Chemical Bonding and Intermolecular Forces

Ionization Energy

• Definition: The energy required to remove an electron from a gaseous atom or ion.

  • Structured as: M(g) + energy → M⁺(g) + e⁻.

• First Ionization Energy: Energy required to remove the first electron.

• Subsequent Ionization Energies (2nd, 3rd, etc.): Increased energy requirements for subsequent electron removal, especially when electrons are removed from lower principal energy levels.

  • Example: Magnesium (Mg):

  • 1st Ionization: 736 kJ/mol

  • 2nd Ionization: 1450 kJ/mol

  • 3rd Ionization: 7740 kJ/mol

• Stability: Formation of stable electron configurations after multiple ionizations (e.g., Mg²⁺).

Electronegativity

• Definition: Measure of an atom's ability to attract electrons.

  • Trends in the Periodic Table:

  • Increases from left to right across a period.

  • Decreases down a group due to increased distance of valence electrons from the nucleus.

  • Key Element: Fluorine has the highest electronegativity value.

  
  

• Electronegativity Values Example:

  Element	EN Value
  H	2.2
  He	-
  Li	1.0
  Na	0.93
  Cl	3.0
  F	4.0
  Rn	-

  • Types of Bonds:

    • Metallic Bonding: Non-directional, delocalized electrons around positive metal ions (e.g., Na, Mg).

    • Ionic Bonding: Electron transfer to form charged ions; typically involves metals losing electrons and non-metals gaining electrons (e.g., NaCl).

    • Covalent Bonding: Electron sharing between non-metals to achieve stable electron configurations (e.g., O₂, H₂O).

  Properties of Solids

  • Metallic Solids:

    • Variable hardness, high melting/boiling points (500°C–5500°C), malleable, ductile, good conductors.

  • Ionic Solids:

    • Hard & brittle, high melting/boiling points (200°C–2000°C), soluble in polar solvents, conduct electricity when molten/aqueous.

  • Covalent Network Solids:

    • Very hard, high melting/boiling points (1000°C–4000°C), non-conductors (except graphite).

  • Covalent Molecular Solids:

    • Soft, low melting/boiling points (-200°C–100°C), non-conductors, often soluble in non-polar solvents.

  Intermolecular Forces

  • Definition: Forces of attraction between individual molecules in covalent substances.

  • Types of Intermolecular Forces:

    1. Dispersion Forces: Weak forces between all molecules; significant in larger molecules.

    2. Dipole-Dipole Forces: Attraction between polar molecules with net dipoles.

    3. Ion-Dipole Forces: Attraction between an ion and a polar molecule (e.g., Na⁺ and H₂O).

    4. Hydrogen Bonding: Strong forces between hydrogen and electronegative atoms (N, O, F).

  Summary of Characteristics of Intermolecular Forces

  • Dispersion Forces: Present in all substances; particularly in non-polar molecules; weak.

  • Dipole Forces: Occur in polar molecules; results from differences in electronegativity leading to dipoles.

  • Hydrogen Bonds: Strongest form of intermolecular forces when H is bonded to electronegative atoms.

  Conclusion

  • Understanding the concepts of ionization energy, electronegativity, types of bonding, properties of solids, and intermolecular forces is crucial for mastering the fundamentals of chemical bonding and interactions in Chemistry for Life Sciences.

  • Prepare adequately for practical applications in upcoming lab sessions and quizzes by reviewing the content thoroughly.