Atoms are the fundamental building blocks of matter, and while all atoms of a given element share the same number of protons, they can differ in the number of neutrons. This variation leads to the formation of isotopes, which are defined as atoms of the same element that have the same atomic number (number of protons) but different mass numbers due to differing numbers of neutrons.
The mass number of an atom is calculated by summing the total number of protons and neutrons in its nucleus. For example, carbon has three isotopes: carbon-12, carbon-13, and carbon-14. All three isotopes have 6 protons, which is the atomic number of carbon. However, they differ in their neutron counts: carbon-12 has 6 neutrons, carbon-13 has 7 neutrons, and carbon-14 has 8 neutrons. Consequently, their mass numbers are 12, 13, and 14, respectively.
To calculate the mass number for each isotope, the formula is:
Mass Number = Number of Protons + Number of Neutrons
For carbon-12: 6 protons + 6 neutrons = 12
For carbon-13: 6 protons + 7 neutrons = 13
For carbon-14: 6 protons + 8 neutrons = 14
While isotopes of an element have the same number of protons and electrons, they can differ in abundance. For instance, carbon-12 is the most abundant isotope, making up about 99% of all carbon atoms, while carbon-13 and carbon-14 are much less common. This abundance affects the atomic mass of the element, which is the weighted average of all isotopes. For carbon, the atomic mass is approximately 12.011, reflecting the predominance of carbon-12 and the minor contributions from the heavier isotopes.
Understanding isotopes is crucial in various fields, including chemistry, physics, and environmental science, as they play significant roles in processes such as radioactive decay and radiometric dating.