Covalent Bond

A covalent bond is a chemical bond formed when two atoms share one or more electron pairs while each atom contributes equally. Covalent bonding is a stable equilibrium of the attractive and repulsive forces between two atoms that occurs when they share electrons. Bonding pairs or sharing pairs are other names for these type of electron pairs. Because electrons are shared among several molecules, each atom can reach the equivalent of a full valence shell, resulting in a stable electronic state. In organic chemistry, covalent bonds are significantly more common than ionic bonds.

Covalent bonding encompasses a wide range of other forms of interactions, including bent bonds, three-center two-electron bonds, agostic interactions, metal-to-metal links, and three-center four-electron bonds, among others. Covalent bonds were first described in 1939. A "co-valent bond" basically signifies that the atoms share "valence," as is covered in valence bond theory, because the prefix co- means jointly, related in action, partnered to a lesser extent, etc.

In order to share their two electrons, the hydrogen atoms in the molecule H₂ establish covalent bonds. The highest covalency occurs when two atoms with similar electronegativities are present for sharing their electrons. For a covalent bond to form, the two atoms merely need to have similar electronegativity; they do not necessarily need to be from the same element. Delocalized covalent bonding is characterised by covalent bonds in which more than two atoms share an electron.

Hydrogen Bond

History of covalent bond

Lewis first postulated the second form of combination in 1916, stating that when each atom contributes equally, some atoms can achieve a noble gas configuration by sharing one or more electron pairs. The electron pair or pair(s) becomes a shared characteristic of both. Atoms with and without similarity can form such a connection. Since the electron or electrons are not totally lost in this link, the atoms are not charged. The shared electrons are positioned between the nuclei of the two atoms where the force of attraction between the two nuclei is greatest. As a result, the link is known as a non-polar bond.

Covalent bonds are defined as chemical bonds generated by sharing one or more electron pairs amongst atoms when each atom participates equally.

A covalent bond can have one, two, or three bonds. Multiple covalent bonds refer to double and triple covalent bonds. One electron pair is shared to create a single covalent bond. A single dash is used to denote this relationship (-). Atoms bound together create double and triple covalent bonds when they share two or three electron pairs, respectively. Double dashes (=) and triple dashes (==) are used to denote these bonds, respectively.

Lewis postulated that an atom creates sufficient covalent bonds to close off (or form a complete) its outer electron shell. According to the octet rule, the carbon atom in this diagram of methane is surrounded by eight electrons (four from the carbon atom itself and four from the hydrogens bound to it) since it has a valence of four. In accordance with the duet rule, each hydrogen atom has a valence of one and is surrounded by two electrons: its own and one from the carbon. According to the atom's quantum theory, the numbers of electrons correspond to full shells. For example, the outer shell of a carbon atom is the n = 2 shell, which can contain eight electrons, whereas the outer (and only) shell of a hydrogen atom is the n = 1 shell, which can only hold two electrons.

The phrase "covalence" was originally used to describe bonding in 1919 by Irving Langmuir in an essay titled "The Arrangement of Electrons in Atoms and Molecules" in the Journal of the American Chemical Society. The number of pairs of electrons that an atom shares with its neighbours will be referred to as covalence, according to Langmuir.

Despite the fact that the concept of shared electron pairs offers a useful qualitative representation of covalent bonding, quantum mechanics is still necessary to comprehend the nature of these connections and foresee the structures and characteristics of small molecules. The first successful quantum mechanical explanation of a chemical bond (molecular hydrogen) is attributed to Walter Heitler and Fritz London in 1927. The valence bond concept, which holds that a chemical bond is created when there is good overlap between the atomic orbitals of participating atoms, served as the foundation for their research.

Types of covalent bonds

Except for s orbitals, atomic orbitals have unique directional characteristics that give rise to several kinds of covalent bonds. Sigma (𝛔) bonds, which result from the head-on overlapping of orbitals on two distinct atoms, are the strongest covalent bonds. Typically, a bond refers to a single bond. Due to lateral overlap between p (or d) orbitals, pi (𝜋) bonds are weaker. One and one bonds make up a double bond, whereas one and two bonds make up a triple bond between any two specified atoms.

Covalent bonds' chemical polarity is also influenced by the electronegativity of the atoms they are attached to. Equivalent electronegativity between two atoms allows them to form nonpolar covalent bonds like H-H. Polar covalent bonds, such as the one with H-Cl, are produced by unequal relationships. However, geometric asymmetry is also necessary for polarity; otherwise, dipoles may cancel out and produce a non-polar molecule.

Covalent structures

Covalent substances can take on a variety of structures, including single molecules, molecular structures, macromolecular structures, and enormous covalent structures. Although the atoms in each molecule have strong bonds holding them together, the forces of attraction between molecules are often very weak. Usually gases, such covalent compounds include HCl, SO2, CO₂, and CH₄. There are weak attraction forces present in molecular structures. Low-melting-temperature solids and low-boiling-temperature liquids (like ethanol) are examples of such covalent compounds (such as iodine and solid CO₂). In macromolecular structures, which include biopolymers like proteins and starch as well as manmade polymers like polyethylene and nylon, several atoms are connected by covalent bonds to form chains. Large numbers of atoms are connected together in sheets, such as in graphite, to form network covalent structures, also known as massive covalent structures or three-dimensional structures (such as diamond and quartz). These materials are usually brittle, have high melting and boiling temperatures, and exhibit significant electrical resistance. Such massive macromolecular structures are frequently formed by elements with strong electronegativity and the capacity to create three or four electron pair bonds.

One and three electron bonds

Radical species, which have an odd number of electrons, can have bonds with one or three electrons. The dihydrogen cation, H+2, is the most straightforward instance of a 1-electron bond. One-electron bonds, sometimes known as "half bonds," frequently have half the bond energy of a two-electron bond. There are, however, certain exceptions. For instance, in the case of dilithium, the link between the 1- and 2-electron Li+2 actually has a stronger binding than the latter. Hybridization and inner-shell effects can be used to explain this outlier.

The helium dimer cation, He+2, is the most straightforward instance of three-electron bonding. Because there is only one shared electron present (as opposed to two), it is referred to as a "half bond"; in terms of molecular orbitals, the third electron is in an anti-bonding orbital, which cancels out half of the bond created by the other two electrons. Nitric oxide, or NO, is another example of a molecule with a 3-electron link in addition to two 2-electron bonds. The paramagnetism and formal bond order of the oxygen molecule, O₂, can also be explained by its two 3-electron bonds and one 2-electron bond. Three-electron bonds are also present in the heavier counterparts of chlorine dioxide, such as bromine dioxide and iodine dioxide.

Odd-electron bond molecules are typically very reactive. These bonds are only durable when formed between atoms with comparable electronegativities..

Electron deficiency

Three atoms share two electrons in a link known as a three-center two-electron bond (3c-2e). Because there are insufficient valence electrons to create localised (2-centre 2-electron) bonds connecting all the atoms, this sort of bonding happens in boron hydrides like diborane (B2H6), which is frequently referred to as electron deficient. The molecules can, however, be categorised as electron-precise as the more recent description employing 3c-2e bonds does give enough bonding orbitals to connect all the atoms.

Each of these bonds—there are two of them per molecule of diborane—contains two electrons that form a banana-shaped connection between the boron atoms. A proton—the nucleus of a hydrogen atom—is located in the centre of the bond and shares electrons with both boron atoms. So-called four-center two-electron bonds have also been proposed for some cluster compounds.