By Chris Deziel, Updated Mar 24, 2022
Every atomic nucleus, except hydrogen, is composed of protons and neutrons. While the nucleus is far too small to be seen, scientists can determine the number of neutrons in any isotope using mass spectrometry and the data provided by the periodic table.
The total mass of an atom is essentially the sum of its protons, neutrons, and, to a negligible degree, electrons. Because electrons weigh only about 1/1836 of a proton, they can be ignored when calculating neutron numbers. Consequently, the atomic mass (in atomic mass units, amu) reflects the combined mass of protons and neutrons alone.
The periodic table lists elements in order of increasing proton count, which is also the atomic number (Z). Directly beneath the element symbol is the atomic number and, next to it, the average atomic mass. This average accounts for all naturally occurring isotopes and often includes a fractional component.
To find the typical neutron count:
Atomic mass ≈ protons + neutrons. Subtract the atomic number (protons) from the rounded atomic mass to get the neutron count.
Uranium is element 92. Its listed atomic mass is 238.039 amu. Rounding gives 238; subtract 92 protons, and you obtain 146 neutrons. The high neutron-to-proton ratio is a key factor in uranium’s radioactivity.
Isotopes are variants of an element that differ only in neutron number. While most elements have multiple isotopes (tin has ten, xenon nine), each isotope’s atomic mass is an integer, representing the exact count of protons and neutrons.
For example, carbon‑14 (C‑14) has a mass of 14 amu. With an atomic number of 6, it contains 8 neutrons—two more than the stable C‑12, which contributes to its radioactive decay.
By following the simple subtraction method above, you can determine the neutron number for any element or isotope you encounter.
