Atomic mass and mass number both describe aspects of an atom’s mass, but they measure different things. Mass number is the total number of protons and neutrons in an atom’s nucleus, while atomic mass is a measure of the atom’s actual mass. The distinction becomes especially important when comparing isotopes, which are different forms of the same chemical element.
Although the two terms are related, they are not interchangeable. Understanding how they differ helps explain the numbers on the periodic table, why atoms of the same element can have different masses, and how scientists identify and study individual atoms.
What is atomic mass?
Atomic mass describes how much matter an atom contains, expressed using a unit called the atomic mass unit, abbreviated as u. This unit is defined in relation to carbon-12: one atomic mass unit equals exactly one-twelfth of the mass of a neutral carbon-12 atom in its ground state.
The masses of individual atoms are extremely small when expressed in grams, so atomic mass units provide a more convenient way to compare them. A single carbon-12 atom has a mass of exactly 12 u by definition, while a typical oxygen-16 atom has a mass of approximately 15.995 u.
An atom’s mass comes primarily from its nucleus, which contains protons and neutrons. Electrons also contribute to the total, but their individual masses are much smaller than those of protons and neutrons.
Atomic mass is not always a whole number because protons and neutrons do not each contribute exactly 1 u, and the total mass of an atom is affected by the energy that binds its nucleus together. In particular, nuclear binding energy reduces the mass of a bound nucleus compared with the combined mass of its separate protons and neutrons.
For an individual isotope, scientists can determine its atomic mass through precise measurements. The result is often close to, but not exactly equal to, its mass number.
What is mass number?
Mass number is a whole-number count of the particles in an atom’s nucleus. It is calculated by adding the number of protons and the number of neutrons.
Mass number = number of protons + number of neutrons
Protons determine which element an atom belongs to. For example, every carbon atom has six protons, and every oxygen atom has eight. The number of neutrons can vary without changing the element’s identity.
Consider carbon-12. Its nucleus contains six protons and six neutrons, giving it a mass number of 12. Carbon-14 also has six protons, but it contains eight neutrons, so its mass number is 14.
The mass number is always a whole number because protons and neutrons are counted as individual particles. It does not represent a precise measurement of an atom’s mass.
Scientists commonly write an isotope’s mass number after the element’s name, as in carbon-12 or uranium-238. In nuclear notation, the mass number appears as a superscript to the left of the element’s chemical symbol. Carbon-12, for example, is written as ¹²C.
The key differences between atomic mass and mass number
The central distinction is that mass number counts particles, whereas atomic mass measures mass.
| Feature | Atomic mass | Mass number |
|---|---|---|
| What it describes | The mass of an atom | The total number of protons and neutrons |
| How it is determined | Measured experimentally or calculated from physical data | Found by counting protons and neutrons |
| Typical value | Often a decimal | Always a whole number |
| Unit | Atomic mass units (u) | No unit |
| Changes between isotopes? | Yes | Yes |
| Affected by nuclear binding energy? | Yes | No, because it is a particle count |
These differences explain why an isotope can have a mass number of 16 but an atomic mass slightly below 16 u. The number 16 counts its protons and neutrons; the measured mass reflects the physical properties of the entire atom.
The two quantities are closely related because protons and neutrons account for most of an atom’s mass. As a result, an isotope’s atomic mass in atomic mass units is often close to its mass number. However, this approximation is not precise enough to treat the values as identical.
How isotopes reveal the difference
Isotopes provide one of the clearest ways to understand the relationship between atomic mass and mass number.
Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Because the proton count remains unchanged, all isotopes of an element have essentially the same chemical behavior, although their masses can influence some physical properties and reaction rates.
Carbon has several isotopes, including carbon-12 and carbon-14. Both contain six protons, but their neutron counts differ.
Carbon-12 contains six neutrons, giving it a mass number of 12. Carbon-14 contains eight neutrons, giving it a mass number of 14. Their atomic masses are approximately 12 u and 14.003 u, respectively.
The difference in mass is related to the additional neutrons, but the atomic masses are not simply equal to their mass numbers. Nuclear binding energy and other small contributions affect the measured values.
Carbon-14 is also radioactive, meaning its nucleus is unstable and eventually decays into another nucleus. Carbon-12 is stable. This illustrates another important point: mass number helps identify an isotope, but it does not by itself tell us whether that isotope is stable or radioactive.
Why the atomic mass on the periodic table is often a decimal
A periodic table usually lists an atomic mass that is not a whole number. For many elements, this value is a weighted average of the masses of the element’s naturally occurring isotopes.
