Heavy water is a form of water that contains a larger than normal amount of the hydrogen isotope deuterium, rather than the common hydrogen-1 isotope that makes up most of the hydrogen in normal water. The presence of deuterium gives the water different nuclear properties, the increase of mass gives it different physical and chemical properties when compared to normal water. Deuterium is a hydrogen isotope with a nucleus containing a proton; the additional neutron makes a deuterium atom twice as heavy as a protium atom. A molecule of heavy water has two deuterium atoms in place of the two protium atoms of ordinary "light" water; the weight of a heavy water molecule, however, is not different from that of a normal water molecule, because about 89% of the molecular weight of water comes from the single oxygen atom rather than the two hydrogen atoms. The colloquial term'heavy water' refers to a enriched water mixture that contains deuterium oxide D2O, but some hydrogen-deuterium oxide and a smaller amount of ordinary hydrogen oxide H2O.
For instance, the heavy water used in CANDU reactors is 99.75% enriched by hydrogen atom-fraction—meaning that 99.75% of the hydrogen atoms are of the heavy type. For comparison, ordinary water contains only about 156 deuterium atoms per million hydrogen atoms, meaning that 0.0156% of the hydrogen atoms are of the heavy type. Heavy water is not radioactive. In its pure form, it has a density about 11% greater than water, but is otherwise physically and chemically similar; the various differences in deuterium-containing water are larger than in any other occurring isotope-substituted compound because deuterium is unique among heavy stable isotopes in being twice as heavy as the lightest isotope. This difference increases the strength of water's hydrogen-oxygen bonds, this in turn is enough to cause differences that are important to some biochemical reactions; the human body contains deuterium equivalent to about five grams of heavy water, harmless. When a large fraction of water in higher organisms is replaced by heavy water, the result is cell dysfunction and death.
Heavy water was first produced in a few months after the discovery of deuterium. With the discovery of nuclear fission in late 1938, the need for a neutron moderator that captured few neutrons, heavy water became a component of early nuclear energy research. Since heavy water has been an essential component in some types of reactors, both those that generate power and those designed to produce isotopes for nuclear weapons; these heavy water reactors have the advantage of being able to run on natural uranium without using graphite moderators that pose radiological and dust explosion hazards in the decommissioning phase. Most modern reactors use enriched uranium with ordinary water as the moderator. Semiheavy water, HDO, exists whenever there is water with light deuterium in the mix; this is because hydrogen atoms are exchanged between water molecules. Water containing 50% H and 50% D in its hydrogen contains about 50% HDO and 25% each of H2O and D2O, in dynamic equilibrium. In normal water, about 1 molecule in 3,200 is HDO, heavy water molecules only occur in a proportion of about 1 molecule in 41 million.
Thus semiheavy water molecules are far more common than "pure" heavy water molecules. Water enriched in the heavier oxygen isotopes 17O and 18O is commercially available, e.g. for use as a non-radioactive isotopic tracer. It is "heavy water" as it is denser than normal water —but is called heavy water, since it does not contain the deuterium that gives D2O its unusual nuclear and biological properties, it is more expensive than D2O due to the more difficult separation of 17O and 18O. H218O is used for production of fluorine-18 for radiopharmaceuticals and radiotracers and for positron emission tomography. Tritiated water contains tritium in place of protium or deuterium, therefore it is radioactive; the physical properties of water and heavy water differ in several respects. Heavy water is less dissociated than light water at given temperature, the true concentration of D+ ions is less than H+ ions would be for a light water sample at the same temperature; the same is true of OD OH − ions. For heavy water Kw D2O = 1.35 × 10−15, must equal for neutral water.
