Average Atomic Mass
Calculates the weighted average of the atomic masses of an element's naturally occurring isotopes, considering their relative abundances.
What is Average Atomic Mass?
Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers and atomic masses.
Key formula / rule: Average Atomic Mass
Key points
- Define average atomic mass.
- Explain the concept of isotopes and their role in determining average atomic mass.
- Calculate the average atomic mass of an element given the masses and relative abundances of its isotopes.
- Relate average atomic mass to the values found on the periodic table.
Common exam trap
Calculating a simple arithmetic mean of isotope masses instead of a weighted average.
Definitions
- Term
Isotopes
- Meaning
Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers and atomic masses.
- Term
Atomic Mass Unit (u)
- Meaning
A unit of mass used to express atomic and molecular masses, approximately equal to the mass of one proton or one neutron. It is defined as 1/12th the mass of an atom of carbon-12.
- Term
Relative Abundance
- Meaning
The proportion of a specific isotope in a natural sample of an element, usually expressed as a percentage or a fraction.
Learning objectives
Define average atomic mass.
Explain the concept of isotopes and their role in determining average atomic mass.
Calculate the average atomic mass of an element given the masses and relative abundances of its isotopes.
Relate average atomic mass to the values found on the periodic table.
Formulae
- Name
Average Atomic Mass
- Note
Where 'i' represents each naturally occurring isotope of the element.
- Expression
Average Atomic Mass = Σ (Mass of Isotopeᵢ × Fractional Abundance of Isotopeᵢ)
- Name
Fractional Abundance
- Note
Used to convert percentage abundance to a decimal for calculations.
- Expression
Fractional Abundance = Percentage Abundance / 100
Prerequisites
Concept of an element and its atomic number.
Understanding of isotopes and their mass numbers.
Basic arithmetic operations, including multiplication and addition.
Concept of percentage and fractional abundance.
Common mistakes
Calculating a simple arithmetic mean of isotope masses instead of a weighted average.
Using percentage abundance directly instead of converting it to fractional abundance.
Confusing average atomic mass with the mass number of a specific isotope.
Keywords
Average Atomic Mass
Isotopes
Relative Abundance
Fractional Abundance
Atomic Mass Unit
Periodic Table
Weighted Average
Stoichiometry
Practice preview
The average atomic mass of a sample of magnesium is 24.305 amu. If magnesium has three isotopes, Mg-24, Mg-25, and Mg-26, with natural abundances of 79.0%, 10.0%, and 11.0% respectively, which of the following statements…
hard
Chlorine has two naturally occurring isotopes, Cl-35 and Cl-37, with relative abundances of 75.77% and 24.23% respectively. Calculate the average atomic mass of chlorine.…
medium
The average atomic mass of an element is 16.2 amu. It has two isotopes with mass numbers 16 and 18. If the relative abundance of the lighter isotope is 75%, what is the relative abundance of the heavier isotope?…
medium