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12.1.1 Basic Concept of Isotopes in Environment
Isotopes are atoms of a chemical element whose nucleus has the same atomic
number, Z, but a different atomic mass, A. The word isotope means in the same
place and comes from the fact that the isotopes are in the same place in the periodic
table of the elements. The isotope ratio is the constant ratio of isotopes in chemical
elements in nature and the characteristic size for each atomic species.
The atomic number corresponds to the number of protons in the atom. Therefore
the isotopes of a given element contain the same number of protons. The difference
in atomic masses stems from the difference in the number of neutrons in the atomic
nucleus. In scientific nomenclature, isotopes are denoted by the names of a given
element followed by a dash and the number of nucleons (protons and neutrons) in
the atomic nucleus (helium-3, carbon-12, carbon-14, iron-57, uranium-238). In
symbolic form, the number of nucleons is expressed above the chemical symbol on
the left (
3
He,
12
C,
14
C,
57
Fe,
238
U).
Isotopes of all elements form a set of nuclides. A nuclide is a definition for a
particular atomic nucleus and is defined by the ordinal number Z and the mass number A: AZHh. Strictly speaking, it should be said that an element such as fluorine
consists of one nuclide rather than one isotope. In a neutral atom, the number of
electrons is equal to the number of protons. Therefore, the isotopes of a given element also have the same number of electrons and the same electronic structure. As
the chemical behaviour of an atom is largely determined by its electronic structure,
isotopes have almost identical chemical properties. The main exception is that, due
to higher masses, heavier isotopes tend to react somewhat more slowly than lighter
isotopes. This ‘mass effect’, or isotopic effect, is most pronounced for protium (1H)
and deuterium (common name for 2H), because deuterium has twice the mass of
protium. With heavier elements, the relative difference in mass is much smaller, so
the effect of mass is usually negligible. Although isotopes have almost identical
electronic and chemical properties, their nuclear properties differ dramatically. The
atomic nucleus consists of protons and neutrons held together by a strong nuclear
force. As the protons are positively charged, they repel each other. Neutrons that are
electroneutral make a certain distance between protons, reduce electrostatic repulsion and stabilize the nucleus. For this reason, neutrons are needed to bind two or
more protons to the nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus, for example, although the neutron/proton
ratio at
3
He is 1/2, the neutron/proton ratio at
238
U is >3/2. However, if too many
neutrons are present, the nucleus becomes unstable. As the isotopes of a given element have different numbers of neutrons, they also have different neutron/proton
ratios. This affects their nuclear stability, which results in the fact that some isotopes
are subject to nuclear decay. The decay of these radioactive isotopes (shorter radioisotopes) is an important topic in nuclear physics. By studying the way this decay
takes place, physicists gain insight into the properties of the atomic nucleus. Usually
several isotopes of the same element can be found in nature. Stable isotopes are the
most common, although significant amounts of unstable isotopes with long half- lives
L. Fan et al.
12.1.1 Basic Concept of Isotopes in Environment
Isotopes are atoms of a chemical element whose nucleus has the same atomic
number, Z, but a different atomic mass, A. The word isotope means in the same
place and comes from the fact that the isotopes are in the same place in the periodic
table of the elements. The isotope ratio is the constant ratio of isotopes in chemical
elements in nature and the characteristic size for each atomic species.
The atomic number corresponds to the number of protons in the atom. Therefore
the isotopes of a given element contain the same number of protons. The difference
in atomic masses stems from the difference in the number of neutrons in the atomic
nucleus. In scientific nomenclature, isotopes are denoted by the names of a given
element followed by a dash and the number of nucleons (protons and neutrons) in
the atomic nucleus (helium-3, carbon-12, carbon-14, iron-57, uranium-238). In
symbolic form, the number of nucleons is expressed above the chemical symbol on
the left (
3
He,
12
C,
14
C,
57
Fe,
238
U).
Isotopes of all elements form a set of nuclides. A nuclide is a definition for a
particular atomic nucleus and is defined by the ordinal number Z and the mass number A: AZHh. Strictly speaking, it should be said that an element such as fluorine
consists of one nuclide rather than one isotope. In a neutral atom, the number of
electrons is equal to the number of protons. Therefore, the isotopes of a given element also have the same number of electrons and the same electronic structure. As
the chemical behaviour of an atom is largely determined by its electronic structure,
isotopes have almost identical chemical properties. The main exception is that, due
to higher masses, heavier isotopes tend to react somewhat more slowly than lighter
isotopes. This ‘mass effect’, or isotopic effect, is most pronounced for protium (1H)
and deuterium (common name for 2H), because deuterium has twice the mass of
protium. With heavier elements, the relative difference in mass is much smaller, so
the effect of mass is usually negligible. Although isotopes have almost identical
electronic and chemical properties, their nuclear properties differ dramatically. The
atomic nucleus consists of protons and neutrons held together by a strong nuclear
force. As the protons are positively charged, they repel each other. Neutrons that are
electroneutral make a certain distance between protons, reduce electrostatic repulsion and stabilize the nucleus. For this reason, neutrons are needed to bind two or
more protons to the nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus, for example, although the neutron/proton
ratio at
3
He is 1/2, the neutron/proton ratio at
238
U is >3/2. However, if too many
neutrons are present, the nucleus becomes unstable. As the isotopes of a given element have different numbers of neutrons, they also have different neutron/proton
ratios. This affects their nuclear stability, which results in the fact that some isotopes
are subject to nuclear decay. The decay of these radioactive isotopes (shorter radioisotopes) is an important topic in nuclear physics. By studying the way this decay
takes place, physicists gain insight into the properties of the atomic nucleus. Usually
several isotopes of the same element can be found in nature. Stable isotopes are the
most common, although significant amounts of unstable isotopes with long half- lives
L. Fan et al.