Chemical Bonding and Molecular Structure: common doubts, answered
The questions students ask most often about Chemical Bonding and Molecular Structure, each with a short answer. For the full chapter, read the Chemical Bonding and Molecular Structure notes.
About the chapter
How is Chemical Bonding and Molecular Structure tested in NEET?
Expect shapes from VSEPR, hybridisation of the central atom, bond angle comparisons, and which molecules have zero or non-zero dipole moments. Molecular orbital questions ask for bond order, stability order and magnetic behaviour of species like O₂⁺, O₂⁻, N₂ and B₂. Hydrogen bonding appears through boiling point comparisons. Most of these reward careful counting of electron pairs and electrons rather than memorised lists.
Kössel-Lewis approach and the octet rule
Read this section in the notes →What are the exceptions to the octet rule?
There are three kinds. Some central atoms end up with fewer than eight electrons, as in BeH₂, BCl₃ and LiCl. Odd-electron molecules such as NO and NO₂ cannot give every atom an octet. And atoms from the third period onward can hold more than eight, as in PF₅, SF₆ and H₂SO₄, an expanded octet. The rule also says nothing about molecular shape.
How do you calculate formal charge in a Lewis structure?
Use formal charge = V − L − ½S, where V is the atom's valence electrons, L the electrons in its lone pairs and S the electrons in its bonds. In ozone the central oxygen gets 6 − 2 − ½(6) = +1 and the singly bonded end oxygen gets 6 − 6 − ½(2) = −1. The preferred structure usually keeps formal charges as small as possible.
Ionic bonding and lattice enthalpy
Read this section in the notes →Why do ionic compounds form if making the ions costs energy?
Because the energy released when the ions pack into a crystal more than repays that cost. Turning Na and Cl atoms into Na⁺ and Cl⁻ needs energy overall, since sodium's ionization enthalpy is larger than the energy chlorine gives out on gaining an electron. The lattice enthalpy of NaCl, 788 kJ mol⁻¹, is far greater, so the ionic solid ends up much more stable.
Bond parameters: length, angle, enthalpy and order
Read this section in the notes →Why does bond length decrease as bond order increases?
More shared electron pairs between two atoms pull the nuclei closer, so the bond gets shorter as bond order rises. Carbon-carbon bonds show this: single 154 pm, double 134 pm and triple 120 pm, with bond enthalpy rising in the same order. So a higher bond order means a shorter and stronger bond, a link used to compare species such as O₂⁺ and O₂⁻.
Resonance structures
Read this section in the notes →What is resonance and why are both O–O bonds in ozone equal?
Resonance is the idea used when a single Lewis structure cannot describe the real bonding. Ozone can be drawn with the double bond on either side, but the actual molecule is one hybrid in which the extra electron pair is spread over both bonds. So the two O–O bonds are identical, with a length between single and double bonds, and the drawn structures never exist separately.
Polarity of bonds and dipole moment
Read this section in the notes →Why does CO₂ have zero dipole moment while H₂O does not?
Dipole moment is a vector, so the bond dipoles must be added along with their directions. CO₂ is linear, and its two C=O dipoles point in opposite directions and cancel exactly. Water is bent at about 104.5°, so its two O–H dipoles add up to a resultant of about 1.85 D. For the same reason BF₃ and CH₄ are non-polar despite having polar bonds.
Why is the dipole moment of NH₃ greater than that of NF₃?
Both molecules are pyramidal with a lone pair on nitrogen, but their bond dipoles point in different directions. Nitrogen is more electronegative than hydrogen, so the N–H dipoles point towards nitrogen, the same way as the lone pair's moment, and they reinforce it. Fluorine is more electronegative than nitrogen, so the N–F dipoles point away and partly cancel the lone pair's moment.
VSEPR theory and molecular shapes
Read this section in the notes →Why is NH₃ called trigonal pyramidal and not tetrahedral?
The four electron pairs around nitrogen are arranged tetrahedrally, but one of them is a lone pair. A molecule's shape describes only where its atoms sit, so nitrogen and its three hydrogens form a pyramid on a triangular base. Water, with two lone pairs out of four, is bent for the same reason. Tetrahedral describes the electron pairs, not the shape of these molecules.
Why is the bond angle in H₂O smaller than in NH₃ and CH₄?
All three have four electron pairs around the central atom, but they carry different numbers of lone pairs. Lone pairs repel more strongly than bond pairs, in the order lp-lp > lp-bp > bp-bp, and push the bonds closer together. CH₄ has no lone pair and keeps 109.5°, NH₃ has one and narrows to about 107°, and H₂O has two and narrows to about 104.5°.
Why is SF₄ see-saw shaped and why does its lone pair sit in the equatorial position?
