Isotopes are variants of a particular chemical element that differ in the number of neutrons. All isotopes of certain elements have the same number of protons in each atom. The term isotope is formed from the Greek root isos (???? "sama") and topos (????? "place"), meaning "the same place"; thus, the meaning behind the name is that different isotopes of one element occupy the same position on the periodic table.
The number of protons in the atomic nucleus is called the atomic number and is equal to the number of electrons in neutral (unionized) atoms. Each atomic number identifies a particular element, but not an isotope; an atom of a particular element may have a wide range in its neutron count. The number of nucleons (both protons and neutrons) in the nucleus is the number of atomic masses, and each isotope of a particular element has a different mass number.
For example, carbon-12, carbon-13 and carbon-14 are three isotopes of carbon element with mass numbers 12, 13 and 14 respectively. The number of carbon atoms is 6, which means that each carbon atom has 6 protons, so the number of neutrons of this isotope are 6, 7 and 8 respectively.
Video Isotope
Isotop vs. nuklida
Nuclides are the atomic species with the number of protons and specific neutrons in the nucleus, eg carbon-13 with 6 protons and 7 neutrons. The concept (referring to individual nuclear species) emphasizes the nuclear nature of chemical properties, whereas the concept isotope (grouping all the atoms of each element) emphasizes chemistry over nuclear. The number of neutrons has a major effect on nuclear properties, but their effects on chemical properties are negligible for most elements. Even in the case of light elements where the ratio of the number of neutrons to atomic numbers varies between isotopes it usually has only a small effect, although it is important in some circumstances (for hydrogen, the lightest element, the isotope effect is large enough to affect biology). The term
, 13
6 C
, 14
6 C
is isotope (nu klida with the same atomic number but different mass number), but 40
line-height: inherit; vertical-align: baseline "> 40
19 K
, 40
20 Ca
isobar (nuklida with the same mass number).But, because isotope is a longer term, it is better known than nuclides , and is sometimes still used in contexts where nuclides may be more appropriate, such as nuclear technology and nuclear medicine. Maps Isotope
Definition of isotope
Isotopes are atoms of the same element that have the same number of protons (atomic number), but the number of different neutrons. They have the same chemical properties due to the same electronic configuration but different physical properties.
Notation
An isotope and/or nuclide is determined by the name of a particular element (this denotes the atomic number) followed by dashes and mass numbers (eg helium-3, helium-4, carbon-12, carbon-14, uranium-235 and uranium-239). When a chemical symbol is used, e.g. "C" for carbon, standard notation (now known as "AZE notation" because A is the mass number, Z atomic number, and E for element) is to indicate the mass number (the number of nucleons) with the superscript at the top left of the chemical symbol and to show the atomic number with the subscript on the bottom left (eg 3
2 He > , 4
12
6 C
, 14
6 C
235
92 U
, and 239
92 U
). Since atomic numbers are given by element symbols, it is common to declare only mass numbers in superscript and leave the atomic number of the subscript (eg 3
He
, 4 < br> He
12
C
, 14 C
235 U , and 239 U ) The letters m are sometimes added after the mass number to denote a nuclear isomer, a metastable nuclear state or an energetic energy (as opposed to the lowest energy state), for example 180m
73 Ta
(tantalum-180m).
The general pronunciation of AZE notation differs from how it was written: 4
2 He
is generally pronounced as helium instead of four-two-helium, and 235
92 U
as uranium two three five (American English) or uranium-two-three-five (English) rather than 235-92-uranium.
Radioactive, primordial, and stable isotopes
Some isotopes/nuclides are radioactive, and are therefore referred to as radioisotopes or radionuclides, while others have never been observed radioactively decomposed and are called stable isotopes or stable nuclides. For example, 14 C > is a radioactive form of carbon, while 12
C
and 13 C
is a stable isotope. There are about 339 nuclides that form naturally on Earth, where 286 are primordial nuclides, meaning that they have existed since the formation of the Solar System.
