Average Atomic Mass Lecture Notes

Introduction

• Atomic Mass on Periodic Table: Represents a weighted average of all the isotopes of an element.
• Carbon Example:
  • Atomic mass = 12.011
  • Major isotopes: Carbon-12 and Carbon-13
  • Carbon-12: 98.93% abundance, mass = 12
  • Carbon-13: 1% abundance, mass = 13.00335
  • Carbon-14: Trace amounts; not a significant contributor to average atomic mass.

Weighted Averages

• Average Closeness: The average atomic mass is closest to the most abundant isotope.
• Symbols Used: AMU (atomic mass unit), U (unified atomic mass unit), or grams per mole.

Calculating Average Atomic Mass

Example 1: Magnesium

• Isotopes and Percentages:

  • Magnesium-24: 78.7%
  • Magnesium-25: 10.0%
  • Magnesium-26: 11.3%

• Calculation Steps:

  1. Convert percentage to decimal:
     • 78.7% -> 0.787
     • 11.3% -> 0.113
     • 10.0% -> 0.100
  2. Multiply by respective masses:
     • 0.787 × 24 = 18.888
     • 0.113 × 26 = 2.938
     • 0.100 × 25 = 2.500
  3. Add contributions: 18.888 + 2.938 + 2.500 = 24.326
  4. Average atomic mass for magnesium = 24.326 U

Example 2: Element X

• Isotopes and Percentages:

  • X-114: 94.2%
  • X-113: 4.8%
  • X-112: 1%

• Calculation Steps:

  1. Convert percentage to decimal:
     • 94.2% -> 0.942
     • 4.8% -> 0.048
     • 1% -> 0.010
  2. Multiply by respective masses:
     • 0.942 × 114 = 107.388
     • 0.048 × 113 = 5.424
     • 0.010 × 112 = 1.120
  3. Add contributions: 107.388 + 5.424 + 1.120 = 113.932
  4. Average atomic mass for element X = 113.932 U

Shortcuts and Tips

• Quick Method:

  • Change percentages to decimals and calculate directly using a calculator.
  • Example for element X: 0.942 × 114 + 0.048 × 113 + 0.010 × 112 = 113.932

• Reminder: Show your work when practicing.

Questions

• Write down any questions to discuss with your teacher in class.

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Note: These notes summarize the process of calculating average atomic mass using isotopic weights and provide examples to illustrate the method.