Lecture Notes: Atomic Mass and Atomic Weight

Key Isotopes of Carbon

• Carbon-12
  • Most common isotope
  • Comprises 98.89% of Earth's carbon
  • Mass: Exactly 12 atomic mass units (by definition)
• Carbon-13
  • Second most common isotope
  • Comprises 1.11% of Earth's carbon
  • Mass: 13.0034 atomic mass units

Understanding Atomic Mass

• Atomic mass refers to the mass of an individual isotope.
• Calculated experimentally for each isotope.

Calculating Atomic Weight

• Atomic Weight: Weighted average of the atomic masses of an element's isotopes.
• Formula:
  • Weighted by the abundance of each isotope.
  • Example calculation for carbon:
    • ( \text{Atomic Weight} = (0.9889 \times 12) + (0.0111 \times 13.0034) )
    • Result: Approximately 12.01 atomic mass units
    • Calculation steps:
      • ( 0.9889 \times 12 = 11.8668 )
      • ( 0.0111 \times 13.0034 = 0.1445374 )
      • Adding results gives 12.01113774
      • Rounded to 12.01 for periodic table usage

Differences Between Carbon-12 and Carbon-13

• Similarities:
  • Both have 6 protons, which define the element as carbon.
• Differences:
  • Carbon-12 has 6 neutrons
  • Carbon-13 has 7 neutrons (1 additional neutron)
  • Difference in atomic mass: +1.0034 atomic mass units

Understanding Neutron Mass

• Adding a neutron generally increases atomic mass by roughly 1 atomic mass unit.
• Protons have a similar effect on atomic mass.

Conclusion

• Atomic Mass: The mass of a specific isotope.
• Atomic Weight: Averages atomic masses based on isotopic abundance on Earth.
• Understanding the calculation of atomic weight helps in appreciating chemical properties and behaviors of elements.