Molecular Geometry and Bonding Theories: Lewis Structures, Formal Charges, and Resonance

Chapter 9: Molecular Geometry and Bonding Theories

Introduction: Chemistry of Vision and Molecular Geometry

The geometry of molecules plays a crucial role in chemical processes, including biological phenomena such as vision. The retinal molecule in our eyes changes geometry upon absorbing visible-light photons, triggering electrical signals that allow us to perceive light.

• Molecular geometry affects the function and reactivity of molecules.

• Example: Retinal undergoes a geometric change when exposed to light, initiating the process of vision.

Lewis Structures for Covalent Compounds

Lewis structures are visual representations of molecules that show all valence electrons, including both bonding and non-bonding (lone pair) electrons.

• Bonds shared between atoms are represented by lines.

• Lone pairs are represented by dots.

• Example: H2 is shown as H–H, and Cl2 as Cl–Cl with lone pairs around each Cl atom.

Writing Lewis Structures: Stepwise Method

Constructing Lewis structures involves a systematic approach to ensure all valence electrons are accounted for and the octet rule is satisfied where applicable.

• Step 1: Sum the valence electrons of all atoms in the molecule or ion.

  • For polyatomic anions, add electrons equal to the negative charge.

  • For polyatomic cations, subtract electrons equal to the positive charge.

• Step 2: Identify the central atom (usually the least electronegative).

• Step 3: Connect outer atoms to the central atom using single bonds.

• Step 4: Complete the octets of the outer atoms with remaining electrons.

• Step 5: Place any remaining electrons on the central atom and ensure its octet is complete.

Example: Phosphorus Trichloride (PCl3)

• Valence electrons: P (5) + 3 × Cl (7) = 26

• Three single bonds use 6 electrons; 20 electrons remain.

• Complete Cl octets (18 electrons); 2 electrons remain for P.

Lewis Structures for Molecules with Multiple Bonds

When the central atom does not achieve an octet after placing all available electrons, form multiple bonds (double or triple) as needed.

• Example: Hydrogen Cyanide (HCN)

  • After distributing electrons, a triple bond is formed between C and N to satisfy the octet rule.

Formal Charges in Lewis Structures

Formal charges help determine the most accurate Lewis structure when multiple possibilities exist. The formal charge of an atom is calculated as:

• Formal charge = (Valence electrons) – (Lone pair electrons + ½ Bonding electrons)

Assign formal charges to minimize their number and place negative charges on the most electronegative atoms.

Atom	Valence electrons	Electrons assigned	Formal charge
O	6	6	0
C	4	4	0
O	6	6	0
O (with extra electron)	6	7	-1

Example: NCO-

• The best Lewis structure has the fewest formal charges and places the negative charge on the most electronegative atom (O).

Resonance Structures

Some molecules cannot be accurately represented by a single Lewis structure. Resonance structures are used to depict delocalized electrons within molecules.

• Example: Ozone (O3)

  • Ozone has two resonance structures, each with a different arrangement of single and double bonds.

  • Experimental data shows both O–O bonds are of equal length, and each outer oxygen atom has a charge of –½.

• Resonance structures are connected by double-headed arrows and represent the true electron distribution as a hybrid of all possible structures.

Summary Table: Key Concepts in Lewis Structures and Resonance

Concept	Description	Example
Lewis Structure	Shows all valence electrons, bonds as lines, lone pairs as dots	H–H, Cl–Cl
Formal Charge	Helps select the most stable structure	NCO- with negative charge on O
Resonance	Multiple structures for delocalized electrons	Ozone (O3)

Key Equations

• Formal charge calculation: $\text{Formal charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2} \times \text{Bonding electrons})$

Additional info: These notes cover foundational concepts in molecular geometry and bonding theories, including the importance of electron arrangement in chemical and biological systems.