Thus pKw D2O = p + p = 7.44 + 7.44 = 14.87, the p of neutral heavy water at 25.0 °C is 7.44. The pD of heavy water is measured using pH electrodes giving a pH value, or pHa, at various temperatures a true acidic pD can be estimated from the directly pH meter measured pHa, such that pD+ = pHa + 0.41. The electrode correction for alkaline conditions is 0.456 for heavy water. The alkaline correction is pD+ = pHa + 0.456. These corrections are different from the differences in p and p of 0.44 from the corresponding ones in heavy water. Heavy water is 10.6% denser than ordinary water, heavy water's physically different properties can be seen without equipment if a frozen sample is dropped into normal water, as it will sink. If the water is ice-cold the higher melting tem
Deuterium is one of two stable isotopes of hydrogen. The nucleus of deuterium, called a deuteron, contains one proton and one neutron, whereas the far more common protium has no neutron in the nucleus. Deuterium has a natural abundance in Earth's oceans of about one atom in 6420 of hydrogen, thus deuterium accounts for 0.0156% of all the occurring hydrogen in the oceans, while protium accounts for more than 99.98%. The abundance of deuterium changes from one kind of natural water to another; the deuterium isotope's name is formed from the Greek deuteros, meaning "second", to denote the two particles composing the nucleus. Deuterium was named in 1931 by Harold Urey; when the neutron was discovered in 1932, this made the nuclear structure of deuterium obvious, Urey won the Nobel Prize in 1934. Soon after deuterium's discovery and others produced samples of "heavy water" in which the deuterium content had been concentrated. Deuterium is destroyed in the interiors of stars faster. Other natural processes are thought to produce only an insignificant amount of deuterium.
Nearly all deuterium found in nature was produced in the Big Bang 13.8 billion years ago, as the basic or primordial ratio of hydrogen-1 to deuterium has its origin from that time. This is the ratio found in the gas giant planets, such as Jupiter. However, other astronomical bodies are found to have different ratios of deuterium to hydrogen-1; this is thought to be a result of natural isotope separation processes that occur from solar heating of ices in comets. Like the water cycle in Earth's weather, such heating processes may enrich deuterium with respect to protium; the analysis of deuterium/protium ratios in comets found results similar to the mean ratio in Earth's oceans. This reinforces theories; the deuterium/protium ratio of the comet 67P/Churyumov-Gerasimenko, as measured by the Rosetta space probe, is about three times that of earth water. This figure is the highest yet measured in a comet. Deuterium/protium ratios thus continue to be an active topic of research in both astronomy and climatology.
Deuterium is represented by the chemical symbol D. Since it is an isotope of hydrogen with mass number 2, it is represented by 2H. IUPAC allows 2H, although 2H is preferred. A distinct chemical symbol is used for convenience because of the isotope's common use in various scientific processes, its large mass difference with protium confers non-negligible chemical dissimilarities with protium-containing compounds, whereas the isotope weight ratios within other chemical elements are insignificant in this regard. In quantum mechanics the energy levels of electrons in atoms depend on the reduced mass of the system of electron and nucleus. For the hydrogen atom, the role of reduced mass is most seen in the Bohr model of the atom, where the reduced mass appears in a simple calculation of the Rydberg constant and Rydberg equation, but the reduced mass appears in the Schrödinger equation, the Dirac equation for calculating atomic energy levels; the reduced mass of the system in these equations is close to the mass of a single electron, but differs from it by a small amount about equal to the ratio of mass of the electron to the atomic nucleus.
For hydrogen, this amount is about 1837/1836, or 1.000545, for deuterium it is smaller: 3671/3670, or 1.0002725. The energies of spectroscopic lines for deuterium and light hydrogen therefore differ by the ratios of these two numbers, 1.000272. The wavelengths of all deuterium spectroscopic lines are shorter than the corresponding lines of light hydrogen, by a factor of 1.000272. In astronomical observation, this corresponds to a blue Doppler shift of 0.000272 times the speed of light, or 81.6 km/s. The differences are much more pronounced in vibrational spectroscopy such as infrared spectroscopy and Raman spectroscopy, in rotational spectra such as microwave spectroscopy because the reduced mass of the deuterium is markedly higher than that of protium. In nuclear magnetic resonance spectroscopy, deuterium has a different NMR frequency and is much less sensitive. Deuterated solvents are used in protium NMR to prevent the solvent from overlapping with the signal, although deuterium NMR on its own right is possible.
