Atomic Mass Calculator
Why Isn't the Periodic Table Simple?
If you look at Carbon on the Periodic Table, its mass is 12.011. Why not just 12?
This is because elements in nature don't exist in just one form. They exist as mixtures of Isotopes.
- Isotope: Atoms of the same element (same number of protons) but with a different number of neutrons (different mass).
For Carbon, most atoms are Carbon-12 (Mass 12). But about 1.1% of all Carbon is Carbon-13 (Mass 13.003). When you average them out based on how common they are, you get 12.011.
The "Weighted Average" Formula
You cannot simply add the masses and divide by 2. That would be a regular average.
Because some isotopes are much more common than others, we must use a Weighted Average. The formula is:
Where "Relative Abundance" is the percentage written as a decimal (e.g., 50% = 0.50).
Example Problem: Chlorine
Chlorine is a classic chemistry problem because it has two very common isotopes.
- Chlorine-35: Mass = 34.969 amu. Abundance = 75.78%
- Chlorine-37: Mass = 36.966 amu. Abundance = 24.22%
Step 1: Convert percents to decimals.
75.78% -> 0.7578
24.22% -> 0.2422
Step 2: Multiply Mass × Abundance.
(34.969 × 0.7578) = 26.50
(36.966 × 0.2422) = 8.95
Step 3: Add them up.
26.50 + 8.95 = 35.45 amu.
This matches the number you see on the Periodic Table!
• Mass Number: A whole number (Protons + Neutrons) for a specific single atom (e.g., Carbon-12).
• Atomic Mass: A decimal number. The average weight of a random sample of that element found in nature.
What is "amu"?
The unit for this calculation is the Atomic Mass Unit (amu), sometimes denoted as u or Daltons (Da).
1 amu is defined as exactly 1/12th the mass of a Carbon-12 atom. It is the standard scale for weighing protons and neutrons.