COVALENT BOND: CHARACTERISTICS, PROPERTIES AND EXAMPLES - CHEMISTRY - 2025

The covalent bonds are a type of bond between atoms forming molecules through sharing electron pairs. These bonds, which represent a fairly stable balance between each species, allow each atom to achieve the stability of its electronic configuration.

These bonds are formed in single, double or triple versions, and have polar and non-polar characters. Atoms can attract other species, thus allowing the formation of chemical compounds. This union can occur by different forces, generating a weak or strong attraction, ionic characters or electron exchange.

Covalent bonds are considered "strong" bonds. Unlike other strong bonds (ionic bonds), covalent ones usually occur in non-metallic atoms and in those that have similar affinities for electrons (similar electronegativities), making the covalent bonds weak and require less energy to break.

In this type of bond, the so-called Octet rule is usually applied to estimate the number of atoms to share: this rule states that each atom in a molecule requires 8 valence electrons to remain stable. Through sharing, they must achieve loss or gain of electrons between species.

characteristics

Covalent bonds are affected by the electronegative property of each of the atoms involved in the interaction of electron pairs; When you have an atom with considerably higher electronegativity than the other atom in the junction, a polar covalent bond will form.

However, when both atoms have a similar electronegative property, a nonpolar covalent bond will form. This happens because the electrons of the most electronegative species will be more bound to this atom than in the case of the one with the least electronegativity.

It is worth noting that no covalent bond is completely egalitarian, unless the two atoms involved are identical (and thus have the same electronegativity).

The type of covalent bond depends on the difference in electronegativity between species, where a value between 0 and 0.4 results in a non-polar bond, and a difference of 0.4 to 1.7 results in a polar bond (the Ionic bonds appear from 1.7).

Non-polar covalent bond

The non-polar covalent bond is generated when electrons are shared equally between atoms. This generally occurs when the two atoms have a similar or equal electronic affinity (same species). The more similar the electron affinity values ​​are between the atoms involved, the stronger the resulting attraction.

This usually occurs in gas molecules, also known as diatomic elements. Non-polar covalent bonds work with the same nature as polar ones (the atom with higher electronegativity will attract more strongly the electron or electrons of the other atom).

However, in diatomic molecules the electronegativities cancel out because they are equal, resulting in a charge of zero.

Non-polar bonds are crucial in biology: they help form the oxygen and peptide bonds that are seen in amino acid chains. Molecules with a high number of nonpolar bonds are usually hydrophobic.

Polar covalent bond

The polar covalent bond occurs when there is an unequal sharing of electrons between the two species involved in the union. In this case, one of the two atoms has a considerably higher electronegativity than the other, and for this reason it will attract more electrons from the junction.

The resulting molecule will have a slightly positive side (the one with the lowest electronegativity), and a slightly negative side (with the atom with the highest electronegativity). It will also have an electrostatic potential, giving the compound the ability to weakly bind to other polar compounds.

The most common polar bonds are those of hydrogen with more electronegative atoms to form compounds such as water (H ₂ O).

Properties

In the structures of covalent bonds, a series of properties are taken into account that are involved in the study of these bonds and help to understand this phenomenon of electron sharing:

Octet rule

The octet rule was formulated by the American physicist and chemist Gilbert Newton Lewis, although there were scientists who studied this before him.

It is a rule of thumb that reflects the observation that the atoms of the representative elements tend to combine in such a way that each atom reaches eight electrons in its valence shell, leading it to have an electronic configuration similar to noble gases. Lewis diagrams or structures are used to represent these junctions.

There are exceptions to this rule, for example in species with an incomplete valence shell (molecules with seven electrons such as CH ₃, and reactive species with six electrons such as BH ₃); it also happens in atoms with very few electrons, such as helium, hydrogen and lithium, among others.

Resonance

Resonance is a tool used to represent molecular structures and represent delocalized electrons where bonds cannot be expressed with a single Lewis structure.

In these cases, the electrons must be represented by various "contributing" structures, called resonant structures. In other words, resonance is that term that suggests the use of two or more Lewis structures to represent a particular molecule.

This concept is completely human, and there is no one or another structure of the molecule at any given time, but can exist in any version of it (or all) at the same time.

Furthermore, the contributing (or resonant) structures are not isomers: only the position of the electrons can differ, but not the atom nuclei.

Aromaticity

This concept is used to describe a cyclic, planar molecule with a ring of resonant bonds that exhibit greater stability than other geometric arrangements with the same atomic configuration.

Aromatic molecules are very stable, since they do not break easily nor do they usually react with other substances. In benzene, the prototype aromatic compound, conjugated pi (π) bonds are formed in two different resonant structures, which form a highly stable hexagon.

Sigma link

It is the simplest bond, in which two "s" orbitals join. Sigma bonds occur in all simple covalent bonds, and they can also occur in "p" orbitals, as long as they are looking at each other.

Bond pi (π)

This bond occurs between two "p" orbitals that are in parallel. They bind side by side (unlike sigma, which binds face to face) and form areas of electron density above and below the molecule.

Covalent double and triple bonds involve one or two pi bonds, and these give the molecule a rigid shape. Pi bonds are weaker than sigma bonds, since there is less overlap.

Types of covalent bonds

Covalent bonds between two atoms can be formed by a pair of electrons, but they can also be formed by two or even three pairs of electrons, so these will be expressed as single, double and triple bonds, which are represented by different types of unions (sigma and pi bonds) for each.

Single bonds are the weakest and triple bonds are the strongest; This occurs because triples have the shortest bond length (greater attraction) and the greatest bond energy (they require more energy to break).

Simple link

It is the sharing of a single pair of electrons; that is, each atom involved shares a single electron. This union is the weakest and involves a single sigma (σ) bond. It is represented by a line between the atoms; for example, in the case of the hydrogen molecule (H ₂):

H H

Double link

In this type of bond, two shared pairs of electrons form bonds; that is, four electrons are shared. This bond involves one sigma (σ) and one pi (π) bond, and is represented by two lines; for example, in the case of carbon dioxide (CO ₂):

O = C = O

Triple bond

This bond, the strongest that exists among covalent bonds, occurs when the atoms share six electrons or three pairs, in a sigma (σ) and two pi (π) bonding. It is represented with three lines and can be seen in molecules such as acetylene (C ₂ H ₂):

HC≡CH

Lastly, quadruple bonds have been observed, but they are rare and are mainly limited to metallic compounds, such as chromium (II) acetate and others.

Examples

For simple bonds, the most common case is that of hydrogen, as can be seen below:

The case of a triple bond is that of nitrogens in nitrous oxide (N ₂ O), as seen below, with the sigma and pi bonds visible:

References

1. Chang, R. (2007). Chemistry. (9th ed). McGraw-Hill.
2. Chem Libretexts. (sf). Retrieved from chem.libretexts.org
3. Anne Marie Helmenstine, P. (nd). Retrieved from thoughtco.com
4. Lodish, H., Berk, A., Zipursky, SL, Matsudaira, P., Baltimore, D., & Darnell, J. (2000). Molecular Cell Biology. New York: WH Freeman.
5. Wikiversity. (sf). Retrieved from en.wikiversity.org