Primordial nuclides include 32 nuclides with very long half-lives (over 100 million years) and 253 formally considered "stable nuclides", because they have not been observed to decompose. In many cases, for obvious reasons, if an element has a stable isotope, the isotope dominates in the abundance of elements found on Earth and in the Solar System. However, in the case of three elements (tellurium, indium, and rhenium) the most abundant isotope found in nature is actually one (or two) very long-lived radioisotopes of the element, although these elements have one or more stable isotopes.
The theory predicts that many seemingly "stable" isotopes/nuclides are radioactive, with very long half-lives (reducing the possibility of proton decay, which will make all nuclides ultimately unstable). Of the 253 nuclides never seen to rot, only 90 of these (all of the first 40 elements) are theoretically stable for all known forms of decay. The 41 element (niobium) is theoretically unstable through spontaneous fission, but this has never been detected. Many other stable nucleons in theory are energetically susceptible to other known forms of decomposition, such as alpha decay or double beta decay, but no decay products have been observed, and this isotope is said to be "observably stable". The estimated half-life for these nuclides often exceeds the estimated age of the universe, and in fact there are also 27 known radionuclides (see primordial nuclides) with a half-life longer than the life of the universe.
Adding radioactive nuclides that have been made artificially, there are 3,339 nuclides that are known today. This includes 905 nuclides that are stable or have a half-life of more than 60 minutes. See the list of nuclides for details.
History
Radioactive isotope
The existence of the first isotope was proposed in 1913 by radiochemist Frederick Soddy, based on studies of a radioactive decay chain that indicated about 40 different species called radioelements (ie, radioactive elements) between uranium and lead, although the periodic table Only allowed for 11 elements of uranium to lead.
Some attempts to separate these new radioelements have chemically failed. For example, Soddy has pointed out in 1910 that the mesothorium (then shown as 228 Ra), radium ( 226 Ra, the longest-lived isotope), and thorium X (< > 224 Ra) can not be separated. Attempts to place radioelements in the periodic table led Soddy and Kazimierz Fajans independently to propose their radioactive switch laws in 1913, stating that alpha decay produces a two-place element to the left in the periodic table, while beta decay produces a one-way element emission to the right. Soddy admits that emission of alpha particles followed by two beta particles causes the formation of a chemical element identical to the initial element but with a mass of four lighter units and with different radioactive properties.
Soddy proposes that some types of atoms (different in radioactive properties) can occupy the same place in the table. For example, the alpha decay of uranium-235 forms thorium-231, while the beta-acting-230 beta forms thorium-230. The term "isotope", Greek for "in the same place", was suggested to Soddy by Margaret Todd, a doctor and a Scottish family friend, during a conversation in which he explained his ideas to him. He won the 1921 Nobel Prize in Chemistry partly because of his work on isotopes.
In 1914 T. W. Richards found a variation between the weight of lead atoms from various mineral sources, caused by variations in the isotopic composition due to different radioactive origins.
Stable isotope
The first evidence for several isotopes of a stable (non-radioactive) element was discovered by J. J. Thomson in 1913 as part of an exploration into the composition of the canal rays (positive ions). Thomson channeled the flow of neon ions through magnetic and electric fields and measured their deflection by placing photographic plates in its path. Each stream creates a radiant glow on the plate at that point struck. Thomson observed two separate patches of light on the photographic plate (see figure), which showed two different deflection parabolas. Thomson has finally concluded that some atoms in neon gas have a higher mass than others.
F. W. Aston later found several stable isotopes for various elements using mass spectrograph. In 1919 Aston studied fluorescents with sufficient resolution to show that two isotopic masses are very close to integers 20 and 22, and that they are not equal to the known molar mass (20.2) of neon gas. This is an example of the rule of the Aston integer for the mass of the isotope, which states that the massive deviation of the elemental molar mass of the integers is primarily due to the fact that the element is a mixture of isotopes. The same aston shows that the chlorine molar mass (35.45) is the weighted average of the nearly integral mass for the two isotopes 35 Cl and 37 Cl.
Variations of properties between isotopes
Chemical and molecular properties
Neutral atoms have the same number of electrons as protons. So different isotopes of a given element all have the same number of electrons and share the same electronic structure. Because the chemical behavior of atoms is largely determined by its electronic structure, different isotopes exhibit almost identical chemical behavior.
