Atoms – Nuclides and Radioisotopes
The basic building block of all matter
All matter in the world begins with atoms - they form elements like oxygen, hydrogen, and carbon.
An atom consists of protons and neutrons, that make up the nucleus, and electrons that orbit the nucleus. The nucleus carries a positive charge; protons are positively charged, and neutrons don't carry a charge. The electrons, which carry a negative charge, move around the nucleus in clouds (or shells). The negative electrons are attracted to the positive nucleus by an electrical force. This attraction is what keeps the electrons orbiting the nucleus.
It is the number of protons in the nucleus, or the atomic number, that distinguishes each element. The number of protons is unique to each element. For example, there are six protons in a carbon nucleus; therefore, its atomic number is 6 on the periodic table.
Atoms are stable when the number of neutrons and protons in the nucleus are balanced. When there is a significant imbalance between the number of neutrons and protons in a nucleus, the atom becomes unstable and in order to achieve stability, the atom may undergo a transformation or radioactive decay.
Atoms from one or more elements combine to form larger compounds, which are called molecules. A molecule of water, for example, is formed of two atoms of hydrogen combined with one atom of oxygen (H2O).
A nuclide is a specific type of atom characterized by the number of protons and neutrons in the nucleus, which approximates the mass of the nuclide. The number that is sometimes given with the name of the nuclide is called its mass number (the total number of protons and neutrons in the nucleus). For example, carbon-12 is a nuclide of carbon with 6 protons and 6 neutrons.
Understanding the isotope
Nuclides of an element that have the same number of protons but not the same number of neutrons are called isotopes of that element. They are a variant of a basic element. For example, there are three isotopes (or variants) of hydrogen: hydrogen-1 (one proton and no neutrons), hydrogen-2 or deuterium (one proton and one neutron), and hydrogen-3 which is called tritium (one proton and two neutrons). Another example is uranium-235 which has 92 protons and 143 neutrons, and uranium 238 which has 92 protons and 146 neutrons. Both uranium-235 and uranium-238 are isotopes of uranium.
The stable isotope
Many isotopes are stable. They will not undergo radioactive decay and give off radiation. Other isotopes are not stable. An isotope is stable when there is a balance between the number of neutrons and protons. When an isotope is small and stable, it contains close to an equal number of protons and neutrons. Isotopes that are larger and stable have slightly more neutrons than protons.
An example of a stable isotope is carbon-12, which has six protons and six neutrons, for a total mass of 12g.
The unstable isotope
When there is an imbalance between protons and neutrons, usually when the ratio of neutrons to protons is too low, the isotope will want to transform itself into a more stable form – a different atom. When this happens, the atom decreases its mass by emitting alpha particles, beta particles, positrons and/or gamma rays, but some may also gain stability through spontaneous fission or electron capture. It is a spontaneous process that is known as radioactive decay.
There are three main types of radioactive decay:
Alpha decay: Alpha decay occurs when the atom ejects a particle from the nucleus which consists of two neutrons and two protons, causing the atomic number to decrease by two and the mass to decrease by four.
Beta decay: In basic beta decay, a neutron is turned into a proton and an electron is emitted from the nucleus. The atomic number increases by one, but the mass only decreases slightly.
Gamma decay: Gamma decay takes place when there is residual energy in the nucleus following either alpha or beta decay, or after neutron capture in a nuclear reactor. The residual energy is released as a photon of gamma radiation. Gamma decay does not generally affect the mass and atomic number of the radioisotope.
When an isotope disintegrates spontaneously, the excess energy that is emitted is a form of ionizing radiation. In other words, the disintegration gives off radiation and this is called activity. The isotope that changes and emits radiation is called a radioisotope.
These disintegrations are expressed or measured in a unit called the becquerel (Bq). One Bq equals one disintegration per second.
Half-life is the time it takes for a radioisotope to decay to half of its starting activity. The symbol is t½. Each radioisotope has a unique half-life that can range from a fraction of a second to billions of years. The decay is exponential.
For example, iodine-131 takes eight days to reach half of its original activity, while plutonium-239 takes 24,000 years.
If the original source of the radioactivity is known, how long it will take to decay can be predicted. Similarly, the reverse is true. If the half-life is known, you can identify the radioisotope.
Specific activity is the activity (the disintegration rate of an isotope that gives off radiation) per unit of mass of a radioisotope. Half-life and specific activity have an inversely proportional relationship:
- Isotopes that have a longer half-life, like uranium-238, will decay slower and give off less radiation per second by gram.
- Isotopes with a shorter half-life, like cobalt-60, will decay faster and give off more radiation per second by gram.
As an example, if you had 1 gram of cobalt-60 and 1 gram of uranium-238, the gram of Cobalt-60 will give off much more radiation for a much shorter time than the Uranium-238.
The following table compares different radioisotopes with different half-lives and shows the inverse relationship between half-life and specific activity. The last column, ‘Miligrams per MBq’ shows how many miligrams of the isotope are required to yield one megabecquerel of activity. Specific activity is expressed in units of becquerels per gram, and the specific activity of each radioisotope will depend on its half-life and atomic mass.
|Radioisotope||Half-life||Specific activity (Bq/g)||Milligrams (mg) per MBq|
|Polonium-210||138.4 days||166.3E12 (166.3 TBq/gr)||0.000006|
|Cobalt-60||5.3 years||41.9E12(41.9 TBq/gr)||0.0000239|
|Caesium-137||30.04 years||3.2E12 (3.2 TBq/gr)||0.0003109|
|Carbon-14||5700 years||165.7E9 (165.7 GBq/gr)||0.006|
|Technetium-99||214 000 years||624.9E6 (624.9 MBq/gr)||1.6|
|Uranium-235||703 800 000 years||80.0E3 (80.0 kBq/gr)||12,000.5|
|Uranium-238||4 468 000 000 years||12.4E3(12.4 kBq/gr)||80,000.4|
Natural versus artificial sources of radioisotopes
Many radioisotopes are naturally occurring. They originated from the formation of the solar system and from the interaction of cosmic rays with molecules in the atmosphere. Tritium, for example, is formed by cosmic ray interaction with atmospheric molecules. Some radioisotopes that were formed when our solar system was created have half-lives of billions of years and continue to be present in our environment. Uranium and thorium are examples.
Radioisotopes are produced as a by-product of nuclear reactors and by radioisotope generators, such as cyclotrons. Many artificial radioisotopes are used in the fields of nuclear medicine and biochemistry, in the manufacturing industry and in agriculture.
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