Covalent Bonding and Electron-Dot Structures: Octet Rule, Resonance, and Exceptions

Covalent Bonding and Electron-Dot Structures

Introduction to Covalent Bonding

Covalent bonding involves the sharing of electron pairs between atoms, typically nonmetals, to achieve stable electron configurations. The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, resembling the electron configuration of noble gases.

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

• Octet rule: Atoms gain, lose, or share electrons to achieve eight electrons in their valence shell.

• Electron-dot (Lewis) structures: Diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

Example: In water (H2O), each hydrogen shares one electron with oxygen, allowing oxygen to complete its octet.

Electron-Dot Structures: Practice and Application

Drawing electron-dot structures helps visualize the arrangement of electrons in molecules and ions. The total number of valence electrons is distributed to satisfy the octet rule for each atom, where possible.

• SO32- (Sulfite ion): Has 26 valence electrons.

• SO3 (Sulfur trioxide): Has 24 valence electrons and exhibits resonance.

• ClO3- (Chlorate ion): Has 26 valence electrons.

• O3 (Ozone): Has 18 valence electrons.

• NO3- (Nitrate ion): Has 24 valence electrons and resonance structures.

Example: The Lewis structure for SO32- distributes 26 electrons among the atoms, with lone pairs shown as dots.

Resonance Structures

Some molecules and ions cannot be represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for a molecule, indicating delocalized electrons.

• Resonance: Occurs when more than one valid Lewis structure can be drawn for a molecule or ion.

• Example: SO3 and NO3- both have resonance structures, with double bonds and lone pairs distributed among different atoms.

Additional info: Resonance stabilizes molecules by delocalizing electrons over multiple atoms.

Exceptions to the Octet Rule

While the octet rule applies to many molecules, there are important exceptions:

• Free radicals: Molecules or ions with an odd number of electrons, resulting in at least one unpaired electron. These do not satisfy the octet rule.

• Electron-deficient molecules: Some atoms, such as Be and B, form compounds where the central atom has fewer than eight electrons.

• Expanded octet: Atoms in period 3 or below (e.g., S, P, Cl, Xe) can have more than eight electrons due to available d orbitals.

Examples of free radicals: ClO2, NO, and NO2 all have an odd number of electrons.

Examples of electron-deficient molecules: BeCl2 and BF3 have central atoms with less than eight electrons.

Examples of expanded octet: SF6, IF5, XeF4, and ClF3 have central atoms with more than eight electrons.

Identifying Free Radicals and Electron-Deficient Molecules

To identify free radicals, count the total number of valence electrons. If the number is odd, the species is a free radical.

• ClO2: $1 \times 7 + 2 \times 6 = 19$ electrons (odd, free radical)

• NO: $5 + 6 = 11$ electrons (odd, free radical)

• NO2: $5 + 2 \times 6 = 17$ electrons (odd, free radical)

Electron-deficient molecules are most common for Be and B compounds.

Expanded Octet: Electron-Dot Structures

Atoms in period 3 or below can accommodate more than eight electrons. Drawing electron-dot structures for these molecules shows the expanded valence shell.

• SF6: $6 + 6 \times 7 = 48$ electrons

• IF5: $7 + 5 \times 7 = 42$ electrons

• XeF4: $8 + 4 \times 7 = 36$ electrons

Example: The Lewis structure for SF6 shows sulfur surrounded by six fluorine atoms, each with a full octet, and sulfur with twelve electrons in its valence shell.

Summary Table: Types of Octet Rule Exceptions

Type	Description	Examples
Free Radicals	Odd number of electrons, at least one unpaired electron	NO, NO2, ClO2
Electron-Deficient	Central atom has fewer than 8 electrons	BeCl2, BF3
Expanded Octet	Central atom has more than 8 electrons	SF6, IF5, XeF4

Key Equations and Concepts

• Counting valence electrons: Sum the valence electrons for all atoms, add electrons for negative charges, subtract for positive charges.

• Lewis structure steps:

  1. Count total valence electrons.

  2. Draw skeletal structure.

  3. Distribute electrons to satisfy octet rule.

  4. Check for resonance and exceptions.

Example equation:

$\text{Total valence electrons} = \sum \text{(valence electrons per atom)} + \text{charge}$

Additional info: Mastery of electron-dot structures and recognition of exceptions is essential for understanding molecular structure, reactivity, and chemical bonding in general chemistry.