Average atomic mass — Science definition

The weighted average mass of all naturally occurring isotopes of an element, weighted by each isotope's abundance.

Where it is taught

In use

The listed value is not the mass of any one atom — it is the average atomic mass, a weighted average of the masses of all the element's naturally occurring isotopes, weighted by how common each one is.

Isotopes & Average Atomic Mass

Questions using average atomic mass

Problem 1

An element has two isotopes: one of mass 20.0 amu at 90.0% abundance and one of mass 22.0 amu at 10.0% abundance. Find the average atomic mass in amu.
Answer and explanation

20.2 amu

Average = 20.0 × 0.900 + 22.0 × 0.100 = 18.0 + 2.2 = 20.2 amu.

Problem 2

Copper has two isotopes: copper-63 (62.93 amu, 69.15%) and copper-65 (64.93 amu, 30.85%). Find copper's average atomic mass in amu.
Answer and explanation

63.55 amu

62.93 × 0.6915 = 43.52 amu; 64.93 × 0.3085 = 20.03 amu; sum = 63.55 amu.

Problem 3

Chlorine's average atomic mass, 35.45 amu, is much closer to 35 than to 37. What does that tell you about the abundances of its two isotopes?
Answer and explanation

Chlorine-35 must be far more abundant than chlorine-37 (roughly 75% versus 25%), because a weighted average is pulled toward the mass of the most common isotope.

Problem 4

An element has just two isotopes, of mass 10.0 amu and 11.0 amu, and its average atomic mass is 10.8 amu. What percent of its atoms are the mass-11 isotope?
Answer and explanation

80 %

Let the mass-10 fraction be x: 10x + 11(1 − x) = 10.8 → 11 − x = 10.8 → x = 0.20. So the mass-11 fraction is 0.80 = 80%.

Taught alongside

Elsewhere

Browse all courses