Lecture on Chemical Bonds

Introduction

• Definition of Molecules: Made up of atoms participating in chemical bonds.
• Types of Bonds: Determined by the difference in electronegativity between atoms.

Types of Chemical Bonds

Ionic Bonds

• Electronegativity Difference: Greater than about 2.
• Example: Sodium (Na) and Chlorine (Cl).
  • Sodium: Loses an electron.
  • Chlorine: Gains an electron, forming chloride ions.
• Result: Formation of ions (Na⁺ and Cl⁻).
• Characteristic: Strong electrostatic attraction between oppositely charged ions.
• Nature: Electrons are transferred completely, not shared.

Covalent Bonds

• Electronegativity Difference: Less than about 1.7.
• Nature of Electrons: Shared between atoms.

Polar Covalent Bonds

• Electronegativity Difference: Between 0.5 and 1.7.
  • Example: Hydrogen and Chlorine.
  • Result: Partial charges develop.
    • More Electronegative Atom: Becomes partially negative (δ-).
    • Less Electronegative Atom: Becomes partially positive (δ+).

Nonpolar Covalent Bonds

• Electronegativity Difference: Less than 0.5.
  • Nature: Electrons shared evenly.
  • Case: Atoms of the same element share electrons precisely evenly, resulting in no partial charges.

Summary of Bond Prediction

• Nonpolar Covalent: Electronegativity difference < 0.5.
• Polar Covalent: Electronegativity difference between 0.5 and 1.7.
• Ionic: Electronegativity difference > 2.

Additional Notes

• Metallic Bonding: Mentioned as a phenomenon beyond typical bonding categories.

Conclusion

• Understanding Bonds: Based on comparing electronegativities.
• Contact Information: Professor Dave - [email protected]
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