Electronic theory of valence:

 Electronic theory of valence:

The electronic theory of valence, proposed by Gilbert N. Lewis and Walther Kossel in 1916, explains chemical bonding based on the arrangement of electrons in atoms. 

The main principle of this theory is that atoms tend to achieve the stable electron configuration of a noble gas, which is an octet (eight electrons) in their outermost or valence shell. 

The only exception is for the first period elements (like hydrogen and helium), which are stable with a duplet (two electrons). Atoms achieve this stability through two primary types of bonding: ionic and covalent.

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Ionic Bonding

Ionic bonds are formed when there is a complete transfer of one or more valence electrons from one atom to another. This typically occurs between a metal and a nonmetal.

• A metal atom, which has a low electronegativity, readily donates its valence electrons to become a positively charged ion, or cation. This process reveals a previously full inner shell, which now serves as the stable octet.

• A nonmetal atom, with high electronegativity, accepts these electrons to fill its valence shell, becoming a negatively charged ion, or anion. The resulting ions are held together by strong electrostatic forces of attraction.

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Covalent Bonding

Covalent bonds are formed when atoms share pairs of valence electrons to achieve a stable octet. This type of bonding usually occurs between nonmetal atoms.

• When two atoms with similar electronegativity approach each other, neither is strong enough to completely pull electrons away from the other.

• Instead, they share one or more pairs of electrons, and these shared pairs are counted toward the valence shells of both atoms, helping each to achieve a stable octet.

• Lewis dot structures are a common way to visually represent this sharing of electrons. A single bond represents one shared pair, a double bond represents two shared pairs, and so on

Limitations of the Theory

While the electronic theory of valence, particularly the octet rule, is a foundational concept in chemistry, it has several limitations:

• It does not explain the shapes or geometries of molecules.

• It cannot fully account for the bond energies and bond lengths of molecules.

• There are many compounds that do not follow the octet rule:

• An incomplete octet is a condition where an atom in a molecule has fewer than eight valence electrons after bonding. This is an exception to the octet rule, which states that atoms tend to form bonds to achieve eight electrons in their outermost shell. Atoms that often exhibit an incomplete octet are typically from Group 2 (like Beryllium) and Group 13 (like Boron and Aluminum) of the periodic table

• Beryllium (Be) is a common example. In beryllium chloride (BeCl2​), the central beryllium atom forms two single bonds with two chlorine atoms. This results in only four valence electrons around the beryllium, an incomplete octet.

• Boron (B) is another classic case. In boron trifluoride (BF3​), the central boron atom forms three single covalent bonds with three fluorine atoms. This leaves the boron with only six valence electrons.

• Hydrogen is also often considered an exception to the octet rule, as it is stable with only two electrons, following the duet rule.

• An expanded octet is a situation where a central atom in a molecule has more than eight valence electrons. This is a significant exception to the octet rule and is also known as hypervalency.

  Why Expanded Octets Occur

  Atoms in the third period and beyond (like Sulfur, Phosphorus, and Chlorine) can have expanded octets because they have available d-orbitals that can be used for bonding. ⚛️ Elements in the second period, such as carbon, nitrogen, and oxygen, do not have accessible d-orbitals and therefore cannot expand their octet.

  Common Examples

  • Phosphorus pentachloride ($PC{l}_5$): In this molecule, the central phosphorus atom forms five single bonds with five chlorine atoms. This gives phosphorus a total of ten valence electrons (five bonding pairs), exceeding the octet.

  • Sulfur hexafluoride ($S{F}_6$): The central sulfur atom is bonded to six fluorine atoms, resulting in twelve valence electrons around the sulfur.

• It also fails to explain about some molecules, An odd electron species is a molecule or ion that has an odd number of valence electrons. This directly violates the octet rule because it is impossible to pair all electrons and give every atom a full octet. These species are also known as free radicals and are typically highly reactive.

• Nitric oxide ($NO$): It has 11 valence electrons (5 from N, 6 from O).

• Nitrogen dioxide ($N{O}_2$): It has 17 valence electrons (5 from N, 12 from 2 O).

• Chlorine dioxide ($Cl{O}_2$): It has 19 valence electrons (7 from Cl, 12 from 2 O).

More advanced theories, like Valence Bond Theory and Molecular Orbital Theory, were developed later to address these limitations by incorporating the principles of quantum mechanics