Covalent Bond Definition and Examples

Covalent Bond Definition and Example
A covalent bond is a type of chemical bond characterized by two atoms sharing valence electrons.

A covalent bond is a chemical bond between two atoms where they share one or more pairs of electrons. Usually, sharing electrons gives each atom a full valence shell and makes the resulting compound more stable than its constituent atoms are on their own. Covalent bonds usually form between nonmetals. Examples of covalent compounds include hydrogen (H2), oxygen (O2), carbon monoxide (CO), ammonia (NH3), water (H2O), and all organic compounds. There are compounds that contain both covalent and ionic bonds, such as potassium cyanide (KCN) and ammonium chloride (NH4Cl).

What Is a Covalent Bond?

Covalent bonding is one of the main types of chemical bonds, along with ionic and metallic bonds. Unlike these other bonds, covalent bonding involves the sharing of electron pairs between atoms. These shared electrons exist in the outer shell of the atom, the so-called valence shell.

The water molecule (H2O) is an example of compound with covalent bonds. The oxygen atom shares one electron with each of the two hydrogen atoms, forming two covalent bonds.

Octet Rule and Covalent Bonding

The concept of covalent bonding ties in with the octet rule. This rule states that atoms combine in such a way that each atom has eight electrons in its valence shell, resembling the electronic configuration of a noble gas. By sharing electrons through covalent bonding, atoms effectively fill their outer shells and satisfy the octet rule.

Covalent Bond vs Ionic and Metallic Bonds

Covalent bonds differ significantly from ionic and metallic bonds. Ionic bonds form when one atom gives up one or more electrons to another atom, forming ions that attract each other due to their opposite charges. Sodium chloride (NaCl) is an example of a compound with ionic bonds.

Metallic bonds, on the other hand, form between metal atoms. In these bonds, electrons are not shared or transferred between atoms but instead move freely in what is sometimes referred to as an “electron sea”. This fluidity of electrons gives metals their unique properties, such as electrical conductivity and malleability.

Types of Covalent Bonds

Covalent bonds are either polar covalent bonds or nonpolar covalent bonds.

A nonpolar covalent bond forms when two atoms with the same electronegativity share electrons equally, as in a molecule of hydrogen gas (H2).

A polar covalent bond, on the other hand, forms when the atoms involved in the bond have different electronegativities, resulting in unequal sharing of electrons. The atom with the higher electronegativity pulls the shared electrons closer, creating a region of slight negative charge, while the other atom becomes slightly positive. An example is water (H2O), where the oxygen atom is more electronegative than the hydrogen atoms.

Electronegativity and the Type of Bonding

Electronegativity is a measure of an atom’s tendency for attracting a bonding pair of electrons. The electronegativity values, proposed by Linus Pauling, range from around 0.7 to 4.0. The higher the electronegativity, the greater an atom’s attraction for bonding electrons.

When considering whether a bond is ionic or covalent, the difference in electronegativity between the two atoms a helpful guideline.

  1. If the electronegativity difference is greater than 1.7, the bond is ionic. This is because the more electronegative atom attracts the electron(s) so strongly that it effectively “steals” them from the other atom.
  2. If the electronegativity difference is less than 1.7 but greater than 0.5, the bond is polar covalent. The atoms do not share electrons equally. The more electronegative atom attracts the electron pair. This leads to a separation of charge, with the more electronegative atom carrying a slight negative charge and the other atom a slight positive charge.
  3. If the electronegativity difference is less than 0.5, the bond is nonpolar covalent. The atoms share the electron pair more or less equally.

However, these are just guidelines and there is no absolute cut-off value that cleanly separates ionic and covalent bonds. In reality, many bonds falling somewhere in between. Also, electronegativity is not the only factor that determines the type of bond formed. Other factors also play a role, including the size of the atoms, the lattice energy, and the overall structure of the molecule.

Single, Double, and Triple Bonds

Covalent bonds exist as single, double, or triple bonds. In a single covalent bond, two atoms share one pair of electrons. Hydrogen gas (H2 or H-H) has a single covalent bond, where each hydrogen atom shares its single electron with the other.

In a double bond, atoms share two pairs of electrons. A typical example is oxygen gas (O2 or O=O), where each oxygen atom shares two electrons with the other. A double bond is stronger than a single bond, but less stable.

Triple bonds involve the sharing of three pairs of electrons, as seen in nitrogen gas (N2 or N≡N). The triple bond is strongest, yet least stable.

Properties of Covalent Compounds

Compounds that have covalent bonds often share several common properties.

  • Low Melting and Boiling Points: Covalent compounds generally have lower melting and boiling points than ionic bonds due to the weaker forces of attraction between molecules.
  • Poor Conductivity: Most covalent compounds do not conduct electricity because they lack free-moving charges (such as ions or delocalized electrons) that are necessary for the flow of electrical current. There are exceptions, such as graphite, which conducts electricity due to the delocalization of its electrons. Thermal conductivity varies widely among covalent compounds. For example, diamond, a form of carbon with each carbon atom covalently bonded to four other carbon atoms, is one of the best known thermal conductors. In contrast, many other covalently bonded substances, like water or polymers, are relatively poor thermal conductors.
  • Insolubility in Water: Many covalent compounds are nonpolar and are not soluble in water. Water and ethanol are examples of polar covalent compounds that do dissolve ionic compounds and other polar compounds.
  • Solubility in Organic Solvents: While nonpolar covalent compounds don’t dissolve well in water, they often dissolve well in organic solvents like benzene or in nonpolar solvents such as carbon tetrachloride. This is due to the ‘like dissolves like’ principle, where polar substances dissolve polar substances, and nonpolar substances dissolve nonpolar substances.
  • Lower Density: Covalent compounds generally have lower densities than ionic compounds. This is because the atoms in covalently bonded substances are not packed as closely together as in ionic substances. As a result, they are lighter for their size.
  • Brittle Solids: When covalent compounds do form solids, they are generally brittle. They are not ductile or malleable. This is due to the nature of their bonds. If a layer of atoms gets shifted, it disrupts the network of covalent bonds and the substance breaks.


  • Atkins, Peter; Loretta Jones (1997). Chemistry: Molecules, Matter and Change. New York: W.H. Freeman & Co. ISBN 978-0-7167-3107-8.
  • Langmuir, Irving (1919). “The Arrangement of Electrons in Atoms and Molecules”. Journal of the American Chemical Society. 41 (6): 868–934. doi:10.1021/ja02227a002
  • Lewis, Gilbert N. (1916). “The Atom and the Molecule”. Journal of the American Chemical Society. 38 (4): 772. doi:10.1021/ja02261a002
  • Pauling, Linus (1960). The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry. ISBN 0-801-40333-2. doi:10.1021/ja01355a027
  • Weinhold, F.; Landis, C. (2005). Valency and Bonding. Cambridge University Press. ISBN 0521831288.