Why comparison is the skill here
A typical NEET bonding question does not ask you to construct a molecular orbital diagram. It asks which of four species has the shortest bond, or which molecule is polar, or how the bond angle in ammonia compares with that in water. Each is a comparison, and each is answered from a small set of principles rather than a calculation.
That means the productive preparation is building the comparison tools - bond order predicts length and strength, lone pairs compress angles, symmetry cancels dipoles - rather than memorising properties molecule by molecule.
How NEET actually asks Chemical Bonding
NTA publishes no chapter-wise weightage for NEET, so figures circulating online are coaching estimates from past papers; confirm the current paper structure in the official NTA information bulletin.
Recurring types include: predict a shape and bond angle, determine hybridisation, compare bond lengths or strengths using bond order, decide whether a molecule is polar, and explain an anomalous boiling point. Assertion-reason questions are common and reward knowing why a fact holds rather than only that it does.
Key concepts, compressed
- ✓Ionic bonding involves electron transfer, covalent bonding electron sharing, and most real bonds lie between the two.
- ✓VSEPR arranges electron pairs to minimise repulsion; the shape is then named from the atom positions alone.
- ✓Hybridisation is determined by the steric number - sigma bonds plus lone pairs.
- ✓Bond order from molecular orbital theory predicts bond length, bond strength and magnetic behaviour.
- ✓Molecular polarity depends on both bond polarity and molecular symmetry.
- ✓Hydrogen bonding explains most boiling-point anomalies and requires H bonded to N, O or F.
Comparisons you need before the questions
| Comparison | Result | Reason |
|---|---|---|
| CH₄, NH₃, H₂O bond angles | 109.5° > 107° > 104.5° | increasing lone pairs |
| Bond order and length | higher order, shorter bond | |
| Bond order and strength | higher order, stronger bond | |
| N₂ bond order | 3, diamagnetic | |
| O₂ bond order | 2, paramagnetic | two unpaired electrons |
| He₂ bond order | 0 | does not exist |
| CO₂ polarity | non-polar | linear, dipoles cancel |
| H₂O polarity | polar | bent, dipoles do not cancel |
| Steric number to hybridisation | 2 sp, 3 sp², 4 sp³ | pi bonds excluded |
| Hydrogen bonding | requires N, O or F | explains high boiling points |
| Lone pair repulsion | lp-lp > lp-bp > bp-bp | compresses bond angles |
The five mistakes that cost the most marks
- ✓Naming the shape from the hybridisation. Ammonia is sp³ but pyramidal, and water is sp³ but bent.
- ✓Assuming polar bonds mean a polar molecule. Carbon dioxide and carbon tetrachloride are both non-polar despite polar bonds.
- ✓Counting pi bonds in the steric number. Only sigma bonds and lone pairs determine hybridisation.
- ✓Describing oxygen as diamagnetic. It has two unpaired electrons and is paramagnetic.
- ✓Invoking hydrogen bonding for hydrogen attached to carbon or chlorine. It requires nitrogen, oxygen or fluorine.
Practice set 1: bond formation and parameters
1. How does an ionic bond form?
By transfer of electrons from a metal of low ionisation energy to a non-metal of high electron affinity, producing oppositely charged ions held by electrostatic attraction. The lattice energy released on assembling the crystal is what makes the process favourable overall, even though ionisation itself costs energy.
2. How does a covalent bond form, and what does a Lewis structure show?
By sharing electron pairs between atoms of comparable electronegativity, so both achieve a stable configuration. A Lewis structure shows the valence electrons as bonding pairs and lone pairs, which is the starting point for predicting shape - you cannot apply VSEPR without first identifying the lone pairs.
3. What are the main exceptions to the octet rule?
Incomplete octets, as in BF₃ where boron has six electrons; expanded octets, as in SF₆ where sulphur has twelve, possible only from period three onward where d orbitals are available; and odd-electron species such as nitric oxide. NEET questions frequently select an exception deliberately.
4. How is formal charge calculated and why is it useful?
Formal charge = valence electrons − non-bonding electrons − half the bonding electrons. It is used to choose between competing Lewis structures: the preferred structure minimises formal charges and places any negative charge on the most electronegative atom.
5. What are the main bond parameters?
Bond length, the distance between bonded nuclei; bond angle, the angle between two bonds at an atom; bond enthalpy, the energy needed to break the bond; and bond order. They are interrelated - higher bond order means shorter length and greater enthalpy, which is what makes comparison questions answerable.