The main exception to this is the kinetic isotope effect: because of their larger mass, heavier isotopes tend to react somewhat slower than the lighter isotopes of the same element. This is most obvious so far for the protium ( 1 H
), deuterium ( 2 H
), and tritium ( 3
), since deuterium has twice the mass of protium and tritium having three times the mass of the protium. This mass difference also affects the behavior of chemical bonds each, by altering the center of gravity (decreasing mass) of the atomic system. However, for heavier elements the relative mass differences between isotopes are much less, so the effect of mass differences on chemicals is usually negligible. (The heavy elements also have relatively more neutrons than lighter elements, so the ratio of nuclear masses to collective electronic mass is slightly larger.)
Similarly, two different molecules only in their isotopic isotopes have identical electronic structures, and therefore physical and chemical properties are almost indistinguishable (again with deuterium and tritium being the main exception). The vibrational mode of the molecule is determined by its shape and by its constituent atomic mass; Different isotopologists have different vibration mode sequences. Because the vibration mode allows molecules to absorb suitable energy photons, isotopologists have different optical properties in the infrared range.
Nuclear properties and stability
The atomic nucleus consists of protons and neutrons bound together by the remaining strong force. Because protons are positively charged, protons repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their attitude pushes protons slightly apart, reducing electrostatic repulsion between protons, and they exert a nuclear force that attracts each other and on protons. For this reason, one or more neutrons are required for two or more protons to bind to the nucleus. As the number of protons increases, so does the ratio of neutrons to the protons necessary to ensure a stable nucleus (see chart on the right). For example, although the ratio of neutrons: proton 3 2 He
is 1: 2, the ratio of neutrons : proton 238
92 U
larger than 3: 2. Some lightweight elements have stable nuclides with a 1: 1 ratio ( Z = N ). Nuclide 40
20 Ca
(calcium-40) is the heaviest stable nuclide observation with number of neutrons and the same proton; (Theoretically, the heaviest stable is sulfur-32). All stable nuclides heavier than calcium-40 contain more neutrons than protons.
Number of isotopes per element
Of the 80 elements with stable isotopes, the largest number of stable isotopes observed for each element is ten (for tin elements). No element has nine stable isotopes. Xenon is the only element with eight stable isotopes. The four elements have seven stable isotopes, eight have six stable isotopes, ten have five stable isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have two stable isotopes (calculate 180m
73 Ta
as stable), and 26 elements have only one stable isotope (from this, 19 is called a mononuclide element, has a stable isotope primordials that dominate and fix the atomic weight of the natural element for high precision; 3 radioactive mononuclidic elements occur as well). In total, there are 253 unrevealed nuclides. For 80 elements that have one or more stable isotopes, the average number of stable isotopes is 253/80 = 3.1625 isotopes per element.
Odd and peculiar number of odd nonsense
Proton: the ratio of neutrons is not the only factor that affects nuclear stability. It also depends on the flatness or strangeness of the atomic number Z , the neutron number N and, consequently, of their number, the mass of A . The oddities of both Z and N tend to decrease nuclear binding energy, making the odd nucleus, generally, less stable. The remarkable difference between nuclear energies between neighboring nuclei, especially the odd isoar - A , has important consequences: unstable isotopes with neutron counts or unoptimal release of protons by beta decay (including positron decay) , capture electrons or other exotic means, such as spontaneous cleavage and group decay.
The majority of stable nuclides are protons-even-neutrons, where all numbers Z , N , and A are even. The stable nuclide that is A is split (roughly evenly) into neutron-neutron necrons, and even-proton-necron neutrons. The neutron-odd-weird nucleus is the most common.
Even atomic number
The 148 even-proton, even neutron (EE) nuclides comprise ~ 58% of all stable nuclides and all have a 0 spin because of the pair. There are also 22 primordial long-lived even-nuclides. As a result, each of the 41 even numbered elements from 2 to 82 has at least one stable isotope, and most of these elements have some primordial isotopes. Half of these even-numbered elements have six or more stable isotopes. The extreme stability of helium-4 because of double pairs of 2 protons and 2 neutrons prevents any nuclides containing five or eight nucleons from long enough to function as a platform for heavier nuclear element buildup. fusion in stars (see process alpha three).