Sulphur in SF₄ has five electron pairs, four bonding and one lone, arranged as a trigonal bipyramid. A lone pair placed equatorially has only two neighbours at 90°, while an axial one would have three, so the equatorial site keeps repulsion lowest. With the lone pair there, the four fluorines form a see-saw. The same rule makes ClF₃ T-shaped and XeF₂ linear.
Why is XeF₄ square planar?
Xenon in XeF₄ has six electron pairs, four bonding and two lone, so the pairs point to the corners of an octahedron. The two lone pairs take opposite positions, 180° apart, which keeps their strong repulsion as small as possible. That leaves the four fluorine atoms in one plane around xenon, a square planar shape. BrF₅, with one lone pair among six, is square pyramidal instead.
Valence bond theory and orbital overlap
Read this section in the notes →Why is a sigma bond stronger than a pi bond?
A σ bond forms by head-on overlap of orbitals along the line joining the two nuclei, whereas a π bond forms by sideways overlap of parallel p orbitals above and below that line. Head-on overlap is greater, so the σ bond is stronger. That is why a double bond contains one σ and one π bond, and why the π part of C=C is the weaker part.
Hybridisation
Read this section in the notes →How do you find the hybridisation of the central atom quickly?
Add the number of σ bonds the central atom forms to the number of its lone pairs: 2 means sp, 3 sp², 4 sp³, 5 sp³d and 6 sp³d². Leave π bonds out, since they use unhybridised p orbitals. In NH₃, nitrogen has three σ bonds and one lone pair, so it is sp³; in CO₂, carbon has two σ bonds and no lone pair, so it is sp.
Why are the axial bonds in PCl₅ longer than the equatorial bonds?
PCl₅ is trigonal bipyramidal, with phosphorus sp³d hybridised, three equatorial bonds and two axial ones. Each axial bond pair lies at 90° to three equatorial pairs, while each equatorial pair has only two neighbours at 90°. The extra repulsion pushes the axial bonds outward, so they are longer and weaker than the equatorial ones, which helps explain why PCl₅ is quite reactive.
Molecular orbital theory
Read this section in the notes →Why is O₂ paramagnetic according to molecular orbital theory?
MO theory places oxygen's last two electrons one each in the two π*2p orbitals of equal energy, with parallel spins, as Hund's rule requires. Two unpaired electrons make O₂ paramagnetic, which is why liquid oxygen is pulled towards a magnet. The Lewis structure O=O pairs every electron and so wrongly predicts a diamagnetic molecule; this was a major success for MO theory.
How do you calculate bond order and use it to compare O₂, O₂⁺, O₂⁻ and O₂²⁻?
Bond order is ½(Nb − Na), with Nb and Na the electrons in bonding and antibonding orbitals. Removing an electron from O₂ takes it out of an antibonding π* orbital, so bond order rises; adding electrons does the reverse. The values are O₂⁺ 2.5, O₂ 2, O₂⁻ 1.5 and O₂²⁻ 1, so stability falls and bond length grows in that order.
Homonuclear diatomic molecules
Read this section in the notes →Why does He₂ not exist?
He₂ would hold four electrons: two in the bonding σ1s orbital and two in the antibonding σ*1s orbital. The antibonding pair cancels the bonding pair, giving a bond order of ½(2 − 2) = 0, so there is no net bond and the molecule is not stable. H₂ has bond order 1 and is stable, and He₂⁺, with bond order ½, can exist.
What is the molecular orbital energy order for N₂ versus O₂, and why does it matter?
For B₂, C₂ and N₂ the σ2pz orbital lies above the two π2p orbitals; for O₂ and F₂ it lies below them. The order decides how the lighter molecules fill: B₂ puts its last two electrons singly into the two π2p orbitals and is paramagnetic, while C₂ fills both π orbitals and is diamagnetic. For O₂, bond order and magnetism come out the same either way.
Hydrogen bonding
Read this section in the notes →Why is water a liquid but H₂S a gas at room temperature?
Water molecules are held together by hydrogen bonds, while H₂S molecules have only weak van der Waals attractions. Hydrogen bonding needs hydrogen joined to a small, highly electronegative atom: fluorine, oxygen or nitrogen. Sulphur is larger and far less electronegative, so S–H bonds are not polar enough. Breaking water's hydrogen bonds needs extra energy, which lifts its boiling point well above that of H₂S.
Why is o-nitrophenol more volatile than p-nitrophenol?
o-Nitrophenol forms a hydrogen bond within each molecule, between the –OH and the neighbouring –NO₂ group, so its molecules hardly attach to one another and escape easily. In p-nitrophenol the two groups are too far apart, so each molecule hydrogen-bonds to its neighbours, and more energy is needed to separate them. The isomers have the same mass, so the difference comes entirely from the bonding.
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