Isotope Meaning - Overview, Structure, Properties & Uses

Isotope Meaning - Overview, Structure, Properties & Uses

What is an isotope/Define isotopes with example/What is an isotope in chemistry/What are isotopes give example/What is meant by isotopes/What are isotopes class 9

Isotope definition chemistry/Isotope definition simple: Isotope meaning is the term isotope refers to two or more types of atoms that share similar atomic numbers (number of protons in their nuclei) and position in the periodic table (so that they belong to the same chemical element), but differ in nucleon numbers (mass numbers) because of isotopes they have different numbers of neutrons. There is no such thing as an identical isotope of an element because of isotopes they have different physical and chemical properties. A single element's different isotopes occupy identical positions on the periodic table, an inference derived from the Greek roots isos and tope meaning "the same place.". A suggestion from Margaret Todd to chemist Frederick Soddy in 1913 led to its coining.

The atomic number is the number of protons during the nuclear process and is equal to an atom's number of electrons during the neutral (non-ionized) phase. It is only the element's atomic number that identifies it, not the type of isotope. Atoms of a given element may vary widely in their neutron count. A nucleus' mass number determines how many nucleons it has (even if they are protons or neutrons), and each isotope of an element has a different mass number. Isotope examples of isotopes of carbon are carbon-12, carbon-13, and carbon-14, which are chemical substances that have mass numbers 12, 13, and 14. Carbon has an atomic number of 6, implying that it contains six protons, so its neutron numbers are six, seven, and eight.

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A brief history

Isotopes of radioactive material

Radiochemist Frederick Soddy first proposed the existence of isotopes in 1913, based on studies of radioactive decay chains that indicated about 40 different species (radioactive elements) between uranium and lead, though the periodic table only classified 11 elements between them.

These new radioelements had been separated chemically several times but without success. Therefore, Soddy demonstrated in 1910 that radium (226Ra, the longest-lived isotope), mesothelium (224Ra), and thorium X (224Ra) cannot be separated. Attempts to properly place the radioelements in the periodic table led Soddy and Kazimierz Faja’s to independently propose the alpha decay and beta decay displacement laws in 1913, indicating that alpha decay produced two alphas and betas in the periodic table. He recognized that the emission of an alpha particle followed by two beta particles produced an element whose chemical properties were identical to the initial element but had a lighter mass than it and different radioactivity.

Nature's occurrence

A mononucleotide (mononuclidic element) or, in the case of naturally occurring isotopes, more than one naturally occurring isotope is composed of a given element. There are two categories of unstable (radioactive) isotopes: new or ancient. Originating from stellar-mass fusion or another type of nucleosynthesis such as cosmic ray spallation, they have been preserved to the present because of isotopes of their slow decay rates (e.g., uranium-238 and potassium-40). A radioactive primordial isotope decays to a radioactive radiogenic isotope daughter (e.g., uranium to radium) or to cosmogenic nuclides under cosmic ray bombardment (e.g., tritium, carbon-14). It is also possible to naturally synthesize a few nucleonic nuclides by another natural nuclear reaction, such as when neutrons are absorbed by another atom during natural nuclear fission.

Isotopes atomic mass

Isotopes (neutralizes) have mass numbers that are determined primarily by their atomic mass (Mr) (i.e., number of nucleons in their nucleus). Some corrections are required due to the physical behaviour of the nucleus (see mass defect), the mass differences between the proton and neutron, and the mass of the electrons associated with the atom. These differences arise because of isotopes of the electron: nucleon ratio. There is no dimension to the mass number. A difference between an atomic mass and an average mass is measured using the atomic mass unit based on the mass of carbon-12. A unified atomic mass unit is defined by "U" (for the unified atomic mass unit) or by "Da" (for Dalton).An element's atomic mass is determined by the atomic masses of its naturally occurring isotopes

Isotope and isobar comparison

A different chemical species, an isobar, contains the same nucleons but different atomic numbers. There is a specific atomic number, number of protons, number of electrons, and number of neutrons for each group of isobars. In spite of this, the number of nucleons will always be the same. The sum of the protons and neutrons in an isobar group will always be the same, because of isotopes of this. As an example, the isobars calcium-40, potassium-40, sulfur-40, and chlorine-40 are isotope examples. As a short summary, isotopes have the same mass numbers but differ in their atomic numbers, while isobars have the same mass numbers but differ in their atomic numbers.

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