Electronic theory of chemical bonding

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The electronic theory of chemical bonding was introduced by W. Kossel and G. N. Lewis in 1916. Just as the electronic configuration of an atom is important in chemistry, such is the fundamental significance of the electronic theory while studying chemical bonding.

Each individual atom in the Periodic Table of elements wants to gain stability, and in their struggle to gain stability, these atoms form different types of chemical bonds. An atom gains stability when its valence shell electronic configuration is the same as that of the nearest noble gas element. The combining power of an atomic element with another element is known as its valency.

So, let’s see through this article how the electronic theory of chemical bonding postulates the formation of a new chemical bond. You will also find in this article interesting facts about the ionic bond, crystal structures and lattice energy. So, continue reading!

What is the electronic theory of chemical bonding -Definition

As per the electronic theory of chemical bonding, atoms tend to achieve a stable noble gas electronic configuration by gaining or losing electrons. A chemical bond is formed between constituent atoms by the complete transference of electrons from one atom to the other.

Postulates of the electronic theory of chemical bonding

The electronic theory of chemical bonding, also known as the electronic theory of valency, is based on Bohr’s atomic structure. According to Bohr’s atomic structure, an atom consists of a dense positively charged nucleus at the center. Negatively charged electrons revolve around the nucleus in specific energy levels called shells. The outermost shell of an atom is known as the valence shell. During chemical bonding electrons are removed or added to this valence shell.

Noble gas elements are present in zero group, i.e., Group VII A (or 18) of the Periodic Table, such as helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), etc. They have a total of 2 or 8 electrons in their outermost shell, which denotes a complete valence shell. 2 electrons in the valence shell of helium represent a complete duplet, while 8 valence electrons in Ne, Ar, Kr, etc., represent a complete octet electronic configuration.

These elements are inert in nature, which means they do not react to form a chemical bond with any other element as they are already stable. Contrarily, all the other Periodic elements are not stable, so they react to form a chemical bond with other elements. 

Hydrogen (H) gains stability by gaining a duplet configuration. However, almost all the other elements present in the Periodic Table tend to gain stability by gaining or losing electrons such that they achieve a complete octet electronic configuration. This is also known as the octet law.

What is an ionic or an electrovalent bond

An ionic or electrovalent bond is defined as an electrostatic force of attraction that develops between two atoms/ions via the complete transference of electrons from one to another.

An ionic bond is formed between a metal atom and a non-metal atom. Metals are electropositive in nature, while non-metals are electronegative. Electronegativity refers to the ability of an atomic element to attract electrons. So highly electronegative elements such as halogens from Group VI A (or 17) of the Periodic Table have the ability to attract electrons. Contrarily, metal atoms from the first two groups of the Periodic Table can easily lose electrons.

The electrons lost by metal atoms are consequently gained by non-metals. Metal atoms change into positively charged ions (cations), while non-metals transform into negatively charged ions (anions).

An electrostatic force of attraction develops between the oppositely charged ions, and an electrovalent bond is formed. Both the parent elements achieve a stable electronic configuration through this chemical bond, i.e., a complete octet.

Examples of ionic compounds are NaCl, MgO, CsCl, CaCl2, ZnCl2, etc.

Energetics of ionic bond formation

The ionic bond formation is an exothermic process. A large amount of energy is released as an ionic compound is formed. The amount of energy released is represented by negative enthalpy change (-∆H).

 The overall potential energy of the system decreases. The total energy released during ionic bond formation is known as bond energy. An equivalent amount of energy needs to be supplied in order to break this ionic bond. The greater the bond energy of an ionic compound, the higher is its stability. An ionic bond is usually very stable and stronger than other chemical bonds, such as covalent or metallic bonds. The strength of an ionic bond strongly dominated the properties of ionic compounds.

Properties of ionic compounds

Ionic compounds are:

  • Solids at r.t.p (having high melting points).
  • Hard and brittle.
  • Rigid.
  • Usually water soluble.
  • Poor conductors of heat and electricity in the solid state.
  • Good conductors of heat and electricity in molten or aqueous form (due to the presence of freely moving ions).

Structures of ionic compounds

A large number of oppositely charged ions arrange in an orderly, three-dimensional pattern and produce a giant crystal lattice.  The cations and anions are arranged in an alternate pattern in the crystal lattice.

The number of nearest neighbors of each ion in an ionic crystal is known as its coordination number.

The smallest structural unit in this three-dimensional arrangement is known as a unit cell.

 X-ray diffraction spectroscopy helps study the different bond lengths and bond angles present in an ionic crystal. Based on the different bond angles and bond lengths, there are seven types of unit cells and, thus, seven different crystallographic arrangements of ionic compounds:

  • Cubic
  • Tetragonal
  • Orthorhombic
  • Monoclinic
  • Triclinic
  • Hexagonal
  • Rhombohedral

Example: Sodium chloride crystal structure

Sodium chloride is the most popular example of an ionic compound. An ionic bond is formed between each Na+ and Cl ion as a Na metal atom loses 1 valence electron, which is then gained by a Cl non-metal atom. This chemical reaction occurs in gaseous form, so [Na+ Cl] is represented as an ion pair.

A large number of [Na+ Cl] ion pairs arrange three-dimensionally to give NaCl crystal lattice. Each unit cell is made up of 4 [Na+ Cl] ion pairs in this lattice structure. Thus, the crystal structure of NaCl is face-centered cubic. Each Na+ ion is surrounded by 6 Cl ions in this structure. Similarly, each Cl ion is surrounded by 6 Na+ ions, so the coordination number of the NaCl crystal is 6.

The ratio of a cation radius to an anion radius in an ionic structure is known as the radius ratio. The radius ratio of NaCl crystal is 0.53.

This close packing ensures that a huge amount of energy is required to break the strong force of attraction between oppositely charged ions and, hence, to melt the ionic crystal. NaCl thus has a high melting point (801°C). As per the thermal energy equation, Q =mc∆T, temperature (T) is directly proportional to heat or energy (Q). So, an extremely high melting point of NaCl signifies a high lattice energy.

Lattice energy

The lattice energy, also known as crystal energy, of a crystal lattice is defined as the amount of energy that evolved in bringing oppositely charged gaseous ions while forming the ionic crystal. It is equal to the amount of energy required to break the crystal into its constituent chemical entities. The lattice energy of a crystal structure such as NaCl can be theoretically calculated via the formula given below or experimentally determined by the Born-Haber cycle.

Lattice energy formula :

U = A\frac{q^{+}.q^{-}}{r^{+} + r^{-}} - B

Where r+ and r are ionic radii, while q+ and q are the charges present on the ions. A= constant (depending upon the type of crystal lattice), and B= small constant of repulsion.

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1. Adams, J. B. 2001. ‘Bonding Energy Models.’ in K. H. Jürgen Buschow, Robert W. Cahn, Merton C. Flemings, Bernhard Ilschner, Edward J. Kramer, Subhash Mahajan and Patrick Veyssière (eds.), Encyclopedia of Materials: Science and Technology (Elsevier: Oxford).

2. Chaudhary, Ghulam Rasool. 2006. Inorganic chemistry, pg. 64-70. 

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