Lewis Structures

A Lewis structure, also called a Lewis dot diagram, is a simple way of showing the valence electrons of atoms and how those electrons are arranged when atoms form molecules.

Lewis structures are particularly useful for understanding covalent bonding, because they show which electrons are shared between atoms and which electrons remain as unshared pairs.

What Does a Lewis Dot Diagram Show?

Lewis diagrams focus only on an atom's valence electrons.

Valence electrons are the electrons in the atom's outermost occupied electron shell. These are the electrons involved in chemical bonding.

In a Lewis diagram:

• the element's chemical symbol represents the nucleus and inner electrons
• dots around the symbol represent valence electrons
• each dot represents one electron

For example, carbon has 4 valence electrons, so its Lewis dot diagram contains four dots around C.

Oxygen has 6 valence electrons, while chlorine has 7.

 

 

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Suggested placement: A diagram comparing the Lewis symbols for C, N, O, F, and Ne works well here because students can see the gradual increase in valence electrons.

Finding Valence Electrons

For the main-group elements, the group number can help us determine the number of valence electrons.

*Helium is an exception because it has only 2 electrons, filling its first shell.

For example:

Carbon → Group 14 → 4 valence electrons

Oxygen → Group 16 → 6 valence electrons

Chlorine → Group 17 → 7 valence electrons

Drawing Lewis Diagrams for Atoms

To draw a Lewis diagram for an atom:

Step 1 – Write the element symbol

For nitrogen:

N

Step 2 – Determine the valence electrons

Nitrogen is in Group 15.

Therefore:

5 valence electrons

Step 3 – Add the electrons as dots

Place one electron on each side of the symbol before pairing electrons.

A simple representation is:

  ••
• N •
  •

Nitrogen therefore has:

• 5 valence electrons
• one electron pair
• three unpaired electrons

Those unpaired electrons can participate in covalent bonding.

What Happens During Covalent Bonding?

A covalent bond forms when atoms share electrons.

The shared electrons are attracted to the nuclei of both atoms, helping hold the atoms together.

A shared pair of electrons forms one covalent bond.

So:

1 shared pair → single bond

2 shared pairs → double bond

3 shared pairs → triple bond

Shared and Unshared Electron Pairs

Not every valence electron has to form a bond.

In a molecule, electron pairs can be classified as either shared pairs or unshared pairs.

Shared Pair

A pair of electrons located between two bonded atoms.

A shared pair forms a covalent bond.

Unshared Pair

A pair of valence electrons that does not participate in bonding.

These are also called lone pairs.

Understanding the difference between bonding pairs and lone pairs is essential when interpreting Lewis structures.

Hydrogen Molecule – H₂

Each hydrogen atom has one valence electron.

Two hydrogen atoms can share their electrons:

H:H

The two electrons between the atoms form one shared pair.

A shared pair can also be represented using a line:

H—H

The line represents one covalent bond.

Each hydrogen now has access to two electrons, filling its first electron shell.

Chlorine Molecule – Cl₂

Each chlorine atom has 7 valence electrons.

Each chlorine needs one more electron to achieve a full outer shell.

Two chlorine atoms can therefore share one pair of electrons:

Cl—Cl

The shared pair forms a single covalent bond.

Each chlorine atom also has three lone pairs.

This demonstrates an important idea:

Atoms do not have to transfer electrons to become stable. They can also share them.

Water – H₂O

Water provides a useful example containing more than one bond.

Oxygen has 6 valence electrons.

Each hydrogen has 1 valence electron.

Oxygen shares one electron pair with each hydrogen:

  ••
H—O—H
  ••

The structure contains:

• 2 O—H covalent bonds
• 2 shared electron pairs
• 2 lone pairs on oxygen

Oxygen now has access to eight electrons in its outer shell, while each hydrogen has access to two.

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Single, Double, and Triple Bonds

Sometimes atoms must share more than one pair of electrons.

Single Bond

One shared pair:

H—H

1 shared pair = 1 covalent bond

Double Bond

Two shared pairs:

O=O

Each oxygen atom needs two additional electrons to complete its outer shell.

They therefore share two pairs of electrons.

This forms a double covalent bond.

Triple Bond

Three shared pairs:

N≡N

Each nitrogen atom has five valence electrons and needs three more.

The two nitrogen atoms share three electron pairs, producing a triple covalent bond.

So:

 

 

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Suggested placement: A side-by-side comparison of H₂, O₂, and N₂ makes the progression from single to double to triple bonding particularly clear.

