What Are Isotopes? Atoms of the Same Element with Different Masses

Isotopes are atoms of the same element that contain different numbers of neutrons. Every carbon atom has six protons, which is what makes it carbon, but carbon atoms can carry six, seven or eight neutrons. Those versions have nearly identical chemistry yet different masses, and some of them are unstable and decay over time.

This article explains how isotopes are named, why the average atomic mass on the periodic table is rarely a whole number, and how isotopes are used in dating, medicine and industry. The figures used below are standard reference values that you can check against any modern chemistry text.

Protons define the element

The atomic number Z counts the protons in the nucleus. Changing Z changes the element. An atom with one proton is hydrogen, an atom with eight is oxygen, and an atom with 92 is uranium. No chemical reaction changes the proton count, which is why the atomic number is the fixed identity of an element.

Neutrons are the other nuclear particle. They have almost the same mass as protons but no electric charge. Neutrons add mass and help hold the nucleus together against the repulsion between positive protons, but they do not change which element an atom is.

Mass number and notation

The mass number A is the total count of protons and neutrons in a particular atom. The number of neutrons N is simply A minus Z.

Isotopes are written with the mass number before the element symbol, for example carbon-12, carbon-13 and carbon-14. The notation ¹²C shows the same information with the mass number as a superscript.

Isotope	Protons	Neutrons	Mass number	Stable?
Carbon-12	6	6	12	Yes
Carbon-13	6	7	13	Yes
Carbon-14	6	8	14	No (radioactive)

Carbon-14 decays by emitting a beta particle, with a half-life of about 5,730 years. That slow decay is the basis of radiocarbon dating, discussed below.

Natural abundance

Most elements occur in nature as a mixture of isotopes. The fraction of each isotope is called its natural abundance, and it is stable across most ordinary samples on Earth, which is why chemists can quote a single average atomic mass for an element.

Typical natural abundances for common elements are shown below.

Element	Isotope	Approximate natural abundance
Hydrogen	Protium (H-1)	about 99.98%
Hydrogen	Deuterium (H-2)	about 0.015%
Carbon	Carbon-12	about 98.9%
Carbon	Carbon-13	about 1.1%
Chlorine	Chlorine-35	about 75.8%
Chlorine	Chlorine-37	about 24.2%

Hydrogen-3, called tritium, is also present in trace amounts but is not listed because it is radioactive and produced mainly by cosmic rays and in nuclear reactors. Its half-life is about 12.3 years.

Average atomic mass

The atomic mass shown on the periodic table is a weighted average of the masses of an element's isotopes. Each isotope's mass is multiplied by its fractional abundance, and the products are added together.

For chlorine, the calculation uses the two stable isotopes:

0.758 × 34.969 + 0.242 × 36.966 ≈ 35.45

That result explains why chlorine's atomic mass is about 35.45 rather than a whole number. No single chlorine atom has a mass of 35.45. The number describes the average over a large sample.

This also explains why some atomic masses look odd. The value is only as precise as the abundance data behind it, and abundances can vary slightly between sources, such as rocks from different regions.

Stable and radioactive isotopes

An isotope is stable when its nucleus does not change over time. Stability depends on the balance between protons and neutrons. Light elements tend to be stable with roughly equal numbers of each, while heavy elements need more neutrons than protons to stay bound together.

When a nucleus has too many or too few neutrons for its size, it becomes radioactive. It decays toward a more stable configuration by emitting radiation. Common decay modes include alpha decay, where a helium-4 nucleus leaves; beta decay, where a neutron turns into a proton and an electron; and gamma decay, where excess energy leaves as a photon.

The half-life is the time for half of a sample to decay. It does not depend on temperature, pressure or chemical form. A sample with a half-life of one day halves in one day, halves again in the next, and so on. After ten half-lives, about one thousandth of the original atoms remain.

Uses of isotopes

Isotopes are among the most practical tools in chemistry and physics.

Radiocarbon dating

Living organisms take in carbon from the air and food, keeping their ratio of carbon-14 to carbon-12 close to that of the atmosphere. After death, the carbon-14 decays without being replaced. Measuring the remaining ratio gives an estimate of how long ago the organism died. The method works for material up to roughly 50,000 years old, because beyond that little carbon-14 remains to measure.

Medical imaging and treatment

Radioactive isotopes such as technetium-99m are widely used in diagnostic imaging, because their gamma emissions can be detected outside the body. Other isotopes are used in cancer treatment, where the radiation is aimed at diseased tissue. The choice of isotope depends on its half-life, the type of radiation it emits and how easily it can be attached to a useful molecule.

Nuclear energy

Uranium-235 is the fissile isotope that powers most nuclear reactors. Natural uranium contains only about 0.7% uranium-235, so it is enriched before use in many reactor designs. The rest of the uranium, mostly uranium-238, is not easily fissioned by slow neutrons.

Tracing reactions

Chemists use isotopic labels, such as deuterium or carbon-13, to follow atoms through a reaction. Because isotopes have the same chemistry, the labelled atom goes wherever its unlabelled partner goes. Mass spectrometry or nuclear magnetic resonance can then show where the label ended up, revealing the reaction pathway.

Common misconceptions

Students sometimes assume that isotopes are different elements or that they react differently in ordinary chemistry. Neither is true. Isotopes have the same number of electrons in a neutral atom, so they have the same electron configuration and very similar chemical behaviour. The one important exception is the kinetic isotope effect, where a heavier isotope reacts somewhat more slowly, because bonds to heavier atoms vibrate with lower energy. That effect is small for most elements but large for hydrogen and deuterium.

Another misconception is that all heavy isotopes are dangerous. Many heavy isotopes are stable and harmless. Radioactivity depends on the specific nucleus, not on mass alone.

Key takeaways

Isotopes share a proton count and therefore an element identity, but they differ in neutron count and so in mass and nuclear stability. The mass number is the sum of protons and neutrons, and the notation shows it clearly.

The average atomic mass on the periodic table is a weighted average of natural isotope abundances. That is why it is rarely a whole number, and why the value reflects a population rather than a single atom.

Radioactive isotopes decay at a fixed rate described by their half-life. That property makes isotopes useful for dating, medicine, energy and chemical research, and it is the reason the topic matters well beyond the classroom.