Atomic Mass vs. Atomic Weight

Atomic Mass vs. Atomic Weight

  • Chemists use several terms to describe the "heaviness" of an element:
    • Atomic Mass (synonymous with Mass Number)
    • Atomic Weight
  • Atomic weight is constant for a given element (reported on the periodic table).
  • Atomic mass/mass number varies from one isotope to another.

Atomic Mass

  • The mass of one proton is approximately 1 AMU (atomic mass unit).
  • Definition of AMU: Exactly 1/121/12 the mass of a Carbon-12 atom (approximately 1.66×10−241.66 \times 10^{-24} grams).
  • Carbon-12 has 6 protons and 6 neutrons, so 1 AMU is approximately the mass of a proton or neutron.
  • The mass difference between protons and neutrons is very small; approximately the mass of an electron.
  • The atomic mass of an atom (in AMU) is nearly equal to its mass number (sum of protons and neutrons).
  • Some mass is lost as binding energy.

Isotopes

  • Isotopes: Atoms of the same element with varying mass numbers.
  • Isotopes differ in their number of neutrons.
  • Referred to by the element name followed by the mass number (e.g., Carbon-12, Carbon-13).
  • Hydrogen isotopes have unique names:
    • Protium: 1 proton, atomic mass of 1 AMU.
    • Deuterium: 1 proton, 1 neutron, atomic mass of 2 AMU.
    • Tritium: 1 proton, 2 neutrons, atomic mass of 3 AMU.
  • Isotopes generally exhibit similar chemical properties due to the same number of protons and electrons.

Atomic Weight

  • Most elements exist as two or more isotopes in nature.
  • These isotopes are usually present in the same proportions in any naturally occurring sample.
  • Atomic weight: The weighted average of these isotopes; this is the number reported on the periodic table.
  • Example: Chlorine has two main isotopes: Chlorine-35 and Chlorine-37.
  • Chlorine-35 is about three times more abundant than Chlorine-37.
  • Therefore, the atomic weight of chlorine is closer to 35 than 37 and is approximately 35.5.
  • Half-lives of isotopes relate to their stability and help determine their relative proportions.

Utility of Atomic Weight

  • Represents both:
    • The mass of the average atom of that element in AMU.
    • The mass of one mole of that element in grams.
  • Mole: A number of things (atoms, ions, molecules) equal to Avogadro's number NA=6.02×1023N_A = 6.02 \times 10^{23}.
  • Example: The atomic weight of carbon is 12 AMU.
  • The average carbon atom has a mass of 12 AMU
  • 6.02×10236.02 \times 10^{23} carbon atoms have a combined mass of 12 grams.

Example Problem

  • Element Q consists of three isotopes: A, B, and C.
    • Isotope A: Atomic mass of 40 AMU, accounts for 60% of naturally occurring Q.
    • Isotope B: Atomic mass of 44 AMU, accounts for 25% of Q.
    • Isotope C: Atomic mass of 41 AMU, accounts for 15% of Q.
  • Question: What is the atomic weight of element Q?
Solution:
  • The atomic weight is the weighted average of the isotopes.
  • Calculation: (0.6×40 AMU)+(0.25×44 AMU)+(0.15×41 AMU)(0.6 \times 40 \text{ AMU}) + (0.25 \times 44 \text{ AMU}) + (0.15 \times 41 \text{ AMU})
  • =24 AMU+11 AMU+6.15 AMU=41.15 AMU= 24 \text{ AMU} + 11 \text{ AMU} + 6.15 \text{ AMU} = 41.15 \text{ AMU}