1.
Radioactive decay
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A material containing such unstable nuclei is considered radioactive. Certain highly excited short-lived nuclear states can decay through neutron emission, or more rarely, however, for a collection of atoms, the collections expected decay rate is characterized in terms of their measured decay constants or half-lives. This is the basis of radiometric dating, the half-lives of radioactive atoms have no known upper limit, spanning a time range of over 55 orders of magnitude, from nearly instantaneous to far longer than the age of the universe. A radioactive nucleus with spin can have no defined orientation. If there are multiple particles produced during a single decay, as in decay, their relative angular distribution. Such a parent process could be a previous decay, or a nuclear reaction, the decaying nucleus is called the parent radionuclide, and the process produces at least one daughter nuclide. Except for gamma decay or internal conversion from an excited state. When the number of changes, an atom of a different chemical element is created. The first decay processes to be discovered were alpha decay, beta decay, alpha decay occurs when the nucleus ejects an alpha particle. This is the most common process of emitting nucleons, but highly excited nuclei can eject single nucleons, or in the case of cluster decay, specific light nuclei of other elements. Beta decay occurs when the nucleus emits an electron or positron, the nucleus may capture an orbiting electron, causing a proton to convert into a neutron in a process called electron capture. All of these result in a well-defined nuclear transmutation. By contrast, there are radioactive decay processes that do not result in a nuclear transmutation, another type of radioactive decay results in products that vary, appearing as two or more fragments of the original nucleus with a range of possible masses. For a summary table showing the number of stable and radioactive nuclides in each category, there are 29 naturally occurring chemical elements on Earth that are radioactive. They are those that contain 34 radionuclides that date before the time of formation of the solar system, well-known examples are uranium and thorium, but also included are naturally occurring long-lived radioisotopes, such as potassium-40. Radioactivity was discovered in 1896 by the French scientist Henri Becquerel and these materials glow in the dark after exposure to light, and he suspected that the glow produced in cathode ray tubes by X-rays might be associated with phosphorescence. He wrapped a photographic plate in black paper and placed various phosphorescent salts on it, all results were negative until he used uranium salts. The uranium salts caused a blackening of the plate in spite of the plate being wrapped in black paper and these radiations were given the name Becquerel Rays
2.
Stable isotope
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The term stable isotope has a similar meaning to stable nuclide, but is preferably used when speaking of nuclides of a specific element. Hence, the plural form stable isotopes usually refers to isotopes of the same element, the relative abundance of such stable isotopes can be measured experimentally, yielding an isotope ratio that can be used as a research tool. Theoretically, such stable isotopes could include the radiogenic daughter products of radioactive decay, however, the expression stable isotope ratio is preferably used to refer to isotopes whose relative abundances are affected by isotope fractionation in nature. This field is termed stable isotope geochemistry, measurement of the ratios of naturally occurring stable isotopes plays an important role in isotope geochemistry, but stable isotopes are also finding uses in ecological and biological studies. Other workers have used oxygen isotope ratios to reconstruct historical atmospheric temperatures, the long history of study of these elements is in part because the proportions of stable isotopes in these light and volatile elements is relatively easy to measure. However, recent advances in isotope ratio mass spectrometry now enable the measurement of isotope ratios in heavier elements, such as iron, copper, zinc, molybdenum. The variations in oxygen and hydrogen isotope ratios have applications in hydrology since most samples will lie between two extremes, ocean water and Arctic/Antarctic snow, stable isotopes of water are also used in partitioning water sources for plant transpiration and groundwater recharge. Another application is in paleotemperature measurement for paleoclimatology, for example, one technique is based on the variation in isotopic fractionation of oxygen by biological systems with temperature. Species of Foraminifera incorporate oxygen as calcium carbonate in their shells, the ratio of the oxygen isotopes oxygen-16 and oxygen-18 incorporated into the calcium carbonate varies with temperature and the oxygen isotopic composition of the water. This oxygen remains fixed in the calcium carbonate when the dies, falls to the sea bed. In forensic science, research suggests that the variation in certain isotope ratios in drugs derived from plant sources can be used to determine the drugs continent of origin and it also has applications in doping control, to distinguish between endogenous and exogenous sources of hormones. Chondrite meteorites are classified using the isotope ratios. In addition, a signature of carbon-13 confirms the non-terrestrial origin for organic compounds found in carbonaceous chondrites
3.
