1. Introduction
I.
Isotopes:
An isotope is one of two or more species of atoms of a chemical element with the same atomic number (same number or protons in the nucleus) and position in the periodic table and nearly identical chemical behavior but with different atomic masses and physical properties. Every chemical element has one or more isotopes.
An atom is first identified and labeled according to the number of protons in its nucleus. This atomic number is ordinarily given the symbol Z. The great importance of the atomic number derives from the observation that all atoms with the same atomic number have nearly, if not precisely, identical chemical properties. A large collection of atoms with the same atomic number constitutes a sample of an element. A bar of pure uranium, for instance, would consist entirely of atoms with atomic number 92. The periodic table of the elements assigns one place to every atomic number, and each of these places is labeled with the common name of the element, as, for example, calcium, radon, or uranium.
I.
Stable and Nonstable Isotopes:
Isotopes utilized in nuclear medicine fall into two broad
categories: Stable and Unstable. Stable isotopes do not undergo
radioactive decay.
A "stable isotope" is any of two or more forms of
an element whose nuclei contains the same number of protons and electrons, but
a different number of neutrons. Stable isotopes remain unchanged indefinitely,
but "unstable" (radioactive) isotopes undergo spontaneous
disintegration.
1.
Radioactivity
and Radioisotopes
I.
Radioactivity:
Radioactive decay occurs in unstable atomic
nuclei – that is, ones that don’t have enough binding energy to hold the
nucleus together due to an excess of either protons or neutrons.
It comes in three main types – named alpha, beta
and gamma for the first three letters of the Greek alphabet.
a)
Alpha
decay:
An alpha particle is identical to a helium
nucleus, being made up of two protons and two neutrons bound together. It
initially escapes from the nucleus of its parent atom, invariably one of the
heaviest elements, by quantum mechanical processes and is repelled further from
it by electromagnetism, as both the alpha particle and the nucleus are
positively charged.
The process changes the original atom from which
the alpha particle is emitted into a different element. Its mass number
decreases by four and its atomic number by two. For example, uranium-238 will
decay to thorium-234. Sometimes one of these daughter nuclides will also be
radioactive, usually decaying further by one of the other processes described
below.
b)
Beta
decay:
Beta decay itself comes in two kinds: β+ and
β-. β- Emission occurs by the transformation of one
of the nucleus’s neutrons into a proton, an electron and an antineutrino.
Byproducts of fission from nuclear reactors often undergo β- decay,
as they are likely to have an excess of neutrons. β+ decays is
a similar process, but involves a proton changing into a neutron, a positron
and a neutrino.
c)
Gamma
decay:
After a nucleus undergoes alpha or beta decay,
it is often left in an excited state with excess energy. Just as an electron
can move to a lower energy state by emitting a photon somewhere in the
ultraviolet to infrared range, an atomic nucleus loses energy by emitting a
gamma ray.
Gamma radiation is the most penetrating of the
three, and will travel through several centimeters of lead. Beta particles will
be absorbed by a few millimeters of aluminum, while alpha particles will be
stopped in their tracks be a few centimeters of air, or a sheet of paper – although
this type of radiation does the most damage to materials it hit.
I.
Radioisotopes:
Radioisotopes are the unstable form
of an element that emit radiation to transform into a more stable form.
Radiation is easily traceable and can cause changes in the substance it falls
upon. These special attributes make radioisotopes useful in medicine, industry
and other areas.
Of
the 118 elements listed in the periodic table, only 94 occur naturally. While
there are 254 stable isotopes, more than 3,000 radioisotopes are known, of
which only about 84 are seen in nature. The radiation emitted is energetic and
can be of different types, most often alpha (a), beta (b) and gamma (g).
Most
radioisotopes are artificially produced in research
reactors and accelerators by exposing a target material to
“intense particles,” such as neutrons or protons, followed by different
chemical processes to bring them into the required chemical form.
Radioisotopes
are an effective tool used in radiopharmaceutical sciences, industrial applications,
environmental tracing and biological studies. Aside from research reactors and
accelerators, they are also obtained from radioisotope generators.
1.
History of Radioisotopes:
I.
