Nothing is perfect! Haven’t you heard it too many times in your life? Yeah, and it’s true! This belief applies to chemistry as well. If you thought that the Lewis theory explained all about compounds and molecules, you are wrong! It failed to explain many concepts and that is why we have the Valence Bond Theory. Here, we will read more about the valence bond theory and also look at its limitations. Yes, even this theory isn’t perfect guys! Let’s learn why.
Why a Need for Valence Bond Theory Arose?
The theory given by Lewis explained the structure of molecules. However, it failed to explain the chemical bond formation. Similarly, VSEPR theory explained the shape of simple molecules. But, it’s application was very limited. It also failed to explain the geometry of complex molecules. Hence, scientists had to introduce the theory of valence bonds to answer and overcome these limitations.
Browse more Topics under Chemical Bonding And Molecular Structure
- Bond Parameters
- Covalent Compounds
- Fundamentals of Chemical Bonding
- Hydrogen Bonding
- Ionic or Electrovalent Compounds
- Molecular Orbital Theory
- Polarity of Bonds
- Resonance Structures
- VSEPR Theory
Valence Bond Theory
Heitler and London introduced this theory. This is primarily based on the concepts of atomic orbitals, electronic configuration of elements, the overlapping of atomic orbitals, hybridization of atomic orbitals. The overlapping of atomic orbitals results in the formation of a chemical bond. The electrons are localized in the bond region due to overlapping.
Valence bond theory describes the electronic structure of molecules. The theory says that electrons fill the atomic orbitals of an atom within a molecule. It also states that the nucleus of one atom is attracted to the electrons of another atom. Now, we move on and look at the various postulates of the valence bond theory.
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Postulates of Valence Bond Theory
- The overlapping of two half-filled valence orbitals of two different atoms results in the formation of the covalent bond. The overlapping causes the electron density between two bonded atoms to increase. This gives the property of stability to the molecule.
- In case the atomic orbitals possess more than one unpaired electron, more than one bond can be formed and electrons paired in the valence shell cannot take part in such a bond formation.
- A covalent bond is directional. Such a bond is also parallel to the region of overlapping atomic orbitals.
- Based on the pattern of overlapping, there are two types of covalent bonds: sigma bond and a pi bond. The covalent bond formed by sidewise overlapping of atomic orbitals is known as pi bond whereas the bond formed by overlapping of atomic orbital along the inter nucleus axis is known as a sigma bond.
Limitations of Valence Bond Theory
As we pointed out earlier, nothing is perfect! In a similar way, the Valence Bond theory is also not perfect. It has its own set of limitations. They are:
- It fails to explain the tetravalency of carbon.
- This theory does not discuss the electrons’ energies.
- The assumptions are about the electrons being localized to specific locations.
Solved Examples for You
Question: Based on the overlapping of orbitals, how many types of covalent bonds are formed and what are they?
Answer: Based on the overlapping of orbitals, two types of covalent bonds are formed. These are known as sigma(σ) and pi(π) bonds.
- Sigma bonds are formed by the end-to-end overlap of atomic orbitals along the inter-nuclear axis known as a head-on or axial overlap. End-on overlapping is of three types, they are s-s overlapping, s-p overlapping and p-p overlapping.
- A pi bond is formed when atomic orbitals overlap in a specific way that their axes remain parallel to each other and perpendicular to the internuclear axis.