3.1.3 Nature of Bond in Terms of Pauling Electronegativity Values

This topic explains how the difference in electronegativity between two bonded atoms determines whether the bond is ionic, polar covalent, or nonpolar covalent. It also provides examples to illustrate these bond types.

Definition

Pauling Electronegativity Values: A numerical scale (designed by Linus Pauling) that measures the tendency of an atom to attract shared electrons in a chemical bond.

Bond Type: The nature of the chemical bond between two atoms, determined by the difference in their electronegativity values.

Main Idea

• Ionic bond: Electronegativity difference greater than 1.8.
• Polar covalent bond: Electronegativity difference between 0.4 and 1.8.
• Nonpolar covalent bond: Electronegativity difference 0.4 or lower.
• Pure covalent bond: When the electronegativity of both atoms is the same (difference = 0), the pair of electrons is shared equally.

Core Concept

When two atoms with different electronegativities form a bond, the more electronegative atom pulls the shared electrons closer to itself. This creates a polar bond. If the difference is large enough, the bond becomes ionic. If the difference is very small or zero, the bond is nonpolar covalent.

Detailed Explanation

Bond Type Based on Electronegativity Difference

Example 3.1: Determine the Bond Types of the Following Substances

**(i) Oxygen molecule (O=O)**

• This molecule consists of two oxygen atoms.
• Oxygen has an electronegativity value of 3.44.
• So each atom in this molecule has an electronegativity of 3.44.
• The difference in electronegativity is 0.
• This means that none of the two atoms in the molecule pulls the electrons to itself more strongly than the other.
• This means that this molecule has a covalent bond (nonpolar covalent).

(ii) Potassium Chloride (KCl)

• This compound has a potassium atom with an electronegativity of 0.82.
• It has a chlorine atom with an electronegativity of 3.16.
• There is a difference in electronegativity of 2.34, which is greater than 1.8.
• Hence the bond will be ionic.

(iii) Hydrogen Chloride (HCl)

• This molecule consists of a hydrogen atom with an electronegativity of 2.20.
• It has a chlorine atom with an electronegativity of 3.16.
• In this molecule, chlorine is more electronegative than hydrogen.
• This means that chlorine pulls the electrons more closely to itself than the hydrogen atom.
• The difference between their electronegativities is 0.96, which is between 0.4 and 1.8.
• Hence the bond will be polar covalent.

Summary of Examples

Additional Examples

Intermediate Level

Bond Type Determination

Example Calculations

• O = 3.44, O = 3.44
• Difference = 0
• Bond type = Nonpolar covalent

• K = 0.82, Cl = 3.16
• Difference = 2.34
• Bond type = Ionic

• H = 2.20, Cl = 3.16
• Difference = 0.96
• Bond type = Polar covalent

Explanation

• Ionic bond: The more electronegative atom takes the electron completely.
• Polar covalent bond: The more electronegative atom pulls the shared electrons closer to itself, creating partial charges (δ+ and δ−).
• Nonpolar covalent bond: Both atoms share the electrons equally.

Important Points

• 1. Ionic bond: Electronegativity difference > 1.8.
• 2. Polar covalent bond: Electronegativity difference between 0.4 and 1.8.
• 3. Nonpolar covalent bond: Electronegativity difference ≤ 0.4.
• 4. Pure covalent bond: Electronegativity difference = 0.
• 5. Example: O₂ (nonpolar), KCl (ionic), HCl (polar covalent).
• 6. Partial charges: δ+ and δ− in polar covalent bonds.

Remember

• - Ionic: > 1.8
• - Polar covalent: 0.4 – 1.8
• - Nonpolar covalent: ≤ 0.4
• - Pure covalent: 0
• - Examples: O₂ (nonpolar), KCl (ionic), HCl (polar covalent)

Common Mistakes

• 1. Thinking all covalent bonds are nonpolar: Polar covalent bonds have a difference between 0.4 and 1.8.
• 2. Forgetting the ranges: Ionic > 1.8; Polar covalent 0.4–1.8; Nonpolar covalent ≤ 0.4.
• 3. Confusing electronegativity with electron affinity: Electronegativity is about attracting shared electrons; electron affinity is about gaining an electron.
• 4. Ignoring partial charges: Polar covalent bonds have δ+ and δ− charges.

Quick Revision

• Ionic bond: Electronegativity difference > 1.8.
• Polar covalent bond: Electronegativity difference between 0.4 and 1.8.
• Nonpolar covalent bond: Electronegativity difference ≤ 0.4.
• Pure covalent bond: Electronegativity difference = 0.
• Example: O₂ (nonpolar), KCl (ionic), HCl (polar covalent).
• Partial charges: δ+ and δ− in polar covalent bonds.

60-Second Revision

• Ionic: > 1.8
• Polar covalent: 0.4 – 1.8
• Nonpolar covalent: ≤ 0.4
• Pure covalent: 0
• Examples: O₂ (nonpolar), KCl (ionic), HCl (polar covalent)

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