Chemical Bonding | Lec # 7 | Valence Bond Theory | organic chemistry | Dr Rizwana

Valence bond theory ties chemical bonding to how atomic orbitals overlap—small, controlled overlap leads to stable molecules, while weaker overlap produces weaker bonds. The framework treats bonding as the result of two half-filled atomic orbitals (with opposite electron spins) coming together so their electron density concentrates between nuclei, forming a more stable molecular orbital. A key quantitative guideline is that effective overlap should be less than 50%—meaning only partial overlap of orbital lobes is needed to create an effective bond; stronger overlap increases bond strength and bond energy while shortening bond length.

Bond strength and bond length are linked directly to overlap and electron density. When orbitals overlap more tightly, electron density around the nuclei increases, producing a stronger bond with higher bond energy and a shorter bond length. When overlap is reduced, electron density drops and the bond becomes weaker, with a longer bond length. The theory also emphasizes that the overlap must occur between half-filled orbitals, and the electrons involved must have opposite spins to allow pairing.

Beyond overlap, the lecture connects bond formation to potential energy minimization. Chemical bonds form when atoms approach each other to reach a lower potential energy state, rather than requiring complete octets on the central atom. The example of hydrogen illustrates this energy profile: as hydrogen atoms come together, potential energy decreases until an equilibrium separation is reached at about 0.74 nm, corresponding to the stable H–H bond. If the atoms move closer or farther from this distance, potential energy rises again, making the bond less stable.

Valence bond theory then breaks bonding into two categories based on the type of orbital overlap: sigma (σ) and pi (π) bonds. Sigma bonds arise from “head-to-head” overlap, where orbital lobes merge along the internuclear axis. The lecture describes head-to-head overlap as involving s–s, s–p, or p–p combinations aligned so that lobes overlap directly between the two nuclei. It also stresses a structural rule: the first bond formed between two atoms is always a sigma bond.

Pi bonds come from “parallel” overlap, where p-orbital lobes overlap sideways (with lobes aligned vertically in the described geometry). The lecture uses oxygen to illustrate how p–p sideways overlap produces a pi bond, while the corresponding head-to-head overlap produces the sigma bond. Nitrogen is presented as having multiple overlap possibilities: one overlap arrangement forms a sigma bond, and the other two arrangements form two pi bonds, yielding the characteristic bonding pattern.

Finally, the lecture uses BF3 and related electron-sharing logic to show how sigma bonding can form even when electron counts initially suggest an octet mismatch. In BF3, boron has three valence electrons and each fluorine has seven; bonding proceeds through orbital participation that leads to sigma bonds, with boron shifting into an excited state to enable the required overlap. Overall, valence bond theory provides a concrete, orbital-based explanation for how sigma and pi bonds form and why bond strength depends on overlap geometry and electron pairing.