Valence Bond Theory (VBT) is a basic concept in chemistry that explains the formation of chemical bonds through the overlap of atomic orbitals. Developed by Linus Pauling, Heitler and London (1927) VBT provides insights in how atoms combine to form molecules, emphasising the role of valence electrons. It is particularly useful for understanding covalent bonding and the shapes of molecules.
Basic Principles of Valence Bond Theory
1.Atomic Orbital Overlap:-
Valence Bond Theory (VBT) put forward that a covalent bond forms when the atomic orbitals of two atoms overlap. The overlapping region allows the sharing of electron pairs, which leads to bond formation.
Greater the overlap, stronger the bond.
2. Electron Pairing:-
The electrons involved in bonding must have opposite spins. This pairing reduces repulsion and stabilises the molecule.
3. Localised Bonds:-
Valence Bond Theory (VBT) assumes that bonds are localised between specific pairs of atoms, unlike Molecular Orbital Theory, which delocalises electrons.
4. Hybridisation:
To explain molecular geometries, Valence Bond Theory introduces the concept of hybridisation. Atomic orbitals mix to form new hybrid orbitals that align in specific geometries favorable to bond formation.
Types of Bonds in Valence Bond Theory
1.Sigma (σ) Bond:-
Formed by the head-on overlap of orbitals.
Found in all single bonds and the first bond of multiple bonds.
2. Pi (π) Bond:-
Arises from the side-by-side overlap of orbitals, such as p-orbitals.
Weaker than sigma bonds and present in double and triple bonds.
Hybridization and Molecular Geometry
One of the significant contributions of Valence Bond Theory is the concept of hybridization, which helps explain the shapes and bond angles in molecules:
1.sp Hybridization:
Linear geometry, bond angle of 180°.
Example: BeCl₂.
2. sp² Hybridization:
Trigonal planar geometry, bond angle of 120°.
Example: BF₃.
3. sp³ Hybridization:
Tetrahedral geometry, bond angle of 109.5°.
Example: CH₄.
4. sp³d Hybridization:
Trigonal bipyramidal geometry, bond angles of 90° and 120°.
Example: PCl
5. sp³d² Hybridization:
Octahedral geometry, bond angle of 90°.
Example: SF₆.
Strengths of Valence Bond Theory
1.Simple and Intuitive:
Valence Bond Theory provides a straightforward explanation for chemical bonding, focusing on the interaction between valence electrons.
2. Explains Molecular Geometry:
Through hybridization, Valence Bond Theory accurately describes the shapes and bond angles of molecules.
3. Localized Bonds:
Helps visualise bonds between specific pairs of atoms, making it easier to predict molecular behavior.
Limitations of Valence Bond Theory
1.Inability to Explain Delocalization:
Lack of Quantitative Predictions: Valence Bond Theory cannot adequately describe phenomena like resonance in bensene, where electrons are delocalised across multiple bonds.
2. Fails for Some Molecules:
Valence bond theory struggles to explain bonding in molecules with unusual geometries or bonding patterns, such as in O₂ or NO.
Applications of Valence Bond Theory
1.Bond Formation:
Used to predict the types of bonds and their orientations in molecule
2. Molecular Geometry:
Explains the spatial arrangement of atoms in molecules.
3. Predicting Reactivity:
Helps understand how molecules interact based on their bonding and structure
Note
Valence Bond Theory remains a keystone in understanding chemical bonding despite its limitations. It laid the foundation for more advanced theories and continues to be a valuable tool in explaining molecular structures and reactions at a fundamental level.
Valence Bond Theory explains the formation of covalent bonds through the overlap of atomic orbitals, where electrons are shared between atoms. It describes how bonds form and their directional nature.
According to VBT, a covalent bond forms when atomic orbitals of two atoms overlap, allowing electrons to pair in the overlapping region. Greater overlap results in stronger bonds.
Sigma (σ) Bond: Formed by the head-on overlap of orbitals, present in all single bonds.
Pi (π) Bond: Formed by the side-by-side overlap of p-orbitals, present in double and triple bonds.
Hybridization is the mixing of atomic orbitals to form new hybrid orbitals that align in specific geometries. It explains molecular shapes and bond angles, such as the tetrahedral structure of CH₄.
In CH₄, the carbon atom undergoes sp³ hybridization, forming four equivalent hybrid orbitals that overlap with hydrogen’s s-orbitals. This results in a tetrahedral geometry with bond angles of 109.5°.
Valence bond theory cannot explain delocalized bonding, as seen in resonance structures (e.g., benzene).
It does not predict magnetic properties of certain molecules like O₂.
It is less effective in providing quantitative insights compared to Molecular Orbital Theory.
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