An isotope, from greek ἴσος (ìsos, “same”) and τόπος (tòpos, “place”), is an atom, of any chemical element, that maintains the same atomic number (Z) but different mass number (A) and therefore different atomic mass (M). The difference in mass number is due to the number of neutrons, present in the nucleus of the atom, with the same atomic number.

Isotopes are denoted in the following way: proper name of the base element followed by the mass number. Depending on the context, it is usual to write them with the mass number superscripted in front of the element abbreviation (e.g. 4H), or with the element abbreviation followed by a hyphen and the mass number (e.g. H-4). In both examples given, the correct way to quote them is “Hydrogen four”.

If two nuclei contain the same number of protons, but a different number of neutrons, the two nuclei will have the same chemical behavior (with minor differences in reaction times and binding energy, collectively called isotopic effects), but will have different physical behaviors, one being heavier than the other.

Same isotopes that differ only in their excited state are called isomers.

  • Atoms of different elements with the same mass number (e.g. 14C and 14N) are called isobars.
  • Atoms of different elements with the same number of neutrons are called isotones (e.g. 56Fe and 58Ni both have 30 neutrons).

Because elements may have several stable isotopes, the average mass number of an element is the atomic weight and is commonly not an integer.

Some isotopes may emit neutrons, protons, and electrons, and attain a more stable atomic configuration (lower level of potential energy); these are radioactive isotopes or radioisotopes. Radioactive decay (carbon-14 decaying to eventually become nitrogen-14) describes the energy loss that occurs when an unstable atom’s nucleus releases radiation.

Isotopes in nature

The elements that can be observed and manipulated on a human scale are not clusters of atoms that are all the same, but contain different isotopes of the same basic element. Chlorine, for example, is a mixture of two isotopes: Cl-35 and Cl-37. Both chlorine atoms have the same number of protons, which by definition is equivalent to the atomic number Z of the element, that is 17, but different mass number A, from which we derive that the first has 18 neutrons while the second has 20.

Always on a very large scale compared to the microscopic world, if we look at a sufficiently large sample of Hydrogen we can see that it is composed of three variants of the basic element: protium, deuterium and tritium. They possess none, one and two neutrons respectively and are the only isotopes to which a proper name has been assigned.

Stable isotopes

Isotopes are divided into stable isotopes (about 252) and non-stable or radioactive isotopes (about 3000 known and another 4000 assumed by theoretical calculations up to element 118). The concept of stability is not clear, in fact there are “almost stable” isotopes. Their stability is due to the fact that, although they are radioactive, they have an extremely long half-life even if compared with the age of the Earth of 4500 Ma. According to recent cosmological theories no isotope is to be considered properly stable.

There are 21 elements (e.g. beryllium-9, fluorine-19, sodium-23, scandium-45, rhodium-103, iodine-127, gold-197 or thorium-232, almost stable) that have in nature only one stable isotope even if in most cases the chemical elements are composed of more than one isotope with a natural isotopic mixture, which in many cases is variable as a result of hydro-geological phenomena (eg: hydrogen and oxygen), radioactive decays (e.g., lead), and human-caused manipulations (e.g., hydrogen/deuterium/tritium and uranium isotopes). Therefore, IUPAC continually updates the recommended average atomic mass values for the various chemical elements taking into account this variability. It is largely conditioned by the geologic site of origin (aquifer, terrestrial, atmospheric), as well as by extraterrestrial or very rarely extrasolar origin (meteorites).

Since the average atomic mass of polyisotopic elements is sometimes variable, its value must have significant digits in appropriate numbers (e.g. 58.933 195(5) u for 59Co which is monoisotopic, 58.6934(2) u for Ni, 207.2(1) u for Pb which is the product of the decay of the natural radioactive chains of 235U, 238U and 232Th).

Stable isotopes are chemical isotopes that may or may not be radioactive, but if they are, they have half-lives too long to measure.

Different isotopes of the same element (both stable and unstable) have almost the same chemical characteristics and therefore behave almost identically in biology (a notable exception are hydrogen isotopes – see Heavy water). Mass differences, due to a difference in the number of neutrons, will result in partial separation of light and heavy isotopes during chemical reactions and during physical processes such as diffusion and evaporation. This process is called isotopic fractionation. For example, the mass difference between the two stable isotopes of hydrogen, 1H (1 proton, no neutron, also known as protium) and 2H (1 proton, 1 neutron, also known as deuterium) is nearly 100%. Therefore, significant fractionation will occur.

Commonly analyzed stable isotopes include oxygen, carbon, nitrogen, hydrogen, and sulfur. These isotopic systems have been under investigation for many years for the purpose of studying isotopic fractionation processes in natural systems because they are relatively simple to measure. Recent advances in mass spectrometry (i.e., plasma mass spectrometry inductively coupled to multiple collectors) now allow the measurement of heavier stable isotopes, such as iron, copper, zinc, molybdenum, etc.

Stable isotopes have been used for many years in botanical and biological investigations of plants, and increasingly ecological and biological studies are discovering the extreme utility of stable isotopes (mostly carbon, nitrogen, and oxygen). Other practitioners have used oxygen isotopes to reconstruct historical atmospheric temperatures, making them important tools for climate research. Measurements of the ratios of one naturally occurring stable isotope to another play an important role in radiometric dating and isotopic geochemistry, and are also useful in determining patterns of rainfall and movement of elements through living organisms, helping to elucidate the dynamics of food webs in ecosystems.

Related keywords

  • Stable nuclide
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