Discovery:
Evidence
for the existence of isotopes emerged from two independent lines of research,
the first being the study of radioactivity. By 1910, it had become clear that
certain processes associated with radioactivity, discovered some years before
by Henri Becquerel, could transform one element into another. In particular,
ores of the radioactive elements uranium and thorium had been found to contain
small quantities of several radioactive substances never before observed. These
substances were thought to be elements and accordingly received special names.
Uranium ores, for example, yielded "ionium," and thorium ores gave
"mesothorium." Painstaking work completed soon afterward revealed,
however, that ionium, once mixed with ordinary thorium, could no longer be retrieved
by chemical means alone. Similarly, mesothorium was shown to be chemically
indistinguishable from radium. As chemists used the criterion of chemical
indistinguishability as part of the definition of an element, they were forced
to conclude that ionium and mesothorium were not new elements after all, but
rather new forms of old ones. Generalizing from these and other data, Frederick
Soddy in 1910 observed that "elements of different atomic weights may
possess identical (chemical) properties" and so belong in the same place
in the periodic table. With considerable prescience, he extended the scope of
his conclusion to include not only radioactive species but stable elements as
well. A few years later, Soddy published a comparison of the atomic weights of
the stable element lead as measured in ores rich in uranium and thorium,
respectively. He expected a difference because uranium and thorium decay into
different isotopes of lead. The lead from the uranium-rich ore had an average
atomic weight of 206.08 compared to 207.69 for the lead from the thorium-rich
ore, thus verifying Soddy's conclusion.
The unambiguous confirmation of isotopes in stable elements not associated directly with either uranium or thorium followed a few years later with the development of the mass spectrograph by Francis William Aston. His work grew out of the study of positive rays (sometimes called canal rays), first discovered in 1886 by Eugen Goldstein and soon thereafter recognized as beams of positive ions. As a student in the laboratory of J.J. Thomson, Aston had learned that the gaseous element neon produced two positive rays. The ions in the heavier ray had masses about 2 units, or 10 percent, greater than the ions in the lighter ray. To prove that the lighter neon had a mass very close to 20 and that the heavier ray was indeed neon and not a spurious signal of some kind, Aston had to construct an instrument that was considerably more precise than any other of the time. By 1919, he had done so and convincingly argued for the existence of neon-20 and neon-22. Information from his and other laboratories accumulated rapidly in the ensuing years and by 1935 the principal isotopes and their relative proportions were known for all but a handful of elements.
The unambiguous confirmation of isotopes in stable elements not associated directly with either uranium or thorium followed a few years later with the development of the mass spectrograph by Francis William Aston. His work grew out of the study of positive rays (sometimes called canal rays), first discovered in 1886 by Eugen Goldstein and soon thereafter recognized as beams of positive ions. As a student in the laboratory of J.J. Thomson, Aston had learned that the gaseous element neon produced two positive rays. The ions in the heavier ray had masses about 2 units, or 10 percent, greater than the ions in the lighter ray. To prove that the lighter neon had a mass very close to 20 and that the heavier ray was indeed neon and not a spurious signal of some kind, Aston had to construct an instrument that was considerably more precise than any other of the time. By 1919, he had done so and convincingly argued for the existence of neon-20 and neon-22. Information from his and other laboratories accumulated rapidly in the ensuing years and by 1935 the principal isotopes and their relative proportions were known for all but a handful of elements.
II.
Radioactivity Units:
The activity of a radioactive sample is measured as
the number of radioactive emissions in a second. The activity of an isotope
defines how quickly (or slowly) it emits radiation. The unit for measuring
disintegrations is called the curie (Ci) for the scientist Marie Curie. The SI
unit for measuring disintegrations is called the Becquerel (Bq) after Henri
Becquerel.
I.
Half lives:
Every radioactive
isotope emits radiation, at a different rate. Unstable isotopes emit radiation
more rapidly. The rate of decay is measured as half-life, the time it takes for
one-half (50%) of the atoms in a sample to decay. Natural radioisotopes have long half-lives.
Radioisotopes used in medicine have short half-lives; radioactivity is
eliminated quickly.
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