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Kossel–Lewis approach to chemical bonding
The Kossel-Lewis approach to chemical bonding is one of the fundamental concepts in understanding how atoms combine to form molecules. This theory highlights the role of electrons in making or breaking bonds and aims to explain how atoms achieve stable electronic configurations, often called octet configurations. Let us discuss this approach in detail.
Historical context
In the early 20th century, chemists Walter Kossel and Gilbert N. Lewis independently presented theories about chemical bonding centered around electrons. They observed that the most stable compounds form when atoms achieve a stable electronic configuration similar to that of the noble gases. This observation was the birth of the octet rule.
Octet rule
The octet rule requires atoms to form bonds in such a way that they have eight electrons in their valence shell, resulting in a stable electronic configuration similar to that of the noble gases. This can be achieved through sharing, gaining, or losing electrons.
General representation of the Octet Rule: Noble Gas Configuration: ns 2 np 6 Where n represents the principal quantum number indicating the outermost shell.
Nature of chemical bonds
According to the Kossel-Lewis theory, different types of chemical bonds arise due to the interaction of electrons between atoms.
Ionic bond
Walter Kossel focused primarily on the formation of ionic bonds. Ionic bonds are formed when there is a complete transfer of electrons from one atom to another, resulting in the formation of charged ions. For example, when a sodium (Na) atom transfers electrons to a chlorine (Cl) atom, they form a sodium ion (Na +
) and a chloride ion (Cl -
), resulting in sodium chloride (NaCl
).
Example of Ionic Bond Formation: Na (1s 2 2s 2 2p 6 3s 1 ) + Cl (1s 2 2s 2 2p 6 3s 2 3p 5 ) ➡ Na + + Cl - ➡ NaCl
This transfer results in electrostatic attraction between the oppositely charged ions, holding them together in an ionic lattice.
Covalent bonds
Gilbert N. Lewis primarily discussed covalent bonds, where atoms share pairs of electrons to satisfy the octet rule. In covalent bonding, the shared electrons hold the atoms together. A classic example of this is the hydrogen molecule (H 2
), where two hydrogen atoms form a stable molecule by sharing their lone electrons.
Example of Covalent Bond Formation: H (1s 1 ) + H (1s 1 ) ➡ H 2
For a visual representation, consider the water molecule (H 2O
), where one oxygen atom shares electrons with two hydrogen atoms:
[ O ]--[ H ] | [ H ]
Lewis structures
Lewis structures are diagrams that show the bonds between the atoms of a molecule and the lone pairs of electrons that are present in the molecule. They are simple ways of showing where and how electrons are being shared or transferred, and they make understanding chemical bonding much easier.
Follow these steps to draw a Lewis structure:
- Identify the total number of valence electrons in the molecule.
- Arrange the atoms, determine the central atom (usually the least electronegative).
- Use pairs of electrons to form bonds between atoms, and try to adjust all the atoms to satisfy the octet rule.
- If any electron is left then place it as a lone pair on the central atom or elsewhere if required.
- Double or triple bonds may be necessary to obtain an octet.
Example: For carbon dioxide (CO 2
):
Total valence electrons = 4 (Carbon) + 6*2 (Oxygen) = 16 O=C=O Oxygen forms double bonds with carbon to satisfy the octet.
.. O=C=O ..
Limitations and exceptions of the octet rule
Although the octet rule is widely applied, there are some notable exceptions, including:
Incomplete octet
Some elements are stable with fewer than eight electrons in their valence shell. Boron in BF 3
is a common example.
BF 3 Lewis Structure: F | B--F | F Boron is stable with six valence electrons.
Extended octet
Elements in the third period and beyond can expand their valence shells to have more than eight electrons. Consider PF 5
:
PF 5 Lewis Structure: F | F--P--F | FF Phosphorus can hold 10 valence electrons.
Free radicals
Molecules with an odd number of electrons have unpaired electrons and are called free radicals. For example, nitric oxide (NO
):
NO Lewis Structure: .. N--O Nitrogen has seven valence electrons leading to an unpaired electron.
Conclusion
The Kossel-Lewis approach to chemical bonding provides important information about the nature of chemical bonds. By considering the transfer and sharing of electrons, it provides a framework for understanding molecular structures and properties. Despite its limitations and exceptions, the octet rule serves as a useful tool for predicting how atoms bond, leading to the formation of molecules with diverse properties.
Overall, Kossel and Lewis's contributions mark an important milestone in the field of chemical bonding, establishing foundational concepts that continue to inform modern chemistry.