Average Atomic Mass Calculator
Our general chemistry calculator computes average atomic mass accurately. Enter measurements for results with formulas and error analysis.
Reviewed by Manoj Kumar, Mathematics Educator
Formula
Average Mass = sum(mass_i * fractional_abundance_i)
The average atomic mass is the sum of each isotope mass multiplied by its fractional natural abundance. This weighted average appears on the periodic table and represents the expected mass of a random atom from a natural sample.
Worked Examples
Example 1: Chlorine Average Atomic Mass
Problem:Cl-35: mass 34.969 amu, abundance 75.77%. Cl-37: mass 36.966 amu, abundance 24.23%
Solution:Average = 34.969(0.7577) + 36.966(0.2423)\n= 26.496 + 8.957\n= 35.453 amu
Result:Average atomic mass = 35.453 amu
Example 2: Silicon with Three Isotopes
Problem:Si-28: 27.977 amu (92.23%), Si-29: 28.976 amu (4.67%), Si-30: 29.974 amu (3.10%)
Solution:Average = 27.977(0.9223) + 28.976(0.0467) + 29.974(0.0310)\n= 25.803 + 1.353 + 0.929\n= 28.085 amu
Result:Average atomic mass = 28.085 amu
Frequently Asked Questions
What is average atomic mass?
Average atomic mass is the weighted average of the masses of all naturally occurring isotopes of an element, taking into account their relative abundances. It is the value reported on the periodic table and represents the expected mass of a randomly selected atom of that element from a natural sample. For example, chlorine has two stable isotopes: Cl-35 (75.77% abundance, 34.969 amu) and Cl-37 (24.23% abundance, 36.966 amu), giving an average atomic mass of 35.453 amu.
How do you calculate average atomic mass from isotope data?
Average atomic mass is calculated by multiplying each isotope mass by its fractional abundance (decimal form) and summing the products. The formula is: Average Mass = (mass_1 * fraction_1) + (mass_2 * fraction_2) + ... If abundances are given as percentages, divide each by 100 first. The fractional abundances must sum to 1.0 (or percentages to 100%). For example, boron has B-10 (19.9%, 10.013 amu) and B-11 (80.1%, 11.009 amu): Average = 10.013(0.199) + 11.009(0.801) = 10.811 amu.
Why is average atomic mass not a whole number?
Average atomic mass is almost never a whole number because it is a weighted average of multiple isotope masses, each of which is also not exactly a whole number due to nuclear binding energy effects. The mass of each isotope differs from its mass number because of mass defect, and the averaging across isotopes with different abundances produces a non-integer result. The only exception is fluorine, which has only one stable isotope (F-19), but even its precise atomic mass is 18.998 amu, not exactly 19. Carbon-12 is defined as exactly 12 amu by convention.
Can you determine isotope abundance from average atomic mass?
Yes, if you know the average atomic mass and the individual isotope masses, you can calculate the relative abundances for an element with two stable isotopes. Set the abundance of one isotope as x and the other as (1-x), then solve: Average Mass = mass_1(x) + mass_2(1-x). Rearranging gives x = (Average Mass - mass_2) / (mass_1 - mass_2). For elements with three or more isotopes, you need additional information such as the abundance of at least one isotope to solve the system of equations.
References
Reviewed by Manoj Kumar, Mathematics Educator · Editorial policy