Chemical Bonding and Molecular Structure Notes for Class 11, JEE and NEET

Chemical bonding explains why atoms combine and why molecules have particular shapes and properties. It is Chapter 4 of NCERT Class 11 Chemistry. It builds on Structure of Atom (Class 11 Chemistry Chapter 2), and it is used again in Coordination Compounds, the p-Block and Organic Chemistry. These Chemical Bonding and Molecular Structure Notes follow the chapter in the same order as the textbook.

India's Best Exam Preparation for Class 11th - Download Now

Chemical Bonding and Molecular Structure Notes

India's Best Exam Preparation for Class 11th - Download Now

Also read, Class 11 Chemistry Chapter 2 Notes

India's Best Exam Preparation for Class 11th - Download Now

India's Best Exam Preparation for Class 11th - Download Now

India's Best Exam Preparation for Class 11th - Download Now

India's Best Exam Preparation for Class 11th - Download Now

India's Best Exam Preparation for Class 11th - Download Now

India's Best Exam Preparation for Class 11th - Download Now

Chemical Bonding and Molecular Structure Notes: Chapter Overview and Syllabus

The chapter covers:

• Kossel–Lewis approach, the octet rule and its limitations
• Ionic bonding and lattice enthalpy
• Covalent bonding: Lewis structures, formal charge, bond parameters and resonance
• Polarity: electronegativity, dipole moment and Fajan’s rules
• VSEPR theory and the shapes of molecules
• Valence bond theory and hybridisation (s, p and d orbitals)
• Molecular orbital theory for homonuclear diatomic molecules
• Hydrogen bonding and an elementary idea of metallic bonding

Check the latest syllabi of CBSE, JEE Main and NEET for the exact topic lists.

Tip from the eSaral Chemistry team: Do not memorise shapes one by one. Count the bond pairs and lone pairs on the central atom, find the steric number, and then read the shape from the VSEPR table. The same method works for every molecule.

Octet Rule, Ionic Bonds and Lattice Enthalpy in Chemical Bonding and Molecular Structure Notes

Octet Rule and Its Limitations

Atoms tend to gain, lose or share electrons to get eight electrons in the outermost shell (two for hydrogen). The octet rule fails for:

• Incomplete octet: BeCl₂ (4 electrons around Be), BF₃ (6 electrons around B)
• Expanded octet: PCl₅, SF₆ and other compounds of the third period and below
• Odd-electron molecules: NO and NO₂

Ionic Bond and Lattice Enthalpy

An ionic bond forms by the transfer of electrons. It is favoured by a low ionisation enthalpy of the metal, a high electron gain enthalpy of the non-metal and a high lattice enthalpy. Lattice enthalpy is the energy released when one mole of an ionic solid forms from its gaseous ions, and it is larger when the ions are smaller and have higher charges. For example, MgO has a much higher lattice enthalpy than NaCl.

Covalent Bonds, Formal Charge and Resonance in Chemical Bonding and Molecular Structure Notes

• Formal charge = (valence electrons of the free atom) − (non-bonding electrons) − ½(bonding electrons). The structure with formal charges closest to zero is the most stable.
• Resonance: when one Lewis structure cannot describe a molecule, the real structure is a hybrid of the contributing structures. In the carbonate ion (CO₃²⁻), all three C–O bonds are equal, with a bond order of 4/3.

Bond Parameters

Polarity, Dipole Moment and Fajan’s Rules in Chemical Bonding and Molecular Structure Notes

• Dipole moment: μ = q × d. The unit is the debye (D), where 1 D = 3.336 × 10⁻³⁰ C·m. It is a vector, and the molecular dipole moment is the vector sum of the bond dipoles. Symmetrical molecules such as CO₂, BF₃ and CH₄ have zero dipole moment.
• Fajan’s rules: the covalent character of an ionic bond increases when the cation is small and has a high charge, the anion is large and has a high charge, and the cation has a pseudo-noble gas configuration (for example Cu⁺) instead of a noble gas configuration. LiCl is more covalent than NaCl.

VSEPR Theory and Shapes of Molecules in Chemical Bonding and Molecular Structure Notes

The VSEPR theory says that electron pairs around the central atom arrange themselves to minimise repulsion. The order of repulsion is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.

Ammonia is the standard trigonal pyramidal example. Its bond angle (about 107°) is smaller than the tetrahedral angle (109.5°), because the lone pair repels the bond pairs more strongly. For its formula, preparation and properties, see the chemical formula of ammonia.

