Covalent Bonding Model (VCE SSCE Chemistry): Revision Notes

Covalent Bonding Model

What is covalent bonding?

When non-metal atoms combine, they form bonds by sharing electrons. This type of bonding is called covalent bonding, and it occurs because non-metallic atoms have relatively high numbers of electrons in their outer shells.

A molecule is a discrete (individually separate) group of atoms with a known formula, bonded together. The molecular formula tells you the number and type of atoms in a molecule. For example, $\text{H}_2\text{O}$ shows that a water molecule contains two hydrogen atoms and one oxygen atom.

The octet rule

Most atoms become more stable when they have eight electrons in their outer shell. This is called the octet rule. When non-metal atoms bond covalently, they share electrons so that each atom achieves eight electrons in its outer shell (or two electrons for hydrogen, which has only one shell).

The covalent bonds that form between atoms within a molecule are called intramolecular bonds (bonds within a molecule).

Single covalent bonds

A single covalent bond forms when two atoms share one pair of electrons (two electrons total), with each atom contributing one electron.

Hydrogen molecules

Hydrogen atoms have just one electron. The outer shell of a hydrogen atom can hold a maximum of two electrons, so a hydrogen atom needs one more electron to achieve a stable configuration.

When two hydrogen atoms approach each other, they can each share their single electron. This creates a molecule of hydrogen gas, $\text{H}_2$.

A molecule containing two atoms is called a diatomic molecule.

You can represent a hydrogen molecule in different ways:

• Molecular formula: $\text{H}_2$
• Structural formula: $\text{H}-\text{H}$ (the line represents the single bond)
• Electron dot diagram: showing the shared electrons

In a hydrogen molecule, each electron is attracted to both protons (one in each nucleus). The two electrons spend most of their time between the two nuclei. Even though the two positively charged nuclei repel each other, the electrostatic attraction to the electrons (which are closer) is stronger and holds the molecule together.

Chlorine molecules

A chlorine atom has the electronic configuration $2,8,7$. It needs one more electron to achieve eight electrons in its outer shell.

When two chlorine atoms approach each other, each can share one electron with the other. This forms a chlorine molecule, $\text{Cl}_2$, with a single covalent bond.

After bonding:

• Both chlorine atoms have eight electrons in their outer shells
• One pair of electrons is shared between them (the bonding pair)
• Each chlorine atom also has six electrons not involved in bonding (three lone pairs)

The Hindenburg disaster

Lewis structures

Lewis structures (also called electron dot structures) are a useful way to represent molecules. They show the valence shell electrons of atoms, which are the only electrons involved in bonding.

For a chlorine molecule:

• There is one bonding pair (the shared electrons)
• Each chlorine atom has three lone pairs (six non-bonding electrons)

You can represent electrons in Lewis structures using dots, crosses, lines, or a combination. For example, these all show a chlorine molecule:

$\text{:Cl:Cl:}$ $\text{:Cl}-\text{Cl:}$ $\text{:Cl} \times \text{Cl:}$

Double covalent bonds

A double covalent bond forms when two pairs of electrons (four electrons total) are shared between two atoms.

Oxygen molecules

An oxygen atom has the electronic configuration $2,6$. It needs two more electrons to achieve eight electrons in its outer shell.

When two oxygen atoms bond, each atom shares two of its electrons with the other. This creates a double covalent bond in the oxygen molecule, $\text{O}_2$.

The Lewis structure of oxygen shows the double bond:

$\text{:O}=\text{O:}$

or

$\text{:}\ddot{\text{O}}::\ddot{\text{O}}\text{:}$

Triple covalent bonds

A triple covalent bond forms when three pairs of electrons (six electrons total) are shared between two atoms.

Nitrogen molecules

A nitrogen atom has the electronic configuration $2,5$. It needs three more electrons to achieve eight electrons in its outer shell.

When two nitrogen atoms bond, each atom shares three of its electrons with the other. This creates a triple covalent bond in the nitrogen molecule, $\text{N}_2$.

The Lewis structure of nitrogen shows the triple bond:

$\text{:N}\equiv\text{N:}$

Nitrogen's strength and importance

The triple covalent bond in $\text{N}_2$ is very strong and difficult to break. This makes nitrogen gas relatively unreactive. Although nitrogen makes up 78% of the air and is essential for life (it's a major component of proteins), very few organisms can use nitrogen gas directly because it's so unreactive.

Molecular compounds

Covalent bonds don't just form between atoms of the same element. They can also form between atoms of different elements to create molecular compounds.

Hydrogen chloride

Hydrogen chloride ($\text{HCl}$) is a simple molecular compound. A hydrogen atom needs one electron to fill its outer shell, and a chlorine atom also needs one electron. They can each share one electron to form a single covalent bond.

In the hydrogen chloride molecule:

• One electron pair is shared between the hydrogen and chlorine atoms
• The hydrogen atom has two electrons in its outer shell (full)
• The chlorine atom has eight electrons in its outer shell (full)
• The chlorine atom also has three lone pairs

Polyatomic molecules

Polyatomic molecules contain more than two atoms. Let's look at three important examples.

Water ($\text{H}_2\text{O}$)

When hydrogen and oxygen combine, one oxygen atom bonds with two hydrogen atoms.

Methane ($\text{CH}_4$)

A carbon atom has the electronic configuration $2,4$. It needs four more electrons to achieve eight electrons in its outer shell. When carbon combines with hydrogen, four hydrogen atoms bond to one carbon atom.

Ethene ($\text{C}_2\text{H}_4$)

Ethene is another compound of carbon and hydrogen. In this molecule:

• Each carbon atom shares two electrons with the other carbon atom, forming a double bond
• Each carbon atom also shares one electron with each of two hydrogen atoms
• There are one double covalent bond and four single covalent bonds

How to draw Lewis structures

Follow these steps to draw a Lewis structure:

Step 1: Write the electronic configuration of each atom in the molecule

Step 2: Determine how many electrons each atom needs for a stable outer shell

Step 3: Draw the Lewis structure, ensuring:

• Each atom has a stable outer shell
• Shared electrons are shown between atoms
• Non-bonding electrons are shown as lone pairs

Worked example: Hydrogen sulfide ($\text{H}_2\text{S}$)

Historical development of bonding theories

Before the twentieth century, scientists had limited understanding of chemical bonds. An early theory suggested atoms had 'hooks' and 'eyes' that allowed them to connect!

By the nineteenth century, chemists knew atoms combined in set proportions. Friedrich Kekule and Archibald Scott Couper proposed that carbon formed four bonds, and Kekule famously proposed the structure for benzene ($\text{C}_6\text{H}_6$).

Remember!