Lewis Structures — Revision Notes

⚡ 30-Second Revision

Valence Electrons: — Outermost electrons, determine bonding.
Octet Rule: — Atoms aim for 8 valence electrons (H for 2).
Steps:

1. Count total valence electrons (add for -, subtract for +). 2. Identify central atom (least electronegative, not H). 3. Form single bonds. 4. Distribute lone pairs to terminal atoms (octets first). 5. Distribute remaining lone pairs to central atom. 6. Form multiple bonds if central atom lacks octet.

Formal Charge: — $FC = V - N - \frac{1}{2}B$. Used for stability (minimize FC, negative on electronegative atoms).
Exceptions:

* Incomplete Octet: $BF_3$, $BeCl_2$. * Expanded Octet: $PCl_5$, $SF_6$, $XeF_4$ (Period 3+ elements). * Odd-Electron: $NO$, $NO_2$.

Resonance: — Multiple equivalent structures (delocalized electrons), indicated by $leftrightarrow$.

2-Minute Revision

Lewis structures are 2D diagrams showing valence electron distribution in molecules and ions. The goal is typically to satisfy the octet rule (8 electrons) for most atoms, and the duplet rule (2 electrons) for hydrogen.

The process begins by summing all valence electrons, remembering to adjust for ionic charges. The least electronegative atom (excluding hydrogen) usually serves as the central atom. Single bonds are first drawn between the central and terminal atoms, consuming two electrons per bond.

Remaining electrons are then distributed as lone pairs, prioritizing the terminal atoms to complete their octets. Any leftover electrons go to the central atom. If the central atom still lacks an octet, lone pairs from terminal atoms are converted into double or triple bonds.

Formal charge calculation, $FC = V - N - \frac{1}{2}B$, is crucial for selecting the most stable Lewis structure, aiming for minimal charges and placing negative charges on more electronegative atoms. Key exceptions to the octet rule include incomplete octets (e.

g., $BF_3$), expanded octets (e.g., $PCl_5$, $SF_6$ for Period 3+ elements), and odd-electron molecules (radicals like $NO_2$). Resonance structures are needed when electron delocalization means a single Lewis structure is insufficient, representing the molecule as a hybrid of multiple contributing forms.

These structures are fundamental for predicting molecular geometry, hybridization, and polarity.

5-Minute Revision

Lewis structures are essential tools for visualizing chemical bonding. They depict valence electrons as dots and covalent bonds as lines, helping us understand electron distribution. The core principle is the octet rule, where atoms strive for eight valence electrons for stability (hydrogen aims for two, the duplet rule).

Steps to draw a Lewis structure:

1
Count Total Valence Electrons: — Sum valence electrons for all atoms. Add one electron for each negative charge, subtract one for each positive charge. For $CO_3^{2-}$: C(4) + 3*O(6) + 2(charge) = 4 + 18 + 2 = 24 electrons.
2
Identify Central Atom: — Usually the least electronegative atom (never H). In $CO_3^{2-}$, Carbon is central.
3
Form Single Bonds: — Connect central atom to terminal atoms with single bonds. $C-O$ (3 bonds = 6 electrons used). Remaining = $24 - 6 = 18$.
4
Distribute Lone Pairs: — First, complete octets of terminal atoms. Each O needs 6 electrons (3 O * 6 e- = 18 electrons used). Remaining = $18 - 18 = 0$. Carbon currently has only 6 electrons.
5
Form Multiple Bonds: — If the central atom lacks an octet, convert lone pairs from terminal atoms into double/triple bonds. For $CO_3^{2-}$, convert one lone pair from an oxygen to form a C=O double bond. Now, Carbon has 8 electrons (1 double, 2 single bonds).