Beta decay
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In nuclear physics, beta decay is a type of radioactive decay in which a beta ray, and a neutrino are emitted from an atomic nucleus. Neither the beta particle nor its associated neutrino exist within the prior to beta decay. By this process, unstable atoms obtain a more stable ratio of protons to neutrons, the probability of a nuclide decaying due to beta and other forms of decay is determined by its binding energy. The binding energies of all existing nuclides form what is called the valley of stability. Beta decay is a consequence of the force, which is characterized by relatively lengthy decay times. Nucleons are composed of up or down quarks, and the force allows a quark to change type by the exchange of a W boson. For example, a neutron, composed of two quarks and an up quark, decays to a proton composed of a down quark. Decay times for many nuclides that are subject to beta decay can be thousands of years, electron capture is sometimes included as a type of beta decay, because the basic nuclear process, mediated by the weak force, is the same. In electron capture, an atomic electron is captured by a proton in the nucleus, transforming it into a neutron. The two types of decay are known as beta minus and beta plus. β+ decay is known as positron emission. Beta decay conserves a number known as the lepton number, or the number of electrons. These particles have lepton number +1, while their antiparticles have lepton number −1, since a proton or neutron has lepton number zero, β+ decay must be accompanied with an electron neutrino, while β− decay must be accompanied by an electron antineutrino. This new element has a mass number A, but an atomic number Z that is increased by one. As in all nuclear decays, the element is known as the parent nuclide while the resulting element is known as the daughter nuclide. The beta spectrum, or distribution of values for the beta particles, is continuous. The total energy of the process is divided between the electron, the antineutrino, and the recoiling nuclide. In the figure to the right, an example of an electron with 0.40 MeV energy from the decay of 210Bi is shown
4.
Proton emission
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Proton emission is a rare type of radioactive decay in which a proton is ejected from a nucleus. For a proton to escape a nucleus, the proton separation energy must be negative - the proton is therefore unbound, Proton emission is not seen in naturally occurring isotopes, proton emitters can be produced via nuclear reactions, usually using linear particle accelerators. Research in the field flourished after this breakthrough, and to more than 25 isotopes have been found to exhibit proton emission. The study of proton emission has aided the understanding of nuclear deformation, masses and structure, in 2002, the simultaneous emission of two protons was observed from the nucleus iron-45 in experiments at GSI and GANIL. In 2005 it was determined that zinc-54 can also undergo double proton decay. Proton drip line Diproton Free neutron Neutron emission Nuclear Structure and Decay Data - IAEA with query on Proton Separation Energy
5.
Neutron emission
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Neutron emission is a type of radioactive decay of atoms containing excess neutrons, in which a neutron is simply ejected from the nucleus. Neutron emission is one of the ways an atom reaches its stability, an atom is unstable, therefore radioactive, when the forces in the nucleus are unbalanced. The instability of the results from the nuclei having extra neutrons or extra protons. Two examples of isotopes that emit neutrons are beryllium-13 and helium-5, commonly, it is abbreviated with a lower case n. As only a neutron is lost in process, the atom does not gain or lose any protons. Instead, the atom will become a new isotope of the original element, Neutron emission usually happens from nuclei that are in an excited state, such as the excited 17O* produced from the beta decay of 17N. The neutron emission process itself is controlled by the force and therefore is extremely fast. This process allows unstable atoms to become more stable, a synonym for such neutron emission is prompt neutron production, of the type that is best known to occur simultaneously with induced nuclear fission. Induced fission happens only when a nucleus is bombarded with neutrons, gamma rays, many heavy isotopes, most notably californium-252, also emit prompt neutrons among the products of a similar spontaneous radioactive decay process, spontaneous fission. Spontaneous fission happens when a nucleus splits into two smaller nuclei and generally one or more neutrons. Most neutron emission outside prompt neutron production associated with fission, is from neutron-heavy isotopes produced as fission products. Thus, the delay in neutron emission is not from the process, but rather its precursor beta decay, which is controlled by the weak force. The beta decay half lives for the precursors to delayed neutron-emitter radioisotopes, are typically fractions of a second to tens of seconds, Neutron radiation Neutron source Neutron drip line Proton emission Why Are Some Atoms Radioactive. The LIVEChart of Nuclides - IAEA with filter on delayed neutron emission decay Nuclear Structure and Decay Data - IAEA with query on Neutron Separation Energy
6.
