Lecture: Difference Between Mass Number and Atomic Mass

Introduction

• Mass Number and Atomic Mass are two distinct concepts in chemistry.
• Although they sound similar, they have different meanings and calculations.

Mass Number

• Represents the sum of protons and neutrons in an atom's nucleus.
• Example:
  • Atom with 5 protons and 5 neutrons -> Mass Number = 10.
  • Atom with 5 protons and 6 neutrons -> Mass Number = 11.
• Mass number is a whole number and gives an estimate of atomic mass in atomic mass units (amu).
• Protons and Neutrons each approximately weigh 1 amu.
• Electrons are negligible in mass for these calculations.

Atomic Mass

• Also known as average atomic mass, relative atomic mass, or atomic weight.
• Reflects the weighted average mass of all isotopes of an element as found in nature.
• Consider isotopes of Boron:
  • Boron-10 (5 protons, 5 neutrons) weighs 10 amu.
  • Boron-11 (5 protons, 6 neutrons) weighs 11 amu.
• Isotopes have the same number of protons (same element) but different numbers of neutrons.
• Example calculation of Atomic Mass:
  • 20% of Boron atoms are Boron-10, and 80% are Boron-11.
  • Atomic Mass = (0.20 * 10 amu) + (0.80 * 11 amu) = 10.8 amu.
• This weighted average corresponds to the atomic mass number on the periodic table.

Key Differences

• Mass Number:
  • Specific to individual atom/isotope.
  • Sum of protons and neutrons.
  • Gives a rough idea of atomic mass in amu.
• Atomic Mass:
  • Average mass considering all isotopes.
  • Considers isotopic abundances.
  • Calculated using weighted average method.

Conclusion

• Understanding these differences is crucial for interpreting data about elements.
• For further learning, additional resources or videos on calculating atomic mass are available.