Kossel–Lewis Approach to Chemical Bonding

Full Notes

Introduction to Kössel and Lewis Concepts

Early models by Kössel and Lewis in 1916 laid the groundwork for understanding how atoms combine to form molecules. Both aimed to explain bond formation based on noble gas configuration, a state of high stability.

• Kössel’s model focused on electron transfer to form ionic bonds.
• Lewis’s model highlighted electron sharing to form covalent bonds.
• Both models introduced the octet rule, proposing that atoms combine to achieve eight electrons in their valence shell, similar to noble gases.

4.1.1 Octet Rule

The Octet Rule states that atoms tend to gain, lose, or share electrons to achieve eight electrons in their outermost shell.

• Atoms of elements like Na, Mg, Cl, O, N, etc., commonly react to achieve noble gas configurations.
• Ionic compounds result from transfer of electrons between atoms, forming charged ions:

Example Formation of sodium chloride

Na → Na⁺ + e⁻, Cl + e⁻ → Cl⁻ → NaCl

• Covalent compounds arise from electron sharing:

Example Formation of hydrogen molecule

Two hydrogen atoms share one electron pair to form H₂. (Note: Hydrogen and helium aim for two electrons in their outermost shell – the duplet rule.)

4.1.2 Covalent Bond

A covalent bond is formed when two atoms share one or more pairs of electrons. It was introduced to explain molecules like H₂, O₂, Cl₂, HCl, CH₄, etc., where atoms achieve an outer shell octet through sharing.

• Single covalent bond: One electron pair shared (e.g., Cl₂).
• Double bond: Two pairs shared (e.g., O₂).
• Triple bond: Three pairs shared (e.g., N₂).

These shared electrons occupy the valence shells of both atoms, helping each attain an octet (or duet in hydrogen’s case). Covalent bonding mainly occurs between non-metals.

4.1.3 Lewis Representation of Simple Molecules (The Lewis Structures)

Lewis structures depict molecules using dots to represent valence electrons around each atom.

Single bonding examples:

Double bonding examples:

Steps for writing Lewis structures:

1. Calculate total valence electrons.
2. Choose central atom (usually the least electronegative).
3. Form single bonds (shared pairs).
4. Distribute remaining electrons to satisfy octets.
5. Use multiple bonds if necessary.

Less structures can be useful for predicting:

• Bonding and non-bonding electrons.
• Molecular shape and resonance structures.

4.1.4 Formal Charge

Formal charge is a tool for evaluating how well electrons are distributed in a Lewis structure.

Where:

• Valence electrons are from group number in periodic table (e.g. 5 for N, 6 for O)
• Non-bonding electrons: lone pair electrons on the atom
• Bonding electrons: total electrons shared in bonds around the atom

Guidelines for Using Formal Charge

• A preferred Lewis structure will have the smallest possible formal charges.
• Avoid like charges on adjacent atoms.
• Place negative formal charges on more electronegative atoms.
• Have a total charge equal to the overall charge on the molecule or ion.

4.1.5 Limitations of the Octet Rule

Despite its usefulness, the octet rule has several exceptions:

For Example Incomplete octet

Some elements (e.g., Be, B) are stable with less than 8 electrons. For Example: BeCl₂ and BF₃.

For Example Expanded octet

Atoms in Period 3 and beyond can have more than 8 electrons. Due to availability of d-orbitals. For Example: SO₂, PCl₅, SF₆.

For Example Odd-electron molecules

Some molecules have odd numbers of electrons, leading to incomplete octets. For Example: NO.

• Transition elements: Their bonding cannot always be explained using octet theory.

Summary

• Kössel–Lewis ideas explain bonding through electron transfer or sharing to attain noble gas configurations.
• Ionic bonds form via electron transfer and covalent bonds form via sharing electron pairs.
• Lewis structures and formal charge help represent bonding and evaluate charge distribution.
• Octet rule has exceptions including incomplete octet, expanded octet and odd-electron species.