Kossel-Lewis Approach to Chemical Bonding — Revision Notes

⚡ 30-Second Revision

Octet Rule — Atoms strive for 8 valence electrons (stable noble gas config).
Duplet Rule — For H, He, stability is 2 valence electrons.
Kossel's Contribution — Ionic bond = complete electron transfer (metal to non-metal).
Lewis's Contribution — Covalent bond = sharing of electron pairs (non-metal to non-metal).
Lewis Dot Structure — Visual representation of valence electrons (dots) and bonds.
Formal Charge ($FC$) — $FC = V - N - \frac{1}{2}B$. Helps select best Lewis structure.
Exceptions to Octet Rule

- Incomplete Octet: < 8 electrons (e.g., $BF_3$, $BeCl_2$). - Expanded Octet: > 8 electrons (e.g., $SF_6$, $PCl_5$, $XeF_4$). Occurs for Period 3+ elements. - Odd-Electron Molecules: Odd total valence electrons (e.g., $NO$, $NO_2$).

Limitations — Doesn't explain shapes, bond strengths, magnetic properties.

2-Minute Revision

The Kossel-Lewis approach is fundamental to understanding chemical bonding, driven by atoms' desire to achieve stable noble gas electron configurations, primarily an octet (8 valence electrons) or a duplet (2 for H/He).

Kossel explained ionic bonds as the complete transfer of electrons from metals to non-metals, forming ions. Lewis described covalent bonds as the sharing of electron pairs between non-metals, using Lewis dot structures to visualize this.

Key to this approach is the octet rule, but it has important exceptions: incomplete octets (e.g., $BF_3$ with 6 electrons on B), expanded octets (e.g., $SF_6$ with 12 electrons on S, possible for Period 3+ elements), and odd-electron molecules (e.

g., $NO$). Formal charge, calculated as $FC = V - N - \frac{1}{2}B$, helps determine the most plausible Lewis structure. Remember, this theory is qualitative and doesn't explain molecular shapes, bond strengths, or magnetic properties, which are limitations addressed by later theories.

5-Minute Revision

The Kossel-Lewis approach, a cornerstone of chemical bonding, explains how atoms achieve stability by mimicking noble gas electron configurations. This often means attaining an 'octet' of eight valence electrons, or a 'duplet' for hydrogen and helium.

Kossel's Insight (Ionic Bonding): He proposed that highly electropositive metals (like Na) transfer electrons to highly electronegative non-metals (like Cl). This complete transfer forms oppositely charged ions ($Na^+$ and $Cl^-$), which are then held together by strong electrostatic forces, forming an ionic bond. The number of electrons transferred defines 'electrovalency'.

Lewis's Insight (Covalent Bonding): Lewis suggested that atoms can achieve stability by sharing electron pairs. Each shared pair constitutes a covalent bond. He introduced 'Lewis dot structures' to represent valence electrons (dots) and shared pairs (lines or dots). For example, in $CH_4$, carbon shares four electron pairs with four hydrogen atoms, achieving an octet, while each hydrogen achieves a duplet.

Steps for Drawing Lewis Structures:

1
Count total valence electrons.
2
Identify the central atom (least electronegative, never H).
3
Form single bonds to terminal atoms.
4
Distribute remaining electrons as lone pairs to terminal atoms to complete octets.
5
Place any leftover electrons on the central atom.
6
If the central atom lacks an octet, convert lone pairs from terminal atoms into multiple bonds.

Formal Charge: A crucial concept for evaluating Lewis structures. $FC = V - N - \frac{1}{2}B$. The most stable structure minimizes formal charges and places negative charges on more electronegative atoms. For example, in $CO_2$ ($O=C=O$), all atoms have a formal charge of 0.