Deuterium is thought to have played an important role in setting the number and ratios of the elements that were formed in the Big Bang. Combining thermodynamics and the changes brought about by cosmic expansion, one can calculate the fraction of protons and neutrons based on the temperature at the point that the universe cooled enough to allow formation of nuclei; this calculation indicates seven protons for every neutron at the beginning of nucleogenesis, a ratio that would remain stable after nucleogenesis was over. This fraction was in favor of protons primarily because the lower mass of the proton favored their production; as the universe expanded, it cooled. Free neutrons and protons are less stable than helium nuclei, the protons and neutrons had a strong energetic reason to form helium-4. However, forming helium-4 requires the intermediate step of forming deuterium. Through much of the few minutes after the big bang during which nucleosynthesis could have occurred
Tritium is a radioactive isotope of hydrogen. The nucleus of tritium contains one proton and two neutrons, whereas the nucleus of protium contains one proton and no neutrons. Occurring tritium is rare on Earth, where trace amounts are formed by the interaction of the atmosphere with cosmic rays, it can be produced by irradiating lithium metal or lithium-bearing ceramic pebbles in a nuclear reactor. Tritium is used as a radioactive tracer, in radioluminescent light sources for watches and instruments, along with deuterium, as a fuel for nuclear fusion reactions with applications in energy generation and weapons; the name of this isotope is derived from Greek, Modern τρίτος, meaning'third'. While tritium has several different experimentally determined values of its half-life, the National Institute of Standards and Technology lists 4,500 ± 8 days, it decays into helium-3 by beta decay as in this nuclear equation: and it releases 18.6 keV of energy in the process. The electron's kinetic energy varies, with an average of 5.7 keV, while the remaining energy is carried off by the nearly undetectable electron antineutrino.
Beta particles from tritium can penetrate only about 6.0 mm of air, they are incapable of passing through the dead outermost layer of human skin. The unusually low energy released in the tritium beta decay makes the decay appropriate for absolute neutrino mass measurements in the laboratory; the low energy of tritium's radiation makes it difficult to detect tritium-labeled compounds except by using liquid scintillation counting. Tritium is produced in nuclear reactors by neutron activation of lithium-6; this is possible with neutrons of any energy, is an exothermic reaction yielding 4.8 MeV. In comparison, the fusion of deuterium with tritium releases about 17.6 MeV of energy. For applications in proposed fusion energy reactors, such as ITER, pebbles consisting of lithium bearing ceramics including Li2TiO3 and Li4SiO4, are being developed for tritium breeding within a helium cooled pebble bed known as a breeder blanket. High-energy neutrons can produce tritium from lithium-7 in an endothermic reaction, consuming 2.466 MeV.
This was discovered. High-energy neutrons irradiating boron-10 will occasionally produce tritium: A more common result of boron-10 neutron capture is 7Li and a single alpha particle. Tritium is produced in heavy water-moderated reactors whenever a deuterium nucleus captures a neutron; this reaction has a quite small absorption cross section, making heavy water a good neutron moderator, little tritium is produced. So, cleaning tritium from the moderator may be desirable after several years to reduce the risk of its escaping to the environment. Ontario Power Generation's "Tritium Removal Facility" processes up to 2,500 tonnes of heavy water a year, it separates out about 2.5 kg of tritium, making it available for other uses. Deuterium's absorption cross section for thermal neutrons is about 0.52 millibarns, whereas that of oxygen-16 is about 0.19 millibarns and that of oxygen-17 is about 240 millibarns. Tritium is an uncommon product of the nuclear fission of uranium-235, plutonium-239, uranium-233, with a production of about one atom per each 10,000 fissions.
The release or recovery of tritium needs to be considered in the operation of nuclear reactors in the reprocessing of nuclear fuels and in the storage of spent nuclear fuel. The production of tritium is not a goal, but rather a side-effect, it is discharged to the atmosphere in small quantities by some nuclear power plants. In June 2016 the Tritiated Water Task Force released a report on the status of tritium in tritiated water at Fukushima Daiichi nuclear plant, as part of considering options for final disposal of this water; this identified that the March 2016 holding of tritium on-site was 760 TBq in a total of 860000 m3 of stored water. This report identified the reducing concentration of tritium in the water extracted from the buildings etc. for storage, seeing a factor of ten decrease over the five years considered, 3.3 MBq/L to 0.3 MBq/L. According to a report by an expert panel considering the best approach to dealing with this issue, "Tritium could be separated theoretically, but there is no practical separation technology on an industrial scale.