Formal Charge: Calculate $FC = V - N - \frac{1}{2}B$. For the double-bonded oxygen in $CO_3^{2-}$: $6 - 4 - \frac{1}{2}(4) = 0$. For single-bonded oxygen: $6 - 6 - \frac{1}{2}(2) = -1$. For carbon: $4 - 0 - \frac{1}{2}(8) = 0$. Sum of FCs = $0 + 0 + (-1) + (-1) = -2$, matching the ion's charge. Structures with minimal formal charges, especially negative charges on more electronegative atoms, are more stable.

Exceptions to Octet Rule:

Incomplete Octet: — $BF_3$ (B has 6 electrons).
Expanded Octet: — $PCl_5$ (P has 10 electrons), $SF_6$ (S has 12 electrons). Occurs for Period 3 and heavier elements due to d-orbitals.
Odd-Electron Molecules (Radicals): — $NO$, $NO_2$ (cannot achieve octet for all atoms).

Resonance: When multiple equivalent Lewis structures can be drawn (e.g., $O_3$, $NO_3^-$, $SO_2$), the actual molecule is a hybrid of these forms, with electrons delocalized. This concept is crucial for understanding bond lengths and stability. Mastering Lewis structures is the gateway to VSEPR theory, hybridization, and molecular polarity.

Prelims Revision Notes

Lewis structures are fundamental for NEET, providing a visual representation of valence electron distribution. The core principle is the octet rule (atoms aim for 8 valence electrons, H for 2), but exceptions are critical.

Key Steps for Drawing Lewis Structures:

1
Total Valence Electrons: — Sum valence electrons for all atoms. For anions, add electrons equal to the charge; for cations, subtract. (e.g., $NO_3^-$: N(5) + 3*O(6) + 1 = 24 e-).
2
Central Atom: — Usually the least electronegative atom (never H). If one atom is unique, it's often central.
3
Single Bonds: — Connect central atom to terminal atoms with single bonds (2 electrons each).
4
Lone Pairs on Terminal Atoms: — Distribute remaining electrons to terminal atoms to complete their octets (or duplets for H).
5
Lone Pairs on Central Atom: — Place any leftover electrons on the central atom as lone pairs.
6
Multiple Bonds: — If the central atom lacks an octet, convert lone pairs from terminal atoms into double or triple bonds to satisfy the central atom's octet.

Formal Charge (FC): $FC = (\text{Valence electrons}) - (\text{Non-bonding electrons}) - \frac{1}{2}(\text{Bonding electrons})$. The most stable Lewis structure has minimal formal charges, with negative charges on more electronegative atoms. The sum of FCs must equal the molecule's/ion's charge.

Octet Rule Exceptions:

Incomplete Octet: — Atoms with fewer than 8 valence electrons (e.g., B in $BF_3$ has 6, Be in $BeCl_2$ has 4). These are often Lewis acids.
Expanded Octet: — Atoms from Period 3 and beyond (P, S, Cl, Br, I, Xe) can accommodate more than 8 electrons due to available d-orbitals (e.g., P in $PCl_5$ has 10, S in $SF_6$ has 12, Xe in $XeF_4$ has 12).
Odd-Electron Molecules (Radicals): — Molecules with an odd total number of valence electrons (e.g., $NO$ (11), $NO_2$ (17)). These are paramagnetic and highly reactive.

Resonance Structures: Occur when a single Lewis structure cannot accurately depict electron delocalization. Multiple equivalent structures are drawn, connected by a double-headed arrow ($leftrightarrow$). The actual molecule is a resonance hybrid (e.g., $O_3$, $CO_3^{2-}$, $SO_2$).

NEET Relevance: Lewis structures are the foundation for VSEPR theory (molecular geometry), hybridization, and molecular polarity. Expect questions on identifying correct structures, formal charges, and octet rule exceptions.

Vyyuha Quick Recall

Very Clever Students Love Learning Molecules Fast!

Valence electrons (Count total)
Central atom (Identify)
Single bonds (Form)
Lone pairs (Distribute to terminal atoms)
Lone pairs (Distribute to central atom)
Multiple bonds (Form if needed)
Formal charges (Calculate for stability)