What are Isotopes? Definition, examples and uses

Every carbon atom in your body has six protons — that is what makes it carbon. But not every carbon atom weighs the same. Most carry six neutrons; a few carry seven, and a rare handful carry eight. They are all carbon, all chemically interchangeable, yet subtly different in mass. These variants are isotopes, and that small difference of a neutron or two turns out to be one of the most useful facts in all of science.

In this article you will learn what isotopes are, how they are written, why some are stable and others radioactive, why they make the atomic masses on the periodic table come out as decimals, and how isotopes are used to date the past, image the body and power reactors.

§ 01 What isotopes are

An element is defined by its number of protons, the atomic number (Z). All carbon atoms have Z = 6; all oxygen atoms have Z = 8. The neutrons, however, are free to vary. Change the neutron count and you do not change the element — you create a different isotope of it.

Because chemistry is governed by electrons (and a neutral atom has as many electrons as protons), isotopes of an element are chemically almost identical. They burn, bond and react the same way. What changes is the mass number (A) — the total of protons plus neutrons — and, for some, whether the nucleus is stable. To explore how this fits into atomic structure, see what an atom is.

§ 02 How isotopes are written

There are two common ways to name an isotope, and both encode the mass number:

• Name or symbol with the mass number — "carbon-14" or "C-14". The number is the mass number (protons + neutrons).
• Nuclear notation — the symbol with the mass number as a superscript and the atomic number as a subscript, both written to the left.

To find the number of neutrons in any isotope, simply subtract the atomic number from the mass number. Carbon-14 has 14 − 6 = 8 neutrons; uranium-235 has 235 − 92 = 143. The isotopes explorer lets you read off these numbers for every element at a glance.

§ 03 Stable versus radioactive isotopes

Whether an isotope lasts forever or decays comes down to the balance of protons and neutrons in its nucleus.

• Stable isotopes have a nucleus that holds together indefinitely. They do not decay or emit radiation. Most elements have one or more stable isotopes — carbon-12 and carbon-13, for instance.
• Radioactive isotopes (radioisotopes) have an unstable nucleus that decays over time, emitting radiation and transforming into another nucleus. Carbon-14 and tritium are examples.

Every radioisotope is characterised by its half-life — the time for half a sample to decay. The full story of how and why these nuclei break down is covered in what is radioactivity; the key point here is that being radioactive is a property of the isotope, not the element. Carbon is not radioactive, but the isotope carbon-14 is.

What tips an isotope from stable to unstable is the ratio of neutrons to protons. Light elements are most stable when the two are roughly balanced, while heavier elements need a growing surplus of neutrons to dilute the electrical repulsion between their many protons. Isotopes whose neutron count strays outside this narrow band of stability decay in order to move back toward it. That is why, for any given element, only a handful of neutron numbers yield stable atoms — and everything beyond them is radioactive.

§ 04 Isotopes and atomic mass

Here is the payoff that explains a puzzle every student notices. Look at the periodic table and the atomic masses are almost never whole numbers: chlorine is 35.45, copper is 63.55. If a single atom's mass is essentially a whole number of protons and neutrons, where do the decimals come from?

The answer is isotopes. The mass listed on the periodic table is the relative atomic mass — a weighted average of all the element's natural isotopes, accounting for how common each one is.

You can practise this calculation yourself: the average is each isotope's mass multiplied by its fractional abundance, all added together. Try it for any element with the molar-mass and average-mass tools in the calculators.

§ 05 Famous examples of isotopes

A few isotope families do an enormous amount of scientific work:

Hydrogen is the only element whose isotopes have their own names. Deuterium combined with oxygen makes "heavy water", used in some nuclear reactors, while the two uranium isotopes differ so little chemically that separating them requires the painstaking process of enrichment.

