Covalent Bonds: Formation, Types, Lewis Structures, Electronegativity, and Polarity

Covalent Bonds

Introduction

Covalent bonds are a fundamental type of chemical bond in which atoms share electrons to achieve stable electron configurations. This topic is essential for understanding molecular structure, chemical reactivity, and the properties of substances in General Chemistry.

Section 1: Formation of Covalent Bonds

Electron Sharing and Molecular Compounds

• Covalent bonds are formed when two nonmetal atoms share one or more pairs of electrons.

• Molecules are groups of atoms held together by covalent bonds.

• In nature, elements like oxygen exist as diatomic molecules (O2), not as ions (O2−) or neutral atoms (O).

• Examples of diatomic molecules: H2, O2, N2; polyatomic molecules: S8 (octasulfur), NH4+ (ammonium).

• Halogens also form diatomic molecules: I2, Br2, F2.

• Molecular compounds are substances formed by covalent bonding. Examples include butane (C4H10), methane (CH4), and carbon dioxide (CO2).

Types of Covalent Bonds

• Two nonmetal atoms can form up to three covalent bonds, depending on the number of electron pairs needed to achieve a noble gas configuration.

• Single bond: Sharing one pair of electrons (e.g., H–H, H–Cl).

• Double bond: Sharing two pairs of electrons (e.g., O=O in oxygen, O=C=O in carbon dioxide).

• Triple bond: Sharing three pairs of electrons (e.g., N≡N in nitrogen, H–C≡C–H in ethyne/acetylene).

• Double and triple bonds are stronger and shorter than single bonds.

Orbital Overlap and Bond Formation

• Valence electron orbitals from each atom overlap to form a shared orbital, resulting in a covalent bond.

• Single covalent bonds are represented by a dash (–).

Lewis Structures of Molecules and Polyatomic Ions

Guidelines for Drawing Lewis Structures

1. Count total valence electrons for all atoms in the molecule or ion. Examples: NH3, NO2, CH4

2. Arrange atoms in a skeleton structure, choosing the central atom (usually the first in the formula or the one with highest covalency).

3. Distribute electrons as dots between and around atoms. Use dashes to represent shared pairs (bonds).

4. Apply the octet rule: If atoms lack eight electrons, shift lone pairs or convert single bonds to double/triple bonds as needed.

5. Verify electron count: Ensure all electrons are accounted for and each atom has the correct number of electrons.

Exceptions to the Octet Rule

• Odd number of electrons: Some molecules (e.g., NO) cannot achieve octets for all atoms.

• Less than an octet: Some atoms (e.g., B in BF3) are stable with fewer than eight electrons.

• More than an octet: Atoms in period 3 or higher (e.g., P in PF5) can have expanded octets.

Practice Exercise: Lewis Structures

• Draw Lewis structures for: PCl3, OF2, C2H2, CHCl3, H2S, SiF4, N2O5, NCl3, NH3, SO42−

Section 2: Electronegativity and Polarity

Electronegativity and Bond Type

• Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond.

• The difference in electronegativity (ΔEN) between two atoms determines the bond type:

Bond Type	Electronegativity Difference (ΔEN)
Ionic	ΔEN ≥ 1.7
Polar Covalent	1.7 > ΔEN > 0.4
Nonpolar Covalent	ΔEN ≤ 0.4

• Examples for calculating ΔEN: Ca–Cl, C–S, Se–P, C–O, K–F.

Polarity of Molecules

• Molecules are classified as polar or nonpolar based on bond polarity and molecular geometry.

• Nonpolar molecules have no polar bonds or their bond dipoles cancel due to symmetry (e.g., CO2, BF3, CCl4).

• Polar molecules have at least one polar bond and an asymmetric shape, so dipoles do not cancel (e.g., NH3, HCl, H2O).

• Example: CO2 has polar C–O bonds, but the linear geometry causes dipoles to cancel, making it nonpolar.

Relationship Between Polarity and Molecular Symmetry

• Molecules with only nonpolar covalent bonds are nonpolar.

• Molecules with polar covalent bonds and symmetric charge distribution are nonpolar.

• Molecules with polar covalent bonds and asymmetric charge distribution are polar.

Practice Exercise: Classifying Molecular Polarity

• Classify the following as polar or nonpolar: CH4, CH2Br2, C2H6, SO2, OF2

Summary Table: Covalent Bond Types and Examples

Bond Type	Electron Pairs Shared	Example
Single Bond	1	H–H, H–Cl
Double Bond	2	O=O, O=C=O
Triple Bond	3	N≡N, H–C≡C–H

Key Equations

• Electronegativity difference: $\Delta EN = |EN_{A} - EN_{B}|$

Additional info: The notes have been expanded to include definitions, examples, and academic context for clarity and completeness. The tables and equations are reconstructed for study purposes.