Defining Isotopes
An isotope refers to atoms of the same chemical element that have an identical number of protons but a different number of neutrons. This variation in neutron count means isotopes of an element have the same atomic number but different mass numbers. For instance, all carbon atoms have six protons, but they can have six, seven, or eight neutrons, leading to different carbon isotopes.
Atomic Structure and Identity
The number of protons defines an element's atomic number and its chemical identity; therefore, isotopes exhibit nearly identical chemical properties as the non-isotopic atoms of that element. The varying number of neutrons primarily affects the atom's atomic mass and, in some cases, its nuclear stability, leading to radioactivity. This difference in mass is significant in physical processes.
A Common Example: Hydrogen Isotopes
A classic example is hydrogen, which has three common isotopes: protium (hydrogen-1) with one proton and no neutrons, deuterium (hydrogen-2) with one proton and one neutron, and tritium (hydrogen-3) with one proton and two neutrons. All three are hydrogen because they each have one proton, but their differing neutron counts give them distinct atomic masses.
Importance and Applications of Isotopes
Isotopes are incredibly important across various fields. Radioactive isotopes (radioisotopes) are used in medicine for diagnostic imaging (like PET scans) and cancer therapy, in carbon dating to determine the age of ancient artifacts, and as tracers in biological and chemical research. Stable isotopes are also used in environmental studies and to understand metabolic pathways.