Alpha decay
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An alpha particle is identical to the nucleus of a helium-4 atom, which consists of two protons and two neutrons. Alpha decay typically occurs in the heaviest nuclides, in practice, this mode of decay has only been observed in nuclides considerably heavier than nickel, with the lightest known alpha emitters being the lightest isotopes of tellurium. Alpha decay is by far the most common form of cluster decay and it is the most common form because of the combined extremely high binding energy and relatively small mass of the alpha particle. Like other cluster decays, alpha decay is fundamentally a quantum tunneling process, unlike beta decay, it is governed by the interplay between both the nuclear force and the electromagnetic force. Alpha particles have a kinetic energy of 5 MeV and have a speed of about 15,000,000 m/s. There is surprisingly small variation around this energy, due to the dependence of the half-life of this process on the energy produced. Approximately 99% of the helium produced on Earth is the result of the decay of underground deposits of minerals containing uranium or thorium. The helium is brought to the surface as a by-product of natural gas production, Alpha particles were first described in the investigations of radioactivity by Ernest Rutherford in 1899, and by 1907 they were identified as He2+ ions. By 1928, George Gamow had solved the theory of alpha decay via tunneling, the alpha particle is trapped in a potential well by the nucleus. The nuclear force holding an atomic nucleus together is very strong, however, the nuclear force is also short range, dropping quickly in strength beyond about 1 femtometre, while the electromagnetic force has unlimited range. A nucleus with 210 or more nucleons is so large that the nuclear force holding it together can just barely counterbalance the electromagnetic repulsion between the protons it contains. Alpha decay occurs in such nuclei as a means of increasing stability by reducing size, one curiosity is why alpha particles, helium nuclei, should be preferentially emitted as opposed to other particles like a single proton or neutron or other atomic nuclei. Single proton emission, or the emission of any particle with an odd number of nucleons would violate this conservation law, the rest of the answer comes from the very high binding energy of the alpha particle. For example, performing the calculation for uranium-232 shows that alpha particle emission would need only 5.4 MeV, while a single proton emission would require 6.1 MeV. Most of this energy becomes the kinetic energy of the alpha particle itself. However, since the numbers of most alpha emitting radioisotopes exceed 210. These disintegration energies however are substantially smaller than the barrier provided by the nuclear force. An alpha particle can be thought of as being inside a potential barrier whose walls are 25 MeV, however, decay alpha particles only have kinetic energies of 4 MeV to about 9 MeV, far less than the energy needed to escape
7.
Spontaneous fission
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Spontaneous fission is a form of radioactive decay that is found only in very heavy chemical elements. The first nuclear fission process discovered was the fission induced by neutrons, because cosmic rays produce some neutrons, it was difficult to distinguish between induced and spontaneous events. Cosmic rays can be shielded by a thick layer of rock or water. Spontaneous fission was identified in 1940 by Soviet physicists Georgy Flyorov, cluster decay was shown to be a superasymmetric spontaneous fission process. The lightest natural nuclides that are subject to spontaneous fission are niobium-93. Spontaneous fission has never observed in the naturally occurring isotopes of these elements. In practice, these are stable isotopes, spontaneous fission is feasible over practical observation times only for atomic masses of 232 amu or more. These are elements at least as heavy as thorium-232 – which has a somewhat longer than the age of the universe. Thorium-232 is the lightest primordial nuclide that has evidence of undergoing spontaneous fission in its minerals. The known elements most susceptible to spontaneous fission are the synthetic high-atomic-number actinides and transactinides with atomic numbers from 100 onwards, hence, the spontaneous fission of these isotopes is usually negligible, except in using the exact branching ratios when finding the radioactivity of a sample of these elements. Within the framework of liquid drop model, the criterion for whether spontaneous fission can occur in a short enough to be observed by present methods, is approximately. Where Z is the number and A is the mass number. Spontaneous fission rates, In practice 239Pu will invariably contain an amount of 240Pu due to the tendency of 239Pu to absorb an additional neutron during production. 240Pus high rate of fission events makes it an undesirable contaminant. Weapons-grade plutonium contains no more than 7. 0% 240Pu, the rarely used gun-type atomic bomb has a critical insertion time of about one millisecond, and the probability of a fission during this time interval should be small. Almost all nuclear bombs use some kind of implosion method, spontaneous fission can occur much more rapidly when the nucleus of an atom undergoes superdeformation. Spontaneous fission gives much the result as induced nuclear fission. However, like other forms of decay, it occurs due to quantum tunneling
8.