These stable nuclides have an even number of protons and an odd number of neutrons. They are a minority compared to even-isotope, which is about 3 times as much. Among the 41 even- Z elements that have stable nuclides, there are only two elements (argon and cerium) that do not have an even strange stable nuclide. One element (lead) has three. There are 24 elements that have one even nucl and 13 which have two odd nuclides. Of the 35 primordial radionuclides there are four even nuclides (see table on the right), including fissile 235
92 U
. Because of the strange neutron numbers, even strange nuclides tend to have a large cross-section of neutrons, due to the energy generated by the effects of neutron pairs. These stable neutron nuclides tend to be infrequent due to abundance in nature, generally because, to form and enter into primordial abundance, they must escape capturing neutrons to form other stable even-stable isotopes, during both process s and r-capture processes neutrons, during the nucleosynthesis of stars. For this reason, just 78 Pt
and 9
4 Be
is the most natural isotope of its element.
Odd atomic number
Forty-eight stable proton-in-neutron neutron nuclides, stabilized by their number of neutrons in pairs, form the most stable isotope of odd-numbered elements; very few strange-proton-odd-neutron nuclides are made up of others. There are 41 odd-numbered elements with Z = 1 to 81, of which 39 have stable isotopes (technetium element (
43 Tc
) and promethium (
style = "font-size: inherit; vertical-align: baseline"> 61 Pm
) does not have an isotope stable). Of these 39 odd elements, 30 elements (including hydrogen-1 where even 0 neutrons) have one odd-even stable isotope, and nine elements: chlorine (
17 Cl
), potassium (
19 K
), copper (
29 Cu
), gallium (
31 Ga < inherit; line-height: inherit; vertical-align: baseline "> 35 Br
), silver (
47 Ag
), antimony (
51 Sb
), iridium ( < span>
77 Ir
), and thallium (
81 Tl
), has two odd isotopes-even each stable.This makes a total of 30 2 (9) = 48 odd-stable isotopes.
There are also five primordial radioactive odd-long-life isotopes, 87
< sub style = "font-size: inherit; line-height: inherit; vertical-align: baseline"> 37 Rb
, 115
49 In
187
75 Re
, 151
63 Eu
, and 209
83 Bi
The last two have just been found to rot, with the half-life of more than 10 18 .
Only five stable nuclides contain an odd number of protons and the odd number of neutrons. The first four "weird" nuclides occur in the low-mass nuclei, which convert the protons into neutrons or vice versa will cause a very oblique proton-neutron ratio ( 2
1 H
6 3 Li , 10 style = "font-size: inherit; line- height: inherit; vertical-align: baseline "> 14
7 N
; spinning 1, 1, 3, 1) The only weird nuclides that Another "stable" is 180m
73 Ta
(spin 9) is considered the rarest of 253 stable isotopes, and is the only Primordial nuclear isomers, which have not yet been observed to rot even when the experiment is carried out.
Many odd radionuclides (such as tantalum-180) with a relatively short half-life are known. Usually, they experience beta decay to nearby nearby isbars that have combined protons and neutrons in pairs. Of the nine primordial strange nuclides (five stable and four radioactive with half life lengths), just 14
7 N
is the most common isotope of common elements. This is the case because this is part of the CNO cycle. Nuklida 6 3 Li
and 10
5 B
is a minor isotope of the elements that itself rarely compare to other light elements, while the other six isotopes form only a small percentage of natural abundance their elements.
Odd neutron number
Actinides with strange neutron numbers are generally fissile (with thermal neutrons), whereas even with the number of neutrons generally not, though fissionable with fast neutrons. All observable observable strange nuclides have a non-zero integer spin. This is because the only unpaired neutrons and unpaired protons have a greater nuclear force attraction to each other if their spins are parallel (resulting in a total rotation of at least 1 unit), not anti-alignment. See deuterium for the simplest case of this nuclear behavior.
