract the number of valence electrons used

from the total number of valence electrons.

N N2 valence electrons used.

Total of 10 valence electrons

– 2 used electrons

8 unused valence electrons

Triple Bond Example5. Place the unused valence electrons as lone

pairs of electrons on each of the two atoms.

N N

8 unused Valene Electrons

6. Make sure all atoms meet the octet rule.

1st Nitrogen

1 bond = 2 electrons

2 lone pairs = 4 electrons

2 + 4 = 6

No

2nd Nitrogen

1 bond = 2 electrons

2 lone pairs = 4 electrons

2 + 4 = 6

No

Triple Bond Example7. If an atom does not meet the octet rule,

convert lone pairs into bonded pairs.

N N

1st Nitrogen

3 bond = 6 electrons

1 lone pair = 2 electrons

8 electrons

Yes

Check Octet rule

again

2nd Nitrogen

3 bond = 6 electrons

1 lone pair = 2 electrons

8 electrons

Yes

Triple Covalent Bonds Triple covalent bonds are the strongest

chemical bonds ever formed.

78% of our air

consists of nitrogen

gas (N2).

We breathe it in, and breathe it right back out

again because our bodies are unable to break

that triple covalent bond.

Nitrogen Cycle Remember when you learned about the nitrogen

cycle?

The only

organisms

capable of

breaking the triple

covalent bond are

bacteria that live

on the roots of

legume plants.

The End