Accordingly, a controlled environmental release is said to be the best way to treat low-tritium-concentration water." Tritium's decay product helium-3 has a large cross section for reacting with thermal neutrons, expelling a proton, hence it is converted back to tritium in nuclear reactors. Tritium occurs due to cosmic rays interacting with atmospheric gases. In the most important reaction for natural production, a fast neutron interacts with atmospheric nitrogen: Worldwide, the production of tritium from natural sources is 148 petabecquerels per year; the global equilibrium inventory of tritium created by natural sources remains constant at 2,590 petabecquerels. This is due to losses proportional to the inventory. According to a 1996 report from Institute for Energy and Environmental Research on the US Department of Energy, only 225 kg of tritium had been produced in the United States from 1955 to 1996. Since it continually de
The atomic number or proton number of a chemical element is the number of protons found in the nucleus of an atom. It is identical to the charge number of the nucleus; the atomic number uniquely identifies a chemical element. In an uncharged atom, the atomic number is equal to the number of electrons; the sum of the atomic number Z and the number of neutrons, N, gives the mass number A of an atom. Since protons and neutrons have the same mass and the mass defect of nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in unified atomic mass units, is within 1% of the whole number A. Atoms with the same atomic number Z but different neutron numbers N, hence different atomic masses, are known as isotopes. A little more than three-quarters of occurring elements exist as a mixture of isotopes, the average isotopic mass of an isotopic mixture for an element in a defined environment on Earth, determines the element's standard atomic weight, it was these atomic weights of elements that were the quantities measurable by chemists in the 19th century.
The conventional symbol Z comes from the German word Zahl meaning number, before the modern synthesis of ideas from chemistry and physics denoted an element's numerical place in the periodic table, whose order is but not consistent with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence that this Z number was the nuclear charge and a physical characteristic of atoms, did the word Atomzahl come into common use in this context. Loosely speaking, the existence or construction of a periodic table of elements creates an ordering of the elements, so they can be numbered in order. Dmitri Mendeleev claimed. However, in consideration of the elements' observed chemical properties, he changed the order and placed tellurium ahead of iodine; this placement is consistent with the modern practice of ordering the elements by proton number, Z, but that number was not known or suspected at the time. A simple numbering based on periodic table position was never satisfactory, however.
Besides the case of iodine and tellurium several other pairs of elements were known to have nearly identical or reversed atomic weights, thus requiring their placement in the periodic table to be determined by their chemical properties. However the gradual identification of more and more chemically similar lanthanide elements, whose atomic number was not obvious, led to inconsistency and uncertainty in the periodic numbering of elements at least from lutetium onward. In 1911, Ernest Rutherford gave a model of the atom in which a central core held most of the atom's mass and a positive charge which, in units of the electron's charge, was to be equal to half of the atom's atomic weight, expressed in numbers of hydrogen atoms; this central charge would thus be half the atomic weight. In spite of Rutherford's estimation that gold had a central charge of about 100, a month after Rutherford's paper appeared, Antonius van den Broek first formally suggested that the central charge and number of electrons in an atom was equal to its place in the periodic table.
This proved to be the case. The experimental position improved after research by Henry Moseley in 1913. Moseley, after discussions with Bohr, at the same lab, decided to test Van den Broek's and Bohr's hypothesis directly, by seeing if spectral lines emitted from excited atoms fitted the Bohr theory's postulation that the frequency of the spectral lines be proportional to the square of Z. To do this, Moseley measured the wavelengths of the innermost photon transitions produced by the elements from aluminum to gold used as a series of movable anodic targets inside an x-ray tube; the square root of the frequency of these photons increased from one target to the next in an arithmetic progression. This led to the conclusion that the atomic number does correspond to the calculated electric charge of the nucleus, i.e. the element number Z. Among other things, Moseley demonstrated that the lanthanide series must have 15 members—no fewer and no more—which was far from obvious from the chemistry at that time.
After Moseley's death in 1915, the atomic numbers of all known elements from hydrogen to uranium were examined by his method. There were seven elements which were not found and therefore identified as still undiscovered, corresponding to atomic numbers 43, 61, 72, 75, 85, 87 and 91. From 1918 to 1947, all seven of these missing elements were discovered. By this time the first four transuranium elements had been discovered, so that the periodic table was complete with no gaps as far as curium. In 1915 the rea
Nuclear transmutation is the conversion of one chemical element or an isotope into another chemical element. Because any element is defined by its number of protons in its atoms, i.e. in the atomic nucleus, nuclear transmutation occurs in any process where the number of protons or neutrons in the nucleus is changed. A transmutation can be achieved either by nuclear reactions or by radioactive decay, where no outside cause is needed. Natural transmutation by stellar nucleosynthesis in the past created most of the heavier chemical elements in the known existing universe, continues to take place to this day, creating the vast majority of the most common elements in the universe, including helium and carbon. Most stars carry out transmutation through fusion reactions involving hydrogen and helium, while much larger stars are capable of fusing heavier elements up to iron late in their evolution. Elements heavier than iron, such as gold and lead, are created through elemental transmutations that can only take place in supernovae - as stars begin to fuse heavier elements less energy is released from each fusion reaction, each fusion reaction that produces elements heavier than iron is endothermic in nature, stars are incapable of carrying this out.
