Atoms are the fundamental building blocks of all matter, and they are composed of three main subatomic particles: protons, neutrons, and electrons. The atom’s core, known as the nucleus, contains the positively charged protons and the neutral neutrons, which account for almost all of the atom’s mass. Negatively charged electrons orbit this dense nucleus in a cloud-like region, balancing the positive charge of the protons to keep the atom electrically neutral.
The identity of an element is defined by the number of protons in its nucleus, known as the atomic number. For instance, every carbon atom possesses six protons, and every oxygen atom contains eight protons. While atoms of the same element are generally considered identical, a variation exists that changes their physical properties without altering their chemical behavior. This difference in atomic structure defines an isotope.
The Defining Difference: Neutron Count
Isotopes are atoms of the same element, meaning they share an identical number of protons (the atomic number). The factor that distinguishes one isotope from another is the count of neutrons found in the atomic nucleus. Because neutrons have no electrical charge, changing their number does not affect the atom’s overall charge or its electron configuration, which governs chemical interactions.
All isotopes of a specific element react in the same way during chemical reactions because they have the same number of valence electrons. The difference in neutron count, however, makes one version of the atom physically heavier than another. For example, hydrogen has three naturally occurring isotopes: Protium (one proton, zero neutrons), Deuterium (one proton, one neutron), and Tritium (one proton, two neutrons).
How Mass Number Changes
The differing neutron count directly results in a change to the Mass Number, a whole-number value representing the total count of protons and neutrons in an atom’s nucleus. The Mass Number (A) is calculated by adding the number of protons (Z) and the number of neutrons (N), expressed as A = Z + N. Adding an extra neutron, as seen when comparing Carbon-12 (six protons, six neutrons) to Carbon-13 (six protons, seven neutrons), increases the Mass Number by one unit.
Isotopes are commonly named by stating the element’s name followed by their specific Mass Number, such as Carbon-14 or Uranium-235. The Mass Number is distinct from the Atomic Weight, which is the decimal number often found on the periodic table. The Atomic Weight represents a weighted average of the Mass Numbers of all the naturally occurring isotopes of that element, factoring in their relative abundance on Earth.
Stability and Radioactive Decay
The particular combination of protons and neutrons within an isotope’s nucleus determines its stability. A nucleus is considered stable when the attractive strong nuclear force is in balance with the repulsive forces between the positively charged protons. Most isotopes found in nature are stable, meaning their nuclei will persist indefinitely without change.
However, some neutron-to-proton ratios create an imbalance, resulting in an unstable nucleus, which defines a radioactive isotope, or radioisotope. To achieve a more stable configuration, this unstable nucleus must spontaneously break down in a process called radioactive decay. This decay involves the nucleus emitting excess energy and particles, which are experienced as radiation.
The emission of radiation allows radioisotopes to be used in various real-world applications. For instance, the predictable decay rate of Carbon-14 is the mechanism behind radiocarbon dating, allowing scientists to determine the age of ancient organic materials. In medicine, certain radioisotopes, such as Technetium-99m, are used as diagnostic tracers because they can be injected into the body and their emitted radiation can be tracked to image organs or identify disease.