Atomic mass is the mass of a single atom in its ground state. It is usually expressed in unified atomic mass units (u or Da) and is one of the most fundamental quantities in chemistry because it connects the microscopic world of atoms to measurable amounts of matter. Atomic mass depends on the numbers of protons, neutrons, and, to a much smaller extent, electrons in an atom. Since most elements occur as mixtures of isotopes, the atomic masses listed on periodic tables are usually weighted averages rather than the mass of any one atom.
Atomic mass plays a central role in chemistry, physics, geology, biology, and materials science. It is essential for identifying isotopes, calculating molar masses, balancing chemical equations, interpreting mass spectra, determining molecular formulas, and converting between the number of atoms and measurable masses in the laboratory.
Key Takeaways: Atomic Mass
- Atomic mass is the mass of a single atom.
- The standard unit is the unified atomic mass unit (u), also called the dalton (Da).
- One atomic mass unit is defined as exactly 1/12 the mass of a carbon-12 atom.
- Atomic mass depends primarily on the numbers of protons and neutrons.
- Different isotopes of the same element have different atomic masses.
- Periodic tables usually list the standard atomic weight, not the atomic mass of an individual atom.
- Atomic mass is used to calculate molar mass, molecular mass, stoichiometric relationships, and chemical formulas.
Why Atomic Mass Is Important
Atomic mass provides the bridge between the atomic scale and the macroscopic world.
It is essential for:
- Identifying isotopes
- Determining molecular formulas
- Calculating molar masses
- Balancing chemical equations
- Performing stoichiometric calculations
- Understanding nuclear reactions
- Interpreting mass spectrometry data
- Predicting physical properties of isotopes
- Measuring abundances of naturally occurring isotopes
Without atomic mass, chemists could not accurately determine how much material participates in a reaction or compare atoms of different elements quantitatively.
History of Atomic Mass
The concept of atomic mass developed gradually as scientists learned that matter consists of atoms.
Early 1800s: Dalton’s Atomic Theory
In 1803, John Dalton proposed that each element consists of identical atoms having characteristic masses. Because actual atomic masses were unknown, Dalton assigned hydrogen a mass of 1 and expressed all other masses relative to hydrogen.
Early Atomic Weight Tables
During the nineteenth century, chemists compared the masses of elements through chemical reactions. These values became known as atomic weights, even though they were relative masses rather than actual weights.
Cannizzaro’s Contribution
In 1858, Stanislao Cannizzaro clarified how Avogadro’s hypothesis could determine reliable atomic masses, bringing consistency to atomic weight measurements.
Oxygen Standard
For many years chemists defined atomic masses relative to oxygen. Unfortunately, physicists and chemists used slightly different oxygen standards, leading to confusion.
Carbon-12 Standard
In 1961, the international scientific community adopted the carbon-12 standard, defining the atomic mass unit as exactly one-twelfth the mass of a neutral carbon-12 atom in its ground state. This standard remains in use today.
Modern Measurements
Today, highly precise atomic masses come from mass spectrometry, allowing measurements accurate to many decimal places.
What Is Atomic Mass?
Atomic mass is the actual mass of one atom of a specific isotope.
It includes the mass of:
- Protons
- Neutrons
- Electrons
However, electrons contribute only about 0.05% of an atom’s total mass, so nearly all atomic mass comes from the nucleus.
For example:
- Carbon-12 atom: exactly 12 u
- Carbon-13 atom: approximately 13.003355 u
- Hydrogen-1 atom: approximately 1.007825 u
Notice that these values are not whole numbers because nuclear binding energy slightly changes the total mass.
Atomic Mass Units (u)
Atomic masses are extremely small.
Rather than kilograms, chemists use the unified atomic mass unit (u).
Definition
One unified atomic mass unit equals:
1 u = 1/12 the mass of one carbon-12 atom
Exactly.
Other accepted names include:
- atomic mass unit (amu)
- unified atomic mass unit (u)
- dalton (Da)
The SI accepts the dalton for use alongside SI units.
