What are Stable Isotopes?

 

 

The nucleus of each atom contains protons and neutrons. While the number of protons defines the element (e.g., hydrogen, carbon, etc.) and the sum of the protons and neutrons gives the atomic mass, the number of neutrons defines the isotope of that element. For example, most carbon (≈ 99 %) has 6 protons and 6 neutrons and is written as 12C to reflect its atomic mass. However, about 1 % of the carbon in the Earth’s biosphere has 6 protons and 7 neutrons (13C) forming the heavy stable isotope of this important element. Stable isotopes do not decay into other elements. In contrast, radioactive isotopes (e.g., 14C) are unstable and will decay into other elements.

The less abundant stable isotope(s) of an element have one or two additional neutrons than protons, and thus are heavier than the more common stable isotope for those elements. Both heavy and light stable isotopes participate freely in chemical reactions and in biological and geochemical processes, but the rate at which heavy and light stable isotopes react during physical or chemical reactions differs. The chemical bonds and attractive forces of atoms with heavy stable isotopes are stronger than those in the more common, lighter isotopes of an element. As a result, the heavier isotopes react more slowly than the lighter isotopes leading to isotopic separation or fractionation between reactant and product in both physical and biological reactions. Fractionation of the heavy and light stable isotopes is important because it a) produces variation in the stable isotope ratio of different element pools and b) establishes an isotope signal that can indicate the existence or magnitude of key processes involved with elemental cycling.