Carbon Dioxide – CO₂

Carbon dioxide contains both carbon and oxygen.

Carbon has 4 valence electrons, while each oxygen has 6.

Carbon needs four more electrons to complete its outer shell. Each oxygen needs two.

The atoms share electrons to form:

   ••   ••
    O=C=O
   ••   ••

Carbon forms a double bond with each oxygen atom.

Therefore, CO₂ contains:

• 2 double bonds
• 4 shared electron pairs
• 2 lone pairs on each oxygen
• no lone pairs on the central carbon

Methane – CH₄

Carbon has four valence electrons and can form four covalent bonds.

In methane, carbon shares one pair of electrons with each of four hydrogen atoms.

    H
    |
  H—C—H
    |
    H

Carbon forms four single covalent bonds.

Each hydrogen obtains a full first shell containing two electrons, while carbon obtains a stable outer arrangement of eight electrons.

Determining the Number of Covalent Bonds

A Lewis structure makes it easy to count covalent bonds.

Remember:

Single line = 1 bond

Double line = 2 bonds

Triple line = 3 bonds

For example:

Notice that CO₂ has two connections between atoms, but each connection is a double bond, giving four covalent bonds when bond order is counted.

How Lewis Structures Explain Molecular Formation

Lewis structures help explain why particular molecules form.

Consider oxygen.

An oxygen atom has:

6 valence electrons

It needs:

2 additional electrons

Therefore, oxygen commonly forms two covalent bonds.

Carbon has:

4 valence electrons

It needs:

4 additional electrons

Therefore, carbon commonly forms four covalent bonds.

Nitrogen has:

5 valence electrons

It needs:

3 additional electrons

Therefore, nitrogen commonly forms three covalent bonds.

This gives us a useful introductory pattern:

These patterns help us predict the structures of many simple molecules.

Drawing a Lewis Structure Step by Step

Consider ammonia, NH₃.

Step 1 – Count the valence electrons

Nitrogen:

5 electrons

Three hydrogen atoms:

3 × 1 = 3 electrons

Total:

5 + 3 = 8 valence electrons

Step 2 – Choose the central atom

Nitrogen is the central atom.

Step 3 – Form bonds

Nitrogen forms a single bond with each hydrogen.

   H
   |
 H—N—H

Three bonds use 6 electrons.

Step 4 – Place the remaining electrons

There are two electrons remaining.

They form one lone pair on nitrogen:

  ••
 H—N—H
   |
   H

The completed Lewis structure contains:

• 3 single bonds
• 3 shared electron pairs
• 1 lone pair

Reading an Unfamiliar Lewis Structure

Suppose you see:

H—O—H

with two lone pairs on oxygen.

You should be able to determine that:

• the molecule contains 3 atoms
• oxygen is the central atom
• there are 2 single covalent bonds
• there are 2 shared electron pairs
• oxygen has 2 unshared electron pairs

Lewis structures therefore allow us to interpret how electrons are arranged within a molecule.

Did You Know?

Lewis structures are named after American chemist Gilbert N. Lewis, who introduced the idea of representing valence electrons using dots in 1916.

Lewis structures are simplified models. They do not show the actual three-dimensional shape of molecules, but they provide an extremely useful starting point for understanding bonding, electron pairs, and molecular structure.

Key Terms

Lewis structure – A diagram showing valence electrons and bonding between atoms.

Valence electron – An electron in the outermost occupied electron shell.

Covalent bond – A chemical bond formed when atoms share electrons.

Shared pair – A pair of electrons shared between two bonded atoms.

Lone pair – A pair of valence electrons not involved in bonding.

Single bond – A covalent bond containing one shared electron pair.

Double bond – A covalent bond containing two shared electron pairs.

Triple bond – A covalent bond containing three shared electron pairs.

Key Takeaways

• Lewis diagrams represent valence electrons using dots.
• The periodic table can be used to determine the number of valence electrons in main-group atoms.
• Covalent bonds form when atoms share electrons.
• One shared electron pair forms a single bond.
• Two shared pairs form a double bond.
• Three shared pairs form a triple bond.
• Electron pairs that are not involved in bonding are called lone pairs.
• Lewis structures can be used to determine the number of covalent bonds in a molecule.
• Lewis structures help explain how atoms share electrons to form stable molecules.
• Simple Lewis structures can be used to predict and explain the bonding in molecules such as H₂, Cl₂, H₂O, NH₃, CH₄, O₂, N₂, and CO₂.