Positron emission
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Positron emission is mediated by the weak force. The positron is a type of particle, the other beta particle being the electron emitted from the β− decay of a nucleus. Positron decay results in nuclear transmutation, changing an atom of one element into an atom of an element with an atomic number that is less by one unit. Positron emission should not be confused with electron emission or beta decay, which occurs when a neutron turns into a proton and the nucleus emits an electron. This was the first example of β+ decay, the Curies termed the phenomenon artificial radioactivity, since 30 15P is a short-lived nuclide which does not exist in nature. The discovery of radioactivity would be cited when the husband. Isotopes which undergo this decay and thereby emit positrons include carbon-11, potassium-40, nitrogen-13, oxygen-15, aluminium-26, sodium-22, fluorine-18, the two most common types of quarks are up quarks, which have a charge of +2/3, and down quarks, with a −1/3 charge. Quarks arrange themselves in sets of three such that they make protons and neutrons, in a proton, whose charge is +1, there are two up quarks and one down quark. Neutrons, with no charge, have one up quark and two down quarks, via the weak interaction, quarks can change flavor from down to up, resulting in electron emission. Positron emission happens when an up quark changes into a down quark, nuclei which decay by positron emission may also decay by electron capture. For low-energy decays, electron capture is energetically favored by 2mec2 =1.022 MeV, as the energy of the decay goes up, so does the branching ratio towards positron emission. However, if the difference is less than 2mec2, then positron emission cannot occur. Certain isotopes are stable in galactic cosmic rays, because the electrons are stripped away, a positron is ejected from the parent nucleus, and the daughter atom must shed an orbital electron to balance charge. Isotopes which increase in mass under the conversion of a proton to a neutron, or which decrease in mass by less than 2me and these isotopes are used in positron emission tomography, a technique used for medical imaging. Note that the energy emitted depends on the isotope that is decaying, the short-lived positron emitting isotopes 11C, 13N, 15O and 18F used for positron emission tomography are typically produced by proton irradiation of natural or enriched targets. The LIVEChart of Nuclides - IAEA with filter on β+ decay
9.
Electron capture
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Electron capture is a process in which the proton-rich nucleus of an electrically neutral atom absorbs an inner atomic electron, usually from the K or L electron shell. This process thereby changes a proton to a neutron and simultaneously causes the emission of an electron neutrino. P + e− → n + ν e The daughter nuclide, if it is in an excited state, usually, a gamma ray is emitted during this transition, but nuclear de-excitation may also take place by internal conversion. Following capture of an electron from the atom, an outer electron replaces the electron that was captured. Following electron capture, the number is reduced by one, the neutron number is increased by one. Simple electron capture results in an atom, since the loss of the electron in the electron shell is balanced by a loss of positive nuclear charge. However, an atomic ion may result from further Auger electron emission. Electron capture is an example of interaction, one of the four fundamental forces. Electron capture is always an alternate mode for radioactive isotopes that do have sufficient energy to decay by positron emission. It is sometimes called Inverse beta decay, though this term can refer to the interaction of an electron antineutrino with a proton. For example, rubidium-83 will decay to krypton-83 solely by electron capture, the theory of electron capture was first discussed by Gian-Carlo Wick in a 1934 paper, and then developed by Hideki Yukawa and others. K-electron capture was first observed by Luis Alvarez, in vanadium-48 and he reported it in a 1937 paper in Physical Review. Alvarez went on to study electron capture in gallium-67 and other nuclides, radioactive isotopes that decay by pure electron capture can be inhibited from radioactive decay if they are fully ionized. Anomalies in elemental distributions are thought to be partly a result of this effect on electron capture, chemical bonds can also affect the rate of electron capture to a small degree depending on the proximity of electrons to the nucleus. For example, in 7Be, a difference of 0. 9% has been observed between half-lives in metallic and insulating environments and this relatively large effect is due to the fact that beryllium is a small atom whose valence electrons are close to the nucleus. Electron capture happens most often in the heavier neutron-deficient elements where the change is smallest. When the loss of mass in a reaction is greater than zero but less than 2m. Some common radioisotopes that decay by electron capture include, For a full list, the LIVEChart of Nuclides - IAEA with filter on electron capture
10.
