Chemical Bonds and Molecular Interactions

Water Molecule

• The water molecule is composed of hydrogen (H) and oxygen (O).

Bohr Model (Electron Shell Diagram)

• Electrons orbit the nucleus at fixed energy levels.
  • 1st energy level: up to 2 electrons.
  • 2nd energy level: up to 8 electrons.
  • 3rd energy level: up to 8 electrons.
• Example: Aluminum (Atomic #13)
  • 13 protons in the nucleus and 13 electrons orbiting the nucleus at fixed energy levels.

Valence Electrons

• Valence electrons are electrons in the outermost shell of an atom.
• When two atoms interact, valence electrons are the first to come into contact.
• Valence electrons can participate in the formation of bonds.
• Carbon has 4 valence electrons.

Chemical Bond

• A chemical bond is a force of attraction.
  • Types:
    • Intramolecular: within a molecule
      • Ionic Bond: Loss or gain of electrons to form ions (cation +, anion -).
      • Covalent Bond: Sharing of electrons.
    • Intermolecular: between molecules
      • Hydrogen Bond.

Intramolecular Force - Ionic Bonds

• Occur between a metal (e.g., Na) and a nonmetal (e.g., Cl).
• Involve the loss and gain of electrons.
  • Na loses an electron to Cl.
    • Na becomes positively charged, forming a cation.
  • Cl gains an electron from Na.
    • Cl becomes negatively charged, forming an anion.
• Ionic bond is the strongest bond.
  • Involves the attraction of full positive and negative charges.

Intramolecular Force - Covalent Bonds

• Occur between nonmetals.
• Involve the sharing of an electron pair between two atoms.
  • Example: Oxygen gas ($O_2$)
    • Two shared pairs of electrons form a double covalent bond.

Covalent Bond Examples

• Methane ($CH_4$)
• Water ($H_2O$)

Covalent and Ionic Bonds

• Water molecule
• Table salt (NaCl)
• Covalent bonds: sharing of electrons between atoms
• Ionic bonds: transfer of electrons between atoms

Ionic vs. Covalent Bonds

• Ionic Bonds
  • Occur between a metal and a nonmetal.
  • Result from the loss or gain of electrons, forming cations and anions with full charges.
  • Strong attraction = Strong bond.
• Covalent Bonds
  • Occur between two nonmetals.
  • Result from the sharing of electrons.
  • Weaker attraction = Weaker bond.

Electronegativity

• Electronegativity is the tendency of an atom to attract electrons towards itself.
• Factors that can affect electronegativity:
  • Nuclear charge.
  • Atomic size.
• Periodic trend:

Polarity

• Polar
  • Unequal distribution of charge.
  • Molecules with partially positive ($\delta+$) and negative ($\delta-$) poles.
  • Result: Can attract and repel each other.
  • Example: Hydrogen Chloride (HCl).
  • Affected by:
    • Electronegativities of the atoms within the molecule.
    • Molecular geometry (shape of the molecule).
• Nonpolar
  • Equal distribution of charge.
  • Affected by the same factors as polarity.

Polarity Examples

• Polar molecules examples include HF, HCl, and H2O
• Non polar: H2, Cl2, CH4

Polarity of Methane and Carbon Dioxide

• Polarity of Methane ($CH_4$)
  • Nonpolar: The net dipole moment equals zero because of the symmetrical molecular structure.
• Polarity of Carbon Dioxide ($CO_2$)
  *Nonpolar: The net dipole moment equals zero because of the symmetrical molecular structure.

Intermolecular Force - Hydrogen Bonds

• Differences in electronegativity polarity hydrogen bonding
• Weaker than covalent bonds.
• Hydrogen Bond (represented by dashed lines).

Hydrogen Bonds

• Hydrogen bonds are the attraction between a hydrogen atom and a highly electronegative atom (such as oxygen or nitrogen) in another molecule or in a different part of the same molecule.