Conversion
1 u ≈ 1.66053906660 × 10⁻²⁷ kg
Why Carbon-12?
Carbon-12 was chosen because it:
- is stable
- is abundant
- is easy to measure
- avoids earlier disagreements involving oxygen standards
- provides a reproducible international reference
Every atomic mass ultimately traces back to this definition.
Isotopes and Atomic Mass
Different isotopes contain the same number of protons but different numbers of neutrons.
For example:
| Isotope | Protons | Neutrons | Atomic Mass (u) |
|---|---|---|---|
| Hydrogen-1 | 1 | 0 | 1.007825 |
| Hydrogen-2 | 1 | 1 | 2.014102 |
| Hydrogen-3 | 1 | 2 | 3.016049 |
Although all three atoms are hydrogen, each has a different atomic mass.
Why Are Atomic Masses Not Whole Numbers?
A common question is why carbon-13 has a mass of 13.003355 u instead of exactly 13 u.
The answer lies in mass defect.
When protons and neutrons bind together in a nucleus, some mass converts into nuclear binding energy according to Einstein’s equation:
E = mc²
As a result,
atomic mass < sum of individual proton and neutron masses
The stronger the nucleus is bound, the larger the mass defect.
Atomic Mass vs. Mass Number
Atomic mass and mass number are related, but they are not the same quantity.
The mass number (A) is simply the total number of protons and neutrons in an atom’s nucleus. Because protons and neutrons are counted as whole particles, the mass number is always a whole number.
The atomic mass, on the other hand, is the actual measured mass of an atom. It includes the masses of the protons, neutrons, and electrons, as well as the effects of nuclear binding energy. As a result, atomic masses are rarely whole numbers.
For example, the most common isotope of carbon, carbon-12, has a mass number of 12 and, by definition, an atomic mass of exactly 12 u. Carbon-13 has a mass number of 13, but its atomic mass is 13.003355 u because of the mass defect resulting from nuclear binding energy.
| Property | Atomic Mass | Mass Number |
|---|---|---|
| What it represents | Actual mass of one atom | Number of protons + neutrons |
| Symbol | m | A |
| Includes electrons? | Yes (although their contribution is very small) | No |
| Includes binding energy effects? | Yes | No |
| Decimal or whole number? | Usually a decimal | Always a whole number |
| Units | u (Da) | None |
| Example (carbon-13) | 13.003355 u | 13 |
Why Do Different Periodic Tables Show Different Atomic Masses?
Many people notice that different periodic tables list slightly different values.
For example, chlorine may appear as:
- 35.45
- 35.453
- [35.446–35.457]
All of these can be correct.
Reasons include:
Different Rounding
Some tables round to two decimal places.
Others show three or four decimal places.
Updated IUPAC Values
Natural isotope abundances are measured more accurately over time.
When these measurements improve, recommended values may change slightly.
Standard Atomic Weight Intervals
Several elements naturally vary enough in isotopic composition that IUPAC now gives an interval rather than a single value.
Examples include:
- hydrogen
- carbon
- nitrogen
- oxygen
- sulfur
- chlorine
Educational periodic tables usually simplify these intervals into one representative value.
Periodic Table With Atomic Mass
The periodic table is one of the easiest places to find atomic masses (more precisely, standard atomic weights). Each element box typically includes a decimal value beneath the element symbol that represents the weighted average for naturally occurring isotopes. Because these values are averages, they often fall between the mass numbers of the most abundant isotopes.
A downloadable periodic table that includes atomic masses provides a convenient reference for solving chemistry problems involving molar mass, molecular mass, and stoichiometry.
How to Find Atomic Mass
For a Single Isotope
Use a mass spectrometry table or isotope data.
Example:
Carbon-13
Atomic mass = 13.003355 u
For an Element Found in Nature
Calculate the weighted average.
The formula is:
Average atomic mass = Σ(fractional abundance × isotopic mass)
Convert percentages into decimals before multiplying.
