Atomic mass comes almost entirely from two subatomic particles: protons and neutrons. These particles sit in the nucleus of every atom and each weighs approximately 1 atomic mass unit (amu). Electrons, the third component of an atom, contribute so little mass that they’re effectively ignored in the calculation.
Protons and Neutrons Do the Heavy Lifting
A proton has a mass of 1.007 amu, and a neutron comes in slightly heavier at 1.009 amu. An electron, by contrast, weighs about 0.0005 amu, roughly 1/1,836th the mass of a proton. That’s why when scientists talk about atomic mass, they’re really talking about what’s happening inside the nucleus. The electrons orbiting outside are essentially rounding errors.
The total count of protons plus neutrons gives you the “mass number,” which is always a whole number. A carbon atom with 6 protons and 6 neutrons has a mass number of 12. A helium atom with 2 protons and 2 neutrons has a mass number of 4. This integer is the simplest way to approximate an atom’s mass, and it’s often all you need for basic chemistry.
Why Atomic Mass Isn’t a Whole Number
If you look at the periodic table, you’ll notice that atomic masses are decimals, not clean integers. Carbon is listed at 12.01, chlorine at roughly 35.45, lithium at 6.94. Two things explain this.
First, most elements exist in nature as a mix of isotopes, atoms that have the same number of protons but different numbers of neutrons. Chlorine, for example, comes in two stable forms: chlorine-35 (with 18 neutrons) and chlorine-37 (with 20 neutrons). About 75.8% of all chlorine atoms on Earth are chlorine-35, and the remaining 24.2% are chlorine-37. The atomic mass on the periodic table is a weighted average of these isotopes based on their natural abundance. For chlorine, that works out to about 35.45, closer to 35 because the lighter isotope is far more common.
The formula is straightforward: multiply each isotope’s mass by its percentage abundance, then add the results together. This is why the number on the periodic table is technically called the “standard atomic weight” or “relative atomic mass” of an element, not the mass of any single atom.
Mass Number vs. Atomic Mass
These two terms sound interchangeable, but they describe different things. Mass number is the simple integer count of protons and neutrons in one specific atom. Atomic mass (as shown on the periodic table) is the weighted average across all naturally occurring isotopes of that element, expressed as a decimal.
For hydrogen, the mass number of the most common isotope is 1 (one proton, zero neutrons), but the periodic table lists its atomic mass as 1.01 because a tiny fraction of hydrogen atoms are deuterium (one proton plus one neutron). For elements with more isotopes spread across a wider range, the difference between the mass number of the most common isotope and the listed atomic mass can be more noticeable.
The Missing Mass: Binding Energy
There’s one more subtlety that affects atomic mass. If you add up the individual masses of all the protons and neutrons in a nucleus, the total is slightly more than the actual measured mass of that nucleus. The difference is called the mass defect.
This isn’t an error. When protons and neutrons come together to form a nucleus, a small amount of mass converts into energy, the energy that holds the nucleus together. This is the principle behind Einstein’s famous equation, E = mc². The converted mass is tiny on an everyday scale, but it’s measurable with precision instruments. It also explains why nuclear reactions release such enormous amounts of energy: even a small amount of mass translates into a large quantity of energy when multiplied by the speed of light squared.
For most chemistry purposes, the mass defect is negligible. But it’s the reason that the precise atomic mass of an isotope is never exactly equal to its mass number. Carbon-12, for instance, is defined as exactly 12.000 amu (it’s the reference standard), but other isotopes always show slight deviations from their expected whole-number masses because their binding energies differ.
The Unit of Measurement
Atomic mass is measured in unified atomic mass units, symbolized as “u” or “Da” (daltons). One atomic mass unit is defined as exactly one-twelfth the mass of a carbon-12 atom. In absolute terms, that’s about 1.66 × 10⁻²⁷ kilograms, a number so small it’s only useful in particle physics. The amu scale exists because it gives convenient, human-readable numbers: hydrogen is about 1, oxygen about 16, iron about 56.
You may see the older abbreviation “amu” used interchangeably with “u.” The original amu was based on oxygen rather than carbon, but that standard was replaced in 1961. Today, when textbooks write “amu,” they almost always mean the carbon-12-based unit.