Covalent Bond
Summary: A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms, formed between non-metal atoms to achieve noble gas electronic configurations. Tags: igcse chemistry bonding covalent Created: 2026-07-18
A covalent bond is formed when a pair of electrons is shared between two atoms, leading both atoms to achieve the stable electronic configuration of a noble gas (a full outer shell). Unlike ionic bonding, where electrons are transferred from a metal to a non-metal, covalent bonding involves the mutual sharing of electrons and occurs exclusively between non-metal atoms. Each shared pair of electrons is attracted to the nuclei of both atoms, creating a strong electrostatic attraction that holds the atoms together. The number of covalent bonds an atom forms is determined by how many additional electrons it needs to complete its outer shell: atoms in Group 17 typically form one bond, Group 16 form two, Group 15 form three, and Group 14 form four. For a broader discussion including simple molecular and giant covalent structures, see Covalent Bonding.
Definition
A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms. It forms between non-metal atoms, with each atom contributing one or more electrons to the shared pair. This sharing allows both atoms to achieve a full outer shell (noble gas configuration). The bond is directional and holds the two atoms at a fixed distance where attraction and repulsion are balanced.
Exam point: Covalent bond = shared pair of electrons + electrostatic attraction to both nuclei.
Bonding Pairs and Lone Pairs
In a covalent bond, the shared electrons are called a bonding pair. Electrons in the outer shell that are not involved in bonding are called lone pairs (or non-bonding pairs). Lone pairs occupy space around the atom and influence molecular shape, but they do not directly hold atoms together.
| Term | Description | Example in H2O |
|---|---|---|
| Bonding pair | A pair of electrons shared between two atoms | The two O—H bonds (2 bonding pairs) |
| Lone pair | A pair of outer-shell electrons belonging to only one atom | The two non-bonded pairs on the oxygen atom |
In a water molecule (H2O), oxygen has six outer electrons. Two are used in bonding pairs (one shared with each hydrogen), and four remain as two lone pairs. The bonding pairs give each hydrogen access to 2 electrons (duplet), while oxygen achieves 8 outer electrons (octet) — its Noble Gas Electronic Configuration.
Single, Double, and Triple Bonds
Covalent bonds can involve one, two, or three shared pairs of electrons. More shared pairs produce a shorter, stronger bond.
| Bond Type | Shared Pairs | Electrons | Bond Strength | Example | Structural Notation |
|---|---|---|---|---|---|
| Single | 1 | 2 | Weakest | Cl—Cl, H—H, C—H | Single line (— ) |
| Double | 2 | 4 | Intermediate | O=O, C=O | Double line (=) |
| Triple | 3 | 6 | Strongest | N≡N | Triple line (≡) |
Chlorine (Cl2): Each Cl atom has 7 outer electrons and needs 1 more for an octet. Sharing one pair gives both an effective count of 8. This is a single covalent bond: Cl—Cl.
Oxygen (O2): Each O atom has 6 outer electrons and needs 2 more. Sharing two pairs forms a double bond: O=O. Each O now has 8 outer electrons.
Nitrogen (N2): Each N atom has 5 outer electrons and needs 3 more. Sharing three pairs forms a triple bond: N≡N. This bond is very strong, which explains why nitrogen gas is so unreactive at room temperature.
Carbon dioxide (CO2): Carbon (4 outer electrons) needs 4 more for an octet. It forms two double bonds, one to each oxygen atom: O=C=O. Each oxygen shares two pairs with carbon, giving all three atoms full outer shells.
The bond strength order — triple > double > single — means more energy is required to break a triple bond than a double bond, and more for a double than a single. This is directly relevant to Bond Energy Calculations.
Covalent Bond Formation: The Octet and Duplet Rules
Covalent bonds form so that each atom achieves a stable noble gas electronic configuration:
- Octet rule: Most atoms (C, N, O, Cl, etc.) achieve 8 electrons in their outer shell, matching the configuration of neon or argon.
- Duplet rule: Hydrogen (and sometimes lithium) achieves 2 electrons in its outer shell, matching the configuration of helium.
