Just four simple elements make up over 96 percent of your body: oxygen, carbon, hydrogen, and nitrogen. Life without these elements would look completely different. They're found in all sorts of essential substances, from simple molecules to branched chains and tangled polymers. For example, carbon is the basis of most organic molecules, whilst oxygen and hydrogen combine to make water, and nitrogen is a fundamental part of all of our proteins. Scientists have proposed life forms based on other molecules - for example, ammonia-based life - but even these would require some of the above elements.
It is fair to say that life revolves around these four elements. But they're never found as single atoms. Instead, oxygen, carbon, hydrogen and nitrogen, alongside many other elements in the periodic table, bond together into the familiar complex structures that make up much of the world around us using covalent bonds.
- This article is about covalent bonds in physical chemistry.
- We'll define covalent bonds before looking at how you represent them in dot and cross diagrams.
- We'll then explore types of covalent bonds, including dative bonds, double and triple bonds, and sigma and pi bonds. This will involve an introduction to orbitals and hybridisation.
- To conclude, we'll learn about covalent bond properties and how this relates to covalent structures.
- Throughout this article, you'll come across plenty of examples of covalent bonds to help consolidate your learning.
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.
A covalent bond is a shared pair of electrons.
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.
One important point to remember is that the atoms involved don’t form ions. You’ll remember that an ion is an atom that has gained or lost electrons to form a charged particle. Covalent molecules share their electrons instead of losing them and so remain neutral particles.
Some metals break the trend - they're able to bond covalently too. For example, beryllium can join to chlorine with covalent bonds to form beryllium chloride, BeCl2.
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.
Draw a dot and cross diagram for a chlorine molecule, Cl2.
A chlorine atom has seven electrons in its outer shell. To achieve a full outer shell of electrons it needs to gain an extra electron. Two chlorine atoms can do this by coming together, and each sharing one of their electrons through the overlap of their atomic orbitals. This forms a covalent bond. Because each chlorine atom now has an extra electron in its outer shell, both atoms have a noble gas electron configuration and are more stable.
To draw this on a dot and cross diagram, we show the outer electron shells of the chorine atoms as circles. Because the atoms are covalently bonded, we overlap the circles slightly. Each chlorine atom has seven of its own electrons; we draw six of them as three lone pairs around the edge of each circle. The remaining electron from each atom is part of the bonding pair and so is drawn in the overlap between the circles. This shows a shared pair of electrons: a covalent bond.
A dot and cross diagram of a chlorine molecule. Vaia Originals
Some elements can form stable molecules with more than eight electrons in their outer shell. An example is xenon, which often has 10 valence electrons. This is known as an expanded octet. Conversely, other elements are stable with fewer than eight electrons in their outer shell. An example is hydrogen, which likes having only two valence electrons.
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.
Draw the displayed formula of a chlorine molecule, Cl2.
We now know the covalent bonding within a chlorine molecule. To show this using displayed formulae, we simply represent the two atoms using their chemical symbols and draw the bond between them using a straight line. The lone pairs of electrons aren't that important here, and so we leave them out:
The displayed formula of a chlorine molecule. Vaia Originals
Check out "Organic Compounds" for more about the different types of formulae used in chemistry.
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.
Oxygen, O2, contains a double bond, whilst nitrogen, N2, contains a triple bond. Show these molecules using both dot and cross diagrams and displayed formulae.
Oxygen has six valence electrons. It can reach a full outer shell by sharing two pairs of electrons with another oxygen atom, forming a double covalent bond:
A dot and cross diagram and the displayed formula of an oxygen molecule. Vaia Originals
Nitrogen, on the other hand, has five valence electrons. It can reach a full outer shell by sharing three pairs of electrons with another nitrogen atom, forming a triple covalent bond:
A dot and cross diagram and the displayed formula for a nitrogen molecule. Vaia Originals
Single, double and triple bonds also vary in length and energy:
The relative energy and length of single, double and triple covalent bonds. Vaia 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.
You can find out about orbitals and hybridisation in the article with the same name, "Orbitals and Hybridisation".
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.
A dative covalent bond is a type of covalent bond, where both of the electrons in the shared pair come from the same atom. It is also known as a coordinate 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.
The ammonium ion, NH4+, contains a dative covalent bond. Show this using both a dot and cross diagram and a displayed formula.
Nitrogen has five valence electrons. It needs to form three covalent bonds in order to achieve a full outer shell, and it does this by forming three single bonds with three different hydrogen atoms. This leaves nitrogen with one lone pair of electrons. It uses these to bond to a hydrogen ion, H+, which has an empty electron orbital. This forms a dative covalent bond. Here, nitrogen provides both of the electrons in the bonded pair:
A dot and cross diagram and the displayed formula of an ammonium ion. Vaia Originals
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. Vaia OriginalsYou can compare simple covalent molecules and giant covalent macromolecules in the article "Physical Properties of Lattice Structures", which also contrasts them with giant ionic and metallic lattices.
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".
Covalent Bond - Key takeaways
A covalent bond is a shared pair of electrons. It typically forms between two non-metals and results in both atoms having full outer shells. It is formed due to the overlapping of electron orbitals.
We can represent covalent bonds using dot and cross diagrams, which show the outer shell of electrons, or with displayed formulae.
Atoms can form single, double or triple bonds. Single bonds are the longest, whilst triple bonds have the most energy.
Sigma bonds are caused by the end-to-end overlapping of s- or sp-hybridised orbitals, whilst pi bonds are caused by the sideways overlapping of p-orbitals.
A dative covalent bond is a type of covalent bond where one species provides both of the shared electrons. It occurs between a species with a lone electron pair and a species with a vacant orbital.
Covalently bonded atoms can form simple covalent molecules or giant covalent macromolecules. They have different properties due to their numbers and arrangements of covalent bonds.
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