Covalent Bond

A covalent bond is a type of chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding.

1.0Formation of Covalent Bonds

Covalent bonding occurs primarily between nonmetals, which have similar tendencies to attract electrons. Rather than losing or gaining electrons (as in ionic bonding), atoms involved in covalent bonding share electrons because both atoms require more electrons to achieve a stable electronic configuration similar to noble gases. 

 Here are a few examples of molecules with covalent bonds-

 Sharing of electrons may occur in three ways –

2.0Characteristics of Covalent Bonds

1. Shared Electrons: In a covalent bond, electrons are shared mutually between the atoms. The shared electrons contribute to each atom's stability by effectively filling their outermost electron shells.
2. Bond Energy: The energy required to break a covalent bond is known as bond dissociation energy, a measure of the bond's strength. Different bonds have different bond energies depending on the atoms involved and the number of shared electrons.
3. Bond Length: This is the average distance between the nuclei of two bonded atoms. The bond length depends on the size of the atoms and the number of shared electrons; more shared electrons generally result in shorter bond lengths.
4. Orientation: Covalent bonds have specific orientations relative to the atomic nuclei involved, which leads to specific molecular shapes. This is described by theories like VSEPR (Valence Shell Electron Pair Repulsion), which predict the 3D geometry of molecules.

3.0Types of Covalent Bonds

1. Single Bonds: Involve one pair of shared electrons (e.g., the bond in H₂).
2. Double Bonds: Involve two pairs of shared electrons, offering greater strength and shorter bond lengths than single bonds (e.g., the bond in O₂).
3. Triple Bonds: Involve three pairs of shared electrons, are even stronger and shorter than double bonds (e.g., the bond in N₂).
4. Polar Covalent Bonds: Occur when electrons are shared unequally between atoms due to a difference in electronegativity. This creates a dipole moment in the molecule, with partial positive and negative charges (e.g., the bond in HCl).
5. Non-polar Covalent Bonds: Occur when electrons are shared equally between atoms, typically between identical atoms (e.g., the bond in H₂ or Cl₂).

4.0Polarity and Covalent Bonding

Polarity in covalent bonds can affect the physical properties of the compound, such as solubility, boiling point, and melting point. Polar covalent compounds tend to dissolve in polar solvents (like water), whereas nonpolar covalent compounds dissolve in nonpolar solvents (like benzene).

5.0Orbital Concept of Covalent Bond

This concept hinges on the behavior and interaction of atomic orbitals when atoms form covalent bonds. According to quantum mechanics:

Maximum Electron Accommodation

Each orbital can accommodate a maximum of two electrons, and these electrons must have opposite spins. This is due to the Pauli Exclusion Principle, which states that no two electrons in an atom can have the same set of four quantum numbers.

Utilisation of Half-Filled Orbitals

Covalent bonding typically occurs when half-filled orbitals (orbitals containing a single, unpaired electron) overlap. Each unpaired electron from one atom pairs with an unpaired electron from another atom to form a shared pair, thus completing the orbitals involved in bonding. This overlap leads to a lowering of the system's energy, providing stability to the molecule.

Electron Pairing as a Driving Force

The tendency for electrons to pair up is a fundamental aspect of covalent bonding. When two atoms approach each other, their half-filled orbitals overlap, allowing each unpaired electron to pair with one from the other atom. This pairing results in a covalent bond, where the shared electron pair holds the two atoms together.

Octet Completion Not Essential

While the completion of an octet (having eight electrons in the valence shell) is a common outcome of covalent bonding and contributes to stability (as per the Octet Rule), it is not an absolute requirement for the formation of a covalent bond. There are several exceptions to the octet rule, such as molecules and ions with odd numbers of electrons (like NO), molecules where one or more atoms possess more than eight electrons (like SF₆), and molecules where one or more atoms have fewer than eight electrons (like BH₃).

6.0Co-valency

Co-valency refers to the number of covalent bonds an atom can form within a molecule, whether the atom is in its ground state or an excited state. The concept of co-valency is central to understanding how atoms share electrons to form molecules.

7.0Variable Valency in Covalent Bonds

Variable valency occurs in elements with empty orbitals in their outer shell, allowing them to form different numbers of covalent bonds depending on conditions. In these elements, lone pair electrons can be excited to higher sub-shells within the same energy level, increasing the number of unpaired electrons available for bonding and thus raising co-valency in the excited state.

Elements in the 3rd to 6th periods can use vacant d-orbitals to show variable valency, especially when bonded to highly electronegative atoms like F, O, or Cl. The valency typically changes by two units between ground and excited states, and the energy required for this electron excitation is called promotion energy. This process follows the promotion rule, where excitation occurs within the same principal energy level.

For example- NCl₃ exists, but NCl₅ does not. Why?

Nitrogen (Ground state) - 1s² 2s²sp³ 

Co-valency 3

Nitrogen does not have an excited state because it lacks vacant d-orbitals. As a result, nitrogen cannot expand its valency beyond its ground state, which limits its bonding capacity. This is why NCl₅ does not exist—nitrogen cannot form five covalent bonds as it cannot access a higher valency state due to the absence of d-orbitals.