6. What is resonance and what does it indicate?
The representation of a molecule by two or more Lewis structures differing only in electron placement, none of which alone describes it correctly. The actual structure is a resonance hybrid, more stable than any contributing form. Resonance also explains why all bonds in species like the carbonate ion are equivalent in length.
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Practice set 2: VSEPR and molecular shape
7. What principle underlies VSEPR theory?
Electron pairs around a central atom arrange themselves to be as far apart as possible, minimising repulsion. Both bonding and lone pairs count toward the arrangement, but only the positions of atoms determine the shape's name - which is the distinction most shape questions turn on.
8. How does the steric number determine the electron geometry?
Two gives linear, three trigonal planar, four tetrahedral, five trigonal bipyramidal and six octahedral. The steric number is sigma bonds plus lone pairs on the central atom. This gives the electron arrangement; removing the lone pairs from consideration then gives the molecular shape.
9. Compare the shapes and bond angles of methane, ammonia and water.
All three have steric number 4 and a tetrahedral electron geometry, but differ in lone pairs. Methane has none and is tetrahedral at 109.5°. Ammonia has one and is pyramidal at about 107°. Water has two and is bent at about 104.5°. The decreasing angle is the clearest evidence that lone pairs repel more strongly.
10. Why do lone pairs compress bond angles?
Because a lone pair is held by only one nucleus and therefore spreads out more, occupying more angular space than a bonding pair shared between two. The repulsion order is lone pair-lone pair, then lone pair-bonding pair, then bonding pair-bonding pair, which is exactly the pattern the methane-ammonia-water sequence shows.
11. What are the shapes of sulphur hexafluoride and phosphorus pentachloride?
Sulphur hexafluoride is octahedral with all angles 90°. Phosphorus pentachloride is trigonal bipyramidal, with three equatorial bonds at 120° and two axial bonds at 90°. Both are expanded octets, and the trigonal bipyramid is notable for having two different bond types in the same molecule.
12. What is the difference between electron geometry and molecular shape?
Electron geometry describes the arrangement of all electron pairs, including lone pairs. Molecular shape describes only the positions of the atoms. They coincide only when there are no lone pairs, which is why methane is tetrahedral in both senses while water is tetrahedral in one and bent in the other.
13. How would you compare the bond angles in NH₃ and NF₃?
The angle in NH₃ is larger. Fluorine is more electronegative than hydrogen, so it draws the bonding electrons further from nitrogen, reducing bonding-pair repulsion near the central atom and allowing the lone pair to compress the angle more. Comparisons of this kind - same shape, different substituents - are a favourite NEET format.
Practice set 3: hybridisation
14. How is hybridisation determined?
From the steric number - sigma bonds plus lone pairs on the central atom. Two gives sp, three sp², four sp³, five sp³d and six sp³d². Pi bonds are excluded, because hybrid orbitals form only sigma bonds and pi bonds come from unhybridised p orbitals.
15. What geometry and bond angle does each of sp, sp² and sp³ give?
sp gives linear at 180°, sp² gives trigonal planar at 120°, and sp³ gives tetrahedral at 109.5°. These are the ideal angles for the electron arrangement; the presence of lone pairs reduces the observed angle, as the ammonia and water cases show.
16. What is the hybridisation of carbon in methane, ethene and ethyne?
sp³ in methane with four sigma bonds, sp² in ethene with three, and sp in ethyne with two. The steric number falls as multiple bonds replace single ones, because the extra bonds are pi bonds and do not count. The geometry changes correspondingly from tetrahedral to planar to linear.
17. When does sp³d or sp³d² hybridisation occur?
When the central atom has five or six regions of electron density, which requires an expanded octet. That is only possible from period three onward, where vacant d orbitals of comparable energy are available. Phosphorus pentachloride is sp³d and sulphur hexafluoride is sp³d².
18. Why are pi bonds not counted in determining hybridisation?
Because hybrid orbitals form sigma bonds only. Pi bonds arise from sideways overlap of unhybridised p orbitals that remain outside the hybridisation scheme. Counting them would give the wrong answer for every multiply-bonded molecule, so the exclusion is essential rather than a technicality.
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Practice set 4: molecular orbital basics
19. What are bonding and antibonding molecular orbitals?
Bonding orbitals arise from constructive overlap of atomic orbitals and lie lower in energy; antibonding orbitals arise from destructive overlap, lie higher, and have a node between the nuclei. Electrons in bonding orbitals stabilise the molecule while those in antibonding orbitals destabilise it.
