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.