Mass number is the total count of protons and neutrons in an atom’s nucleus. It’s always a whole number, written with the symbol A, and calculated with a simple formula: mass number = number of protons + number of neutrons. If you know any two of those three values, you can find the third.
How Mass Number Is Calculated
Every atom has a nucleus packed with two types of particles: protons (which carry a positive charge) and neutrons (which carry no charge). The mass number simply adds these two together. A helium atom, for example, has 2 protons and 2 neutrons, giving it a mass number of 4.
The most common use of this formula is finding how many neutrons an atom contains. You rearrange the equation: number of neutrons = mass number minus atomic number (the number of protons). Take tungsten-186 as an example. Tungsten has 74 protons, so 186 minus 74 gives you 112 neutrons. For uranium-238, subtract its 92 protons and you get 146 neutrons. The process works the same way for any element.
Mass Number vs. Atomic Mass
These two terms sound almost identical but refer to different things. Mass number is a simple count of nuclear particles, so it’s always a whole number: 12, 14, 238. Atomic mass, on the other hand, is a weighted average that accounts for all naturally occurring versions of an element and their relative abundance. That’s why the atomic mass of carbon on the periodic table reads 12.011 rather than a clean 12. The decimal reflects the tiny contributions of heavier, rarer forms of carbon mixed in with the dominant form.
When a problem or textbook gives you a number after a dash (like carbon-12 or potassium-40), that number is the mass number, not the atomic mass.
How It Appears in Chemical Notation
Standard nuclear notation places the mass number as a superscript to the upper left of the element’s symbol, with the atomic number as a subscript below it. For carbon-12, you’d see the symbol C with 12 above and 6 below on the left side. The 12 tells you the total protons plus neutrons; the 6 tells you it’s carbon. This compact format lets you identify the exact version of an element at a glance.
A shorthand alternative skips the subscript and just writes the element name followed by a dash and the mass number: carbon-12, uranium-238, potassium-40. Both formats convey the same information.
Isotopes: Same Element, Different Mass Numbers
Isotopes are atoms of the same element that differ in their number of neutrons, which means they have the same atomic number but different mass numbers. Carbon is a classic example. Every carbon atom has exactly 6 protons, but the number of neutrons varies from six to eight, producing three isotopes: carbon-12 (6 neutrons), carbon-13 (7 neutrons), and carbon-14 (8 neutrons).
Carbon-12 is by far the most abundant. Carbon-13 and carbon-14 exist in much lower amounts. Carbon-14 is radioactive, which is why it’s useful for dating ancient organic materials. The key point is that adding or removing neutrons changes the mass number and can change the atom’s stability, but it doesn’t change what element the atom is. Protons determine the element; neutrons determine the isotope.
How Radioactive Decay Changes Mass Number
When unstable atoms undergo radioactive decay, the mass number may or may not change depending on the type of decay involved.
- Alpha decay ejects a cluster of 2 protons and 2 neutrons from the nucleus. This drops the mass number by 4. Uranium-238, for instance, emits an alpha particle and becomes thorium-234: 238 minus 4 equals 234.
- Beta decay converts a neutron into a proton inside the nucleus, releasing an electron in the process. Because the total count of nuclear particles stays the same (one neutron lost, one proton gained), the mass number does not change. Thorium-234 undergoes beta decay to become protactinium-234, keeping its mass number at 234.
- Gamma emission releases pure energy from the nucleus with no particles ejected. The mass number stays completely unchanged.
These rules make mass number a useful bookkeeping tool for tracking what happens during nuclear reactions. Both sides of a nuclear equation must have the same total mass number, which is how scientists verify that an equation is balanced correctly.
Why Mass Number Matters in Practice
Mass number is foundational to how scientists identify and distinguish atoms. In mass spectrometry, a laboratory technique used to analyze unknown substances, instruments separate atoms and molecules based on their mass. The concept traces back to the early 1900s, when researchers first discovered isotopes by observing that atoms of the same element could have different masses. Modern mass spectrometers can determine atomic masses with accuracy better than one part in ten million, and those measurements rely on knowing which isotopes are present and in what proportions.
On a more everyday level, mass number is what you’re referencing any time you hear terms like “carbon-14 dating” or “uranium-235 fuel.” The number after the element name tells you exactly which isotope is involved, which determines its behavior, stability, and practical uses. Without mass number as a concept, there would be no clear way to distinguish between isotopes that look chemically identical but have very different nuclear properties.