Chlorine occurs naturally as two main isotopes:
| Isotope | Natural abundance | Isotope mass (amu) |
|---|---|---|
| Chlorine-35 | 75.77% (0.7577) | 34.96885 |
| Chlorine-37 | 24.23% (0.2423) | 36.96590 |
Step 1: Convert the percentages to decimals.
75.77% = 0.7577
24.23% = 0.2423
Step 2: Multiply each isotope’s mass by its fractional abundance.
Contribution from chlorine-35:
0.7577 × 34.96885 = 26.50 amu
Contribution from chlorine-37:
0.2423 × 36.96590 = 8.96 amu
Step 3: Add the contributions.
26.50 + 8.96 = 35.46 amu
This value matches the atomic mass listed on the periodic table (35.45 amu, after rounding).
Example: Boron
Boron has two naturally occurring isotopes.
| Isotope | Natural abundance | Isotope mass (amu) |
|---|---|---|
| Boron-10 | 19.9% (0.199) | 10.01294 |
| Boron-11 | 80.1% (0.801) | 11.00931 |
Multiply each isotope’s mass by its abundance:
10.01294 × 0.199 = 1.99 amu
11.00931 × 0.801 = 8.82 amu
Add the contributions:
1.99 + 8.82 = 10.81 amu
This is the average atomic mass of boron.
Atomic Mass vs Atomic Weight
Although people often use the terms interchangeably, they are not identical.
| Atomic Mass | Atomic Weight |
|---|---|
| Mass of one specific atom | Weighted average relative mass of naturally occurring atoms |
| Refers to a particular isotope | Refers to an element in a specified material or natural terrestrial sample |
| Changes between isotopes | Depends on isotope abundances |
| Units: u (Da) | Dimensionless (although commonly reported numerically like atomic mass values) |
| Example: carbon-13 = 13.003355 u | Carbon = 12.011 |
In most introductory chemistry courses, the number shown on the periodic table is called the atomic mass, even though it is technically the standard atomic weight.
Atomic Mass Compared with Related Mass Terms
Many chemistry terms sound similar but describe different quantities. Understanding the distinctions helps avoid confusion when working with periodic tables, isotopes, molecular formulas, and stoichiometric calculations.
| Quantity | Definition | Typical Notation | Units |
|---|---|---|---|
| Atomic mass | Mass of a single atom of a specific isotope. | m (sometimes mₐₜₒₘ) | u (Da) |
| Relative atomic mass (atomic weight) | Weighted average relative mass of the atoms of an element compared with 1/12 the mass of a carbon-12 atom. | Aᵣ | Dimensionless |
| Standard atomic weight | IUPAC-recommended relative atomic mass for a normal terrestrial sample of an element. | Aᵣ(E) | Dimensionless |
| Isotopic mass | Mass of a single atom of a particular isotope. | m | u (Da) |
| Relative isotopic mass | Ratio of the mass of a specific isotope to 1/12 the mass of a carbon-12 atom. | Aᵣ | Dimensionless |
| Mass number | Total number of protons and neutrons in a nucleus. | A | None |
| Atomic number | Number of protons in a nucleus. | Z | None |
| Integer mass | Atomic mass rounded to the nearest whole number; commonly used in introductory isotope problems. | None (approximate value) | u |
| Average atomic mass | Weighted average of the isotopic masses of an element based on natural abundances. | Σ(fraction × isotopic mass) | u (Da) |
| Molecular mass | Sum of the atomic masses of all atoms in a molecule. | Usually written as the sum of atomic masses | u (Da) |
| Formula mass | Sum of the atomic masses of all atoms in one formula unit of an ionic compound or network solid. | Usually written as the sum of atomic masses | u (Da) |
| Molar mass | Mass of one mole of atoms, molecules, or formula units. Numerically equal to the corresponding atomic, molecular, or formula mass expressed in u. | M | g/mol |
Atomic Mass and Molar Mass
One of the most elegant relationships in chemistry is that the numerical value of an atomic mass in u equals the numerical value of the molar mass in grams per mole.
