A certain element consists of two stable isotopes. The first has a mass of 62.9 amu and a percent natural abundance of 69.1 %. The second has a mass of 64.9 amu and a percent natural abundance of 30.9 %. What is the atomic weight of the element?

63.518

Explanation:

The following data were obtained from the question:

Mass of Isotope A = 62.9 amu

Abundance of isotope A (A%) = 69.1%

Mass of isotope B = 64.9 amu

Abundance of isotope B (B%) = 30.9%

Atomic weight of the element = ... ?

The atomic weight of the element can be obtained as follow:

Atomic weight = [ (Mass of A x A%) / 100] + [ (Mass of B x B%) / 100]

Atomic weight = [ (62.9 x 69.1) / 100] + [ (64.9 x 30.9) / 100]

Atomic weight = 43.4639 + 20.0541

Atomic weight = 63.518

Therefore, the atomic weight of the element is 63.518.