The number of covalent bonds an atom forms is typically 8 minus its group number (for Groups 14—17):
| Element | Group | Outer Electrons | Electrons Needed | Covalent Bonds Formed |
|---|---|---|---|---|
| Carbon | 14 | 4 | 4 | 4 |
| Nitrogen | 15 | 5 | 3 | 3 |
| Oxygen | 16 | 6 | 2 | 2 |
| Chlorine | 17 | 7 | 1 | 1 |
| Hydrogen | (not 4—17) | 1 | 1 (duplet) | 1 |
Occurrence: Non-Metal with Non-Metal
Covalent bonding occurs between atoms of non-metal elements. This is because non-metals have relatively high electronegativity values — neither atom can fully remove an electron from the other (as in ionic bonding). Instead, they share.
Common non-metal pairs that form covalent bonds:
- Two identical non-metal atoms: H2, Cl2, O2, N2 (diatomic elements)
- Two different non-metal atoms: HCl, CO2, SO2, H2O, NH3, CH4
- Extended non-metal networks: the C—C and C—H bonds in organic compounds
When a non-metal bonds with a metal, the bonding is typically ionic (electron transfer), not covalent.
Representing Covalent Bonds: Dot-and-Cross Diagrams
A Dot-and-Cross Diagram shows the outer-shell electrons of atoms, using dots for one atom’s electrons and crosses for the other. Shared pairs appear in the overlapping region between the two atoms. The diagram must show each atom (except hydrogen) surrounded by 8 electrons when bonded, counting both shared and lone pairs. The hydrogen atom must be surrounded by 2 electrons (the duplet).
Electron Pair
An electron pair is two electrons occupying the same molecular orbital but with opposite spins. In a covalent bond, the electron pair is shared between the two bonded atoms. Electron pairs in a molecule are either bonding pairs (shared between two atoms, forming the covalent bond) or lone pairs (not shared, residing on a single atom).
Exam point: A single covalent bond = one shared pair of electrons; a double bond = two shared pairs; a triple bond = three shared pairs.
Electron Sharing
Electron sharing is the fundamental mechanism by which covalent bonds form between non-metal atoms. Each atom contributes one or more electrons to a shared pair, and both atoms are attracted to the shared electrons. This allows each atom to achieve a noble gas electronic configuration (a full outer shell). Equal sharing produces a non-polar covalent bond; unequal sharing produces a polar covalent bond (not required for IGCSE Core).
Exam point: Covalent bonding = sharing of electron pairs between non-metal atoms to achieve full outer shells.
Lone Pair
A lone pair is a pair of electrons in the outer shell of an atom that is not involved in bonding — it is not shared with another atom. Lone pairs occupy space around the central atom and influence the shape of the molecule (though molecular shape is not required at IGCSE level). Lone pairs are shown in dot-and-cross diagrams as paired dots or crosses on one atom only.
Exam point: In H₂O, oxygen has 2 lone pairs; in NH₃, nitrogen has 1 lone pair; in CH₄, carbon has 0 lone pairs.
Sources
- Cambridge IGCSE Chemistry 0620 Syllabus (2023-2025) — Section 2.5 (Simple molecules and covalent bonds), Cambridge Assessment International Education
Common Misconceptions
| Misconception | Reality |
|---|---|
| ”A covalent bond is the same as an intermolecular force” | A covalent bond is the strong attraction within a molecule (between atoms). An Intermolecular Force is the weak attraction between molecules. |
| ”Covalent bonds are weak because covalent substances have low melting points” | Covalent bonds themselves are strong. The low melting points of simple molecular substances are due to weak intermolecular forces, not weak covalent bonds. |
| ”All covalent bonds are single bonds” | Double bonds (O=O, C=O) and triple bonds (N≡N) involve two or three shared pairs of electrons respectively. |
| ”Lone pairs do not matter” | Lone pairs influence the shape of a molecule (e.g., the bent shape of H2O is due to the two lone pairs on oxygen repelling bonding pairs). |
| ”Atoms ‘own’ the shared electrons” | Shared electrons are attracted to the nuclei of both atoms simultaneously. Neither atom gains or loses them permanently. |