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Isotope


What is an isotope? Definition and explanation by means of an example:

Atoms with identical atomic number (proton number), but different number of neutrons, are called isotope designated.
As an example for understanding, an example: The two uranium isotopes uranium-235 and uranium-238 each have 92 protons in their nucleus. The number of protons in the nucleus always determines the chemical element of the atom. In the periodic table, the number of protons of any element can easily be read off the atomic number. This is identical to the atomic number, ie the number of protons.
235 or 238 indicates the mass number or number of nucleons. This number tells you how many protons and neutrons are in the nucleus. Uranium 235 has 143 neutrons (sample calculation: 92 + 143 = 235), uranium 238 has 146 neutrons (sample calculation: 92 + 146 = 238).
The whole again to the overview:
Uranium-235 92 protons + 143 neutrons = 235 (mass number)
Uranium-238 92 protons + 146 neutrons = 238 (mass number)
What should be apparent from this example: Isotopes from an element series (in this case uranium) are always a single element. The distinguishing feature is the different numbers of neutrons! So if you know the isotope, you can easily calculate the corresponding number of neutrons:
Mass number = atomic number + neutron number
Of course, not only the chemical element has uranium isotopes. In total, over 3000 different isotopes are known. Each element has a greater or lesser number of isotopes.

Interesting facts about isotopes:

The term isotope is derived from the Greek 'isos' for 'equal' and 'topos' for 'place'. In the periodic table no distinction between the isotopes of the respective elements takes place. They are therefore at the same place in the periodic table, that is, their root element.
Basically, one can categorize isotopes on two levels. These are the one hand stable vs. unstable Isotopes. Stable isotopes basically have no radioactive decay and thus remain in their 'form'. In contrast, unstable isotopes break down into other isotopes or elements, depending on the respective half-lives. The vast majority of the known isotopes is unstable, z.T. the half-lives are only a few seconds.
In addition, isotopes may still be after natural vs. artificial Categorized isotopes. The latter are practically not observable in nature, that is to say purely theoretical considerations or have been deliberately brought about in the laboratory by scientists. The human-relevant isotopes (e.g. 12Carbon) are almost invariably non-radioactive and therefore stable. Otherwise, survival would hardly be possible if the tiny components of our body cells decay into other isotopes with the release of radioactive radiation.

Isotopes of hydrogen:


Finally, a brief look at the three most important isotopes of hydrogen.
Protium: 99.9% of global hydrogen is in protium
deuterium: also called "heavy hydrogen"
tritium: Also referred to as "superheavy hydrogen", radioactive
The figure to the right shows the notation in the case of deuterium. Top left is the mass number (number of protons + neutrons), bottom left the atomic number of the element, in this case the 1 for hydrogen (H).