Covalent bonding: what is a covalent bond?

What is covalent bonding?

Covalent bonding is the sharing of electrons between atoms. It is one of the three types of strong chemical bond, alongside ionic and metallic bonding, as the AQA GCSE specification puts it. Atoms are most stable with a full outer shell of electrons. Hydrogen's first shell holds 2, and the next shells of most small atoms hold 8. Atoms bond to reach that.

Take two chlorine atoms. Each has 7 outer electrons and needs 1 more. Neither can take an electron from the other, because both want one. So each puts one electron into a shared pair. Now both atoms count that pair as part of their outer shell, so both have 8, and the two atoms are joined as a Cl₂ molecule.

Why the shared pair holds atoms together

Electrons are negative and nuclei are positive. The shared pair sits between the two nuclei and is attracted to both of them at once. That electrostatic attraction is the bond. It is the same force that holds ionic compounds together, just arranged differently: in a covalent bond, the attraction is between two nuclei and the electrons they share.

How many covalent bonds does an atom form?

Usually one bond for each electron the atom needs to fill its outer shell. You can read this from the group number of the periodic table.

This is why water is H₂O: oxygen needs 2 electrons, so it bonds to 2 hydrogen atoms. Ammonia is NH₃ and methane is CH₄ for the same reason. You can build these in our Molecule Builder and see which arrangements work.

Dot and cross diagrams

A dot and cross diagram shows how the outer electrons are shared. One atom's electrons are drawn as dots and the other's as crosses, and the shared pair sits where the circles of the two outer shells overlap. Only outer electrons are drawn, because inner electrons take no part in bonding.

Hydrogen chloride, HCl (outer shells only)

         x x
   H  •x  Cl  x
         x x

•x  = the shared pair: one electron from H (•), one from Cl (x)
x x = Cl's other six outer electrons, in three lone pairs

Outer electrons that are paired up but not shared are called lone pairs. GCSE courses expect you to draw these eight molecules:

Other ways to show covalent bonds

A displayed formula draws each bond as a line, so water is H−O−H. A ball and stick model shows atoms as balls and bonds as sticks, and a 3D model shows the real shape. Each has limits: dot and cross diagrams don't show shape, displayed formulas don't show lone pairs, and ball and stick models make the atoms look far apart when the electron clouds really overlap.

Single, double and triple covalent bonds

Atoms can share more than one pair. An oxygen atom needs 2 electrons, so two oxygen atoms share 2 pairs: a double bond, written O=O. A nitrogen atom needs 3, so N₂ has a triple bond, three shared pairs, written N≡N. Carbon dioxide has two double bonds, O=C=O, because carbon needs 4 electrons and each oxygen supplies 2.

More shared pairs pull the atoms closer and hold them more tightly. A triple bond is shorter and stronger than a double bond between the same two atoms, which is stronger than a single bond. The triple bond in N₂ is the reason nitrogen gas is so unreactive.

Covalent vs ionic bonds

Try it: this question gives you a substance and asks which type of bonding it has.

Polar and nonpolar covalent bonds

Sharing isn't always equal. Electronegativity is "the power of an atom to attract the pair of electrons in a covalent bond", in the words of the AQA A-level specification. When two identical atoms bond, as in Cl₂, they pull equally and the bond is nonpolar. When one atom is more electronegative, the shared pair sits closer to it, and the bond is polar.

In water, oxygen is much more electronegative than hydrogen. The oxygen end of each bond gets a small negative charge, written δ−, and each hydrogen a small positive charge, δ+. Water's bent shape means these charges don't cancel out, so the whole molecule is polar, which explains why it dissolves salt and has an unusually high boiling point for its size. If the difference in electronegativity is very large, one atom takes the electrons completely, and the bond is ionic instead.

Shapes of covalent molecules

Electron pairs around an atom repel each other and spread out as far as they can, which sets the shape. Methane's four bonding pairs point to the corners of a tetrahedron at 109.5°. Ammonia's lone pair pushes its three bonds a little closer, to about 107°, and water's two lone pairs push its bonds to about 104.5°, giving a bent shape (OpenStax Chemistry covers this as VSEPR theory). Our Chemical Bonding course lets you predict the shapes step by step.

Dative covalent bonds

In a normal covalent bond, each atom gives one electron to the shared pair. In a dative covalent bond, also called a coordinate bond, one atom gives both. The AQA A-level specification defines it as a bond that "contains a shared pair of electrons with both electrons supplied by one atom". The atom giving the pair must have a lone pair, and the other atom must have room for two more electrons.

The standard example is the ammonium ion, NH₄⁺. Ammonia's nitrogen has a lone pair. A hydrogen ion, H⁺, has no electrons at all. The nitrogen shares its lone pair with the H⁺, forming a fourth N−H bond. Once formed, all four N−H bonds are identical; the dative bond only describes where the electrons came from. The oxonium ion, H₃O⁺, forms the same way from water. A dative bond is drawn as an arrow pointing from the atom that gave the pair: H₃N→H⁺.

Covalent structures and their properties

Covalent bonds are strong, but covalent substances can still melt at very low temperatures. It depends on the structure.

Simple molecular substances

Water, carbon dioxide, methane and oxygen are made of small, separate molecules. The bonds inside each molecule are strong, but the forces between molecules (intermolecular forces) are weak. Melting or boiling only has to overcome those weak forces, not the covalent bonds, so these substances have low melting and boiling points and are usually gases or liquids at room temperature. Bigger molecules have stronger intermolecular forces, so they boil at higher temperatures. They don't conduct electricity, because the molecules have no overall charge.

Giant covalent structures

In a giant covalent structure, every atom is joined to its neighbours by covalent bonds in a huge network. To melt it, you have to break the covalent bonds themselves, so these substances are solids with very high melting points.

• Diamond: each carbon atom forms 4 covalent bonds to other carbon atoms. Diamond is very hard, has a very high melting point and doesn't conduct electricity.
• Graphite: each carbon atom forms 3 bonds, making flat layers of hexagons with no covalent bonds between the layers. The layers slide over each other, so graphite is soft and slippery. The fourth electron from each carbon is delocalised and free to move, so graphite conducts electricity.
• Silicon dioxide (silica): the main compound in sand, a giant network of silicon and oxygen atoms.

Polymers such as poly(ethene) are a third kind: very long molecules held together by covalent bonds along the chain.

Common mistakes

• Saying covalent bonds are weak. They are strong. It is the forces between simple molecules that are weak.
• Saying covalent bonds break when water boils. Boiling separates whole molecules; the O−H bonds stay intact.
• Drawing inner electrons. Dot and cross diagrams for bonding show outer shells only.
• Forgetting lone pairs. Chlorine in HCl still has 3 lone pairs, and they should be drawn.
• Giving hydrogen 8 electrons. Hydrogen's shell is full with 2.
• Saying graphite conducts because of ions. It conducts because of delocalised electrons.

Read next

Chemistry · Bonding and structureWhat is an ionic bond? Ionic bonding, lattices, propertiesWhat an ionic bond is, how ions form, dot and cross diagrams, working out formulae, giant ionic lattices and the properties of ionic compounds.