Most elements are a quiet blend of several isotopes you never notice. Oxygen, for instance, is overwhelmingly oxygen-16 but carries traces of oxygen-17 and oxygen-18 — and the exact ratio of those heavier forms, locked into ancient ice and seashells, is one of the most important records climate scientists have of the Earth's past temperatures. Even the air you are breathing is an isotopic mixture.

§ 06 How isotopes are used

The slight differences between isotopes — in mass and in stability — make them powerful tools across science, medicine and industry.

Dating the past

Radiocarbon dating uses carbon-14's steady decay to determine the age of once-living material up to around 50,000 years old. For rocks and the age of the Earth, geologists use uranium–lead and other long-lived isotope clocks reaching back billions of years.

Medicine

Radioactive isotopes are the backbone of nuclear medicine. Technetium-99m and fluorine-18 trace blood flow and metabolism in scans; iodine-131 treats thyroid disease; and targeted isotope therapies attack tumours from within.

Tracing and authentication

Because organisms and materials take up isotopes in characteristic ratios, stable-isotope analysis can reveal where food was grown, reconstruct ancient diets and climates, trace pollution, and provide evidence in forensic investigations. Isotope labels also let chemists follow exactly which atom ends up where in a reaction.

§ 07 How isotopes are separated and measured

Because isotopes of an element are chemically identical, telling them apart — or pulling one out from the rest — depends entirely on their small difference in mass.

The instrument that reads that difference is the mass spectrometer. It turns a sample into ions and accelerates them through a magnetic field; lighter isotopes are deflected more sharply than heavier ones, so the ions fan out by mass and a detector counts how many of each arrive. This one technique measures the abundances behind every relative atomic mass, dates archaeological finds, screens athletes for banned substances and reconstructs ancient climates.

Separating isotopes in bulk is far harder. Natural uranium is more than 99% uranium-238 and under 1% of the fissile uranium-235 that reactors need, yet no chemical reaction can tell the two apart. Enrichment instead exploits their tiny mass gap — most often by spinning uranium gas through cascades of high-speed centrifuges until the lighter isotope slowly concentrates. The sheer difficulty of that step is one reason nuclear fuel is so tightly controlled around the world.

§ 08 Interesting facts

• Tin is the champion. Tin has ten stable isotopes — more than any other element.
• Some elements have none. Technetium and promethium have no stable isotopes at all; every version of them is radioactive.
• Heavy water really is heavier. Water made with deuterium is about 11% denser than ordinary water — enough that ice cubes of heavy water sink.
• Carbon-12 sets the standard. The entire scale of atomic mass is defined so that one atom of carbon-12 weighs exactly 12 units.
• Your breath is a clock. The carbon-14 you take in while alive stops being replenished the moment you die — which is what makes radiocarbon dating possible.

§ 09 Common misconceptions

"Isotopes are different elements."

No — isotopes of an element are the same element, with the same chemistry and the same spot on the periodic table. Only the neutron count, and therefore the mass, differs.

"All isotopes are radioactive."

Most isotopes of most elements are perfectly stable and last forever. Only some isotopes are radioactive, and "radioisotope" is the specific term for those.

"Isotopes and ions are the same thing."

They are different. Isotopes differ in neutrons (mass); ions differ in electrons (charge). An atom can be both at once.

"The atomic mass is the mass of one atom."

The periodic-table value is an average over an element's natural isotope mix, weighted by abundance — which is why it rarely lands on a whole number.

§ 10 Why it matters today

Isotopes sit quietly behind a surprising amount of modern life. Climate scientists read past temperatures from oxygen-isotope ratios locked in ice cores and shells. Doctors diagnose and treat disease with carefully chosen radioisotopes. Archaeologists and forensic scientists reconstruct histories from isotope signatures. Energy policy turns on the difference between uranium-235 and uranium-238, and the safety of nuclear waste depends on the half-lives of the isotopes it contains — the same isotopes, such as caesium-137, that drove the Chernobyl contamination.

Understanding isotopes means understanding that "an element" is really a family of closely related atoms — and that the tiny differences between family members are precisely what make them so useful.