Cartesian coordinate system
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Each reference line is called a coordinate axis or just axis of the system, and the point where they meet is its origin, usually at ordered pair. The coordinates can also be defined as the positions of the projections of the point onto the two axis, expressed as signed distances from the origin. One can use the principle to specify the position of any point in three-dimensional space by three Cartesian coordinates, its signed distances to three mutually perpendicular planes. In general, n Cartesian coordinates specify the point in an n-dimensional Euclidean space for any dimension n and these coordinates are equal, up to sign, to distances from the point to n mutually perpendicular hyperplanes. The invention of Cartesian coordinates in the 17th century by René Descartes revolutionized mathematics by providing the first systematic link between Euclidean geometry and algebra. Using the Cartesian coordinate system, geometric shapes can be described by Cartesian equations, algebraic equations involving the coordinates of the points lying on the shape. For example, a circle of radius 2, centered at the origin of the plane, a familiar example is the concept of the graph of a function. Cartesian coordinates are also tools for most applied disciplines that deal with geometry, including astronomy, physics, engineering. They are the most common system used in computer graphics, computer-aided geometric design. Nicole Oresme, a French cleric and friend of the Dauphin of the 14th Century, used similar to Cartesian coordinates well before the time of Descartes. The adjective Cartesian refers to the French mathematician and philosopher René Descartes who published this idea in 1637 and it was independently discovered by Pierre de Fermat, who also worked in three dimensions, although Fermat did not publish the discovery. Both authors used a single axis in their treatments and have a length measured in reference to this axis. The concept of using a pair of axes was introduced later, after Descartes La Géométrie was translated into Latin in 1649 by Frans van Schooten and these commentators introduced several concepts while trying to clarify the ideas contained in Descartes work. Many other coordinate systems have developed since Descartes, such as the polar coordinates for the plane. The development of the Cartesian coordinate system would play a role in the development of the Calculus by Isaac Newton. The two-coordinate description of the plane was later generalized into the concept of vector spaces. Choosing a Cartesian coordinate system for a one-dimensional space – that is, for a straight line—involves choosing a point O of the line, a unit of length, and an orientation for the line. An orientation chooses which of the two half-lines determined by O is the positive, and which is negative, we say that the line is oriented from the negative half towards the positive half
11.
Number of neutrons
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The neutron number, symbol N, is the number of neutrons in a nuclide. Atomic number plus neutron number equals mass number, Z+N=A, the difference between the neutron number and the atomic number is known as the neutron excess, D = N - Z = A - 2Z. Neutron number is rarely written explicitly in nuclide symbol notation, in order of increasing explicitness and decreasing frequency of usage, Nuclides that have the same neutron number but a different proton number are called isotones. This word was formed by replacing the p in isotope with n for neutron, Nuclides that have the same mass number are called isobars. Nuclides that have the same neutron excess are called isodiaphers, chemical properties are primarily determined by proton number, which determines which chemical element the nuclide is a member of, neutron number has only a slight influence. Neutron number is primarily of interest for nuclear properties, for example, actinides with odd neutron number are usually fissile while actinides with even neutron number are usually not fissile. Only 57 stable nuclides have an odd number, compared to 200 with an even neutron number. No odd-neutron-number isotope is the most naturally abundant isotope in its element, except for beryllium-9 which is the only stable isotope, nitrogen-14. Only two stable nuclides have fewer neutrons than protons, hydrogen-1 and helium-3
12.
Atomic number
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The atomic number or proton number of a chemical element is the number of protons found in the nucleus of an atom of that element. It is identical to the number of the nucleus. The atomic number identifies a chemical element. In an uncharged atom, the number is also equal to the number of electrons. The atomic number Z, should not be confused with the mass number A and this number of neutrons, N, completes the weight, A = Z + N. Atoms with the atomic number Z but different neutron numbers N. Historically, it was these atomic weights of elements that were the quantities measurable by chemists in the 19th century. Only after 1915, with the suggestion and evidence that this Z number was also the nuclear charge, loosely speaking, the existence or construction of a periodic table of elements creates an ordering of the elements, and so they can be numbered in order. Dmitri Mendeleev claimed that he arranged his first periodic tables in order of atomic weight, however, in consideration of the elements observed chemical properties, he changed the order slightly and placed tellurium ahead of iodine. This placement is consistent with the practice of ordering the elements by proton number, Z. A simple numbering based on periodic table position was never entirely satisfactory and this central charge would thus be approximately half the atomic weight. This proved eventually to be the case, the experimental position improved dramatically after research by Henry Moseley in 1913. 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 and this led to the conclusion that the atomic number does closely correspond to the calculated electric charge of the nucleus, i. e. the element number Z. Among other things, Moseley demonstrated that the series must have 15 members—no fewer. After Moseleys death in 1915, the 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, from 1918 to 1947, all seven of these missing elements were discovered. By this time the first four transuranium elements had also been discovered, in 1915 the reason for nuclear charge being quantized in units of Z, which were now recognized to be the same as the element number, was not understood
13.
