Section 1.3 - Molecules

1.3 Introduction to Atoms and Molecules

  • Atoms: Basic units of matter; consist of protons, neutrons, and electrons.

  • Molecules: Formed when two or more atoms bond together.

  • Oxygen Molecule (O2): Diatomic molecule; consists of two oxygen atoms bonded together.

2. Key Concepts in Chemistry

2.1 Ionization of Atoms and Molecules

  • Definition: The process by which an atom or a molecule loses an electron.

  • Ionization Energy: The amount of energy required to remove an electron from an atom.

2.2 Formal Charge and Resonance

  • Formal Charge: Representation of electron distribution in a molecule; helps determine stability.

  • Resonance Structures: Multiple valid Lewis structures for the same molecule, indicating delocalized electrons.

3. Ionic Bonds vs. Covalent Bonds

3.1 Characteristics of Ionic Compounds

  • Formed through the transfer of electrons resulting in ions.

  • Tend to form crystalline structures.

  • Properties: High melting and boiling points; conduct electricity when dissolved in water.

3.2 Comparison with Covalent Bonds

  • Covalent bonds involve sharing electrons between atoms.

  • Ionic bonds result from electrostatic attraction between oppositely charged ions.

4. Ion Sizes and Periodicity

4.1 Ionic Radius Trends

  • Size of ions increases down a group on the periodic table.

  • In a series of isoelectronic ions, ionic size decreases with increasing atomic number.

5. Common Ionic Compounds

5.1 Naming Conventions

  • Monatomic Ions: Simple ions formed from a single atom.

    • Examples: Na+ (sodium), Cl- (chloride).

  • Polyatomic Ions: Ions composed of two or more atoms.

    • Examples: OH- (hydroxide), NH4+ (ammonium).

5.2 Examples of Ionic Compounds

  • NaCl: Sodium chloride

  • KI: Potassium iodide

  • MgO: Magnesium oxide

6. Lewis Structures

6.1 Drawing Lewis Structures

  • Step 1: Count total number of valence electrons.

  • Step 2: Connect atoms with single bonds.

  • Step 3: Distribute remaining electrons to complete octets.

6.2 Formal Charge Calculation

  • To assign formal charges:

    • Formal Charge = (Valence electrons) - (Non-bonding electrons) - 0.5(Bonding electrons)

7. Molecular Geometry and VSEPR Theory

7.1 Predicting Molecular Shape

  • VSEPR Theory: Valence Shell Electron Pair Repulsion Theory; predicts geometry based on electron pair repulsion.

  • Steric Number: Total number of atoms and lone pairs attached to a central atom.

  • Common Geometries:

    • Linear (AX2): 180°

    • Trigonal Planar (AX3): 120°

    • Tetrahedral (AX4): 109.5°

8. Electronegativity and Molecular Polarity

8.1 Electronegativity Differences

  • Defines the ability of an atom to attract electrons in a bond.

  • Polar Bonds: Bonds between atoms of differing electronegativities lead to unequal electron distribution.

8.2 Dipole Moments

  • Dipole Moment: A measure of polarity in a molecule caused by differences in electronegativity.

  • Molecules can have polar bonds but may have no overall dipole moment due to symmetry.