Chemical Bonding and Molecular Structure

Chemical Bonding and Molecular Structure
Chemistry
June 07, 2023

Chemical Bonding and molecular structure are among the most important chapters of chemistry. Several topics revolve around chemical bonding, for example, the formation of ionic bonds, VSPER, hybridization, etc. These topics are crucial as they form the foundational base for the high classes. The chapter mainly deals with the study of attraction between atoms and molecules. This blog post comprises each concept included in the chapter, from the definition to the theories, in more simple and easy language.

Definition of Chemical Bonding

Chemical bonding refers to forming the chemical bond between atoms or molecules and generating the chemical compound. These bonds keep the atoms together and produce compounds. These compounds rely on the strength of the chemical bonds between their constituents. The stronger the bond between the constituents, the more stable the compound is and vice versa.

A force is imposed on one by the other whenever matter interacts with another kind of matter. The energy in nature reduces when the forces are attractive. The energy increases when the forces are repellent. The chemical bond is the attractive force that connects two atoms.

Theories of Chemical Bonding and Molecular Structure

Albrecht Kossel and Gilbert Lewis correctly described the production of chemical bonding in the year 1916. Both of them explained the reason for noble gases' chemical bonding inertness.

Kossel’s Theory

  • Noble gas has been used to separate the very electronegative halogens from the strongly electropositive alkali metals.
  • Halogens may form negatively by receiving an electron, whereas alkali metals can form positively by losing an electron.
  • Both negatively and positively charged electrons have a noble configuration of eight electrons in the outermost shell.
  • The basic electrical configuration of noble gases is ns2np6.
  • The electrostatic attraction created a strong force that kept the oppositely charged particles together. As an example, MgCl2, the magnesium ion, and the chlorine ions are held together by electrostatic attraction. An electrovalent bond is a chemical link between two oppositely charged particles.

Lewis Theory

  • Atoms can be considered as a positively charged kernel in Lewis's theory.
  • The outer shell can only hold a maximum of eight electrons.
  • The eight electrons in the outer shell surrounding the kernel occupy the corners of a cube.
  • An octet configuration is a stable arrangement in which the outermost shell includes 8 electrons.
  • Atoms make chemical connections with other atoms to obtain this stable structure.
  • These chemical bonds can be established by either losing or gaining an electron (NaCl, MgCl2), with sharing of an electron being particularly important in some circumstances.
  • Valence electrons participate in creating chemical bonds, and valence electrons are the only electrons present in the outer shell.
  • Gilbert utilized Lewis symbols as a distinct notation.
  • The valency of an element corresponding to the Lewis symbol, or 8 minus the number of dots, equals the number of dots.

The Valence Shell Electron Pair Repulsion (VSEPR) Theory

The hypothesis of valence shell electron pair repulsion explains molecular shapes. It is based on the atoms' valence shells' repulsive interactions between electron pairs. Sidgwick and Powell proposed this hypothesis in 1940, and Nyholm and Gillespie developed it in 1957. The following are the main postulates of VSEPR theory:

  • The amount of valence shell electron pairs bound or non-bonded surrounding the core atom determines the molecule's structure.
  • Because their electron clouds are negatively charged, pairs of electrons in the valence shell repel one another.
  • These electron pairs want to be at positions in space that reduce repulsion and increase the distance between them.
  • The electron pairs localize on the spherical surface at maximum distance from one another, as if the valence shell were a sphere.
  • The two or three electron pairs of multiple bonds are considered a single super pair, and the multiple bonds are treated as if they were a single electron pair.
  • The VSEPR model may be used for any structure where two or more resonance structures can represent a molecule.

Molecular Orbital Theory

Another way to describe chemical bonding in molecules is to use molecular orbital theory. Mulliken and Hund presented it in 1932. The molecular orbital theory treats the entire molecule as a single entity, with all electrons moving under the influence of all of the nuclei. The following are some of the most important aspects of the molecular orbital theory:

  • There are molecular orbitals around the nuclei of molecules, much like there are atomic orbitals around the nucleus of an atom.
  • The atomic orbitals from which they are produced are completely different from the molecular orbitals. Molecular orbitals are formed when atomic orbitals combine. Conditions for the formation of molecular orbitals from atomic orbitals –
  • The energies of the combining atomic orbitals should be equal.
  • It must overlap to a large extent, as greater the overlap, the more stable the molecular form.
  • In the molecular orbital, the constituent atoms' valence electrons are thought to be moving under the influence of the nuclei of participating atoms.
  • Molecule orbitals have multiple energy levels like atomic orbitals in an isolated atom.
  • The geometries of the atomic orbitals from which molecular orbitals are created determine their shape.
  • Molecular orbitals, like atomic orbitals, are organized in order to increase energy.
  • The number of atomic orbitals combined equals the number of molecular orbitals produced in bond formation.
  • The Aufbau principle, Hund's rule, and Pauli's exclusion principle regulate the filling of electrons in molecular orbitals, much as they do in atomic orbitals.

