Chemical Bonding Lecture Notes

Introduction to Compounds and Chemical Bonds

• Compound: A substance made from two or more different elements chemically combined.
• Chemical Bonds: Attractive forces that hold different elements together.

Types of Chemical Bonds

• Ionic Bonding: Involves the transfer of electrons.
• Covalent Bonding: Involves the sharing of electrons.

Octet Rule

• Noble gases are inert due to having 8 electrons in their outer energy level (except Helium).
• Octet Rule: Atoms tend to reach an electron arrangement with eight electrons in the outermost energy level.
• Exceptions:
  • Transition metals usually do not follow the octet rule.
  • Elements near helium (e.g., hydrogen, lithium, beryllium) aim to achieve the helium configuration.

Ionic Bonding

• Ion: A charged atom or group of atoms.
  • Cations: Positive ions (elements in Groups 1 & 2 tend to lose electrons).
  • Anions: Negative ions (elements in Groups 16 & 17 tend to gain electrons).
• Ionic Bond: Attraction between oppositely charged ions.
• Crystal Lattice: 3D arrangement of ions in ionic compounds.
• Examples: NaCl, FeCl2, FeCl3, Cu2O, CuO.

Covalent Bonding

• Molecule: A group of atoms joined together, the smallest particle of an element or compound that can exist independently.
• Valency indicates the number of atoms an element can combine with:
  • Example: HCl, H2O.
• Bond Types:
  • Single Bond: One pair of electrons shared
  • Double Bond: Two pairs of electrons shared
  • Triple Bond: Three pairs of electrons shared
  • Example: Nitrogen molecule with a triple bond.
• Sigma and Pi Bonds:
  • Sigma bonds from head-on orbital overlaps.
  • Pi bonds from sideways overlaps of p orbitals.

Characteristics of Ionic and Covalent Bonds

• Ionic Compounds:
  • Hard to cut, high melting and boiling points, conduct electricity when melted or dissolved.
• Covalent Compounds:
  • Usually soft, low melting and boiling points, do not conduct electricity.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• Determines molecular shapes based on electron pair repulsions around central atoms.

Electronegativity

• Electronegativity: Measure of an atom's ability to attract electrons.
• Polar Covalent Bond: Electrons not shared equally (e.g., HCl).
• Pure Covalent Bond: Electrons shared equally.
• Predicting bond type through electronegativity difference:

  • 1.7: Ionic bonding

  • ≤1.7: Covalent bonding

  • 0.4 but <1.7: Polar covalent

  • ≤0.4: Non-polar

Intermolecular Forces

• Van der Waals Forces: Weak attractions due to temporary electron shifts.
• Dipole-Dipole Forces: Attraction between poles of polar molecules.
• Hydrogen Bonding: Strong dipole-dipole attraction involving hydrogen bonded to F, O, or N.

Additional Concepts

• Transition Metals: Form colored compounds, used as catalysts, have variable valency due to small energy differences between 4s and 3d sublevels.
• Solubility: Polar covalent compounds dissolve in polar solvents like water.