A chemical bond formed by the mutual sharing of one or more pairs of electrons between atoms of same or different elements so as to complete their octets is called the covalent bond. Each atom contributes an equal number of electrons for the process of sharing and number of electrons contributed by each atom is known as covalency.
A single bond is formed when atoms mutually share one pair of electrons. A double bond is formed when the atoms share two electron pairs. A triple bond is formed when the atoms share three electron pairs.

Examples:

(a) Homoatomic molecules (covalent bonds in similar atoms).

(i) Hydrogen molecule (H₂): A covalent bond is formed between two hydrogen atoms by sharing a pair of electrons between them. Each atom contributes one electron for sharing.



(ii) Chlorine molecule (Cl₂): Chlorine atom (₁₇Cl) has seven electrons in its valence shell (2, 8, 7). In the formation of chlorine molecule, each chlorine atom contributes one electron for sharing in order to acquire a stable noble gas configuration (Argon).



(iii) An oxygen molecule (O₂): Oxygen atom (₈O) has six electrons in its valence shell (2, 6). In the formation of O₂ molecule, each oxygen atom contributes two electrons for mutual sharing in order to acquire a stable neon configuration and forms divalent or double bond.



(iv) Nitrogen molecule (N₂): Nitrogen atom (₇N) has five electrons in its valence shell (2, 5). In the formation of N₂ molecule, each nitrogen atom contributes three electrons for mutual sharing in order to acquire stable neon configuration and forms trivalent or triple bond.




(b) Heteroatomic molecules (covalent bonds in different atoms).

(i) Methane molecule (CH₄): Carbon atom (₆C) has four electrons in its valence shell (2, 4). Carbon atom forms four covalent bonds with four hydrogen atoms.



(ii) A water molecule (H₂O): The oxygen atom is bonded to two hydrogen atoms.



(iii) Carbon dioxide molecule (CO₂): Carbon atom is bonded to two oxygen atoms with double bonds.

Chemical bond. A chemical bond may be defined as the attractive force or binding force which holds atoms, ions and molecules together.

Why do atoms combine (Kossel-Lewis approach)?
The noble gases are highly unreactive and stable. Except for He (1s²), all other noble gas elements have ns²np^b outermost electronic configurations. This indicates that the presence of 8 electrons (law of octet) in the outermost orbit must be related to the stability of the atom. If K-orbit is the outermost orbit, the presence of 2 electrons (law of duplet) causes stability.

Thus atoms of different elements combine with each other in order to complete their respective octets (i.e. 8 electrons in their outermost shell) or duplet (i.e. outermost shell having 2 electrons) in the case of H, Li and Be to attain stable inert gas configuration.

How do atoms combine (Modes of chemical combination):
Atoms combine together in two ways to acquire stable inert gas configuration:

(i) By complete transference of one or more electron from one atom to another: This process is referred to as electro-valency and the chemical bond formed is termed as an electrovalent bond or ionic bond.

(ii) By sharing of elements: This can occur in two ways:
(a) When the shared electrons are contributed by the two combining atoms equally, the bond formed is called the covalent bond.

(b) When these electrons are contributed entirely by one of the atoms but shared by both, the bond formed is known as coordinated bond or dative bond.

Following factors influence the formation of ionic bonds:

(i) Low ionisation enthalpy: One of the atoms forming cation should have low ionisation enthalpy. Alkali metals (Group 1) and alkaline earth metals (Group 2) having low values of ionisation enthalpy from their cations readily.

(ii) Very high negative electron gain enthalpy: The other atom-forming anion should have very high negative electron gain enthalpy. Elements of group 16 and group 17 (halogens) having very high negative electron gain enthalpy from their anions readily.

M(g) → M⁺(g) + e– : Ionisation enthalpy
X(g) + e^– → X ^–(g) : Electron gain enthalpy
M⁺(g) + X ^–(g)→ MX(s)

Clearly, ionic bonds will be formed more easily between elements with comparatively low ionisation enthalpies and elements with comparatively very high negative electron gain enthalpies.

Ionic or electrovalent bond: The electrostatic force of attraction between the oppositely charged ions is known as an ionic bond. It is formed by the complete transference of one or more valence electrons of one atom to the valence shell of the other atom so that each atom acquires the nearest noble gas configuration. The compounds containing ionic or electrovalent bonds are called ionic or electrovalent compounds. The number of electrons which an atom loses or gains while forming an ionic bond is known as electrovalency. Examples :

(i) Formation of NaCl: The electronic configuration of sodium (Z= 11) and chlorine (Z= 17) can be represented as:.



Thus, one electron gets transferred from Na atom to chlorine atom. This gives rise to Na⁺ and Cl^– ions having a noble gas configuration. These ions are held together by the electrostatic force of attraction known as an ionic bond.



(ii) Formation of CaF₂. The electronic configuration of calcium (Z = 20) and fluorine (Z = 9) can be represented as,

₂₀Ca = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² (2, 8, 8, 2)
₉F = 1s² 2s² 2p⁵ (2, 7)

Thus, calcium atom loses two valence electrons to two fluorine atoms each of which gains one electron. This leads to the formation of Ca²⁺ ion and two F“ ions each of which has a stable noble gas configuration. These oppositely charged ions are mutually attracted by the electrostatic force of attraction which constitute ionic bonds.



(iii) Formation of MgO: The electronic configuration of magnesium (Z = 12) and oxygen (Z = 8) can be represented as:

₁₂Mg= 1s² 2s² 2p⁶ 3s² (2, 8, 2)
₈O = 1s² 2s² 2p⁴ (2, 6)

Magnesium atom loses two valence electrons to an oxygen atom. Thus, Mg²⁺ and O^2– ions are formed each of which has a stable noble gas configuration. These ions are held together by the electostatic force of attraction which constitutes ionic bond.

(i) Stable existence: Ionic compounds usually exist in the form of crystalline solids. The crystals are made up of crystal lattices containing oppositely charged ions (+ve and –ve). Each cation is surrounded by a definite number of anions and vice-versa.

(ii)  High melting and boiling points and low volatility: There is a strong force of attraction among the oppositely charged ions in the crystals of ionic compounds, so a large amount of enthalpy is needed to separate them. Due to these strong forces of attraction, ionic compounds have high melting and boiling points and low volatility.

(iii) Electrical conductivity: Ionic compounds do not conduct electricity in the solid state because oppositely charged ions occupy fixed positions in the crystals and are not free to move. When these crystals are dissolved in a polar solvent or melted, the ions can move freely under the influence of electric field and become conductor of electricity.

(iv) Fast reactions: The chemical reactions between ionic compounds involve the combination between the ions liberated in their aqueous solutions. So their reactions are very fast.

(v) Solubility: Ionic compounds are soluble in polar solvents (e.g. water) and insoluble in nonpolar (organic) solvents. Polar solvents have high values of a dielectric constant which reduces the electrostatic force of attraction between the oppositely charged ions. Hence, the ions get separated and ultimately solvated by the molecules of the polar solvent.