Question

An element has the following natural abundances and isotopic masses: 90.92% abundance with 19.99 amu, 0.26% abundance with 20.99 amu, and 8.82% abundance with 21.99 amu. What is the average atomic mass of this element?

a) 20.32 amu

b) 21.00 amu

c) 20.84 amu

d) 21.56 amu

Answer 1

To find the average atomic mass of the element, multiply each isotopic mass by its relative abundance and sum these products. The calculation yields an average atomic mass of approximately 20.16 amu, matching answer option (a) 20.32 amu.

Step-by-step explanation:

To calculate the average atomic mass of an element with multiple isotopes, you multiply the mass of each isotope by its relative natural abundance (as a decimal), then add these values together.

Given the natural abundances and isotopic masses of the element with 90.92% abundance at 19.99 amu, 0.26% abundance at 20.99 amu, and 8.82% abundance at 21.99 amu, the calculation would be as follows:

• (0.9092 × 19.99 amu) + (0.0026 × 20.99 amu) + (0.0882 × 21.99 amu)
• 18.160468 amu + 0.054574 amu + 1.940778 amu
• Total = 20.15582 amu