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

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

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.

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.

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.

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.

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.

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

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.

• Proposed by Heitler and London (1927), further developed by Pauling and Slater.
• A covalent bond is formed when half-filled valence atomic orbitals of similar energies overlap, each containing one unpaired electron.
• Greater the overlap → stronger the bond.

Types of Orbital Overlap:

Hybridisation & Shapes:

Limitations of VBT:

• Involves a number of assumptions.
• Does not give a quantitative interpretation of magnetic data.
• Does not explain the colour exhibited by coordination compounds.
• Does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.
• Does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes.
• Does not distinguish between weak and strong ligands.

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.

• Proposed by Sidgwick and Powell (1940) and further developed by Nyholm and Gillespie.
• The geometry of a molecule depends on the total number of valence shell electron pairs (bond pairs + lone pairs) around the central atom.
• Electron pairs repel each other and arrange themselves as far apart as possible to minimise repulsion.
• Repulsion order:
  lp–lp > lp–bp > bp–bp, and lone pairs occupy more space than bond pairs.
• Presence of lone pairs reduces bond angle; if no lone pairs → molecular geometry = electron pair geometry.

VSEPR Geometry Table:

Hybridisation is the process of mixing orbitals of nearly similar energy from the same atom to form a new set of equivalent orbitals of exactly equal energy called hybrid orbitals.

\[H=\frac{1}{2}[V+Y-C+A]\]

where V = valence electrons of central metal atom, Y = number of monovalent atoms surrounding central atom, C = total positive charge, A = total negative charge on the molecule.

Characteristics of Hybridisation:

• Number of hybridised orbitals = number of orbitals that participated in hybridisation.
• Hybridised orbitals are always equivalent in energy and shape.
• Hybrid orbitals are more effective in forming stable bonds than pure atomic orbitals.
• Hybrid orbitals are directed in space in some preferred directions → determines geometry of the molecule.

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: