Chemistry · Chemical Bonding · NEET
Valence Bond Theory (VBT) was given by Heitler and London (1927) and developed by Pauling. It says: when two atoms come very close, a half-filled orbital of one atom overlaps (partly merges) with a half-filled orbital of the other atom. The two electrons inside, having opposite spins, pair up. This shared, overlapped region is the covalent bond. For NEET, remember VBT explains the ENERGY and STRENGTH of a bond, which the Lewis dot and VSEPR models do not.
Think of two hydrogen atoms coming closer. Each has a half-filled 1s orbital with one electron. As they approach, the orbitals start to interpenetrate (overlap). At one exact distance the energy is lowest and most stable, and the two electrons pair up in the overlap region. That low-energy, stable state IS the H-H bond. If atoms come too close, repulsion rises and energy goes up again, so there is one best bond distance.
A bond needs two electrons of OPPOSITE spin to pair. If an orbital already has 2 electrons (fully filled), it cannot take another electron to pair. If it is empty, it has no electron to share. So a normal covalent bond in VBT comes from two HALF-FILLED orbitals, one from each atom, each giving one electron. (A coordinate bond is the exception, where one atom donates both electrons.)
Head-on / axial overlap happens along the line joining the two nuclei (the internuclear axis). This gives a SIGMA (sigma) bond. Sidewise / lateral overlap happens when p-orbitals lie parallel and overlap above and below the axis. This gives a PI (pi) bond. Sigma = end-to-end along the axis. Pi = side-by-side, perpendicular to the axis.
Bond strength depends on the EXTENT of overlap. In a sigma bond the orbitals overlap head-on along the axis, so the overlap is LARGE and the electron density between the nuclei is high, so it is strong. In a pi bond the orbitals only overlap sidewise, which is a SMALLER overlap, so it is weaker. NEET trick: a sigma bond can exist alone, but a pi bond always comes on TOP of a sigma bond (in double and triple bonds).
These are the three ways a sigma bond can form. s-s overlap: two half-filled s-orbitals overlap along the axis (as in H2). s-p overlap: a half-filled s-orbital of one atom overlaps a half-filled p-orbital of another. p-p overlap: two half-filled p-orbitals overlap end-to-end along the axis. All three are AXIAL, so all three give SIGMA bonds. Sidewise p-p overlap is different and gives a PI bond.
Simple rule: every SINGLE bond = 1 sigma. A DOUBLE bond = 1 sigma + 1 pi. A TRIPLE bond = 1 sigma + 2 pi. Also count every C-H, N-H, O-H etc. as 1 sigma. So the first line drawn between any two atoms is always a sigma; extra lines are pi bonds. This exact rule solves most NEET sigma/pi counting questions.
Match List-I (Molecule) with List-II (Number and types of bonds between the two carbon atoms). A. ethane; B. ethene; C. carbon molecule C2; D. ethyne. List-II: I. one sigma-bond and two pi-bonds; II. two pi-bonds; III. one sigma-bond; IV. one sigma-bond and one pi-bond.
The number of sigma and pi bonds in pent-2-en-4-yne is:
The number of sigma bonds, pi bonds and lone pairs of electrons in pyridine (C5H5N), respectively, are:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
Valence Bond Theory was introduced by Heitler and London in 1927 and was developed further by Linus Pauling and others. NEET sometimes asks this directly.
Extent of overlap means HOW MUCH the two orbitals merge in space. The greater the overlap, the more electron density sits between the two nuclei, and the stronger the bond. This is why a head-on sigma overlap (large) gives a stronger bond than a sidewise pi overlap (small).
Yes. The first bond between any two atoms is always a sigma bond (head-on overlap). A pi bond forms only when there is already a sigma bond, as the second or third bond in double and triple bonds. So you can have a sigma bond alone, but never a pi bond alone between two atoms in normal molecules.
Lewis structures show which electrons pair up, and VSEPR predicts shape, but neither explains the ENERGY of bond formation. VBT explains why the bond is stable, why there is a best bond length, and how orbital overlap decides bond strength. Hybridisation (the next topic) is built on top of VBT to explain real molecular shapes.
When the overlap is head-on (along the internuclear axis), s-s, s-p and p-p overlaps all give SIGMA bonds. Only when p-orbitals overlap sidewise (parallel, above and below the axis) do you get a PI bond.