A weighted average accounts for both the mass of each isotope and how common it is in nature. An isotope that occurs more frequently contributes more to the average than one that is rare.
Chlorine is a useful example. Naturally occurring chlorine consists mainly of chlorine-35 and chlorine-37. Chlorine-35 has 17 protons and 18 neutrons, while chlorine-37 has 17 protons and 20 neutrons.
Their mass numbers are 35 and 37, respectively. Their atomic masses are approximately 34.969 u and 36.966 u. Because chlorine-35 is more abundant in nature, it contributes more heavily to the average atomic mass.
The resulting standard atomic weight of chlorine is approximately 35.45. This decimal does not mean that an individual chlorine atom has 35.45 protons and neutrons. No atom can contain a fraction of a proton or neutron. Instead, the value represents the average mass associated with a natural mixture of chlorine isotopes.
For this reason, the atomic mass shown on a periodic table is generally not the mass number of any single isotope. It describes the element’s isotopic mixture rather than one specific atom.
There is an important distinction between the terms atomic mass and average atomic mass. Atomic mass can refer to the mass of a particular atom or isotope, while the value commonly listed on a periodic table generally represents an isotope-weighted average. For elements whose natural isotopic composition varies, the standard atomic weight may be reported as an interval rather than a single fixed value.
How to calculate mass number and identify an isotope
Calculating mass number is straightforward when the numbers of protons and neutrons are known. Add the two quantities together, and the result is the mass number.
For example, an atom with 11 protons and 12 neutrons has a mass number of 23. An atom with 17 protons and 18 neutrons has a mass number of 35.
The number of protons also identifies the element. An atom with 11 protons is sodium, so an atom containing 11 protons and 12 neutrons is sodium-23.
If the mass number and proton count are known instead, the neutron count can be found by rearranging the equation:
Number of neutrons = mass number − number of protons
A magnesium-24 atom has 12 protons because magnesium’s atomic number is 12. Subtracting 12 from its mass number of 24 shows that it contains 12 neutrons.
Atomic number and mass number therefore provide complementary information. Atomic number identifies the element, while mass number distinguishes its isotopes. Neither number alone gives the atom’s exact measured mass.
If an isotope’s atomic mass is provided, it generally cannot be used as an exact substitute for its mass number. The mass number is a whole-number count, so identifying an isotope from a measured mass requires interpreting that measurement in the context of known isotopes and their properties.
Why atomic mass and mass number matter in science
The distinction between atomic mass and mass number is important in chemistry, physics, and nuclear science.
In chemistry, average atomic masses allow scientists to calculate molar masses and determine how much of a substance is needed for a reaction. For example, the atomic masses listed on the periodic table help determine the mass of a mole of a compound, which connects the microscopic world of atoms to quantities that can be measured in a laboratory.
In nuclear physics, mass numbers help scientists describe nuclear reactions and track how protons and neutrons are distributed before and after a reaction. In a typical nuclear reaction, the total number of nucleons—the collective term for protons and neutrons—is conserved, so the sum of the mass numbers is the same on both sides of the nuclear equation.
Precise atomic masses also reveal information about nuclear binding energy. When nucleons combine to form a nucleus, energy is released, and the resulting bound system has less mass than the same particles would have when separated. The relationship between mass and energy, expressed by Einstein’s equation E = mc², explains why this mass difference corresponds to binding energy.
Isotopes have additional scientific applications. Carbon-14 is used in radiocarbon dating to estimate the age of certain once-living materials, while stable isotopes can help researchers investigate environmental processes, biological pathways, and the origins of materials.
Each application relies on distinguishing an isotope’s identity and nuclear composition from its measured mass and its abundance in a sample.
The essential distinction
Atomic mass and mass number describe related but different properties of atoms. Mass number is the whole-number total of protons and neutrons in a nucleus. Atomic mass is the actual mass of an atom, expressed in atomic mass units, and is influenced by the masses of its constituent particles and the energy associated with their binding.
For a specific isotope, the atomic mass is usually close to its mass number when expressed in atomic mass units, but the two values are not identical. On the periodic table, the decimal value usually represents a weighted average of naturally occurring isotopes rather than the mass of one particular atom.
Remembering the difference between a count and a measurement makes both concepts easier to understand: mass number tells you how many protons and neutrons an atom contains, while atomic mass tells you how much that atom weighs on the atomic scale.
