How to calculate the atomic mass of an element with two isotopes?

To calculate the atomic mass of an element with two isotopes, you determine the weighted average of their masses based on their natural abundance. This "atomic mass" seen on the periodic table is an average, reflecting the different isotopes of an element.

Understanding Atomic Mass and Isotopes

An element's atomic mass listed on the periodic table isn't just the mass of a single atom. Most elements exist as a mixture of isotopes, which are atoms of the same element that have the same number of protons but different numbers of neutrons. This difference in neutron count leads to slight variations in their atomic mass. The atomic mass displayed for an element is a weighted average of the masses of all its naturally occurring isotopes, taking into account their relative abundances.

Steps to Calculate Average Atomic Mass

The process for calculating the average atomic mass, applicable to elements with any number of isotopes including two, involves a straightforward series of steps. According to Study.com, the calculation involves three key steps:

Step 1: Identify Isotope Information

The first crucial step is to gather the necessary data for each isotope. This includes:

• The mass of each isotope: Typically measured in atomic mass units (amu).
• The natural abundance (percentage) of each isotope: This represents how commonly each isotope occurs in nature.

Step 2: Calculate Weighted Contribution for Each Isotope

For each isotope, you need to determine its individual contribution to the overall average atomic mass. To do this:

• Convert the percentage abundance to a decimal: Divide the percentage by 100 (e.g., 75% becomes 0.75).
• Multiply the isotope's mass by its decimal abundance: This gives you the weighted mass contribution for that specific isotope.

Step 3: Sum the Contributions

Once you have calculated the weighted contribution for each of the isotopes (in the case of two isotopes, you will have two such values):

• Add the results from Step 2 together: The sum will be the average atomic mass of the element. This final value is the atomic mass typically found on the periodic table.

Practical Example: Element with Two Isotopes

Let's consider a hypothetical element, "Element X," which has two isotopes: X-20 and X-22.

• Isotope 1: X-20
  • Mass = 19.9924 amu
  • Natural Abundance = 90.51%
• Isotope 2: X-22
  • Mass = 21.9914 amu
  • Natural Abundance = 9.49%

Now, let's apply the steps:

Finally, add the weighted contributions:

Average Atomic Mass of Element X = 18.0954 amu (from X-20) + 2.0870 amu (from X-22)
Average Atomic Mass of Element X = 20.1824 amu

This calculated value represents the average atomic mass of Element X, reflecting the natural proportions of its two isotopes.