w2 Chemical bonding I- Lewis structures- Molecular geometry (VSEPR)- bond polarity___CFM__

Chemical Bonding Overview

  • Learning Outcomes:

    • Explain the octet rule for molecular bonding.

    • Apply the octet rule to draw Lewis structures.

    • Understand limitations of the octet rule.

Types of Bonds

  • Ionic Bonds:

    • Form between metals and non-metals.

    • Electrons are transferred, creating charged ions (e.g., NaCl).

    • Stability achieved via electrostatic attraction between oppositely charged ions.

  • Covalent Bonds:

    • Form between non-metals.

    • Electrons are shared for stability (e.g., H2).

    • The bond is maintained by the attraction between shared electrons and nuclear positive charges.

The Octet Rule

  • Covalent Bonds: Each atom aims for 8 valence electrons.

    • Exception: Hydrogen (H) seeks 2 electrons (duplet).

  • Ionic Bonds: Atoms will lose or gain electrons to achieve a fully filled outer shell.

    • Example: Sodium (Na) loses 1 electron, Chlorine (Cl) gains 1 electron.

Energy Changes in Ionic Bonding

  • Formation of ions may have positive energy changes due to ionization and electron gain, but the subsequent attraction between ions (Coulomb attraction) is exothermic, leading to a net energy decrease that is thermodynamically favorable.

    • Net energy change for NaCl formation: -444 kJ/mol.

Polar Bonds

  • Bonds between atoms with differing electronegativities result in unequal sharing of electrons.

    • Polar bond defined when electronegativity difference is >0.4 and <1.7.

Lewis Structures

  1. Draw skeleton structure (atoms and bonds).

  2. Count total valence electrons (adjust for charge).

  3. Determine required electrons for full shells.

  4. Calculate bonding electrons: total - valence.

  5. Assign bonding electrons to bonds.

  6. Create double/triple bonds if needed.

  7. Assign lone pairs to achieve octets.

  8. Calculate formal charge (FC = V - L - 0.5B).

Formal Charge and Structure Stability

  • Lower absolute values of formal charge indicate more stable structures.

  • Example molecules involve determining FC to assess optimal bonding arrangements.

Breakdown of the Octet Rule

  • Case 1: Odd number of valence electrons (radicals).

  • Case 2: Octet deficient (e.g., BF3).

  • Case 3: Expanded octets (e.g., PCl5).

VSEPR Theory and Molecular Geometry

  • Use VSEPR to predict molecular shapes based on electron pair repulsion.

  • Molecules with polar bonds (e.g., H2O) act differently based on geometry, affecting polarity (H2O is polar, CO2 is non-polar).

VSEPR Steps

  1. Draw Lewis structure and identify central atom.

  2. Count atoms and lone pairs around the central atom to determine steric number.

Bond Angles and Lone Pair Repulsion

  • Lone pairs create increased repulsion compared to bonding pairs, affecting bond angles (e.g., NH3 has smaller angles due to lone pair repulsion).

Summary of Key Concepts

  • Ionic and covalent bonds are fundamental to molecular structure.

  • The octet rule guides electron arrangement but has exceptions.

  • Understanding molecular geometry through VSEPR enhances comprehension of chemical behavior.