Only
78 Pt
, 9
4 Be in and 14
7 N The has a strange neutron number and is the most natural isotope of its element.
Genesis in nature
The elements are composed of one nuclide (mononucidal element) or more naturally occurring isotopes. The unstable isotope (radioactive) is primordial or postprimal. The primordial isotope is a product of the nucleosynthesis of stars or other types of nucleosynthesis such as cosmic ray splitting, and has survived to date because its decay rate is very slow (eg uranium-238 and potassium-40). Post-primordial isotopes are created by the bombardment of cosmic rays as cosmogenic nuclides (eg, tritium, carbon-14), or by decay of radioactive primitive isotopes into radioactive radiogenic nuclide princons (eg uranium to radium). Some isotopes are naturally synthesized as nucleic nuclides, by some other natural nuclear reactions, such as when neutrons from natural nuclear fission are absorbed by other atoms.
As discussed above, only 80 elements have stable isotopes, and 26 of these have only one stable isotope. Thus, about two-thirds of stable elements occur naturally in the Earth in stable isotopes, with the largest number of stable isotopes for the ten elements, for tin (
50 Sn
). There are about 94 elements found naturally on Earth (up to the inclusion of plutonium), although some are only detected in very small amounts, such as plutonium-244. Scientists estimate that the naturally occurring elements on Earth (some just as radioisotopes) occur as 339 isotopes (nuclides) in total. Only 253 of these natural nuclides are stable in the sense of never being observed to rot at this point. An additional 35 primordial nuclides (up to a total of 289 primordial nuclides), are radioactive with half the life known, but have a half-life of over 80 million years, allowing them to live from the beginning of the Solar System. See the list of nuclides for details.
All known stable nuclides occur naturally on Earth; Other natural nuclides are radioactive but occur on Earth because of its relatively long half life, or because of other means of ongoing natural production. These include the previously mentioned cosmogenic nucleicides, nucleobic nuclides, and radiogenic nuclides formed by continuous decay of primordial radioactive nuclides, such as radon and radium from uranium.
An additional ~ 3000 radioactive nuclides not found in nature have been made in nuclear reactors and in particle accelerators. Many short-lived nuclides not found naturally on Earth have also been observed by spectroscopic analysis, naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but is found in abundance on an astronomical scale.
The tabulated atomic mass of the elements is the average that explains the existence of several isotopes with different masses. Prior to the discovery of isotopes, empirically determined noninteger values ​​of atomic masses puzzled scientists. For example, chlorine samples contain 75.8% chlorine-35 and 24.2% chlorine-37, giving an average atomic mass of 35.5 atomic mass units.
According to generally acceptable cosmological theories, only the isotopes of hydrogen and helium, traces of several isotopes of lithium and beryllium, and possibly some boron, were created in the Big Bang, while all other nuclides were then synthesized, in stars and supernovas, and in the interaction between energetic particles such as cosmic rays, and previously produced nuclides. (See nucleosynthesis for details of various processes considered responsible for the production of isotopes.) The abundance of each isotope on Earth is generated from the amount formed by this process, its spreading through the galaxy, and the decay rate for the unstable isotope. After the initial blend of the Solar System, the isotopes are redistributed by mass, and the elemental isotope composition varies somewhat from planet to planet. It is sometimes possible to trace the origin of the meteorite.
Isotope atomic mass
The atomic mass ( m r ) of the isotope (nuclide) is determined primarily by the sum of its mass (ie the number of nucleons in the nucleus). The small correction is due to the binding energy of the nucleus (see mass defect), the small difference in mass between protons and neutrons, and the mass of electrons associated with atoms, the latter due to electrons: the ratio of nucleons differs among isotopes.
The mass number is a dimensionless quantity. The atomic mass, on the other hand, is measured using units of atomic mass based on the mass of carbon-12 atoms. This is denoted by the symbol "u" ​​(for unified atomic mass unity) or "Da" (for dalton).
Massa atom isotop yang terbentuk secara alamiah dari suatu unsur menentukan massa atom dari unsur tersebut. Ketika elemen mengandung N isotop, express bawah ini diterapkan untuk massa atom rata-rata :
where m 1 , m 2 ,..., m N is the atomic mass of each individual isotope, and x 1 ,..., x N is the relative abundance of this isotope.
Applications isotope
Isotope refining
Source of the article : Wikipedia
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