Covalent Bonding and Electron-Dot Structures: Formal Charge and Resonance

Covalent Bonding and Electron-Dot Structures

Introduction to Covalent Bonds and Lewis Structures

Covalent bonding involves the sharing of electron pairs between atoms, resulting in the formation of molecules. Lewis structures (electron-dot structures) are a visual representation of the arrangement of valence electrons among atoms in a molecule, showing both bonding pairs and lone pairs.

• Covalent bond: A chemical bond formed by the sharing of one or more pairs of electrons between atoms.

• Lewis structure: A diagram that shows the bonding between atoms and the lone pairs of electrons in a molecule.

• Valence electrons: Electrons in the outermost shell of an atom, involved in chemical bonding.

Drawing Electron-Dot Structures

To draw a Lewis structure, follow these steps:

1. Determine the total number of valence electrons for all atoms in the molecule.

2. Arrange the atoms, usually with the least electronegative atom in the center (except hydrogen).

3. Connect atoms with single bonds (lines), then distribute remaining electrons as lone pairs to complete octets.

4. If necessary, form double or triple bonds to satisfy the octet rule.

Example: N2 (Nitrogen Gas)

• Total valence electrons: 10 (5 per N atom)

• Structure: $:N \equiv N:$ (triple bond between N atoms, each with one lone pair)

Example: XeF2 (Xenon Difluoride)

• Total valence electrons: 22 (8 from Xe, 7 from each F)

• Structure: $F: - Xe: - :F$ (Xe in the center with three lone pairs, each F with three lone pairs)

Resonance Structures

Some molecules cannot be adequately represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for a molecule, differing only in the placement of electrons.

• Resonance: The phenomenon where two or more Lewis structures can be written for a molecule, and the actual structure is a hybrid of these forms.

• Example: CO2 can be represented as $O=C=O$ or $:O-C\equiv O:$

Formal Charge

Formal charge is a bookkeeping tool used to determine the most stable Lewis structure. It is calculated for each atom in a molecule using the following formula:

• Formal charge formula:

$\text{Formal charge} = \text{Number of valence electrons in free atom} - \left(\frac{1}{2} \times \text{Number of bonding electrons}\right) - \text{Number of nonbonding electrons}$

Alternatively, for simple structures:

$\text{Formal charge} = \text{Valence electrons} - \text{Number of lines (bonds)} - \text{Number of dots (lone electrons)}$

Example: H2O (Water)

• Oxygen: $6 - 2 - 4 = 0$

• Hydrogen: $1 - 1 - 0 = 0$

Guidelines for Selecting the Best Lewis Structure

When multiple Lewis structures are possible, use formal charge to select the most stable structure:

• Structures with more zero formal charges are preferred.

• Negative (or less positive) formal charges should be placed on the most electronegative atoms.

• Avoid structures with like charges (both positive or both negative) on adjacent atoms.

Application: Resonance in CO2

Consider two resonance structures for CO2:

Structure	Formal Charge on C	Formal Charge on O (left)	Formal Charge on O (right)
A: $O=C=O$	0	0	0
B: $:O-C\equiv O:$	+1	-1	0

Structure A is preferred because all formal charges are zero.

Practice: Formal Charge Calculation for OF

Structure	Formal Charge on O	Formal Charge on F
A: $:O-F:$	-1	+1
B: $O-F$	0	0

Structure B is preferred because both atoms have zero formal charge.

Summary

• Lewis structures help visualize covalent bonding and electron arrangement.

• Formal charge is essential for evaluating the stability of possible structures.

• Resonance structures represent delocalized electrons and the true electronic structure is a hybrid.

• Always prefer structures with minimal formal charges and negative charges on more electronegative atoms.