20. How is bond order calculated and what does it predict?
Bond order is the number of bonding electrons minus the number of antibonding electrons, divided by two. It predicts bond length and strength - higher order means shorter and stronger - and a bond order of zero means the species does not exist. This single relation answers most comparison questions in the chapter.
21. What are the bond orders of N₂ and O₂?
Nitrogen has bond order 3 from 14 electrons, giving 10 bonding and 4 antibonding. Oxygen has bond order 2 from 16 electrons, giving 10 bonding and 6 antibonding. Nitrogen's higher bond order is why its bond is shorter and stronger, and why nitrogen gas is so unreactive.
22. Why is oxygen paramagnetic?
Because two of its electrons occupy degenerate antibonding pi orbitals singly, with parallel spins, following Hund's rule. Its Lewis structure suggests all electrons are paired and predicts diamagnetism, which experiment contradicts. This discrepancy is the standard demonstration of why molecular orbital theory is needed.
23. Why does He₂ not exist?
Because its bond order is zero. Four electrons fill both the bonding and antibonding sigma orbitals equally, so the stabilisation from the bonding pair is exactly cancelled by the destabilisation from the antibonding pair. No net bond forms, which is why helium is monatomic.
24. How do you compare bond lengths using bond order?
Higher bond order means a shorter bond. So for the oxygen series, O₂⁺ with bond order 2.5 has the shortest bond, then O₂ at 2, then O₂⁻ at 1.5, then O₂²⁻ at 1 with the longest. Bond strength follows the same ordering, which makes one calculation answer two questions.
Practice set 5: polarity and intermolecular forces
25. What is a dipole moment and what does it depend on?
The product of the charge separation and the distance between the charges, measured in debye and directed from positive to negative. For a molecule it is the vector sum of all bond dipoles, so both the individual bond polarities and the molecular geometry determine it.
26. Why is carbon dioxide non-polar while water is polar?
Because of geometry. Carbon dioxide is linear, so its two bond dipoles point in opposite directions and cancel exactly. Water is bent, so its dipoles have a resultant. Both molecules have strongly polar bonds - the difference is entirely in shape, which is why geometry must be settled before polarity.
27. How does electronegativity difference relate to bond character?
A small difference gives a non-polar covalent bond, a moderate difference a polar covalent bond, and a large difference an ionic bond. The boundaries are gradual rather than sharp, so the commonly quoted value of about 1.7 is a guideline rather than a rule.
28. What conditions are required for hydrogen bonding?
Hydrogen must be covalently bonded to nitrogen, oxygen or fluorine, and there must be a lone pair on an electronegative atom to accept it. These three elements are small and electronegative enough to create the required charge concentration. Hydrogen bonded to carbon or chlorine does not qualify.
29. Why does water have an unexpectedly high boiling point?
Because of extensive hydrogen bonding. Extrapolating from the other group 16 hydrides, water should boil far below zero; it boils at 100 °C instead. Each water molecule can form up to four hydrogen bonds, which requires substantial energy to overcome. Ammonia and hydrogen fluoride show the same anomaly for the same reason.
30. What are van der Waals forces and what affects their strength?
Weak attractions arising from instantaneous and induced dipoles, present between all molecules. They strengthen with increasing molecular size and surface area, which is why boiling points rise down a homologous series and why branched isomers boil lower than straight-chain ones - branching reduces the contact area between molecules.
How to study this chapter efficiently
- ✓For every shape question follow three steps: steric number, electron geometry, then remove the lone pairs to name the shape.
- ✓Learn the methane-ammonia-water angle sequence as one fact - it demonstrates lone pair repulsion in a single line.
- ✓Memorise the bond orders of N₂, O₂, He₂ and the oxygen ions. They answer most comparison questions directly.
- ✓Settle geometry before judging polarity, since symmetry is what cancels bond dipoles.
- ✓Check that any hydrogen bond you invoke involves nitrogen, oxygen or fluorine.
- ✓Stay close to NCERT. NEET's version of this chapter is comparative rather than computational, and the textbook covers what is asked.
Turn this into active practice
Comparison questions are quick when the tools are in place and impossible when they are not - there is no way to reason out a bond order under time pressure if the electron counts are unfamiliar. That makes this chapter unusually well suited to short repeated testing.
The NEET Chemical Bonding quiz on QUFF generates fresh questions across shapes, hybridisation, bond parameters and intermolecular forces, marks them instantly and explains each answer. Do short spaced sets - recognition of this kind consolidates better with spacing than with volume.