Lecture on Ionic and Covalent Bonding

Overview

• Chapter 9 focuses on ionic and covalent bonding.
• Examines how atoms attach or attract each other to form compounds.
• Important concepts:
  • Ionic bonding
  • Covalent bonding
  • Electronegativity
  • Partial charges
  • Formal charges
  • Resonance
  • Core and valence electrons
  • Exceptions to the octet rule
  • Lewis dot structures
  • Bond strength and length

Core and Valence Electrons

• Core electrons: Shells that are filled, typically the noble gas core.
  • E.g., Sodium: Neon core
• Valence electrons: Electrons beyond the core, involved in bonding.
  • E.g., Hydrogen, Lithium, Sodium each have one valence electron.
• Group 1 Metals: One valence electron.
• Group 6 Elements: Oxygen, sulfur, selenium have six valence electrons.
• Pseudonoble Gas Core: Concept for elements like selenium with full inner shells.

Importance of Valence Electrons

• Participate in bonding and chemical reactions.
• Core electrons occupy space but do not participate in reactions.
• Valence electrons determine bonding capacity and reactivity.

Examples of Electron Configurations

• Carbon: 4 valence electrons (Group 4)
• Aluminum: 3 valence electrons (Group 3)
• Beryllium: 2 valence electrons (Group 2)

Ionic Bonds

• Formation involves transfer of electrons from one atom to another.
• Example: Sodium (Na) and Chlorine (Cl)
  • Sodium gives up an electron, becoming Na+
  • Chlorine gains an electron, becoming Cl-
  • Result: NaCl, an ionic compound with cation and anion
• General Principle: Metals lose electrons, non-metals gain electrons.

Covalent Bonds

• Formed by sharing of electrons between atoms.
• Example: Carbon and Fluorine
  • Carbon shares its 4 electrons with 4 fluorine atoms.
  • Each line in a Lewis structure indicates a shared pair of electrons.
• Indicators:
  • When two non-metals bond, it’s often covalent.
  • Two shared electrons per bond.

Identifying Bond Types

• Ionic bonds: Formed between metals and non-metals.
• Covalent bonds: Formed between non-metals.

Lewis Dot Structures

• Illustrate bonding and lone pairs of electrons.
• Steps:
  1. Count total valence electrons.
  2. Arrange atoms, typically the less electronegative in the center.
  3. Distribute electrons to satisfy the octet rule (8 electrons per atom, except hydrogen which is 2).
  4. Adjust for double bonds or rings if electron count doesn’t match.
  5. Check for correct total electron count.

Mixed Bonding in Compounds

• Some compounds have both ionic and covalent bonds.
• Example: Sodium acetate
  • Covalent bonds between carbon, oxygen, and hydrogen.
  • Ionic bond between sodium (metal) and acetate (non-metal).

Conclusion

• Understanding the nature of ionic and covalent bonds is crucial for predicting compound formation and reactivity.
• Practice drawing and interpreting Lewis structures for various compounds.