Exceptions to the Octet Rule: These are vital for NEET:

1
Incomplete Octet: — Central atom has < 8 electrons. Common for Group 2 and 13 elements. E.g., $BF_3$ (B has 6), $BeCl_2$ (Be has 4).
2
Expanded Octet (Hypervalent): — Central atom has > 8 electrons. Occurs for Period 3 and beyond (due to d-orbitals). E.g., $PCl_5$ (P has 10), $SF_6$ (S has 12), $XeF_4$ (Xe has 12).
3
Odd-Electron Molecules: — Total valence electrons are odd, so not all atoms can achieve an octet. These are free radicals. E.g., $NO$ (11 electrons), $NO_2$ (17 electrons).

Limitations: The Kossel-Lewis approach is qualitative. It fails to explain molecular shapes (VSEPR), bond strengths/lengths, magnetic properties ($O_2$ paramagnetism needs MOT), and the precise nature of coordinate bonds or electron delocalization (resonance). Despite these, it's a fundamental starting point for chemical bonding.

Prelims Revision Notes

The Kossel-Lewis approach is a foundational theory explaining chemical bonding based on the stability of noble gas electron configurations. The core principle is the Octet Rule: atoms tend to achieve eight electrons in their valence shell. For hydrogen and helium, it's the Duplet Rule (two electrons).

Kossel's Contribution (Ionic Bonding):

Involves complete transfer of electrons from a metal (low ionization enthalpy) to a non-metal (high electron gain enthalpy).
Forms cations (positive ions) and anions (negative ions).
Ions are held by strong electrostatic forces.
Electrovalency — is the number of electrons lost or gained.
Example: $Na \rightarrow Na^+ + e^-$, $Cl + e^- \rightarrow Cl^-$. $Na^+Cl^-$.

Lewis's Contribution (Covalent Bonding):

Involves sharing of electron pairs between non-metal atoms.
Shared pairs are bond pairs; unshared pairs are lone pairs.
Covalency — is the number of electron pairs shared.
Lewis Dot Structures — Represent valence electrons as dots. Steps:

1. Count total valence electrons. 2. Identify central atom (least electronegative, never H). 3. Form single bonds. 4. Complete octets of terminal atoms with lone pairs. 5. Place remaining electrons on central atom. 6. Form multiple bonds if central atom lacks octet.

Example: $H_2O$: O is central, 2 H single bonds, 2 lone pairs on O. O has 8, H has 2.

Formal Charge (FC):

$FC = (\text{Valence electrons in free atom}) - (\text{Non-bonding electrons}) - \frac{1}{2}(\text{Bonding electrons})$.
Sum of FCs = 0 for neutral molecule, = charge for ion.
Most stable Lewis structure has minimized FCs, with negative FCs on more electronegative atoms.

Exceptions to the Octet Rule (CRITICAL for NEET):

1
Incomplete Octet — Central atom has < 8 electrons.

* Examples: $BF_3$ (B has 6), $BeCl_2$ (Be has 4), $AlCl_3$ (Al has 6).

1
Expanded Octet (Hypervalent Molecules) — Central atom has > 8 electrons.

* Occurs for elements from Period 3 onwards (due to available d-orbitals). * Examples: $PCl_5$ (P has 10), $SF_6$ (S has 12), $IF_7$ (I has 14), $XeF_4$ (Xe has 12).

1
Odd-Electron Molecules (Free Radicals) — Total valence electrons are odd.

* Examples: $NO$ (11 electrons), $NO_2$ (17 electrons).

Limitations:

Does not explain molecular shapes (VSEPR theory).
Does not explain bond strengths, bond lengths, or bond angles.
Does not explain magnetic properties (e.g., paramagnetism of $O_2$ needs MOT).
Does not explicitly explain coordinate bonds or resonance (electron delocalization).

Focus on practicing Lewis structures, formal charge calculations, and identifying octet rule exceptions.

Vyyuha Quick Recall

For Octet Rule Exceptions, remember: Incomplete Expanded Odd.

Incomplete: Boron, Beryllium (e.g., $BF_3$, $BeCl_2$) Expanded: Period 3 and beyond (e.g., $PCl_5$, $SF_6$) Odd: Nitrogen Oxides (e.g., $NO$, $NO_2$)