Covalent Bonding and Electronegativity

Types of Covalent Bonds

Covalent bonds are formed when atoms share electrons. The nature of the bond depends on the difference in electronegativity between the bonded atoms.

• Nonpolar Covalent Bond: Electrons are shared equally between atoms. Occurs when the electronegativity difference is less than 0.5.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges. Occurs when the electronegativity difference is between 0.5 and 1.9.

• Ionic Bond: Electrons are transferred from one atom to another. Occurs when the electronegativity difference is greater than 1.9.

Electronegativity Values: Carbon (C): 2.6 Oxygen (O): 3.5 Chlorine (Cl): 3.0

Examples

• CO2: Each C–O bond is polar covalent (), but the molecule is linear and symmetric, so the dipoles cancel and the molecule is nonpolar.

• CCl4: Each C–Cl bond is polar covalent (), but the molecule is tetrahedral and symmetric, so the overall molecular dipole is zero (nonpolar molecule).

Molecular Dipoles and Polarity

Bond Dipole vs. Molecular Dipole

A bond dipole is the separation of charge across a single bond due to differences in electronegativity. A molecular dipole is the vector sum of all bond dipoles in a molecule, determining the overall polarity.

• Bond Dipole: Pair of equal and opposite charges separated by distance.

• Molecular Dipole: Overall dipole moment of a molecule, resulting from the vector sum of all individual dipoles.

Examples

• CO2: Each C–O bond is polar, but the molecule is linear and the dipoles cancel, so CO2 is nonpolar.

• Formaldehyde (H2CO): The molecule is polar because the bond dipoles do not cancel.

Lewis Structures

Definition and Purpose

A Lewis structure is a diagram that represents the valence electrons of atoms within a molecule. Dots represent electrons, and lines represent covalent bonds. Carbon and hydrogen atoms are explicitly shown.

Steps to Draw a Lewis Structure

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

2. Determine the arrangement of atoms (generally, the least electronegative atom is central; hydrogen is never central).

3. Connect the atoms by single bonds (each bond represents two shared electrons).

4. Arrange the remaining electrons so that each atom has a complete valence shell (octet for most atoms, duet for hydrogen). There are exceptions.

Examples

• Methane (CH4): Step 1: Count valence electrons (C: 4, H: 1 × 4 = 4) Step 2: Arrange atoms (C central, H surrounding) Step 3: Connect by single bonds Step 4: Ensure full valence shells

• Formaldehyde (H2CO): Step 1: Count valence electrons (C: 4, O: 6, H: 1 × 2 = 2) Step 2: Arrange atoms (C central, O and H attached) Step 3: Connect by single and double bonds Step 4: Assign lone pairs and check octet

Additional info:

• Lewis structures are foundational for understanding molecular geometry, reactivity, and formal charge assignment.

• Formal charge can be calculated as:

Summary Table: Bond Types and Electronegativity