The physical world is built from fundamental units called atoms, the smallest particles that retain the properties of a chemical element. Every atom is composed of three subatomic particles: the proton, the neutron, and the electron. Protons carry a positive charge, electrons carry a negative charge, and neutrons are electrically neutral. Protons and neutrons reside tightly packed together in the atom’s central nucleus, while electrons orbit this nucleus in a cloud. Every known element, such as oxygen or gold, is defined entirely by the number of protons in its nucleus, which is known as the atomic number.
Defining Isotopes
The identity of an element is set by its proton count, but the number of neutrons can vary, creating different versions of the element known as isotopes. Isotopes are essentially the same element with a slightly different mass. Because the number of protons and electrons remains the same, isotopes of a single element have virtually identical chemical behavior. These variants will react, bond, and behave chemically in almost the same way. However, their physical properties, such as density and stability, can differ.
The Role of Neutrons and Mass
The difference between isotopes lies entirely in the number of neutrons contained within the nucleus. The number of protons is fixed for any given element; for instance, any atom with six protons is automatically a carbon atom. The total number of protons and neutrons in the nucleus determines the atom’s mass number.
For carbon, the most common form is Carbon-12, which contains six protons and six neutrons, summing to a mass number of 12. Another naturally occurring isotope is Carbon-14, which still has six protons but gains two extra neutrons, resulting in a mass number of 14. This difference in the neutron count means that Carbon-14 is measurably heavier than Carbon-12. The additional neutrons in some isotopes can make their nuclei unstable, resulting in a process called radioactive decay.
Practical Uses of Isotopes
The unique physical properties of isotopes, particularly their stability or lack thereof, have led to important practical applications in science and medicine. One recognizable use is Carbon-14 dating, which relies on the unstable nature of the Carbon-14 isotope. Since Carbon-14 decays at a predictable rate, with a half-life of approximately 5,730 years, scientists can measure the remaining amount in ancient organic materials to estimate their age.
Isotopes also play a major role in healthcare for both diagnosis and treatment. For example, the short-lived isotope Technetium-99m is used in about 80% of all nuclear medicine diagnostic scans, helping to image organs like the heart, brain, and bones. Another isotope, Iodine-131, is used to both diagnose and treat thyroid conditions because the thyroid gland naturally absorbs iodine. By delivering a radioactive form of the element, doctors can target and destroy diseased cells.