ABSTRACT:
In this article, we will discuss about the valence bond theory. Chemical bonding is a fundamental concept in chemistry that explains how atoms come together to form molecules. One of the most widely accepted theories to explain chemical bonding is the Valence Bond Theory (VBT). This theory provides a detailed understanding of the nature of chemical bonds by considering the overlapping of atomic orbitals. We will explore the Valence Bond Theory in detail, along with examples of molecular geometries and references.
INTRODUCTION OF VALENCE BOND THEORY:
In the field of chemistry, understanding the structure and geometry of molecules is crucial for predicting their properties and behavior. Valence Bond Theory (VBT) is a fundamental concept that helps explain the formation of chemical bonds and molecular geometries. This theory provides a framework for understanding the spatial arrangement of atoms in a molecule based on the overlapping of atomic orbitals. Valence Bond Theory, proposed by Linus Pauling in the 1930s. It describes the formation of chemical bonds as a result of the overlapping of atomic orbitals. According to this theory, when two atoms approach each other, their atomic orbitals overlap. It leads to the formation of a bond. The overlapping orbitals must have the same or similar energy levels and opposite spins for a stable bond to form.
TYPES OF OVERLAPPING IN VALENCE BOND THEORY:
There are three main types of overlapping that can occur between atomic orbitals: sigma (σ), pi (π), and delta (δ) bonds. Sigma bonds formed by the head-on overlap of orbitals along the internuclear axis. Pi bonds, on the other hand, formed by the sideways overlap of p orbitals. Delta bonds formed by the overlap of d orbitals and are relatively rare.
EXAMPLES OF BONDING IN VALENCE BOND THEORY:
1. HYDROGEN (H2):
let’s consider the example of the formation of a hydrogen molecule (H2). Each hydrogen atom has one electron in its 1s orbital. When two hydrogen atoms come close together, their 1s orbitals overlap, resulting in the formation of a covalent bond. In this case, the two electrons are shared between the two hydrogen atoms, creating a stable H2 molecule.
2. WATER (H2O):
let’s consider the Lewis structure of water (H2O). Oxygen has six valence electrons, while hydrogen has one valence electron each. The Lewis structure of water shows that oxygen shares two electrons with each hydrogen atom, resulting in a stable molecule. The structure can be represented as H-O-H, with two lines representing the shared electron pairs.
3. METHANE:
The Valence Bond Theory also explains the concept of hybridization, which occurs when atomic orbitals mix to form new hybrid orbitals. Hybridization is necessary to explain the observed shapes of molecules, as well as the bonding in molecules with multiple bonds. For instance, in the case of methane (CH4), carbon undergoes sp3 hybridization, where one 2s orbital and three 2p orbitals combine to form four new sp3 hybrid orbitals. These hybrid orbitals then overlap with the 1s orbitals of four hydrogen atoms, resulting in the formation of four C-H covalent bonds. The Lewis structure of methane shows that each hydrogen atom shares one electron with carbon, forming a stable molecule.
CONCLUSION:
Valence Bond Theory provides a valuable framework for understanding the formation of chemical bonds and the resulting molecular geometries. By considering the overlapping of atomic orbitals, this theory helps explain the spatial arrangement of atoms in a molecule. The examples of molecular geometries discussed in this article illustrate the diverse range of structures that can be explained using Valence Bond Theory. By understanding molecular geometries, chemists can gain insights into the properties and behavior of different compounds, enabling them to make informed predictions and design new molecules for various applications.
REFERENCES:
Pauling, L. (1931). The nature of the chemical bond. Application of results obtained from the quantum mechanics and from a theory of paramagnetic susceptibility to the structure of molecules. Journal of the American Chemical Society, 53(4), 1367-1400. https://pubs.acs.org/doi/10.1021/ja01355a027
Huheey, J. E., Keiter, E. A., & Keiter, R. L. (1993). Inorganic chemistry: principles of structure and reactivity. HarperCollins College Publishers. https://www.worldcat.org/title/26974220
Miessler, G. L., & Tarr, D. A. (2013). Inorganic chemistry. Pearson Education. https://celqusb.files.wordpress.com/2017/12/inorganic-chemistry-g-l-miessler-2014.pdf
Housecroft, C. E., & Sharpe, A. G. (2012). Inorganic chemistry. Pearson Education.
