Covalent Bonding: Sharing Electrons Explained

Most of the substances that make up living things — water, sugar, proteins, DNA — are held together by covalent bonds. So are plastics, fuels, medicines and the gases in the air. Unlike ionic bonding, where one atom hands electrons to another, covalent bonding is about sharing. This guide explains how and why atoms share electrons, how to predict the number of bonds each atom forms, and how covalent structures explain the properties of substances.

What is a covalent bond?

A covalent bond is a shared pair of electrons between two atoms. Both nuclei are attracted to the shared pair, and that mutual attraction holds the atoms together.

Covalent bonds usually form between non-metal atoms. Non-metals all attract electrons fairly strongly (they have high electronegativities), so none can easily take electrons from another. Sharing is the lower-energy option.

The simplest example: hydrogen

A hydrogen atom has one electron. When two hydrogen atoms approach, their electron clouds overlap, and the two electrons pair up in the region between the nuclei:

H• + •H → H:H, written H–H

In the H₂ molecule, each nucleus is attracted to both electrons. The bond is strong: about 436 kJ of energy must be supplied to break one mole of H–H bonds. Each hydrogen now effectively has two electrons around it — the same as helium.

Why sharing lowers the energy

In a single hydrogen atom, the electron is attracted to one nucleus. In H₂, each electron is attracted to two nuclei and spends much of its time between them, shielding the nuclei from each other. The attraction outweighs the repulsions between the nuclei and between the electrons, up to a point. The atoms settle at the distance where the energy is lowest — the bond length, 74 pm for H₂ (see what is a chemical bond?).

How many bonds does each atom form?

For many non-metals, the number of covalent bonds equals the number of electrons needed to reach a full outer shell (eight, or two for hydrogen):

Element	Group	Outer electrons	Electrons needed	Usual number of bonds
Hydrogen	1	1	1	1
Carbon	14	4	4	4
Nitrogen	15	5	3	3
Oxygen	16	6	2	2
Halogens (F, Cl, Br, I)	17	7	1	1

This is called the valency of the atom. It lets you predict formulas:

• Carbon (4 bonds) + hydrogen (1 bond each) → CH₄, methane.
• Nitrogen (3) + hydrogen → NH₃, ammonia.
• Oxygen (2) + hydrogen → H₂O, water.
• Hydrogen + chlorine → HCl, hydrogen chloride.

Lone pairs

Not every outer electron takes part in bonding. In ammonia, nitrogen has five outer electrons: three are shared with hydrogen atoms, and the remaining two form a lone pair (a non-bonding pair). In water, oxygen has two lone pairs.

Lone pairs matter a lot:

• They affect the shape of molecules (see VSEPR and molecular geometry). Water is bent, not straight, because of its two lone pairs.
• They can be donated to form dative (coordinate) bonds — for example, when ammonia accepts H⁺ to form NH₄⁺ (see dative covalent bonds).
• They make molecules act as bases and nucleophiles in reactions.

Double and triple bonds

Atoms can share more than one pair:

• Double bond (two shared pairs): O=O in oxygen; C=O twice in carbon dioxide, O=C=O.
• Triple bond (three shared pairs): N≡N in nitrogen; H–C≡C–H in ethyne.

More shared pairs pull the atoms closer and hold them more tightly:

Bond	Shared pairs	Bond length (pm)	Average bond energy (kJ mol⁻¹)
C–C	1	154	347
C=C	2	134	614
C≡C	3	120	839

A double bond isn’t simply twice as strong as a single bond, because the second and third pairs are shared in a different way (π bonds, which overlap less effectively — see sigma and pi bonds). More in single, double and triple bonds compared.

Drawing covalent molecules

There are three common ways to represent covalent molecules:

1. Dot-and-cross diagrams show where each electron came from, with the electrons of one atom as dots and the other as crosses (see how to draw dot-and-cross diagrams).
2. Lewis structures show bonding pairs as lines and lone pairs as pairs of dots (see Lewis dot structures).
3. Displayed formulas show every bond as a line but usually leave out lone pairs.

Polar and non-polar covalent bonds

When two identical atoms share electrons (H–H, Cl–Cl), the pair is shared equally: a non-polar bond.

When two different atoms share, the more electronegative one pulls the pair closer. In H–Cl, chlorine (electronegativity 3.16) attracts the pair more than hydrogen (2.20), so chlorine carries a small negative charge (δ−) and hydrogen a small positive charge (δ+). This is a polar covalent bond (see bond polarity). Polar bonds help explain why water is such a good solvent and why HCl dissolves to form an acid.

Two kinds of covalent structure

Simple molecular substances

Most covalent substances are made of small, separate molecules — water, carbon dioxide, methane, iodine, sugar. Inside each molecule, atoms are held by strong covalent bonds. But between molecules there are only weak intermolecular forces (see intermolecular forces).

So when a simple molecular substance melts or boils, only the weak forces between molecules are overcome — no covalent bonds break. That’s why these substances usually have:

• Low melting and boiling points (many are gases or liquids at room temperature).
• No electrical conductivity, because the molecules have no overall charge and there are no free electrons.

Larger molecules have stronger intermolecular forces, so their boiling points are higher: methane (CH₄) boils at −162 °C, while octane (C₈H₁₈), a component of petrol, boils at about 126 °C.

Giant covalent structures

In some substances, covalent bonds link huge numbers of atoms into a continuous network:

• Diamond — each carbon atom bonded to four others in a rigid 3D network. Extremely hard; melts (or sublimes) only above about 3,500 °C.
• Graphite — carbon atoms in flat layers of hexagons, each bonded to three others. The fourth electron from each atom is delocalised along the layers, so graphite conducts electricity. Weak forces between layers let them slide, making graphite soft and slippery.
• Silicon dioxide (quartz, sand) — each silicon bonded to four oxygens and each oxygen to two silicons.

Melting a giant covalent structure means breaking many strong covalent bonds, so the melting points are very high.

Covalent vs ionic at a glance

Feature	Covalent	Ionic
Elements involved	Non-metals	Metal + non-metal
Electrons	Shared	Transferred
Particles	Molecules (or giant networks)	Ions in a lattice
Typical melting point	Low (simple) / very high (giant)	High
Conducts when molten?	No (usually)	Yes

See ionic vs covalent bonds for a fuller comparison, and ionic bonding explained for the other side of the story.

Common misconceptions

• “Covalent bonds are weak.” Individual covalent bonds are strong; it’s the forces between simple molecules that are weak.
• “Melting ice breaks the bonds in water.” It breaks hydrogen bonds between molecules; the O–H bonds stay intact.
• “All covalent substances have low melting points.” Giant covalent structures like diamond have extremely high melting points.
• “Atoms share electrons equally.” Only identical atoms share exactly equally; most bonds are polar to some degree.

Key takeaways

• A covalent bond is a shared pair of electrons, attracted to both nuclei.
• Covalent bonds form mainly between non-metals.
• The usual number of bonds: H 1, O 2, N 3, C 4, halogens 1.
• Atoms can share one, two or three pairs; more pairs mean shorter, stronger bonds.
• Simple molecular substances have low melting points (weak intermolecular forces); giant covalent substances have very high ones.