Schooling · Class 11 · Chemistry · Chemical Bonding and Molecular Structure · Chapter on a page
CBSE Class 11 · Chemistry · NCERT Chemistry Part-I
Chemical Bonding and Molecular Structure: chapter on a page
A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).
About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.
Fact-checked on 21 Oct 2026: an AI-assisted check of this note against cited sources; the errors it found were corrected that day. Some errors may remain.
Mind map
Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.
The mind map as a list
- Key ideas
- Octet rule
- Ionic bond
- Covalent bond
- Coordinate (dative) bond
- VSEPR theory
- Hybridisation
- Molecular Orbital Theory (MOT)
- Formulas and facts
- Bond order (MOT)
- Formal charge
- Dipole moment
- Hybridisation and geometry
- sp
- sp²
- sp³
- Worked examples
- Drawing the Lewis structure of CO₂
- Determining hybridisation and shape…
- Bond order of O₂ using MOT
- Watch out for
- Thinking all sp³ molecules are…
- Ignoring lone pairs when predicting…
- Confusing molecular polarity with…
- Assuming higher bond order always…
- Thinking a negative bond order is…
What this chapter is about
This chapter explains why and how atoms combine to form molecules and compounds. You have already studied atomic structure and the arrangement of electrons in shells and subshells. Now you learn the driving force behind bond formation: atoms tend to achieve a stable, low-energy configuration, often resembling the electron arrangement of noble gases.
Key ideas
- Octet rule: Atoms tend to gain, lose or share electrons until their valence shell has eight electrons, achieving noble-gas stability. Hydrogen is an exception; it needs only two electrons (duplet).
- Ionic bond: Formed when one atom transfers electrons to another, creating cations and anions held together by electrostatic attraction. Example: Na loses one electron to Cl, forming Na⁺Cl⁻.
- Covalent bond: Formed when two atoms share one or more pairs of electrons. A single bond involves one shared pair, a double bond two pairs, and a triple bond three pairs.
- Coordinate (dative) bond: A covalent bond where both shared electrons come from one atom. In NH₄⁺, the nitrogen of NH₃ donates its lone pair to H⁺.
- VSEPR theory: Electron pairs around a central atom repel each other and arrange themselves to minimise repulsion. Lone pairs repel more than bonding pairs, distorting bond angles.
- Hybridisation: Atomic orbitals of similar energy mix to form new, equivalent hybrid orbitals. sp gives linear geometry (180°), sp² gives trigonal planar (120°), sp³ gives tetrahedral (109.5°).
- Molecular Orbital Theory (MOT): Atomic orbitals combine to form bonding and antibonding molecular orbitals. Electrons fill these according to Aufbau, Pauli and Hund's rules. Bond order = (bonding electrons − antibonding electrons) / 2.
- Polarity: A bond is polar if the two atoms differ in electronegativity; a molecule is polar if the vector sum of bond dipoles is non-zero.
Formulas and facts to remember
- Bond order (MOT) = (Nb − Na) / 2, where Nb = electrons in bonding MOs, Na = electrons in antibonding MOs. Higher bond order means stronger, shorter bond.
- Formal charge = Valence electrons of free atom − [Non-bonding electrons + ½(Bonding electrons)].
- Dipole moment μ = q × d, where q is charge magnitude and d is separation distance. SI unit is coulomb-metre; often given in debye (1 D ≈ 3.336 × 10⁻³⁰ C·m).
- Hybridisation and geometry:
- sp → linear, 180° (BeCl₂, CO₂)
- sp² → trigonal planar, 120° (BF₃, C₂H₄)
- sp³ → tetrahedral, 109.5° (CH₄, NH₄⁺)
- sp³d → trigonal bipyramidal, 90° and 120° (PCl₅)
- sp³d² → octahedral, 90° (SF₆)
- VSEPR repulsion order: Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
- Sigma (σ) bond: Formed by head-on overlap of orbitals along the internuclear axis.
- Pi (π) bond: Formed by sideways overlap of p orbitals above and below the internuclear axis.
- Resonance: When a single Lewis structure cannot represent actual electron distribution, we draw multiple structures; the real molecule is a hybrid.
Common mistakes
- Thinking all sp³ molecules are tetrahedral → Shape depends on how many positions are lone pairs; NH₃ is sp³ but pyramidal, H₂O is sp³ but bent.
- Ignoring lone pairs when predicting geometry → Always count lone pairs; they occupy space and affect angles.
- Confusing molecular polarity with bond polarity → CO₂ has polar C=O bonds but is non-polar overall because dipoles cancel.
- Assuming higher bond order always means more bonds shown in Lewis structure → MOT bond order can be fractional (e.g., 2.5 in O₂⁺) even when Lewis structure shows integer bonds.
- Thinking a negative bond order is impossible → If antibonding electrons exceed bonding electrons, bond order (Nb − Na)/2 is negative; a negative or zero bond order means the molecule is unstable.
Quick revision
- Octet rule: atoms share, gain or lose electrons to reach eight valence electrons.
- VSEPR: electron pairs repel; shape is decided by bonding pairs; angles shrink when lone pairs are present.
- Hybridisation = mixing of orbitals; sp³ (109.5°), sp² (120°), sp (180°).
- Bond order (MOT) = (Nb − Na) / 2; higher bond order means stronger and shorter bond.
- Polar molecule: net dipole ≠ 0; symmetric molecules often have zero net dipole.
- Sigma bonds form by head-on overlap; pi bonds form by sideways overlap and exist in double/triple bonds.