One type of natural transmutation observable in the present occurs when certain radioactive elements present in nature spontaneously decay by a process that causes transmutation, such as alpha or beta decay. An example is the natural decay of potassium-40 to argon-40, which forms most of the argon in the air. On Earth, natural transmutations from the different mechanisms of natural nuclear reactions occur, due to cosmic ray bombardment of elements, occasionally from natural neutron bombardment. Artificial transmutation may occur in machinery that has enough energy to cause changes in the nuclear structure of the elements; such machines include particle accelerators and tokamak reactors. Conventional fission power reactors cause artificial transmutation, not from the power of the machine, but by exposing elements to neutrons produced by fission from an artificially produced nuclear chain reaction. For instance, when a uranium atom is bombarded with slow neutrons, fission takes place; this releases, on 3 neutrons and a large amount of energy.
The released neutrons cause fission of other uranium atoms, until all of the available uranium is exhausted. This is called a chain reaction. Artificial nuclear transmutation has been considered as a possible mechanism for reducing the volume and hazard of radioactive waste; the term transmutation dates back to alchemy. Alchemists pursued the philosopher's stone, capable of chrysopoeia – the transformation of base metals into gold. While alchemists understood chrysopoeia as a metaphor for a mystical, or religious process, some practitioners adopted a literal interpretation, tried to make gold through physical experiment; the impossibility of the metallic transmutation had been debated amongst alchemists and scientists since the Middle Ages. Pseudo-alchemical transmutation was outlawed and publicly mocked beginning in the fourteenth century. Alchemists like Michael Maier and Heinrich Khunrath wrote tracts exposing fraudulent claims of gold making. By the 1720s, there were no longer any respectable figures pursuing the physical transmutation of substances into gold.
Antoine Lavoisier, in the 18th century, replaced the alchemical theory of elements with the modern theory of chemical elements, John Dalton further developed the notion of atoms to explain various chemical processes. The disintegration of atoms is a distinct process involving much greater energies than could be achieved by alchemists, it was first consciously applied to modern physics by Frederick Soddy when he, along with Ernest Rutherford, discovered that radioactive thorium was converting itself into radium in 1901. At the moment of realization, Soddy recalled, he shouted out: "Rutherford, this is transmutation!" Rutherford snapped back, "For Christ's sake, don't call it transmutation. They'll have our heads off as alchemists."Rutherford and Soddy were observing natural transmutation as a part of radioactive decay of the alpha decay type. The first artificial transmutation was accomplished in 1925 by Patrick Blackett, a research fellow working under Rutherford, with the transmutation of nitrogen into oxygen, using alpha particles directed at nitrogen 14N + α → 17O + p. Rutherford had shown in 1919 that a proton was emitted from alpha bombardment experiments but he had no information about the residual nucleus.
Blackett's 1921-1924 experiments provided the first experimental evidence of an artificial nuclear transmutation reaction. Blackett identified the underlying integration process and the identity of the residual nucleus. In 1932, a artificial nuclear reaction and nuclear transmutation was achieved by Rutherford's colleagues John Cockcroft and Ernest Walton, who used artificially accelerated protons against lithium-7 to split the nucleus into two alpha particles; the feat was popularly known as "splitting the atom," although it was not the modern nuclear fission reaction discovered in 1938 by Otto Hahn, Lise Meitner and their assistant Fritz Strassmann in heavy elements. In the twentieth century the transmutation of elements within stars was elaborated, accounting for the relative abundance of heavier elements in the universe. Save for the first five elements, which were produced in the Big Bang and other cosmic ray processes, stellar nucleosynthesis accounted for the abu
Table of nuclides
A table of nuclides or chart of nuclides is a two-dimensional graph in which one axis represents the number of neutrons and the other represents the number of protons in an atomic nucleus. Each point plotted on the graph thus represents the nuclide of a real or hypothetical chemical element; this system of ordering nuclides can offer a greater insight into the characteristics of isotopes than the better-known periodic table, which shows only elements instead of each of their isotopes. A chart or table of nuclides is a simple map to the nuclear, or radioactive, behaviour of nuclides, as it distinguishes the isotopes of an element, it contrasts with a periodic table, which only maps their chemical behavior, since isotopes do not differ chemically to any significant degree, with the exception of hydrogen. Nuclide charts organize nuclides along the X axis by their numbers of neutrons and along the Y axis by their numbers of protons, out to the limits of the neutron and proton drip lines; this representation was first published by Kurt Guggenheimer in 1934 and expanded by Giorgio Fea in 1935, Emilio Segrè in 1945 or G. Seaborg.