Valence Bond Theory and Hybridisation in Chemical Bonding and Molecular Structure Notes

Hybridisation is the mixing of atomic orbitals of similar energies to form new hybrid orbitals. The number of hybrid orbitals equals the number of orbitals mixed.

Steric number (H) = ½[valence electrons of the central atom + number of monovalent atoms − charge on a cation + charge on an anion].

A sigma (σ) bond forms by the head-on overlap of orbitals, and a pi (π) bond forms by the sideways overlap of p orbitals. A double bond has one σ and one π bond, and a triple bond has one σ and two π bonds.

Molecular Orbital Theory in Chemical Bonding and Molecular Structure Notes

Atomic orbitals combine (LCAO) to form bonding and antibonding molecular orbitals.

Bond order = ½ (N_b − N_a), where N_b and N_a are the numbers of electrons in the bonding and antibonding orbitals. A molecule is stable if its bond order is greater than zero.

• Energy order for O₂, F₂ and Ne₂: σ1s < σ1s < σ2s < σ2s < σ2p_z < (π2p_x = π2p_y) < (π2p_x = π2p_y) < σ*2p_z.
• Energy order for B₂, C₂ and N₂: the π2p orbitals lie below σ2p_z.

A higher bond order means a shorter bond length and a higher bond enthalpy.

Hydrogen Bonding and Metallic Bonding in Chemical Bonding and Molecular Structure Notes

Hydrogen bonding occurs when hydrogen attached to a highly electronegative atom (F, O or N) is attracted to another electronegative atom.

• Intermolecular hydrogen bonding (HF, H₂O, alcohols) increases the boiling point and the solubility in water.
• Intramolecular hydrogen bonding (o-nitrophenol) lowers the boiling point and makes the compound steam volatile, while p-nitrophenol has intermolecular hydrogen bonding and a higher boiling point.
• Hydrogen bonding explains why water has a high boiling point for its molecular mass, and why ice is less dense than water.

Metallic bonding (elementary idea): the metal ions sit in a sea of delocalised electrons, which explains the conductivity, malleability and lustre of metals.

Solved Examples from Chemical Bonding and Molecular Structure Notes

Example 1 (hybridisation and shape): Find the hybridisation and shape of SF₄.

H = ½(6 + 4) = 5, so the hybridisation is sp³d. There are 4 bond pairs and 1 lone pair, so the shape is see-saw.

Example 2 (bond order): Find the bond order and the magnetic behaviour of O₂⁻.

O₂⁻ has 17 electrons: N_b = 10 and N_a = 7. Bond order = ½(10 − 7) = 1.5, and there is 1 unpaired electron, so it is paramagnetic.

Example 3 (dipole moment): Two charges of +1.6 × 10⁻¹⁹ C and −1.6 × 10⁻¹⁹ C are 1.0 Å apart. Find the dipole moment in debye.

μ = q × d = 1.6 × 10⁻¹⁹ × 10⁻¹⁰ = 1.6 × 10⁻²⁹ C·m = 1.6 × 10⁻²⁹/3.336 × 10⁻³⁰ ≈ 4.8 D.

Example 4 (formal charge): Find the formal charge on nitrogen in NH₄⁺.

Formal charge = 5 − 0 − ½(8) = +1.

Example 5 (bond angle): Arrange CH₄, NH₃ and H₂O in the decreasing order of bond angle.

CH₄ (109.5°) > NH₃ (about 107°) > H₂O (about 104.5°). The lone pairs repel the bond pairs more strongly than bond pairs repel each other.

Example 6 (Fajan’s rules): Which is more covalent, LiCl or NaCl? Which has the higher lattice enthalpy, NaCl or MgO?

LiCl is more covalent, because Li⁺ is smaller than Na⁺ and polarises the anion more. MgO has the higher lattice enthalpy, because its ions are smaller and carry higher charges.

Common Mistakes in Chemical Bonding and Molecular Structure Notes Problems

These are the errors that students make most often when they apply the Chemical Bonding and Molecular Structure Notes to problems.

How to Study Chemical Bonding and Molecular Structure Notes

1. Read the NCERT chapter first, and learn the VSEPR and hybridisation tables. Download the textbook from eSaral’s NCERT books page.
2. Revise Structure of Atom first if electronic configurations feel weak, because hybridisation and MOT depend on them.
3. Practise shapes, hybridisation and bond order for 10 species every day for a week.
4. Revise the chapter with eSaral’s Class 11 Chemistry notes, and then solve previous year questions.

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