For example:
| Species | Atomic or Molecular Mass | Molar Mass |
|---|---|---|
| Hydrogen atom | 1.008 u | 1.008 g/mol |
| Carbon atom | 12.011 u | 12.011 g/mol |
| Water molecule | 18.015 u | 18.015 g/mol |
The numbers are the same because the mole is defined using the Avogadro constant, linking microscopic masses to macroscopic quantities.
How Atomic Mass Is Measured
Modern atomic masses are determined almost entirely by mass spectrometry.
In a mass spectrometer:
- Atoms are ionized.
- Electric fields accelerate the ions.
- Magnetic or electric fields separate ions according to their mass-to-charge ratio.
- Detectors measure their abundances.
Using highly precise instruments, scientists determine isotope masses to many significant figures.
Applications of Atomic Mass
Atomic mass is important in many scientific fields.
Examples include:
- Stoichiometric calculations
- Pharmaceutical chemistry
- Nuclear chemistry
- Geochemistry
- Radiometric dating
- Environmental chemistry
- Mass spectrometry
- Forensic science
- Astronomy
- Materials science
Common Misconceptions
Atomic mass and atomic weight are exactly the same.
Not quite. Atomic mass refers to one atom (or one isotope), while atomic weight is a weighted average relative to carbon-12.
The periodic table lists atomic masses.
Technically, most periodic tables list standard atomic weights, although the values are often labeled “atomic mass” in educational settings.
Atomic mass equals the mass number.
No. Mass number is a whole number counting protons and neutrons. Atomic mass includes mass defect and is rarely an integer.
Electrons contribute most of an atom’s mass.
No. Nearly all atomic mass comes from protons and neutrons.
Every atom of an element has the same atomic mass.
Only if all atoms are the same isotope. Most elements occur as mixtures of isotopes.
FAQs
What is the difference between atomic mass and mass number?
Mass number counts protons plus neutrons and is always an integer. Atomic mass is the measured mass of an atom and includes the effects of nuclear binding energy, so it is usually not a whole number.
Why is chlorine’s atomic mass 35.45 instead of 35 or 37?
Natural chlorine is a mixture of chlorine-35 and chlorine-37. The periodic table value is the weighted average of both isotopes.
Why is carbon-12 exactly 12?
The atomic mass unit is defined so that a neutral carbon-12 atom has a mass of exactly 12 u.
Is atomic mass measured or calculated?
Both. Scientists measure the masses of individual isotopes using mass spectrometry. The average atomic masses commonly used in chemistry are calculated from isotope masses and their natural abundances.
Why do some periodic tables list ranges instead of single values?
Some elements naturally have variable isotope abundances depending on their source. For these elements, IUPAC recommends an interval rather than a single standard atomic weight.
Is the dalton the same as the atomic mass unit?
Yes. The dalton (Da) and the unified atomic mass unit (u) are numerically equivalent. The dalton is widely used in biochemistry and molecular biology, especially for proteins and nucleic acids.
Can an atom’s atomic mass change?
Yes, if the atom undergoes a nuclear reaction that changes its isotope. Ordinary chemical reactions do not change atomic mass because they affect only electrons, not the nucleus.
References
- De Bievre, P.; Peiser, H. S. (1992). “‘Atomic weight’: The name, its history, definition, and units”. Pure Appl. Chem. 64 (10): 1535. doi:10.1351/pac199264101535
- International Bureau of Weights and Measures (2019). The International System of Units (SI) (9th ed.). ISBN 978-92-822-2272-0.
- IUPAC (2025). Compendium of Chemical Terminology (5th ed.). International Union of Pure and Applied Chemistry.
- Meija, J.; Coplen, T.B.; et al. (2016). “Isotopic compositions of the elements 2013 (IUPAC Technical Report)”. J. Pure Appl. Chem. 88(3): 293-306. doi:10.1515/pac-2015-0503
- Prohaska, T.; Irrgeher, J.; et al. (2022). “Standard atomic weights of the elements 2021 (IUPAC Technical Report)”. J. Pure Appl. Chem. 9(5): 573-600 doi:10.1515/pac-2019-0603