Types of Bonding

When the substance participates in chemical bonding, the stability of the resulting compound is gauged by the type of chemical bond it comprises. The form and types of chemical bonds vary due to their strength and properties. The following are the four types of chemical bonds:

  1. Ionic bonding
  2. Covalent bonding
  3. Hydrogen bonding
  4. Polar bonding

This bonding has been formed due to the loss, gain, or sharing of electrons.

Ionic Bonding

Ionic bonding is a form of chemical bonding that involves the transfer of electrons from one atom or molecule to another. An atom loses an electron in this state, acquired by another atom. When an electron transfer occurs, the atom develops a negative charge, known as an anion, and the surviving atom develops a positive charge, known as a cation. The strength of the ionic connection is determined by the charge difference between the two atoms.

Covalent bonding

A covalent bond is formed when an atom shares its electron with another atom. This sort of chemical bonding occurs when a compound includes its carbon. When an atom shares its pair of electrons with the nucleus of another atom, a molecule is formed.

Hydrogen bonding

Compared to ionic and covalent bonding, hydrogen bonding is the weakest type of chemical connection. When the polar covalent connection between oxygen and hydrogen occurs, a positive particle charge emerges in the hydrogen. This means the electrons are attracted to the more electronegative oxygen atom. This link causes the hydrogen to be drawn towards the negative charges of any nearby atom. The qualities of water can be attributed to the hydrogen bond.

Polar bonding

Covalent bonding can be polar or nonpolar. In covalent chemical bonding, the more electrons are shared unequally, and the more electronegative atom pulls the electron pair towards itself and away from the less electronegative atom. Water is an example of polar covalent bonding.

There is a charge differential in different parts of the atom due to the irregular spacing of electrons between the atoms.

Both ends of the molecule are usually distinct, with one end being positively charged and the other being negatively charged.

What is Ionic Bond?

The strong electrostatic forces of attraction between the positive and negative charged species are known as the electrovalent or ionic bond.

Factors affecting the formation of Ionic bonds:

The following factors are responsible for the formation of the ionic bond:

  • Ionization Enthalpy: removing the electron will be easy if the ionization enthalpy is less as it leads to cation formation.
  • Electron Gain Enthalpy: The higher the electron gain enthalpy, the more energy is released, forming an anion.
  • Lattice Energy: the amount of energy generated when a certain number of gaseous positive and negative ions combine to create one mole of an ionic molecule.

Bond Characteristics

The following are the characteristics of the bond:

  1. Bond Length is the equal distance between the nuclei of two bonded atoms in molecules.

Factors affecting the bond length:

  • Size of the atom: As the size of the atom grows larger, the bond length grows longer. HI > HBr > HCl > HF HI > HBr > HCl HF HI > HBr >
  • Bond's multitudinous personality: The bond length falls as the bond order increases.
  • Hybridization type: A smaller orbital means a shorter bond length. The larger the character, the shorter the bond length.
  1. Bond Angle is considered the angle between the orbital, which contains the electronic bonding pair around the central atom in a molecule.

  1. Bonds Enthalpy: the energy required to break one mole of the bond of a specific type between two atoms in a gaseous state.

Factors affecting the bond angle:

  • Atomic Size - The larger the atoms, the longer the bond length, and the lower the bond dissociation enthalpy, i.e., the lower the bond strength.
  • Bond Multiplicity - The bond dissociation enthalpy is proportional to the multiplicity of the bond between the same two atoms. This is because, first, atoms are closer together, and second, the number of bonds to be broken is greater. The bond dissociation enthalpies of H2, O2, and N2, for example, are in the following order: H–H < O = O < N N

  • Number of Lone Pairs of Electrons Present - The more lone pairs of electrons present on bonded atoms, the stronger the repulsion between them and, thus, the lower the bond dissociation enthalpy.
  • Bond Order: As per the Lewis description of the covalent bond, the number of the bonds present between the atoms in a molecule. The greater the bond order, the greater the bond enthalpy, and the shorter the bond length.

Do you know all the macromolecules in our bodies, such as proteins, DNA, RNA, and others, are bonded together by chemical bonding? These help all the structure together. The stability of the structure depends upon the strength of the bonds. The key qualities such as melting point, boiling point, and others are mainly restricted by the strength of the chemical bonds. This chapter is essential as it carries heavy weightage in the board's eczema and the competitive exam such as JEE, NEET, and others. If you have difficulties understanding the chapters, you can take our online chemistry classes and ace your preparation. The following are the additional key benefits you can avail of:

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