Atomic nucleus
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After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. Almost all of the mass of an atom is located in the nucleus, protons and neutrons are bound together to form a nucleus by the nuclear force. The diameter of the nucleus is in the range of 6985175000000000000♠1.75 fm for hydrogen to about 6986150000000000000♠15 fm for the heaviest atoms and these dimensions are much smaller than the diameter of the atom itself, by a factor of about 23,000 to about 145,000. The branch of physics concerned with the study and understanding of the nucleus, including its composition. The nucleus was discovered in 1911, as a result of Ernest Rutherfords efforts to test Thomsons plum pudding model of the atom, the electron had already been discovered earlier by J. J. Knowing that atoms are electrically neutral, Thomson postulated that there must be a charge as well. In his plum pudding model, Thomson suggested that an atom consisted of negative electrons randomly scattered within a sphere of positive charge, to his surprise, many of the particles were deflected at very large angles. This justified the idea of an atom with a dense center of positive charge. The term nucleus is from the Latin word nucleus, a diminutive of nux, in 1844, Michael Faraday used the term to refer to the central point of an atom. The modern atomic meaning was proposed by Ernest Rutherford in 1912, the adoption of the term nucleus to atomic theory, however, was not immediate. In 1916, for example, Gilbert N, the nuclear strong force extends far enough from each baryon so as to bind the neutrons and protons together against the repulsive electrical force between the positively charged protons. The nuclear strong force has a short range, and essentially drops to zero just beyond the edge of the nucleus. The collective action of the charged nucleus is to hold the electrically negative charged electrons in their orbits about the nucleus. The collection of negatively charged electrons orbiting the nucleus display an affinity for certain configurations, which chemical element an atom represents is determined by the number of protons in the nucleus, the neutral atom will have an equal number of electrons orbiting that nucleus. Individual chemical elements can create more stable electron configurations by combining to share their electrons and it is that sharing of electrons to create stable electronic orbits about the nucleus that appears to us as the chemistry of our macro world. Protons define the entire charge of a nucleus, and hence its chemical identity, neutrons are electrically neutral, but contribute to the mass of a nucleus to nearly the same extent as the protons. Neutrons explain the phenomenon of isotopes – varieties of the chemical element which differ only in their atomic mass. They are sometimes viewed as two different quantum states of the particle, the nucleon
14.
Nuclide
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A nuclide is an atomic species characterized by the specific constitution of its nucleus, i. e. by its number of protons Z, its number of neutrons N, and its nuclear energy state. The word nuclide was proposed by Truman P. Kohman in 1947, Kohman originally suggested nuclide as referring to a species of nucleus defined by containing a certain number of neutrons and protons. The word thus was originally intended to focus on the nucleus, nuclide refers to a nucleus rather than to an atom. Identical nuclei belong to one nuclide, for each nucleus of the carbon-13 nuclide is composed of 6 protons and 7 neutrons. The nuclide concept emphasizes nuclear properties over chemical properties, while the isotope concept emphasizes chemical over nuclear, the neutron number has large effects on nuclear properties, but its effect on chemical properties is negligible for most elements. Since isotope is the term, it is better known than nuclide. A set of nuclides with equal number, i. e. of the same chemical element but different neutron numbers, are called isotopes of the element. Particular nuclides are still often loosely called isotopes, but the term nuclide is the one in general. In similar manner, a set of nuclides with equal mass number A, but different atomic number, are called isobars, and isotones are nuclides of equal neutron number but different proton numbers. The name isotone has been derived from the isotope to emphasize that in the first group of nuclides it is the number of neutrons that is constant. See Isotope#Notation for an explanation of the used for different nuclide or isotope types. Nuclear isomers are members of a set of nuclides with equal number and equal mass number. An example is the two states of the single isotope 99 43Tc shown among the decay schemes, each of these two states qualifies as a different nuclide, illustrating one way that nuclides may differ from isotopes. The most long-lived non-ground state nuclear isomer is the nuclide tantalum-180m and this nuclide occurs primordially, and has never been observed to decay to the ground state. There are about 254 nuclides in nature that have never observed to decay. They occur among the 80 different elements that have one or more stable isotopes, see stable isotope and primordial nuclide. Unstable nuclides are radioactive and are called radionuclides and their decay products are called radiogenic nuclides. About 254 stable and about 85 unstable nuclides exist naturally on Earth, natural radionuclides may be conveniently subdivided into three types
15.
