Lecture Notes: Chapter 3 - Composition of Compounds and Solutions

Section 3.1: Formula Mass and the Mole Concept

Key Concepts:

• Connecting Microscopic to Macroscopic:
  • Understanding atoms and molecules by connecting atomic mass units (AMU) to grams.
  • Use of moles to bridge the gap between microscopic and macroscopic worlds.

Atomic Mass:

• Carbon-12 atom = 12 AMU
• Hydrogen atom ≈ 1.008 AMU
• Oxygen-16 atom = 16.00 AMU

The Mole:

• Defined as a unit to count particles (like a dozen for eggs).
• Avogadro's Number: 6.022 x 10^23 (counts atoms/molecules).
• Molar Mass: Mass of one mole in grams (same numerical value as atomic mass in AMU).

Conversion Factors:

1. Molar Mass (grams/mole):
   • Found on periodic table or given in problems.
2. Avogadro's Number:
   • Used when dealing with atoms, molecules, formula units.

Calculation Examples:

• Finding Mass of 0.89 moles of Calcium:

  • Use periodic table to find molar mass of calcium (40.08 g/mol).
  • Calculation: 0.89 moles x 40.08 g/mol = 36 grams.

• From Atoms to Grams:

  • Zinc example using Avogadro's number and molar mass.
  • 1.65 x 10^23 atoms of Zn to grams.

Molecule and Formula Mass:

• Molecular Mass: Sum of atomic masses in a molecule.
• Formula Mass: Sum of atomic masses in an ionic compound formula.

Practical Problems and Examples:

• Calcium Carbonate Example:
  • Convert grams to moles using molar mass.
• Propylene Example (C3H6):
  • Find molar mass, convert 25g to moles.
• Potassium Bromide (KBr):
  • Convert grams to moles using molar mass.

Practical Applications:

• Hemoglobin in Blood:
  • Calculating number of molecules and mass.
• Cullinan Diamond:
  • Calculations involving carats, grams, and carbon atoms.

Conclusion:

• Grams to atoms conversions involve large numbers.
• Avogadro's number often cumbersome, used mainly for understanding large collections of small entities like atoms.
• Molar mass is a critical tool for conversions in chemistry.