Isotope





Isotopes of a chemical element are atoms whose nuclei have the same atomic
number, Z, but different atomic weights, A. The word isotope, meaning at the
same place, comes from the fact that isotopes are located at the same place
on the periodic table.

The atomic number corresponds to the number of protons in an atom. Thus,
isotopes of a particular element contain the same number of protons. The
difference in atomic weights results from differences in the number of
neutrons in the atomic nuclei. In scientific nomenclature, isotopes are
specified by the name of the particular element followed by a hyphen and the
number of nucleons (protons and neutrons) in the atomic nucleus (e.g.,
iron-57, uranium-238, helium-3). In symbolic form, the number of nucleons is
denoted as a superscripted prefix to the chemical symbol (e.g., 57Fe, 238U, 3He).

Collectively, the isotopes of the elements form the set of nuclides. A
nuclide is a particular type of nucleus (characterised by A and Z). Strictly
speaking, we should say that an element such as fluorine consists of one
nuclide rather than that it has one isotope. Similarly, the tables at the
foot of this article are tables of nuclides.

In a neutral atom, the number of electrons equals the number of protons.
Thus, isotopes of a given element also have the same number of electrons and
the same electronic structure. Because the chemical behavior of an atom is
largely determined by its electronic structure, isotopes exhibit nearly
identical chemical behavior. The primary exception is that, due to their
larger masses, heavier isotopes tend to react somewhat more slowly than
lighter isotopes of the same element. This "mass effect" is most pronounced
for hydrogen and deuterium (the common name of 2H), because deuterium has
twice the mass of hydrogen. For heavier elements the relative mass
difference between isotopes is much less, and the mass effect is usually negligible.

Although isotopes exhibit nearly identical electronic and chemical behavior,
their nuclear behavior varies dramatically. Atomic nuclei consist of protons
and neutrons bound together by the strong nuclear force. Because protons are
positively charged, they repel each other. Neutrons, which are electrically
neutral, allow some separation between the positively charged protons,
reducing the electrostatic repulsion and stabilizing the nucleus. For this
reason neutrons are necessary for two or more protons to be bound into a
nucleus. As the number of protons increases, additional neutrons are needed
to form a stable nucleus, for example, although the neutron/proton ratio of
3He is 1/2, the neutron/proton ratio of 238U is >3/2. However, if too many
neutrons are present, the nucleus becomes unstable.

Because isotopes of a given element have different numbers of neutrons they
also have different neutron/proton ratios. This affects the nuclear
stability, with the result that some isotopes are subject to nuclear decay.
The decay of these radioactive isotopes (radioisotopes for short) is an
important topic in nuclear physics. By studying the manner in which this
decay occurs, physicists gain insight into the properties of the atomic nucleus.

In general, several isotopes of each element can be found in nature. Stable
isotopes are by far the most abundundant; however, significant quantities of
long-lived unstable isotopes, such as uranium-238, can also be found. Small
amounts of short-lived radioactive isotopes are also present in nature.
These arise as products of the decay of larger long-lived radioactive
nuclei. The atomic mass for an element in the periodic table is the average
of the natural abundance of the isotopes of that element.

This part probably needs correction and refinement: The distribution of the
various isotopes on earth is ultimately the result of the initial product
distribution of atoms generated in stars and supernovae during the creation
of the solar system and the subsequent decay patterns of the heavier nuclei.
Someone with some knowledge on the subject could also add more about
creative processes in the universe.

Applications of isotopes

Several applications exist that capitilize on properties of the various
isotopes of a given element. One of the most common applications is as a
tracer or marker in a technique called isotopic labeling. Normally, atoms of
a given element are indistinguishable from each other. However, by using
isotopes of different masses, they can be distinguished by mass spectroscopy
or infrared spectroscopy, which is mass sensitive because heavier atoms
vibrate at different frequencies than lighter atoms.

An example of the use of isotopic labeling is the study of phenol (C6H5OH)
in water. Upon adding phenol to deuterated water (water containing D2O in
addition to the usual H2O), researchers observed the substition of deuterium
for the hydrogen in the hydroxyl group (C6H5OD), indicating that phenol
readily undergoes hydrogen-exchange reactions with water. Only the hydroxyl
group was affected, indicating that the other 5 hydrogen atoms did not
participate in these exchange reactions.

In addition to isotopic labeling, several forms of spectroscopy rely on the
unique nuclear properties of specific isotopes. For example, nuclear
magnetic resonance (NMR) spectroscopy can be used only for isotopes with a
nonzero nuclear spin. The most common isotopes used with NMR spectroscipy
are 1H, 2D, 13C, and 31P. Mossbauer spectroscopy also relies on the nuclear
transitions of specific isotopes, such as 57Fe.

Radioactive isotopes also have important uses. Nuclear power and nuclear
weapons development require relatively large quantities of specific
isotopes. The process of isotope separation represents a significant
technological challenge.

Radioisotopes are also frequently used in medicine, biochemistry, and
chemistry as tracers. Small quantities of the radioisotopes can be readily
detected due to characteristic emissions by the decaying nuclei.