Chapter 4.2: Covalent Bonding

Learning Objectives

• Define electronegativity.
• Assess the polarity of covalent bonds.

Types of Covalent Bonds

Pure Covalent Bond

• Perfectly equal sharing of electrons.
• Equal electron density around both atoms.

Polar Covalent Bond

• Unequal distribution of electrons.
• Characterized by positive and negative ends (poles).
• Partial charges indicated by Greek delta (δ):
  • Partial negative charge (δ-) on more electronegative atom.
  • Partial positive charge (δ+) on less electronegative atom.
• Example: HCl (hydrogen chloride)
  • Chlorine has a partial negative charge.
  • Hydrogen has a partial positive charge.

Electronegativity

• Defined as the tendency of an atom to attract electron density towards itself in a bond.
• Created by Linus Pauling.
• Measured on an arbitrary scale from 0 to 4.
• Not a measurable physical quantity.
• Trend:
  • Increases across the periodic table (left to right).
  • Decreases down the periodic table.
  • Fluorine is the most electronegative element.

Electronegativity vs Electron Affinity

• Electronegativity: Calculated, unitless, related to bond electrons.
• Electron Affinity: Measurable physical quantity, energy change during electron capture.

Types of Bonds Based on Electronegativity Difference (ΔEn)

• Pure Covalent: No difference in electronegativity.
• Polar Covalent: Medium difference in electronegativity.
• Ionic Bonds: Large difference in electronegativity.
  • Example: Metals (low electronegativity) often form ionic bonds with non-metals (high electronegativity).

Example: Water (H₂O)

• Oxygen is more electronegative than hydrogen.
• Electron density is higher around oxygen.
• Bonds in water are polar covalent.

Polyatomic Ions and Compounds

• Ionic compounds can contain covalent components.
• Example: Sodium nitrate (NaNO₃)
  • Contains covalent bonds within the nitrate ion.
  • Forms ionic bonds between sodium ions and nitrate ions.