Chemical element
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A chemical element or element is a species of atoms having the same number of protons in their atomic nuclei. There are 118 elements that have identified, of which the first 94 occur naturally on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have exclusively radioactive isotopes, iron is the most abundant element making up Earth, while oxygen is the most common element in the Earths crust. Chemical elements constitute all of the matter of the universe. The two lightest elements, hydrogen and helium, were formed in the Big Bang and are the most common elements in the universe. The next three elements were formed mostly by cosmic ray spallation, and are rarer than those that follow. Formation of elements with from 6 to 26 protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis, the high abundance of oxygen, silicon, and iron on Earth reflects their common production in such stars. The term element is used for atoms with a number of protons as well as for a pure chemical substance consisting of a single element. A single element can form multiple substances differing in their structure, when different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds. Only a minority of elements are found uncombined as relatively pure minerals, among the more common of such native elements are copper, silver, gold, carbon, and sulfur. All but a few of the most inert elements, such as gases and noble metals, are usually found on Earth in chemically combined form. While about 32 of the elements occur on Earth in native uncombined forms. For example, atmospheric air is primarily a mixture of nitrogen, oxygen, and argon, the history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur, copper and gold. Later civilizations extracted elemental copper, tin, lead and iron from their ores by smelting, using charcoal, alchemists and chemists subsequently identified many more, almost all of the naturally occurring elements were known by 1900. Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, almost all other elements found in nature were made by various natural methods of nucleosynthesis. On Earth, small amounts of new atoms are produced in nucleogenic reactions, or in cosmogenic processes. Of the 94 naturally occurring elements, those with atomic numbers 1 through 82 each have at least one stable isotope, Isotopes considered stable are those for which no radioactive decay has yet been observed. Elements with atomic numbers 83 through 94 are unstable to the point that radioactive decay of all isotopes can be detected, the very heaviest elements undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized
16.
Periodic table
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The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configurations, and recurring chemical properties. This ordering shows periodic trends, such as elements with similar behaviour in the same column and it also shows four rectangular blocks with some approximately similar chemical properties. In general, within one row the elements are metals on the left, the rows of the table are called periods, the columns are called groups. Six groups have names as well as numbers, for example, group 17 elements are the halogens, and group 18, the noble gases. The periodic table can be used to derive relationships between the properties of the elements, and predict the properties of new elements yet to be discovered or synthesized, the periodic table provides a useful framework for analyzing chemical behaviour, and is widely used in chemistry and other sciences. The Russian chemist Dmitri Mendeleev published the first widely recognized periodic table in 1869 and he developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of elements that would be expected to fill gaps in this table. Most of his predictions were proved correct when the elements in question were subsequently discovered, Mendeleevs periodic table has since been expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour. The first 94 elements exist naturally, although some are only in trace amounts and were synthesized in laboratories before being found in nature. Elements with atomic numbers from 95 to 118 have only been synthesized in laboratories or nuclear reactors, synthesis of elements having higher atomic numbers is being pursued. Numerous synthetic radionuclides of naturally occurring elements have also produced in laboratories. Each chemical element has an atomic number representing the number of protons in its nucleus. Most elements have differing numbers of neutrons among different atoms, with variants being referred to as isotopes. Isotopes are never separated in the table, they are always grouped together under a single element. Elements with no stable isotopes have the masses of their most stable isotopes. In the standard periodic table, the elements are listed in order of increasing atomic number, a new row is started when a new electron shell has its first electron. Columns are determined by the configuration of the atom, elements with the same number of electrons in a particular subshell fall into the same columns. Thus, it is easy to predict the chemical properties of an element if one knows the properties of the elements around it
17.
Neutron
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The neutron is a subatomic particle, symbol n or n0, with no net electric charge and a mass slightly larger than that of a proton. Protons and neutrons, each with approximately one atomic mass unit, constitute the nucleus of an atom. Their properties and interactions are described by nuclear physics, the nucleus consists of Z protons, where Z is called the atomic number, and N neutrons, where N is the neutron number. The atomic number defines the properties of the atom. The terms isotope and nuclide are often used synonymously, but they are chemical and nuclear concepts, the atomic mass number, symbol A, equals Z+N. For example, carbon has atomic number 6, and its abundant carbon-12 isotope has 6 neutrons, some elements occur in nature with only one stable isotope, such as fluorine. Other elements occur with many stable isotopes, such as tin with ten stable isotopes, even though it is not a chemical element, the neutron is included in the table of nuclides. Within the nucleus, protons and neutrons are bound together through the nuclear force, neutrons are produced copiously in nuclear fission and fusion. They are a contributor to the nucleosynthesis of chemical elements within stars through fission, fusion. The neutron is essential to the production of nuclear power, in the decade after the neutron was discovered in 1932, neutrons were used to induce many different types of nuclear transmutations. These events and findings led to the first self-sustaining nuclear reactor, free neutrons, or individual neutrons free of the nucleus, are effectively a form of ionizing radiation, and as such, are a biological hazard, depending upon dose. A small natural background flux of free neutrons exists on Earth, caused by cosmic ray showers. Dedicated neutron sources like neutron generators, research reactors and spallation sources produce free neutrons for use in irradiation, neutrons and protons are both nucleons, which are attracted and bound together by the nuclear force to form atomic nuclei. The nucleus of the most common isotope of the atom is a lone proton. The nuclei of the hydrogen isotopes deuterium and tritium contain one proton bound to one. All other types of nuclei are composed of two or more protons and various numbers of neutrons. The most common nuclide of the chemical element lead, 208Pb has 82 protons and 126 neutrons. The free neutron has a mass of about 1. 675×10−27 kg, the neutron has a mean square radius of about 0. 8×10−15 m, or 0.8 fm, and it is a spin-½ fermion
18.
