Revision: Chemical Bonding Chemistry HSC Science (General) 11th Standard Maharashtra State Board

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

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

or

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.

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.

It is the angle between bonded orbitals containing bonding electron pairs around the central atom in a molecule or complex ion.

The equilibrium distance between two nuclei bonded to each other is known as bond length. It is expressed in Å or pm or nm.

1 pm = 10^-12 m, 1 Å = 10^-10 m, 1 nm = 10^-9 m

The minimum amount of energy required to break a bond into one mole of gaseous molecule is known as bond enthalpy.

Bond enthalpy ∝ `\"1"/"Size of atoms"\`

∝ `\"1"/"Number of lone pair of electrons"\` ∝ Multiplicity of bond

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

Number of covalent bond between the two atoms is known as bond order. Integral bond order values of 1, 2 and 3 correspond to single, double and triple bonds, respectively. Isoelectronic molecules and ions have identical bond order.

Bond order ∝ Bond enthalpy ∝`\"1"/"Bond length"\`

Bond enthalpy is defined as the amount of energy required to break one mole of a bond of one type, present between two atoms in a gaseous state.

Dipole moment is defined as the product of the magnitude of charge (q) and distance (d) separating the centres of positive and negative charges.

Its direction is from positive end to negative end. 

µ = q × d

Its unit in CGS system is debye (D).

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

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.

• 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.

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.

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:

• 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:

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:

Resonance occurs when a single Lewis structure cannot adequately represent the actual structure of a molecule — multiple valid Lewis structures (called canonical forms) can be drawn.

• The actual molecule does not switch between these structures; it is a resonance hybrid — a weighted average of all canonical forms.

• The energy of the resonance hybrid is always lower than the energy of any single canonical form (this energy difference is called resonance energy or resonance stabilisation energy).

• All canonical forms must have: the same positions of atoms, the same number of paired and unpaired electrons, and similar energy.

Classic examples:

• CO₃²⁻ (carbonate ion): 3 equivalent resonance structures — each C − O bond is neither single nor double, but intermediate (~1.33 bond order)
• Ozone (O₃): 2 resonance structures with bond length ~128 pm (intermediate between O − O single bond ~148 pm and O = O double bond ~121 pm)