lecture 9 molec compounds formulas

Learning Objectives

• Understand the symbolism in chemical formulas

• Define the mole and Avogadro's number

• Relate mass, moles, and number of atoms/molecules

Chemical Formulas

• Molecular Formula: Represents the types and numbers of atoms in a molecule.

  • Components: 1) Chemical symbols indicate atom types. 2) Subscripts denote quantity of each atom.

  • Structural Formula: Shows how atoms are connected.

Visualization of Methane (Figure 2.16)

• Different representations for CH₄:

  • Molecular formula: CH₄

  • Structural formula: Shows connectivity

  • Ball-and-stick model: Represents atoms and bonds

  • Space-filling model: Represents the volume occupied by atoms

Elements as Molecules

• Discrete Atoms: Elements like noble gases exist as single atoms.

• Diatomic Molecules: Nonmetals found as molecules with two atoms, e.g., H₂, N₂, O₂.

• Allotropes: Different forms of the same element, like:

  • O₂ (oxygen) vs O₃ (ozone)

  • S₈ (sulfur) at room temperature vs S₂ at high temperature

  • Graphite, diamond, C₆₀ (allotropes of carbon)

Empirical and Molecular Formulas (Page 7)

• Empirical Formula: Simplest ratio of atoms in a compound.

  • Example: Benzene C₆H₆ (molecular) vs CH (empirical)

• Molecular Formula: Actual number of atoms in a molecule.

  • Example: Acetic Acid: C₂H₄O₂ (molecular) vs CH₂O (empirical)

Isomers (Page 9)

• Isomers: Same chemical formula but different structures and properties.

  • Example: Acetic acid vs methyl formate (both C₂H₄O₂, different structures)

Formal Charge (Page 100)

• Definition: Hypothetical charge on an atom in a covalently bonded molecule.

• Formula: FC = (valence e- of atom) - (lone pair e-) - ½(bonding pair e-)

• Key Principles:

  1. Low formal charges preferred in plausible structures.

  2. Negative charges on more electronegative atoms.

  3. Sum of all formal charges in neutral molecules = 0; for ions, equal to net charge.

Lewis Structures and Electron Geometry (Page 119)

• Lewis Structure Rules:

  1. Lowest electronegativity at the center

  2. Count valence electrons, account for charge if an ion

  3. Use pairs of electrons for single bonds

  4. Arrange for octets around terminal atoms; leftover electrons go to central atom

  5. Form multiple bonds if central atom lacks an octet

VSEPR Theory (Figure 4.16)

• Concept: Electron-pair geometry predicts molecular shapes based on electron repulsion.

• Example: Methane (CH₄) has a tetrahedral arrangement.

Polarity of Molecules (Figure 4.27)

• Dipole Moment: Derived from bond dipole moments and their arrangement.

  • Example: CO₂ has polar bonds but is nonpolar overall due to cancellation.

  • Water (H₂O) is polar because OH bond moments do not cancel out.