Lecture on Calculating Average Atomic Mass

Introduction

• Discusses the calculation of average atomic mass using percent abundance data.
• Reminder of isotopes and atomic mass variations among elements.

Key Concepts

Isotopes

• Example: Carbon
  • Carbon with mass 12, 13, and 14.
  • All have 6 protons, differing in neutron count.
  • Definition of isotopes.

Atomic Number and Mass

• Atomic number: whole number (e.g., Carbon = 6).
• Atomic mass on the periodic table: decimal number (e.g., Carbon = 12.011).
• Decimal represents the average mass of all isotopes.

Percent Abundance

• Definition: Shows how common each isotope is.
• Example for Carbon:
  • Carbon-12: 98.9% abundant.
  • Carbon-13: 1.1%.
  • Carbon-14: <0.001%.
  • Carbon-12 is the most abundant isotope.

Calculating Average Atomic Mass

Weighted Average

• Necessary due to different percent abundances.
• Example: Copper
  • Isotopes: Copper-63 and Copper-65.
  • Percent abundances: 69.09% for Copper-63, 30.91% for Copper-65.

Steps for Calculation

1. List isotopic masses in a vertical line.
2. Multiply each mass by its percent abundance (convert percent to decimal by moving two places to the left).
3. Example Calculation for Copper:
   • Copper-63: 63 * 0.6909 = 43.527.
   • Copper-65: 65 * 0.3091 = 20.092.
4. Add results to get weighted average: 63.62 (average atomic mass).

Practice Example

• Isotopes: Magnesium-24, Magnesium-25, Magnesium-26.
• Calculate similarly by listing masses, converting percentages, multiplying, and adding results.
• Result: Average atomic mass = 24.46.

Conclusion

• Successful calculation of average atomic masses using weighted averages and percent abundances.
• Encouragement to review the video recap for reinforcement.