A hypothetical element has two main isotopes
with mass numbers of 80 and 83. If 68.00% of
the isotopes have a mass number of 80 amu,
what atomic weight should be listed on the
periodic table for this element?
It's a weighted average.
(80*0.68) + (83*32) = ?
To find the atomic weight of an element, we need to calculate the average mass of all its isotopes, taking into account their relative abundance.
In this case, we have two isotopes: one with a mass number of 80 amu (68.00% abundance) and another with a mass number of 83 amu (remaining abundance of 100 - 68.00 = 32.00%).
To calculate the atomic weight, we'll multiply the mass of each isotope by its relative abundance and then add them together.
Step 1: Calculate the contribution from the isotope with a mass number of 80 amu:
Contribution of isotope 80 = (mass of isotope 80) * (abundance of isotope 80)
= (80 amu) * (68.00%)
= 54.40 amu
Step 2: Calculate the contribution from the isotope with a mass number of 83 amu:
Contribution of isotope 83 = (mass of isotope 83) * (abundance of isotope 83)
= (83 amu) * (32.00%)
= 26.56 amu
Step 3: Calculate the atomic weight:
Atomic weight = Contribution of isotope 80 + Contribution of isotope 83
= 54.40 amu + 26.56 amu
= 80.96 amu
Therefore, the atomic weight that should be listed on the periodic table for this hypothetical element is 80.96 amu.