Isotope Calculator
Find the number of protons, neutrons, and electrons in any isotope from its atomic number and mass number, or compute a weighted average atomic mass from two isotopes and their natural abundances. Two 3D diagrams compare isotopes with different neutron counts and show how abundance-weighting pulls the average mass toward the more common isotope, with charts illustrating both ideas.
Reviewed by the ToolNestr Editorial Team — July 2026
Same element, different neutrons
1. Two isotopes of the same element
Same number of protons (red) in both nuclei, but the isotope on the right carries extra neutrons (grey) — same element, more mass.
2. Abundance-weighted average
The larger sphere represents the more abundant isotope — the average mass sits closer to it than to the rarer, smaller isotope.
Isotope graphs
How it works
The core idea in one line: an element's identity is fixed entirely by its proton count, so isotopes of the same element can carry different numbers of neutrons — and therefore different masses — without ever becoming a different element.
neutrons = A − Z
mass number minus atomic number
electrons = Z − charge
for a neutral atom, charge = 0, so electrons = Z
average mass = Σ(mass_i × fraction_i)
weighted average across all naturally occurring isotopes
Since mass number A is simply the total count of protons and neutrons in the nucleus, subtracting the atomic number (protons) leaves neutrons = A − Z directly. For a neutral atom, the electron count always matches the proton count to balance charge, so electrons = Z (or Z − charge for an ion). When an element occurs in nature as a mixture of isotopes, each contributes to the periodic table's listed atomic mass in direct proportion to how abundant it is — a more common isotope pulls the average closer to its own mass.
Worked example 1 — Carbon-14 composition
Given: Carbon-14 has atomic number Z = 6 and mass number A = 14, with no net charge.
Carbon-14 has 2 more neutrons than the common Carbon-12 isotope — this extra mass and its radioactive instability is exactly what makes carbon-14 useful for radiocarbon dating.
Worked example 2 — chlorine's average atomic mass
Given: Natural chlorine is a mix of Cl-35 (mass 34.969 amu, 75.77% abundance) and Cl-37 (mass 36.966 amu, 24.23% abundance).
This weighted average of 35.45 amu matches chlorine's listed atomic mass on the periodic table — closer to Cl-35's mass because it's the more abundant isotope.
Hydrogen isotopes — same protons, different neutrons
All three are chemically hydrogen (Z=1), but their differing neutron counts give them very different masses and, in tritium's case, radioactivity.
| Isotope | Protons | Neutrons | Notes |
|---|---|---|---|
| Protium (¹H) ★ | 1 | 0 | Most common (~99.98%), stable |
| Deuterium (²H) | 1 | 1 | Stable, used in heavy water |
| Tritium (³H) | 1 | 2 | Radioactive, half-life ~12.3 years |
★ All three have exactly one proton (Z=1), which is what makes them all hydrogen — only the neutron count changes between isotopes.
Where isotopes actually matter
🕰️ Radiocarbon dating
Carbon-14 decays at a known, steady rate, letting archaeologists estimate the age of once-living organic material by measuring how much Carbon-14 remains relative to stable Carbon-12.
🏥 Nuclear medicine tracers
Radioactive isotopes like Technetium-99m are injected into patients in tiny amounts and tracked with imaging equipment to diagnose conditions in bones, organs, and blood flow without invasive surgery.
🔬 Mass spectrometry
A mass spectrometer separates isotopes by their mass-to-charge ratio, letting chemists directly measure the abundance of each isotope in a sample — the exact data used to calculate an element's average atomic mass.
⚛️ Nuclear reactors
Nuclear power plants rely on specific isotopes like Uranium-235 (fissile) rather than the far more abundant Uranium-238, since only certain isotopes of an element can sustain a nuclear chain reaction.
Common misconceptions
"Isotopes of the same element have different atomic numbers."
Isotopes always share the exact same atomic number (Z) — that's precisely what makes them the same element. Only their neutron count, and therefore mass number, differs between isotopes.
"All isotopes are radioactive."
Most elements have at least one stable isotope that never decays. Radioactive (unstable) isotopes exist alongside stable ones — for example, Carbon-12 and Carbon-13 are stable, while only Carbon-14 is radioactive.
"The atomic mass on the periodic table is the mass of the most common isotope."
It's a weighted average across every naturally occurring isotope, weighted by how abundant each one is — not simply the mass of whichever isotope happens to be most common.
"Adding or removing neutrons changes what element an atom is."
Element identity is defined entirely by the number of protons (Z), never by neutrons. Adding or removing neutrons only creates a different isotope of the same element, not a new element.
Formula sources & further reading
The formulas here are standard, traceable to:
- • OpenStax, Chemistry 2e — Chapter 2, "Atoms, Molecules, and Ions" (free, peer-reviewed). openstax.org
- • Brown, LeMay & Bursten, Chemistry: The Central Science — Chapter 2, Atoms, Molecules, and Ions.
- • Zumdahl & Zumdahl, Chemistry — Atomic structure and isotopes.
neutrons = A − Z; electrons = Z − charge; average mass = Σ(mass_i × fraction_i). Isotope masses and natural abundances are standard reference (IUPAC/NIST) values. Results are rounded for display.
How to use this calculator
Pick the mode
"Isotope composition" for protons/neutrons/electrons; "Average atomic mass" for the weighted-average calculation.
Enter the values
Type atomic number, mass number, and charge — or two isotope masses with their percent abundances.
Read the result
Particle counts or the weighted average mass solve instantly.
Related tools
Frequently asked questions
What is an isotope?
Isotopes are atoms of the same element (same number of protons, Z) that have different numbers of neutrons, and therefore different mass numbers (A). Carbon-12, Carbon-13, and Carbon-14 are all isotopes of carbon.
How do you find the number of neutrons in an isotope?
Subtract the atomic number from the mass number: neutrons = A − Z. The mass number A is the total count of protons plus neutrons in the nucleus.
What is the difference between mass number and atomic mass?
Mass number (A) is a whole number — the specific count of protons plus neutrons in one isotope's nucleus. Atomic mass (shown on the periodic table) is a weighted average across all naturally occurring isotopes of an element, which is why it usually isn't a whole number.
Are all isotopes radioactive?
No — most elements have one or more stable isotopes that never decay, alongside any unstable (radioactive) isotopes. Carbon-12 is stable and makes up about 98.9% of natural carbon, while Carbon-14 is radioactive and used for dating organic material.
How is average atomic mass calculated from isotopes?
Multiply each isotope's exact mass by its fractional natural abundance, then add the results together: average mass = Σ(isotope mass × fractional abundance). This weighted average is what appears on the periodic table.