Electronegativity, Bond Polarity, and Molecular Structure

Electronegativity and Bond Polarity

Learning Objectives

• Use electronegativity to determine the polarity of a bond.

• Predict the three-dimensional structure of a molecule.

• Classify molecules as polar or nonpolar based on their structure.

• Describe the intermolecular forces between ions, polar covalent molecules, and nonpolar covalent molecules.

Electronegativity

Definition and Periodic Trends

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a dimensionless quantity, with values assigned based on experimental data (Pauling scale).

• Electronegativity increases from left to right across a period in the periodic table.

• Electronegativity increases from the bottom to the top of a group.

• Nonmetals have higher electronegativity values than metals.

• Fluorine is the most electronegative element (value = 4.0).

Example: In the periodic table, elements like oxygen, nitrogen, and chlorine have high electronegativity, while alkali and alkaline earth metals have low values.

Bond Polarity

Polarity in Covalent Bonds

The polarity of a bond depends on the difference in electronegativity between the two atoms involved:

• Nonpolar covalent bond: Electrons are shared equally between atoms (e.g., H2, Cl2).

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges (e.g., HCl).

Example: In H2, both atoms have the same electronegativity, so the bond is nonpolar. In HCl, chlorine is more electronegative, so the shared electrons are closer to Cl, making the bond polar.

Nonpolar Covalent Bonds

Characteristics and Examples

• Occurs between nonmetals with equal or nearly equal electronegativity.

• Electrons are shared equally or almost equally.

• Very small electronegativity difference (ΔEN ≈ 0.0–0.4).

Atoms	Electronegativity Difference	Type of Bond
N – N	3.0 – 3.0 = 0.0	Nonpolar covalent
Cl – Br	3.0 – 2.8 = 0.2	Nonpolar covalent
H – Si	2.1 – 1.8 = 0.3	Nonpolar covalent

Polar Covalent Bonds

Characteristics and Examples

• Occurs between nonmetal atoms with moderate electronegativity differences.

• Electrons are shared unequally, creating partial positive (δ+) and partial negative (δ−) charges.

• Electronegativity difference (ΔEN ≈ 0.5–1.8).

Atoms	Electronegativity Difference	Type of Bond
O – Cl	3.5 – 3.0 = 0.5	Polar covalent
Cl – C	3.0 – 2.5 = 0.5	Polar covalent
O – S	3.5 – 2.5 = 1.0	Polar covalent

Bond Dipole and Bond Polarity

Dipole Moments

• A dipole is created when there is a separation of charge in a polar bond.

• The positive and negative ends of the dipole are indicated by the lowercase Greek letter delta (δ+ and δ−).

• An arrow is used to show the direction of the dipole, pointing from the positive to the negative end.

Examples of Dipoles in Polar Covalent Bonds:

• C — O: Cδ+ — Oδ−

• N — O: Nδ+ — Oδ−

• Cl — O: Clδ+ — Oδ−

Predicting Type of Chemical Bond

Electronegativity Difference and Bond Type

Electronegativity Difference	Type of Bond	Electron Bonding
0.0 to 0.4	Nonpolar covalent	Electrons shared equally
0.5 to 1.8	Polar covalent	Electrons shared unequally
1.9 to 3.3	Ionic	Electrons transferred

Formula:

$\Delta EN = |EN_{atom1} - EN_{atom2}|$

Where $\Delta EN$ is the electronegativity difference.

Determining Molecular Polarity

Steps to Determine Polarity

1. Determine if the bonds are polar covalent or nonpolar covalent using electronegativity values.

2. If the bonds are polar covalent, draw the Lewis structure and determine if the dipoles cancel due to molecular symmetry.

Nonpolar molecules: All dipoles cancel due to symmetry (e.g., CO2).

Polar molecules: Dipoles do not cancel, resulting in a molecule with a positive and negative end (e.g., H2O).

Intermolecular Forces

Types of Attractive Forces

• Ionic bonds: Strongest attractive forces, present in ionic compounds, usually solid at room temperature.

• Dipole-dipole attractions: Present in polar covalent compounds; positive end of one molecule is attracted to the negative end of another.

• Hydrogen bonds: Special type of dipole-dipole attraction; occurs when hydrogen is bonded to N, O, or F. Strongest force between molecules, important in biological molecules like DNA.

• Dispersion forces (London forces): Weak attractions present in all molecules, but dominant in nonpolar molecules due to temporary dipoles.

Summary Table: Types of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength
Ionic bonds	Ions	Strongest
Hydrogen bonds	Polar molecules with H bonded to N, O, or F	Very strong (for molecular compounds)
Dipole-dipole	Polar molecules	Moderate
Dispersion forces	All molecules (especially nonpolar)	Weakest

Practice Problems (Learning Checks)

Sample Questions

• Use the electronegativity difference to identify the type of bond (nonpolar covalent, polar covalent, or ionic) between the following pairs:

  • K – N

  • N – O

  • Cl – Cl

  • H – Cl

Additional info: Students should refer to the periodic table for electronegativity values and apply the ΔEN ranges to classify each bond.