Atomic Mass Calculator
Calculate the weighted average atomic mass from isotopic abundances. Load accurate data for common elements or enter your own.
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Enter isotopic data or load a common element to see the weighted average here.
Understanding Average Atomic Mass
If you look at the periodic table, you'll notice that the atomic mass of most elements is rarely a whole number. For example, Carbon is 12.011, not just 12. Why is that?
This is because most elements naturally exist as a mixture of different isotopes—atoms of the same element that have the same number of protons but a different number of neutrons. The value you see on the periodic table is a weighted average of all these naturally occurring isotopes, taking into account how abundant each one is.
The Formula
Sum of (Isotope Mass × Fractional Abundance) for all isotopes.
Real-World Applications
Nuclear Medicine
Specific isotopes like Technetium-99m are used in medical imaging to diagnose heart disease and cancer. Calculating precise doses requires knowing the exact isotopic properties.
Carbon Dating
Archaeologists measure the ratio of unstable Carbon-14 to stable Carbon-12 in fossils to determine their age. This relies entirely on our understanding of isotopic decay and abundance.
Energy Production
Nuclear power plants rely on Uranium-235, which is only 0.7% of natural Uranium. Enrichment processes increase this percentage to make fuel efficient enough for reactors.
How to Calculate Manually
Let's say we have an element "X" with two isotopes:
- Isotope A: Mass = 10 amu, Abundance = 20%
- Isotope B: Mass = 11 amu, Abundance = 80%
(10 × 0.20) = 2.0
(11 × 0.80) = 8.8
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Frequently Asked Questions
What is the difference between Mass Number and Atomic Mass?
Mass Number (A) is always a whole number (protons + neutrons) for a specific single atom. Atomic Mass is the weighted average mass of all naturally occurring isotopes and is usually a decimal number.
Why do we calculate the weighted average?
Because a simple average would assume all isotopes are equally common, which is false. For example, Chlorine-35 is 75% abundant, so it influences the atomic mass much more than Chlorine-37.
What unit is used for Atomic Mass?
The standard unit is the Unified Atomic Mass Unit (u) or dalton (Da). One 'u' is defined as exactly 1/12th the mass of a Carbon-12 atom.
Can atomic mass ever be a whole number?
For a specific isotope, it is very close to a whole number but not exact (except Carbon-12, which is exactly 12). For an element (average), it is almost never a whole number due to the mix of isotopes.
How do I find the percent abundance?
Abundance is determined experimentally using a device called a Mass Spectrometer. In textbook problems, these percentages are usually given to you.
What is the most stable isotope?
Generally, the isotope with an atomic mass closest to the average atomic mass on the periodic table is the most abundant and stable one.
Does temperature affect atomic mass?
No. The mass of the nucleus is constant regardless of temperature, pressure, or chemical state.
Why is the mass of a proton not exactly 1 amu?
A proton is 1.007276 amu and a neutron is 1.008665 amu. Also, some mass is converted into 'binding energy' to hold the nucleus together (E=mc²), which changes the final mass slightly.
How many decimal places should I use?
Isotopic masses are very precise. When calculating, use all provided decimal places and only round your final answer to the appropriate number of significant figures.
Is this calculator accurate for all elements?
Yes, as long as you input the correct isotopic masses and abundances. We have pre-loaded data for common elements like Hydrogen, Carbon, Nitrogen, and Oxygen.
What is a 'natural abundance'?
It refers to the percentages of isotopes found naturally on Earth. These ratios can actually vary slightly depending on the source (e.g., water from different oceans), but standard values are set by IUPAC.
Can I use this for artificial isotopes?
Yes. You can input any mass and percentage you effectively want to average, even for theoretical or synthetic mixtures.