3.6 Valence Bond Theory

This topic explains the Valence Bond Theory (VBT), which describes how covalent bonds are formed by the overlapping of atomic orbitals. It covers the main points of VBT, sigma and pi bonds, and applications of the theory.

Definition

Valence Bond Theory (VBT): A theory proposed by Heitler and London (1927) and later developed by Pauling. It states that covalent bonds are formed by the overlapping of partially filled atomic orbitals of two atoms.

Main Idea

According to Valence Bond Theory, covalent bonds are formed by the overlapping of partially filled atomic orbitals of one atom with a partially filled atomic orbital of the other. The two overlapping atomic orbitals retain their identity. The term overlap means that the two orbitals share the same common region in space. This theory has successfully explained bond energies, bond lengths, and shapes of covalent molecules.

Core Concept

A covalent bond is formed when two half-filled atomic orbitals overlap. The greater the overlap, the stronger the bond. The overlapping orbitals must have the same symmetry with respect to the bond axis. VBT explains the formation of sigma (σ) and pi (π) bonds.

Detailed Explanation

Main Points of VBT

1. A bond between two atoms is formed by the overlap of half-filled atomic orbitals of two atoms. These atomic orbitals retain their identity.

2. Electrons of overlapping orbitals have opposite spin.

3. The number of bonds formed is equal to the number of unpaired electrons present in the outermost shell of the atom.

4. By the overlapping of two orbitals, a single bond which is a sigma bond is formed. By the overlapping of additional orbitals, multiple bonds are formed (double and triple).

5. In order to form a bond, the overlapping orbitals must have the same symmetry with respect to the bond axis. A line joining the nuclei of two bonded atoms is called the bond axis.

6. Energy is released by the overlapping of orbitals. The greater the overlap between the orbitals, the greater is the energy released and the stronger will be the bond formed.

Types of Covalent Bonds

The two main types of covalent bonds obtained by overlapping are:

• Sigma bond (σ)
• Pi bond (π)

Sigma Bond (σ)

• When there is a single bond between two atoms, the bond will be a sigma bond.
• Sigma bond is formed by the head-on overlapping of atomic orbitals.
• It is formed between:
• s-s overlap
• s-px, s-py, s-pz overlap
• px-px overlap

Pi Bond (π)

• Pi bond is formed by the parallel overlapping of p orbitals.
• It is formed between:
• py-py overlap
• pz-pz overlap

Multiple Bonds

• When there is a double bond between two atoms, one bond is sigma and the other is pi.
• When there is a triple bond between two atoms, there is one sigma and two pi bonds.

Strength of Sigma and Pi Bonds

• Sigma bond is stronger than pi bond because in sigma bond, electron density is in between the line joining the two nuclei.
• In pi bond, electron density is above and below the plane.
• Hence, pi bond is more diffuse, and less amount of energy is required to break this bond.

Applications of Valence Bond Theory

Single Bond Formation:

• Electronic configuration of hydrogen atom is 1s¹.
• Each half-filled s orbital of hydrogen atoms overlaps to form H−H bond.
• Electron density is in between the line joining the two nuclei, so the bond formed is a sigma (σ) bond.

• Electronic configuration of fluorine is 1s² 2s² 2p⁵.
• Half-filled 2pz orbital of each fluorine atom head-on overlaps to form F−F bond, which is a sigma bond.

• Electronic configurations:
• Fluorine: 1s² 2s² 2p⁵
• Hydrogen: 1s¹
• The half-filled 1s orbital of H atom overlaps with the half-filled 2pz orbital of F to form H−F sigma bond.

• In oxygen molecule, each oxygen atom contains two half-filled 2py and 2pz orbitals.
• According to VBT, two bonds are formed:
• One due to head-on overlap of two half-filled p orbitals → sigma bond
• The other due to parallel overlap of half-filled p orbitals → pi bond

• Nitrogen molecule contains two nitrogen atoms.
• Each nitrogen atom contains three half-filled p orbitals.
• According to VBT, three bonds are formed:
• Half-filled 2px orbitals head-on overlap → sigma bond
• Half-filled 2py and 2pz orbitals parallel overlap → two pi bonds
• Hence, a triple bond is formed between two nitrogen atoms.

Intermediate Level

Main Points of VBT

• Bond formed by overlap of half-filled atomic orbitals.
• Overlapping orbitals retain their identity.
• Electrons have opposite spin.
• Number of bonds = number of unpaired electrons.
• Sigma bond: head-on overlap.
• Pi bond: parallel overlap.
• Greater the overlap, stronger the bond.

Sigma vs. Pi Bonds

Applications

Important Points

• 1. VBT was proposed by Heitler and London (1927) and developed by Pauling.
• 2. Covalent bonds are formed by overlapping of half-filled atomic orbitals.
• 3. Overlapping orbitals retain their identity.
• 4. Electrons in overlapping orbitals have opposite spin.
• 5. Number of bonds = number of unpaired electrons.
• 6. Sigma bond: head-on overlap; stronger.
• 7. Pi bond: parallel overlap; weaker.
• 8. Double bond = 1 sigma + 1 pi.
• 9. Triple bond = 1 sigma + 2 pi.
• 10. Examples: H₂ (sigma), O₂ (sigma + pi), N₂ (sigma + 2 pi).

Remember

• - VBT: Overlap of half-filled atomic orbitals
• - Sigma bond: Head-on overlap; stronger
• - Pi bond: Parallel overlap; weaker
• - Double bond: 1 sigma + 1 pi
• - Triple bond: 1 sigma + 2 pi
• - Examples: H₂ (sigma), O₂ (sigma + pi), N₂ (sigma + 2 pi)

Common Mistakes

• 1. Thinking all covalent bonds are sigma bonds: Double and triple bonds contain pi bonds.
• 2. Confusing sigma and pi bonds: Sigma = head-on; Pi = parallel.
• 3. Forgetting that overlapping orbitals must have opposite spins: This is essential for bond formation.
• 4. Thinking pi bonds are stronger than sigma bonds: Sigma bonds are stronger.
• 5. Ignoring the role of unpaired electrons: The number of bonds equals the number of unpaired electrons.

Quick Revision

• VBT was proposed by Heitler and London (1927) and developed by Pauling.
• Covalent bonds are formed by overlapping of half-filled atomic orbitals.
• Overlapping orbitals retain their identity.
• Electrons in overlapping orbitals have opposite spin.
• Number of bonds = number of unpaired electrons.
• Sigma bond: head-on overlap; stronger.
• Pi bond: parallel overlap; weaker.
• Double bond = 1 sigma + 1 pi.
• Triple bond = 1 sigma + 2 pi.
• Examples: H₂ (sigma), O₂ (sigma + pi), N₂ (sigma + 2 pi).

60-Second Revision

• VBT: Overlap of half-filled atomic orbitals
• Sigma bond: Head-on overlap; stronger
• Pi bond: Parallel overlap; weaker
• Double bond: 1 sigma + 1 pi
• Triple bond: 1 sigma + 2 pi
• Examples: H₂ (sigma), O₂ (sigma + pi), N₂ (sigma + 2 pi)

QUALITY VALIDATION

• Topic: 3.6 Valence Bond Theory
• Chapter: 3 (Chemical Bonding)

• Source re-read: PASS (Pages 51–55)
• Source scope followed: PASS
• Knowledge level matched: PASS
• Important concepts covered: PASS
• Definitions checked: PASS
• Equations checked: N/A
• Derivations checked: N/A
• Examples checked: PASS
• Unsupported knowledge added: NO
• Unnecessary repetition: NO
• Notes format maintained: PASS
• Important source content omitted: NO

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