Chemistry Essentials - Covalent Bonding

Overview of Bonding Types

• Metallic Bonding
• Ionic Bonding
• Covalent Bonding

Electronegativity

• Definition: How much an atom desires electrons.
• Increases as you move up and across the periodic table.
  • Shell Model and Coulomb's Law: Smaller atoms are closer to protons in the nucleus, increasing attraction.
  • Moving right: Positive values increase, enhancing attractive force.

Types of Covalent Bonds

• Nonpolar Covalent Bonds
  • Equal electronegativities.
  • Example: Oxygen molecules (O₂), Carbon-Hydrogen bonds.
• Polar Covalent Bonds
  • Different electronegativities forming a dipole (partial positive and negative charges).
  • Example: Water (H₂O), Hydrochloric acid (HCl).

Graphical Representation of Bonds

• Potential Energy vs. Inter-nuclear Distance
  • Bond Length: Distance between atoms.
  • Bond Energy: Energy required to break the bond.

Bond Characteristics

• Nonpolar Bonds: Equal sharing of electrons.
• Polar Bonds: Unequal sharing leading to partial charges.
  • Higher electronegativity atom gains a partial negative charge.

Transition to Ionic Bonds

• Continuum (gradient) from covalent to ionic based on electronegativity.

Examples

• Hydrogen Gas (H₂): Covalent bond by sharing electrons.
• Methane (CH₄): Nonpolar covalent bond.

Electronegativity Differences

• Nonpolar Covalent: Difference < 0.5
• Polar Covalent: Difference 0.5 - 1.7
• Ionic: Difference > 1.7

Identifying Bond Types

• Covalent Bonds: Usually between two nonmetals.
• Ionic Bonds: Typically involve a metal and a nonmetal.
• Properties:
  • Covalent: Gases, liquids, solids; low melting/boiling points; poor conductors.
  • Ionic: Crystalline solids; high melting/boiling points; good conductors when dissolved.

Recap

• Ranking polarity based on periodic table location.
• Use of Coulomb's law and shell model to explain electronegativity trends.