Chemical Bonding and Molecular Structure Class 11 Notes: The Ultimate Guide for NEET

Mastering chemical bonding is critical for scoring high in the NEET Chemistry section. These chemical bonding class 11 notes cover everything from octet rules to Molecular Orbital Theory. Understanding how atoms combine to form molecules is the foundation for both Inorganic and Organic Chemistry preparation.

01
Introduction to Chemical Bonding

Chemical bonding is the attractive force that holds constituent particles (atoms, ions, or molecules) together in different chemical species. Atoms combine to achieve a stable electronic configuration, specifically to attain the nearest noble gas configuration with a complete valence shell.

PRIMARY DRIVER Atoms combine to minimize potential energy and achieve maximum stability.
OCTET RULE Most atoms tend to have eight electrons in their outermost shell (ns2 np6).
TIP
Remember that bonding is always an exothermic process (energy is released) because the system moves from a higher energy state to a lower, more stable energy state.

02
Kossel–Lewis Approach to Chemical Bonding

According to this approach, atoms achieve the stable octet by either transferring electrons (ionic bond) or sharing electrons (covalent bond). While this rule is fundamental in chemical bonding class 11 notes, it has several notable exceptions that frequently appear in NEET exams.

REPRESENTATION
Lewis Dot Structures: Valence electrons are represented as dots around the atomic symbol.

Limitations of the Octet Rule

  • Incomplete Octet: BeCl2 (Be has 4e−), BF3 (B has 6e−).
  • Expanded Octet: Elements in the 3rd period and beyond (PCl5, SF6, IF7).
  • Odd Electron Molecules: NO and NO2 have unpaired electrons.

03
Ionic Bonding (Electrovalent Bond)

The electrostatic force of attraction between a cation and an anion is called an ionic bond. Its formation depends on several energetic factors including lattice and ionization enthalpies.

Factors favoring Ionic Bond: Low Ionization Enthalpy (Cation) + High Negative Electron Gain Enthalpy (Anion) + High Lattice Enthalpy.
Property Ionic Compounds Covalent Compounds
Physical State Crystalline Solids Gases, Liquids, or Soft Solids
Conductivity Only in Molten/Aqueous State Generally Insulators
Boiling Point Very High Low to Moderate
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04
Covalent Bonding and Lewis Structures

In covalent bonding, atoms share electron pairs to satisfy their octets. A key concept here is the Formal Charge, which helps in identifying the most stable Lewis structure for a molecule.

FORMAL CHARGE CALCULATION
F.C. = [Total valence e−] − [Non-bonding e−] − 1/2[Bonding e−]
WARN
Formal charge does not represent the real charge on the atom; it is a bookkeeping method to compare the stability of different resonance structures.

05
Bond Parameters

Quantitative values describing a bond are essential for solving theoretical NEET questions on chemical bonding class 11 notes.

BOND ORDER Number of bonds between two atoms. For N2, Bond Order = 3.
RELATIONSHIP Bond Order ∝ 1/Bond Length ∝ Bond Enthalpy.

06
VSEPR Theory: Predicting Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the shape of molecules based on the repulsion between electron pairs.

Order of Repulsion: Lone Pair – Lone Pair (lp-lp) > Lone Pair – Bond Pair (lp-bp) > Bond Pair – Bond Pair (bp-bp).
Hybridization Geometry Example Bond Angle
sp Linear BeCl2, CO2 180°
sp2 Trigonal Planar BF3 120°
sp3 Tetrahedral CH4, NH3 (1 lp) 109.5°
sp3d Trigonal Bipyramidal PCl5 90° & 120°

07
Valence Bond Theory (VBT) and Overlap

VBT explains bonding through the overlap of atomic orbitals. Two types of bonds are formed: Sigma (σ) and Pi (π). Sigma bonds result from head-on overlap, whereas Pi bonds result from lateral overlap. Sigma bonds are stronger than Pi bonds.

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08
Hybridization

Hybridization is the intermixing of atomic orbitals of slightly different energies to produce a new set of equivalent orbitals.

STERIC NUMBER FORMULA
H = 1/2 [V + M − C + A]

09
Molecular Orbital Theory (MOT)

MOT is the most advanced concept in chemical bonding class 11 notes. It explains the magnetic properties of molecules (like why O2 is paramagnetic).

Bond Order = 1/2 [Nb − Na]
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Quick Revision Summary

  • Dipole Moment (μ) = Charge (Q) × Distance (r).
  • Polarity depends on the vector sum of individual bond moments.
  • H-bonding: Strongly electronegative elements (F, O, N) attached to Hydrogen.
  • Lattice Enthalpy ∝ (Charge 1 × Charge 2) / (Radius 1 + Radius 2).
  • Paramagnetic substances have unpaired electrons in MOs.
  • Diamagnetic substances have all paired electrons.
  • % Ionic Character = 16|Δχ| + 3.5|Δχ|2.
  • VSEPR predicts Shape; Hybridization predicts Geometry.
Download Formula Sheet (PDF)

10
Frequently Asked Questions

Why is BF3 non-polar despite polar B-F bonds?
BF3 has a symmetrical trigonal planar geometry. The vector sum of the three B-F bond dipoles is zero, making the net dipole moment μ = 0.
How does a lone pair affect the bond angle?
Lone pairs occupy more space and exert greater repulsion on bond pairs. This causes the bond pairs to push closer together, reducing the bond angle (e.g., CH4 is 109.5° while NH3 is 107°).
What is the bond order of O2 and O2+?
The bond order of O2 is 2. When it loses an electron to form O2+, the electron is removed from an anti-bonding orbital, increasing the bond order to 2.5.
What are the conditions for Hydrogen bonding?
1. Hydrogen must be bonded to a highly electronegative atom (F, O, or N). 2. The electronegative atom must be small in size to provide concentrated charge density.
Is every molecule with an expanded octet unstable?
No. Molecules like SF6 and PCl5 are very stable because the central atom can use its d-orbitals to accommodate more than 8 electrons.
What is the main difference between Sigma and Pi bonds?
Sigma bonds are formed by end-to-end overlap and allow free rotation, while Pi bonds are formed by lateral overlap and prevent rotation around the bond axis.

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Table of Contents — Chemistry Class 11

Table of Contents

Chemistry — Class 11

01Some Basic Concepts of ChemistryGo to page
02Structure of AtomGo to page
03Classification of Elements and PeriodicityGo to page
04Chemical Bonding and Molecular StructureGo to page
05ThermodynamicsGo to page
06EquilibriumGo to page
07Redox ReactionsGo to page
08Organic Chemistry — Basic PrinciplesGo to page
09HydrocarbonsGo to page

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