Revision: Chemical Bonding Chemistry HSC Science Class 11 Tamil Nadu Board of Secondary Education

A chemical bond may be defined as the force of attraction between any two atoms in a molecule to maintain stability.

or

A chemical bond may be defined as the force of attraction between any two atoms, in a molecule, to maintain stability.

Reduction is defined as the phenomenon in which an atom gains an electron to form a negatively charged ion called an anion.
During the formation of ionic bond one atom undergoes oxidation while another atom undergoes reduction.

Oxidation is the loss of electrons during a reaction by a molecule, atom or ion. In terms of electron transfer, oxidation is defined as the phenomenon in which an atom loses an electron to form a positively charged cation.
During the formation of ionic bond one atom undergoes oxidation while another atom undergoes reduction.

The cation and the anion being oppositely charged attract each other and form a chemical bond. Since this chemical bond formation is due to the electrostatic force of attraction between a cation and an anion, it is called an electrovalent (or an ionic) bond.

A non-metallic atom, which gains electron(s), becomes a negatively charged ion and is known as an anion.

A metallic element, whose one atom readily loses electron(s) to form a positively charged ion, is an electropositive element.

A non-metallic element, whose atom readily accepts electron(s) to form a negatively charged ion, is an electronegative element.

An ion is a charged particle which is formed due to the gain or the loss of one or more electrons by an atom.

A metallic atom, which loses electron(s), becomes a positively charged ion and is known as a cation.

The chemical compounds formed as a result of the transfer of electrons from one atom of an element to one atom of another element are called ionic (or electrovalent) compounds.

The number of electrons that an atom of an element loses or gains to form a electrovalent bond is called its electrovalency.

The number of bonds formed between the two bonded atoms in a molecule is called the bond order.

Bond order = `("N"_"b" - "N"_"a")/2`

The bond enthalpy is defined as the minimum amount of energy required to break one mole of a particular bond in molecules in their gaseous state. The unit of bond enthalpy is kJ mol^-1.

Hybridisation is the process of mixing of atomic orbitals of the same atom with comparable energy to form an equal number of new equivalent orbitals with the same energy.

\[\mathrm{Bond~Order}=\frac{N_b-N_a}{2}\]

where N_b = number of electrons in bonding MOs, N_a = number of electrons in antibonding MOs.

• Bond order > 0 → molecule is stable

• Bond order = 0 or negative → molecule is unstable (does not exist)

\[\mu=\sqrt{n(n+2)}\text{BM (Bohr Magneton)}\]

where n = number of unpaired electrons. If any unpaired electron is present → paramagnetic; if none → diamagnetic.

An ionic bond is formed by the complete transfer of one or more electrons from an electropositive atom to an electronegative atom, resulting in oppositely charged ions that attract each other.

Key conditions for ionic bond formation:

• One atom must have low ionisation enthalpy (easily loses electron) — typically a metal

• The other must have high electron affinity (easily gains electron) — typically a non-metal

• Large difference in electronegativity between the two atoms

Example: Na + Cl → Na⁺ + Cl⁻ → NaCl

• Sodium (2,8,1) loses 1 electron → Na⁺ (2,8)

• Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8)

Ionic solids are crystalline structures containing cations and anions held together by strong electrostatic ionic bonds.

Postulates:

• All valence shell electron pairs (bonding + lone pairs) around the central atom arrange themselves to be as far apart as possible to minimise repulsion.
• Lone pair–lone pair repulsion > lone pair–bond pair repulsion > bond pair–bond pair repulsion (lp−lp > lp−bp > bp−bp).
• Lone pairs occupy more space than bonding pairs, which compresses bond angles.
• When a molecule has no lone pairs, molecular geometry = electron pair geometry.

VSEPR Geometry Table:

Postulates:

• A covalent bond forms when orbitals of two atoms, each containing one unpaired electron, overlap with each other.
• The overlapping of two atomic orbitals produces a localised molecular orbital called a bonding orbital, occupied by two electrons with opposite spins.
• Maximum electron density lies between the two nuclei (internuclear region).
• The greater the overlap, the stronger the bond and the greater the bond energy — but the shorter the bond length.
• Electrostatic attraction between the nuclei and the accumulated electron cloud between them provides the binding force of the covalent bond.

Types of Orbital Overlap

Axial (head-on) overlap → forms σ (sigma) bond:

• s-s overlap: Two half-filled s orbitals overlap along the internuclear axis
• s-p overlap: Half-filled s and half-filled p orbital overlap along internuclear axis
• p-p overlap (axial): Two half-filled p orbitals overlap along the internuclear axis

Lateral (side-on) overlap → forms π (pi) bond:

• p-p overlap (lateral): Two half-filled p orbitals overlap side-ways (perpendicular to internuclear axis)

Molecular orbitals (MOs) are formed by the linear combination of atomic orbitals (LCAO).

Two types of MOs form:

• Bonding MOs — lower energy than the original atomic orbitals; electrons here stabilise the molecule (σ, π)
• Antibonding MOs — higher energy; electrons here destabilise the molecule (σ*, π*)

Energy Order of MOs for Diatomic Molecules:

For O₂, F₂ (electrons > 14):

For B₂, C₂, N₂ (electrons ≤ 14):

Electronic Configurations and Bond Properties of Diatomic Molecules: