Mass number is a simple count of the protons and neutrons inside a single atom, always expressed as a whole number. Atomic mass is the weighted average mass of all naturally occurring isotopes of an element, which is why it almost always appears as a decimal on the periodic table. That distinction, whole number count versus weighted average, is the core difference.
What Mass Number Tells You
Mass number is the total number of protons and neutrons in one specific atom’s nucleus. The formula is straightforward: mass number equals the number of protons (also called the atomic number) plus the number of neutrons. Because you’re counting whole particles, mass number is always a whole number.
Carbon-12, for example, has 6 protons and 6 neutrons, giving it a mass number of 12. Carbon-14 has 6 protons and 8 neutrons, so its mass number is 14. Both are carbon because they share the same number of protons, but they’re different isotopes because they have different numbers of neutrons. Every isotope of an element has its own mass number.
What Atomic Mass Tells You
Atomic mass (sometimes called atomic weight or relative atomic mass) accounts for the fact that most elements exist in nature as a mixture of isotopes. Rather than describing one specific atom, it represents the average mass across all of an element’s naturally occurring isotopes, weighted by how common each one is. This is the decimal number printed beneath each element’s symbol on the periodic table.
The calculation works by multiplying each isotope’s mass by its natural abundance (as a decimal fraction), then adding the results together. Chlorine is a clean example. It naturally occurs as two isotopes: chlorine-35, which makes up about 75% of all chlorine atoms, and chlorine-37, which makes up about 25%. The weighted average comes out to roughly 35.45, not 35 or 37 but somewhere between, pulled closer to 35 because that isotope is three times more abundant.
This is why you’ll never find an atom that actually weighs 35.45 units. That number is a statistical average, like saying the average household has 2.3 children. No real household has 2.3 children, but the number accurately describes the population as a whole.
Why One Is a Whole Number and the Other Isn’t
Mass number will always be a whole number because protons and neutrons come in whole units. You can’t have half a neutron in a nucleus. Atomic mass, on the other hand, is almost always a decimal for two reasons. First, averaging across isotopes with different abundances naturally produces non-integer results. Second, the actual measured mass of an atom doesn’t perfectly equal the sum of its parts. A small fraction of mass is converted into the binding energy that holds the nucleus together, so even a single isotope’s precise mass isn’t a perfectly round number. Nitrogen-14, for instance, has a measured mass of 14.00643 atomic mass units rather than exactly 14.
The unit used for these measurements, the atomic mass unit (amu), is defined as exactly one-twelfth the mass of a carbon-12 atom. That’s an extraordinarily small quantity: about 1.66 × 10⁻²⁴ grams. The same unit is also called a dalton.
Side-by-Side Comparison
- What it describes: Mass number applies to a single atom of a specific isotope. Atomic mass applies to an element as a whole, averaging all its isotopes.
- Type of value: Mass number is always a whole number. Atomic mass is nearly always a decimal.
- How it’s calculated: Mass number is found by adding protons and neutrons. Atomic mass is found by weighting each isotope’s mass by its natural abundance.
- Where you’ll see it: Mass number appears in isotope notation (the superscript in ¹²C or ²³⁵U). Atomic mass appears on the periodic table beneath the element symbol.
- Units: Mass number has no units because it’s a count. Atomic mass is expressed in atomic mass units (amu).
Carbon: A Quick Example
Carbon has three naturally occurring isotopes. Carbon-12 (mass number 12) has 6 protons and 6 neutrons. Carbon-13 (mass number 13) has 6 protons and 7 neutrons. Carbon-14 (mass number 14) has 6 protons and 8 neutrons. Each of those mass numbers is a clean integer.
But on the periodic table, carbon’s atomic mass is listed as approximately 12.011. That value is dominated by carbon-12, which accounts for about 98.9% of all carbon in nature, with carbon-13 making up most of the remaining ~1.1%. Carbon-14 is so rare that it barely affects the average. The result is a decimal just slightly above 12.
Why Atomic Mass Varies for Some Elements
For most elements, the atomic mass listed on the periodic table is a single number. But for twelve elements, IUPAC (the international body that standardizes chemical data) lists a range instead of a fixed value. This happens when the isotope mix varies depending on where a sample comes from. Hydrogen’s atomic mass, for instance, falls between 1.00784 and 1.00811 depending on the source. Lead shows even wider variation, ranging from 206.14 to 207.94, because different geological processes concentrate different lead isotopes.
Mass number, by contrast, never varies for a given isotope. Lead-208 always has a mass number of 208 regardless of where the atom originated. It’s the proportion of lead-206, lead-207, and lead-208 in a natural sample that shifts, which is what causes the atomic mass to fluctuate.
How to Tell Them Apart on a Test
If a question asks for the mass number, you’re being asked to count protons and neutrons in one specific atom. The answer will be a whole number. If a question asks for atomic mass (or atomic weight), you’re being asked for the weighted average across isotopes. The answer will typically be a decimal, and you’ll need to know each isotope’s mass and its percent abundance to calculate it.
A common source of confusion is that the two values are sometimes close. Fluorine, for example, has only one stable isotope (fluorine-19), so its atomic mass (about 18.998) is very close to its sole mass number of 19. Elements with one dominant isotope will show a similar pattern. The gap between the two values becomes more obvious for elements like chlorine, where two isotopes exist in significantly different proportions.