Chemical Bonding is where chemistry stops being about individual atoms and starts being about the shapes and forces that hold them together — and molecular shape, in particular, ends up explaining an enormous amount of the chemistry covered in every later chapter.
The chapter builds in layers. Lewis structures give a first rough picture of how atoms share electrons. VSEPR theory then predicts actual molecular shape by treating electron pairs, bonding and lone, as mutually repelling regions that arrange themselves as far apart as possible. Hybridisation explains how an atom's orbitals reorganise themselves to match that predicted shape. Molecular orbital theory goes one level deeper still, explaining bond order and magnetic behaviour that valence bond theory alone cannot account for, most famously the paradox of oxygen's paramagnetism.
Hybridisation and molecular shape
| Hybridisation | Shape / example |
|---|---|
| sp | linear (180°); BeCl₂ |
| sp² | trigonal planar (120°); BF₃ |
| sp³ | tetrahedral (109.5°); CH₄ |
| sp³d | trigonal bipyramidal; PCl₅ |
| sp³d² | octahedral; SF₆ |
| Bond order (MOT) | (bonding electrons − antibonding electrons)/2 |
| O₂ bond order | 2 (with two unpaired electrons — paramagnetic) |
Topics covered
- Kossel-Lewis approach — and octet rule as the earliest, simplified model of bonding
- Ionic bonding — lattice energy, and the Born-Haber cycle as an application of Hess's law
- Covalent bonding and Lewis structures — and formal charge as a tool for choosing between possible structures
- VSEPR theory — predicting molecular geometry from the number of bonding and lone electron pairs
- Valence bond theory and hybridisation — sp, sp², sp³ and beyond, and how each connects to a specific predicted shape
- Molecular orbital theory — bonding and antibonding orbitals, and bond order calculation
- Bond parameters — bond length, bond angle and bond enthalpy, and what each reveals about bond strength
- Resonance — and why some molecules cannot be accurately described by a single Lewis structure
- Hydrogen bonding — intermolecular versus intramolecular, and its dramatic effect on boiling point
Why this chapter matters
In NEET and JEE, hybridisation and molecular shape identification for a given molecule is one of the single most frequently repeated question types across the entire chemistry syllabus, since this same skill underlies organic chemistry reaction mechanisms taught in later chapters. Molecular orbital theory bond order and magnetic behaviour questions, especially involving oxygen, nitrogen and their ions, are also a reliable NEET favourite.
For Class 11 boards, expect a VSEPR-based shape prediction question and a numerical or conceptual question on molecular orbital theory bond order.
Mistakes students make in this chapter
- Forgetting lone pairs count toward electron geometry even though they don't appear in the final molecular shape name. Water is tetrahedral in electron geometry but bent in molecular shape, precisely because two of the four electron pairs are lone pairs that repel more strongly and compress the bond angle.
- Assuming hybridisation always matches the number of bonded atoms alone. It matches the total number of electron domains — bonding pairs plus lone pairs — not just the visible bonds, which is exactly why ammonia (three bonds, one lone pair) is still sp³ hybridised like methane.
- Believing valence bond theory can explain oxygen's paramagnetism. It cannot — this is precisely the famous case where molecular orbital theory succeeds where valence bond theory fails, correctly predicting oxygen's two unpaired electrons.
- Confusing resonance structures with an equilibrium between different molecules. Resonance structures are not different molecules interconverting; they're different ways of drawing the same single molecule's actual electron distribution, which is really an average (hybrid) of all the contributing structures.
- Forgetting hydrogen bonding requires hydrogen bonded directly to a highly electronegative atom. Specifically nitrogen, oxygen or fluorine — hydrogen bonded to other elements does not form this same strong intermolecular attraction, regardless of overall molecular polarity.
Worked example: why ClF₃ is T-shaped, not triangular
- Count the electron domains around the central atom using the standard VSEPR count. Chlorine brings 7 valence electrons and forms 3 bonds:domains = ½(V + M) = ½(7 + 3) = 5
- Assign hybridisation and base geometry. Five electron domains always mean sp³d hybridisation, and the parent shape is trigonal bipyramidal — three equatorial positions in a plane, two axial positions perpendicular to it:3 bond pairs + 2 lone pairs → sp³d, trigonal bipyramidal
- Place the lone pairs where repulsion is least. This is the step that decides the answer. Repulsion strength runs lp–lp > lp–bp > bp–bp, so lone pairs seek the roomiest positions. An equatorial position has only two neighbours at 90°, while an axial position has three:lone pairs → equatorial (2 neighbours at 90°, not 3)
- Read off the molecular shape. Shape is named from the atom positions only, ignoring the invisible lone pairs. With two equatorial sites occupied by lone pairs, the three fluorines sit at one equatorial and two axial positions:shape = T-shaped
- Correct the bond angles. The two equatorial lone pairs squeeze the axial bonds inward, so the angle falls below the ideal 90°:F–Cl–F ≈ 87.5° (slightly less than 90°)
What the notes contain
The PDF works through VSEPR-based shape prediction for a wide range of molecules with lone pair effects explained, derives molecular orbital diagrams for O₂, N₂ and their ions with bond order calculated, and covers resonance and hydrogen bonding with worked examples showing their effect on physical properties.
How to use this: build the concept from these notes first, then do one quick NCERT pass before the exam for exact wording and diagram labels, and spend whatever time is left on previous year questions.
Questions
Why is water bent rather than linear if it has a tetrahedral electron geometry?
Water's central oxygen atom has two bonding pairs and two lone pairs, all four repelling each other into a roughly tetrahedral arrangement, but only the two bonding pairs (to hydrogen) are visible in the final molecular shape, which is therefore bent. Lone pairs additionally repel more strongly than bonding pairs, compressing the H-O-H angle slightly below the ideal 109.5°.
Why does molecular orbital theory succeed where valence bond theory fails for oxygen?
Valence bond theory predicts all of oxygen's electrons should be paired, implying no magnetism. Molecular orbital theory, by filling molecular orbitals (rather than treating bonds as simple shared electron pairs) correctly predicts two unpaired electrons in antibonding orbitals, matching oxygen's experimentally observed paramagnetism — the full molecular orbital diagram is worked through in the PDF.
What determines whether a molecule shows resonance?
Resonance occurs when more than one valid Lewis structure can be drawn for the same molecule by only moving electrons (not atoms), typically involving alternating single and double bonds or delocalised charge — the actual molecule is a hybrid of all such structures, not any single one alone.
Does this cover why HF has a higher boiling point than HCl despite F being smaller than Cl?
Yes — hydrogen bonding in HF (since hydrogen is bonded directly to the highly electronegative fluorine) creates strong intermolecular attraction absent in HCl, more than compensating for fluorine's smaller size, and this comparison is worked through directly.
All other chapters
Every chapter of Class 11 and Class 12 Physics, Chemistry and Biology — each with its own notes and PDF.