Covalent Bond

Covalent bond definition

In "Bonding", we learned that atoms like being in the lowest energy state possible. This is when they are at their most stable. For the majority of atoms, this involves having a full outer shell of electrons - or to be precise, eight valence electrons. There are different ways for atoms to achieve a full outer shell, but non-metals find it easiest to reach this stable state by sharing electrons with each other. This is known as a covalent bond.

Covalent bonds form between two non-metal atoms through the overlap of some of their outer shell electrons. It usually results in the atoms having full outer shells. This gives them the electron configuration of a noble gas, which is a more stable electron arrangement. The atoms are held together by strong electrostatic attraction between the positive nuclei and the shared pair of electrons, which is also known as the bonded pair. In contrast, electron pairs that aren't involved in covalent bonding are known as lone pairs.

Covalent bond diagrams

There are two ways of representing covalent bonds. For a more detailed view of the bonding, we use dot and cross diagrams. For a simpler representation, we use displayed formulae.

Dot and cross diagrams

We can show covalent bonds using dot and cross diagrams. They have some important features:

• Dot and cross diagrams show the outer shell of each atom and the electrons within it.
• Electrons from one atom are shown using dots, and electrons from the second are shown using crosses. If we have additional atoms, we might choose to use other symbols like triangles or stars to avoid confusion.
• Electrons are typically drawn in pairs, equally spaced around the atom.
• We show covalent bonds by overlapping the electron shells of two atoms and drawing a pair of electrons within the overlap.

Let’s work through an example together.

Displayed formulae

Dot and cross diagrams can become quite time-consuming to draw for larger molecules. We can show covalent bonding much more easily by simply drawing the chemical symbol of each atom and showing the covalent bonds between them using straight lines. Lone pairs of electrons are generally omitted, but they can be included if they are particularly relevant to the species. This style of representing molecules is known as displayed formulae.

Types of covalent bonds

Covalent bonds all have one thing in common: a shared pair of electrons. But within the field of covalent bonding, there are a few different types of bonding. These include:

• Dative covalent bonds.
• Single, double and triple bonds.
• Sigma and pi bonds.

Let's explore them now.

Single, double and triple bonds

Some atoms need just one shared pair of electrons in order to complete their outer shell. An example is the chlorine molecule, Cl₂, which we explored earlier. Chlorine atoms have seven valence electrons, but if two chlorine atoms share an electron with each other they’ll both have eight, completing their outer shells.

However, some atoms require more than one shared pair of electrons in order to complete their outer shell. To do this, they can form multiple single bonds with different atoms, or a double or triple bond with the same atom.

• In a single covalent bond, two atoms share one pair of electrons. One electron comes from each atom.
• In a double covalent bond, two atoms share two pairs of electrons. Two electrons come from each atom; in total there are four shared electrons.
• In a triple covalent bond, two atoms share three pairs of electrons. Three electrons come from each atom; in total there are six shared electrons.

To show double and triple bonds in dot and cross diagrams, you simply increase the number of electrons found within the overlapping electron shells. For example, a double bond contains two dots and two crosses, making a total of four electrons. To show double and triple bonds in displayed formulae, you simply draw a double or triple line respectively.

Single, double and triple bonds also vary in length and energy:

The relative energy and length of single, double and triple covalent bonds. StudySmarter Originals

Sigma and pi bonds

Do you remember how we said that covalent bonds involve an overlap of atomic orbitals? Well, the orbitals can overlap in different ways, and this creates two further ways of classifying covalent bonds:

• Sigma bonds are formed by the end-to-end overlapping of s-orbitals or sp-hybridised orbitals. All single covalent bonds are sigma bonds.
• Pi bonds are formed by the sideways overlapping of p-orbitals. The second and third bonds in double and triple bonds are pi bonds.

Pi bonds are much weaker than sigma bonds. However, because a double bond contains a sigma bond and a pi bond, it is overall significantly stronger than a single sigma bond on its own. Likewise, because triple bonds contain a sigma bond and two pi bonds, they are stronger still.

Dative covalent bonds

Finally, let's consider dative covalent bonds.

In our example of Cl₂ above, two atoms came together and formed a covalent bond by each sharing one electron. But sometimes both electrons in the bonded pair can come from the same atom. This is known as a dative covalent bond.

To form a dative bond, you need a species with a lone pair of electrons and a species with an empty electron orbital. The species with the lone pair of electrons offers up both of its electrons to form the bonded pair, whilst the species with the empty orbital doesn't share any of its electrons.

You can identify dative covalent bonds in dot and cross diagrams as both electrons come from the same atom - instead of one dot and one cross, you'll see two dots, or two crosses. We show these bonds in displayed formulae using an arrow drawn from the donor species towards the receiving species. However, dative bonds are exactly the same as regular covalent bonds in all other regards - they are the same length and have the same properties.

Covalent bond properties

Covalent bonds are very strong. They're held together by strong electrostatic attraction between the shared pair of electrons and the two atomic nuclei, which requires a lot of energy to overcome. But different covalent species contain different numbers and arrangements of covalent bonds, and this changes their properties. For example, the two most abundant elements that make up the Earth, oxygen and silicon, both contain covalent bonds, but in their elemental form they are structured in very different ways. Whilst oxygen atoms go around in pairs as simple covalent molecules, silicon atoms make up huge crystal structures of indeterminate size known as a giant covalent macromolecule. Their contrasting structures give them both different properties:

• Simple covalent molecules are made up of a small number of atoms covalently bonded together. Although the covalent bonds themselves are strong, the forces between the individual molecules are weak and don’t require much energy to overcome. This gives simple covalent molecules low melting and boiling points.
• Macromolecules, also known as giant covalent structures, are huge lattices, made of many atoms joined together by multiple covalent bonds in all directions. They have high melting and boiling points as all of their covalent bonds are extremely strong and require a lot of energy to overcome. For this same reason, they are hard and strong.

An example of a simple covalent molecule and a giant covalent macromolecule. StudySmarter Originals

Covalent bond examples

Throughout this article, we've included lots of examples of covalent bonds - from single, double and triple bonds, to dative covalent bonds and different covalent structures. But if you want to see more covalent bonds in action, head over to "Examples of Covalent Bonding", where you'll find plenty more molecules. You can also discover some giant covalent macromolecules in the article "Carbon Structures".