Average Atomic Mass and Molar Concepts

Key Concepts

Average Atomic Mass

• Useful for thinking about atomic/molecular mass.
• Seldom deal with individual atoms in a chemistry lab—typically deal with grams of a substance.

Connecting Atomic Mass to Lab-Scale Masses

• Chemistry uses a tool to scale up from atomic to lab-scale masses.
• Example: Lithium has an average atomic mass of 6.94 unified atomic mass units per atom.
• A sample containing a specific number of lithium atoms weighs 6.94 grams.
• This specific number of atoms is 6.02214076 × 10^23, known as Avogadro's Number.

Avogadro's Number

• Named after Amadeo Avogadro, a 19th-century Italian chemist.
• Commonly approximated as 6.022 × 10^23.

Mole Concept

• Introduced by German chemist Wilhelm Ostwald in the late 19th century.
• A mole represents 6.02214076 × 10^23 units of a substance (analogous to a dozen, which represents 12 units).

Practical Example: Germanium

• Given: 15.4 milligrams of Germanium.
• Steps to find the number of atoms:
  • Convert milligrams to grams: 15.4 mg * (1 g / 1000 mg) = 0.0154 g.
  • Convert grams to moles using molar mass (72.63 g/mol for Germanium):
    • 0.0154 g * (1 mol / 72.63 g) ≈ 0.000212 mol.
  • Convert moles to atoms using Avogadro's Number:
    • 0.000212 mol * 6.022 × 10^23 atoms/mol ≈ 1.28 × 10^20 atoms.*

Important Notes

• Dimensional analysis helps ensure correct unit conversions.
• Always pay attention to significant digits in calculations.