In 1958, Walter Seelmann-Eggebert and Gerda Pfennig published the first edition of the Karlsruhe Nuclide Chart. Its 7th edition was made available in 2006. Today, there are several nuclide charts, four of which have a wide distribution: the Karlsruhe Nuclide Chart, the Strasbourg Universal Nuclide Chart, the Chart of the Nuclides from the JAEA and the Nuclide Chart from Knolls Atomic Power Laboratory, it has become a basic tool of the nuclear community. The nuclide table below shows nuclides, including all with half-life of at least one day, they are arranged with increasing atomic numbers from left to right and increasing neutron numbers from top to bottom. Cell color denotes the half-life of each nuclide. In graphical browsers, each nuclide has a tool tip indicating its half-life; each color represents a certain range of length of half-life, the color of the border indicates the half-life of its nuclear isomer state. Some nuclides have multiple nuclear isomers, this table notes the longest one.
Dotted borders mean that a nuclide has a nuclear isomer, their color is represented the same way as for their normal counterparts. Isotopes are nuclides with the same number of protons but differing numbers of neutrons. Isotopes neighbor each other vertically, e.g. carbon-12, carbon-13, carbon-14 or oxygen-15, oxygen-16, oxygen-17. Isotones are nuclides with the same number of neutrons but differing number of protons. Isotones neighbor each other horizontally. Example: carbon-14, nitrogen-15, oxygen-16 in the sample table above. Isobars are nuclides with the same number of nucleons, i.e. mass number, but different numbers of protons and different number of neutrons. Isobars neighbor each other diagonally from lower-left to upper-right. Example: carbon-14, nitrogen-14, oxygen-14 in the sample table above. Isodiaphers are nuclides with the same difference between protons. Like isobars, they at right angles to the isobar lines. Examples: boron-10, carbon-12, nitrogen-14 where N−Z=0. Beyond the neutron drip line along the lower left, nuclides decay by neutron emission.
Beyond the proton drip line along the upper right, nuclides decay by proton emission. Drip lines have only been established for some elements; the island of stability is a hypothetical region of the table of nuclides that contains isotopes far more stable than other transuranic elements. There are no stable nuclides having an equal number of protons and neutrons in their nuclei with atomic number greater than 20 as can be "read" from the chart. Nuclei of greater atomic number require an excess of neutrons for stability; the only stable nuclides having an odd number of protons and an odd number of neutrons are hydrogen-2, lithium-6, boron-10, nitrogen-14 and tantalum-180m. This is because the mass-energy of such atoms is higher than that of their neighbors on the same isobaric chain, so most of them are unstable to beta decay. There are no stable nuclides with mass numbers 5 or 8. There are stable nuclides with all other mass numbers up to 208 with the exceptions of 147 and 151. With the possible exception of the pair tellurium-123 and antimony-123, odd mass numbers are never represented by more than one stable nuclide.
This is because the mass-energy is a convex function of atomic number, so all nuclides on an odd isobaric chain except one have a lower-energy neighbor to which they can decay by beta decay. There are no stable nuclides having atomic number greater than Z=82, although bismuth is stable for all practical human purposes. Elements with atomic numbers from 1 to 82 all have stable isotopes, with the exceptions of technetium and promethium. Interactive Chart of Nuclides app for mobiles: Android or Apple - for PC use The Live Chart of Nuclides - IAEA Another example of a Chart of Nuclides from Korea Data up to Jan 1999 only