A radioactive isotope is a version of a chemical element whose nucleus is unstable and releases energy to reach a more stable state. Every element on the periodic table can have multiple isotopes, meaning atoms with the same number of protons but different numbers of neutrons. When that combination of protons and neutrons creates an imbalanced nucleus, the atom becomes radioactive and sheds particles or energy over time in a process called decay.
Why Some Isotopes Are Unstable
The stability of any atom’s nucleus comes down to the ratio of neutrons to protons packed inside it. For lighter elements, a roughly 1:1 ratio keeps things stable. As elements get heavier, they need proportionally more neutrons to offset the electrical repulsion between all those positively charged protons, and the stable ratio climbs to about 1.5 neutrons for every proton in the heaviest elements.
Isotopes that fall outside this narrow band of stability, whether they have too many neutrons or too few, are radioactive. Their nuclei carry excess energy and will eventually release it by emitting particles or radiation. Above an atomic mass of 208 (the heaviest stable isotope, a form of lead), no isotope of any element is stable at all. Every atom heavier than that is radioactive to some degree.
Three Types of Radioactive Decay
When a radioactive isotope decays, it can release energy in three main forms, each with very different properties.
- Alpha particles are relatively heavy clusters of two protons and two neutrons ejected from the nucleus. They carry a lot of energy but burn through it quickly, traveling only a few centimeters in air. A sheet of paper can stop them. They come from the decay of the heaviest elements like uranium, radium, and polonium.
- Beta particles are fast-moving electrons fired out of the nucleus. They travel farther than alpha particles and can penetrate skin, but a layer of clothing or a thin sheet of aluminum is enough to block them.
- Gamma rays are pure energy with no mass at all, similar to light but far more powerful. They often accompany alpha or beta decay and can pass through the human body entirely. Stopping them requires several inches of lead or a few feet of concrete.
These differences matter because they determine how dangerous an isotope is in different situations. An alpha emitter sitting on a table poses little external threat, but inhaling or swallowing it can cause serious internal damage. A gamma emitter, on the other hand, can be hazardous from a distance.
Half-Life: How Long Decay Takes
Every radioactive isotope decays at its own fixed rate, measured by its half-life. This is the time it takes for exactly half the atoms in a sample to decay into other atoms. After one half-life, half the original radioactive material remains. After two, a quarter. After three, an eighth, and so on.
Half-lives vary enormously. Iodine-131, used in thyroid treatments, has a half-life of about eight days, meaning its radioactivity is essentially gone within a few weeks. Carbon-14, the isotope used in archaeological dating, has a half-life of 5,730 years. Uranium-238, found in rocks and soil, has a half-life of about 4.5 billion years, roughly the age of Earth itself. An isotope with a short half-life is intensely radioactive but fades quickly. One with a long half-life decays so slowly that each individual atom rarely emits anything in a human lifetime.
How Radioactive Isotopes Are Made
Some radioactive isotopes exist naturally. Uranium, thorium, and radium have been present in Earth’s crust since the planet formed. Carbon-14 is continuously created in the upper atmosphere when cosmic rays from space collide with nitrogen atoms. But many of the isotopes used in medicine and industry are manufactured deliberately using one of three methods.
Nuclear reactors produce the majority of artificial radioisotopes, either by bombarding stable materials with neutrons or by extracting useful isotopes from the fission products the reactor generates. Particle accelerators called cyclotrons slam protons into target materials at high speed to create specific short-lived isotopes, particularly those used in PET scans. A third approach uses portable generators that contain a longer-lived “parent” isotope, which gradually decays into the shorter-lived isotope needed at the hospital or lab.
Medical Imaging and Diagnosis
Radioactive isotopes are central to modern diagnostic medicine. The basic idea is simple: attach a tiny amount of a radioactive substance to a molecule that naturally travels to a specific organ, inject it into the patient, and then use a camera that detects the emitted radiation to build an image of what’s happening inside the body.
Technetium-99m is the most widely used radioisotope in diagnostic medicine. Different chemical preparations of it target different organs, making it useful for imaging the brain, bones, liver, spleen, and kidneys, as well as for studying blood flow. Fluorine-18 is the workhorse of PET scanning, where it’s attached to a sugar molecule to reveal areas of high metabolic activity, a hallmark of many cancers. Iodine-123 is used to evaluate thyroid disorders and brain function. Thallium-201 helps cardiologists assess heart blood flow, and xenon-133, an inhaled gas, lets doctors image lung ventilation.
The isotopes chosen for imaging typically have short half-lives, so the radioactivity in a patient’s body fades within hours or days.
Cancer Treatment
Beyond diagnosis, radioactive isotopes can also destroy diseased tissue. Iodine-131 is the best-known example. The thyroid gland naturally absorbs iodine from the bloodstream, so when a patient with thyroid cancer swallows a capsule or liquid containing iodine-131, the radioactive iodine concentrates in the thyroid and delivers a targeted dose of radiation that kills cancer cells while largely sparing the rest of the body.
Other treatment approaches place tiny radioactive sources directly inside or next to a tumor, delivering high doses of radiation over a small area. The principle is always the same: get the radiation as close to the target as possible and choose an isotope whose decay characteristics match the clinical goal.
Dating Ancient Materials
Radiocarbon dating, developed at the University of Chicago in the late 1940s, relies on carbon-14’s predictable decay to determine the age of organic materials. While alive, plants and animals continuously absorb carbon-14 from the environment. Once they die, the intake stops and the carbon-14 already present begins decaying at a known rate. By measuring how much carbon-14 remains in a sample compared to what would be expected in a living organism, scientists can calculate when the organism died.
With a half-life of 5,730 years, carbon-14 dating works reliably on materials up to about 60,000 years old. Beyond that, too little carbon-14 remains for accurate measurement. For older geological samples, scientists use isotopes with much longer half-lives, such as uranium and potassium, which can date rocks billions of years old.
Measuring Radioactivity
Radioactivity is measured in units that describe how many atoms in a sample are decaying per second. The international standard unit is the becquerel, equal to one decay event per second. The older unit still common in the United States is the curie, which equals 37 billion decays per second, originally based on the activity of one gram of radium.
Exposure to radiation and its biological effects are measured separately, in units called sieverts (or millisieverts, abbreviated mSv). International guidelines set the occupational exposure limit for radiation workers at 20 mSv per year averaged over five years, with no single year exceeding 50 mSv. For context, a single chest X-ray delivers roughly 0.02 mSv, and the average person absorbs about 2 to 3 mSv per year from natural background radiation sources like radon gas, cosmic rays, and radioactive minerals in soil.