Revision: Chemical Bonding and Molecular Structure Chemistry Science (English Medium) Class 11 CBSE

Octet rule: Atoms of elements combine with each other in order to complete their respective octets so as to acquire the stable gas configuration.

The octet rule or the electronic theory of chemical bonding was developed by Kossel and Lewis. According to this rule, atoms can combine either by transfer of valence electrons from one atom to another or by sharing their valence electrons in order to attain the nearest noble gas configuration by having an octet in their valence shell.

The octet rule successfully explained the formation of chemical bonds depending upon the nature of the element.

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 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.

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.

Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule.

Bond lengths are expressed in terms of Angstrom (10^–10 m) or picometer

(10^–12 m) and are measured by spectroscopic X-ray diffractions and electron-diffraction techniques.

In an ionic compound, the bond length is the sum of the ionic radii of the constituting atoms (d = r₊ + r_–). In a covalent compound, it is the sum of their covalent radii (d = r_A+ r_B).

Bond-length: It is the equilibrium distance between the nuclei of two bonded atoms in a molecule. Bond-lengths are measured by spectroscopic methods

Electronegativity is the ability of an atom in a chemical compound to attract a bond pair of electrons towards itself.

Electronegativity of any given element is not constant. It varies according to the element to which it is bound. It is not a measurable quantity. It is only a relative number

Electronegativity is the tendency of an atom to attract shared pair of electrons. It is the property of bonded atom.

Anything that has mass and occupies space is called matter.

The process by which matter changes from one state to another and back to the original state, without any change in its chemical composition.

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

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.

In 1916, Kossel and Lewis independently proposed a theory of chemical combination.

• Atoms of different elements take part in chemical combination to complete their octet (8 electrons) or duplet (2 electrons) in the outermost shell.
• All valence shell (outer-shell) electrons of atoms are represented in Lewis structures using dots surrounding the element symbol.
• Lewis structures show only valence electrons of each atom — inner shell electrons are not shown.

Carbon, nitrogen, oxygen, and fluorine always obey the octet rule in their stable compounds. However:

• Second-row elements like B and Be often have fewer than 8 electrons (incomplete octet).
• Third-row elements can exceed 8 electrons (expanded octet) using d-orbitals.

• Octet Rule: Atoms tend to gain, lose, or share electrons so that their outermost shell attains 8 electrons (like a noble gas configuration).
• Duplet Rule: Hydrogen and lithium attain stability with only 2 electrons (like helium).
• The octet rule explains why most main-group elements form bonds in fixed ratios.

• A covalent bond is formed by the mutual sharing of electron pairs between two atoms.
• A single bond shares 1 pair (2 electrons); a double bond shares 2 pairs; a triple bond shares 3 pairs.
• Co-ordinate (Dative) Bond: A special covalent bond in which both electrons are donated by one atom (the donor) to another (the acceptor). Shown as X:→Y.
• Bond angle: The angle between orbitals containing bonding pairs around the central atom. E.g., bond angle of water = 104.5°.
• Bond enthalpy (Bond energy): Energy required to break one mole of a bond in the gaseous state. E.g., H–H bond enthalpy = 435.8 kJ mol⁻¹.

Lewis structures use dots (lone pairs) and dashes (bonds) to represent all valence electrons in a molecule. Key examples:

Steps to draw Lewis structure:

1. Count total valence electrons (add electrons for negative charge; subtract for positive)
2. Arrange atoms — least electronegative atom is usually in the centre
3. Connect atoms with single bonds
4. Complete octets on outer atoms first, then on central atom
5. If central atom has deficit, form multiple bonds

Formal charge is a bookkeeping tool — it is the hypothetical charge on an atom in a Lewis structure assuming electrons in bonds are equally shared. It helps identify the most stable (lowest energy) Lewis structure.

\[F.C.=V.E.-N.E.-\frac{B.E.}{2}\]

where:

• V.E. = Total number of valence electrons of the atom in a free state
• N.E. = Total number of non-bonding (lone pair) electrons on that atom
• B.E. = Total number of bonding (shared) electrons around that atom

Key rules:

• The sum of formal charges in a neutral molecule = 0
• The sum of formal charges in an ion = charge of that ion
• The most stable Lewis structure has formal charges as close to zero as possible
• Negative formal charge should be on the more electronegative atom

Example — CO₃²⁻ and Ozone (O₃): Both have multiple valid Lewis structures (resonance), and formal charges help identify the preferred one.

The octet rule is a useful guideline but not universal. Three important exceptions:

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.

• Distance between nuclei of two bonded atoms at equilibrium.
• Inversely proportional to bond order.
• Trend: Triple < Double < Single (shortest → longest).
• Depends on the size of atoms (larger atoms → longer bond length).
• Measured in picometres (pm) or Ångström (Å) (1 Å = 100 pm).
• Greater bond length → weaker bond.

• Angle between two adjacent bonds at the central atom.
• Determined by the shape of the molecule (VSEPR theory).
• Repulsion order: LP–LP > LP–BP > BP–BP.
• More lone pairs → smaller bond angle.
• Examples: CH₄ (109.5°), NH₃ (107°), H₂O (104.5°).
• Affected by electronegativity and hybridisation.

• Energy required to break one mole of bonds in the gaseous state.
• Directly proportional to bond strength.
• Increases with an increase in bond order.
• Triple bond > Double bond > Single bond (enthalpy trend).
• Depends on bond length (shorter bond → higher enthalpy).
• Measured in kJ mol⁻¹.

• Number of bonds between two atoms.

• Formula (MO theory): \[\frac{N_b-N_a}{2}\]

• Higher bond order → stronger and shorter bond.
• Determines the stability of a molecule.
• Examples: F₂ = 1, O₂ = 2, N₂ = 3.
• CO and NO⁺ have a bond order = 3.

• Resonance occurs when a single Lewis structure cannot accurately describe a molecule.
• Multiple structures (canonical structures) differing only in the arrangement of electrons (not atoms) are drawn.
• The actual molecule is a resonance hybrid — a weighted average of all canonical structures.
• Resonance energy = Actual bond energy − Energy of most stable resonating structure.

Conditions for Writing Resonance Structures:

• Same atomic positions in all structures.
• Same number of unpaired electrons.
• Nearly the same energy.
• Negative charge on the more electronegative atom; positive charge on the electropositive atom.
• Like charges should not reside on adjacent atoms.

Examples:

• 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:

The electrostatic force of attraction between a hydrogen atom (covalently bonded to a highly electronegative atom) and an electronegative atom (F, O, or N) of the same or different molecule.

• Hydrogen bonding is possible when H is attached to N, O, or F (highly electronegative and small atoms).
• A hydrogen bond is weaker than a covalent bond but stronger than van der Waals forces.

• H is bonded to a highly electronegative atom (F, O, N) → H becomes highly positive (δ+).
• This δ+ H is attracted to the lone pair of an electronegative atom of another molecule.