Isotopes
In chemistry, isotopes are defined as variations of a chemical element that share the same atomic number—the number of protons in their nuclei—but possess different mass numbers due to varying numbers of neutrons. For instance, hydrogen has three isotopes:
- Protium
- Deuterium
- Tritium
Each of these isotopes of hydrogen has one proton but differs in neutron count—protium has none, deuterium has one, and tritium has two. This difference in neutron number leads to distinct physical properties without significantly changing their chemical behavior, which is largely determined by the number of protons and electrons.
The significance of isotopes extends beyond hydrogen; many elements exist as mixtures of isotopes. In these cases, the average atomic mass of an element, especially in the presence of isotopes, is calculated based on the relative abundance of each isotope. For example, chlorine naturally occurs as two isotopes, with masses of 35 u and 37 u, present in a 3:1 ratio.
Additionally, isobars—atoms of different elements that have the same mass number—are briefly introduced, illustrating the vast diversity of atomic structures in chemistry. Isotopes have practical applications across various fields, including medicine (for example, certain isotopes are used in cancer treatment or medical imaging) and energy (uranium isotopes in nuclear reactors). Understanding isotopes is fundamental to comprehending atomic theory and the behavior of elements in various chemical contexts.