Proton
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A proton is a subatomic particle, symbol p or p+, with a positive electric charge of +1e elementary charge and mass slightly less than that of a neutron. Protons and neutrons, each with masses of one atomic mass unit, are collectively referred to as nucleons. One or more protons are present in the nucleus of every atom, the number of protons in the nucleus is the defining property of an element, and is referred to as the atomic number. Since each element has a number of protons, each element has its own unique atomic number. The word proton is Greek for first, and this name was given to the nucleus by Ernest Rutherford in 1920. In previous years, Rutherford had discovered that the nucleus could be extracted from the nuclei of nitrogen by atomic collisions. Protons were therefore a candidate to be a particle, and hence a building block of nitrogen. In the modern Standard Model of particle physics, protons are hadrons, and like neutrons, although protons were originally considered fundamental or elementary particles, they are now known to be composed of three valence quarks, two up quarks and one down quark. The rest masses of quarks contribute only about 1% of a protons mass, the remainder of a protons mass is due to quantum chromodynamics binding energy, which includes the kinetic energy of the quarks and the energy of the gluon fields that bind the quarks together. At sufficiently low temperatures, free protons will bind to electrons, however, the character of such bound protons does not change, and they remain protons. A fast proton moving through matter will slow by interactions with electrons and nuclei, the result is a protonated atom, which is a chemical compound of hydrogen. In vacuum, when electrons are present, a sufficiently slow proton may pick up a single free electron, becoming a neutral hydrogen atom. Such free hydrogen atoms tend to react chemically with other types of atoms at sufficiently low energies. When free hydrogen atoms react with other, they form neutral hydrogen molecules. Protons are spin-½ fermions and are composed of three quarks, making them baryons. Protons have an exponentially decaying positive charge distribution with a mean square radius of about 0.8 fm. Protons and neutrons are both nucleons, which may be together by the nuclear force to form atomic nuclei. The nucleus of the most common isotope of the atom is a lone proton
19.
Nuclear drip line
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An arbitrary combination of protons and neutrons does not necessarily yield a stable nucleus. One can think of moving up and/or to the right across the table of nuclides by adding one type of nucleon to a given nucleus. However, adding one at a time to a given nucleus will eventually lead to a newly formed nucleus that immediately decays by emitting a proton. Colloquially speaking, the nucleon has leaked or dripped out of the nucleus, drip lines are defined for protons, neutrons, and alpha particles, and these all play important roles in nuclear physics. The nucleon drip lines are at the extreme of the ratio, at p, n ratios at or beyond the driplines. The location of the drip line is not well known for most of the nuclear chart, whereas the proton. Nuclear stability is limited to those combinations of protons and neutrons described by the chart of the nuclides, the boundaries of this valley are the neutron drip line on the neutron rich side, and the proton drip line on the proton-rich side. The alpha drip line is more difficult to visualize as it also branches down through the center of the chart. These limits exist because of decay, whereby an exothermic nuclear transition can occur by the emission of one or more nucleons. To understand the concept, one needs to apply the principle of conservation of energy to nuclear binding energy. For example, to determine if 12C, the most common isotope of carbon, can undergo proton emission to 11B, one finds that about 16 MeV must be added to the system for this process to be allowed. While Q-values can be used to any nuclear transmutation, for particle decay, the particle separation energy quantity S, is also used. In other words, the proton separation energy Sp indicates how much energy should be added to a nucleus to remove a single proton. Thus, the drip lines defined the boundaries where the particle separation energy is less than or equal to zero. Of the three types of naturally occurring radioactivities, only alpha decay is a type of resulting from the nuclear strong force. The other proton and neutron decays occurred much earlier in the life of the atomic species, thus, alpha-decay can be considered either a form of particle decay or, less frequently, as a special case of nuclear fission. For alpha decay, the timescale can be longer than for proton or neutron emission owing to the high Coulomb barrier seen by an alpha-cluster in a nucleus. Such particle decays are not commonly known because particle decay is governed by the strong force, as well as the Coulomb force in the case of charged particles
20.
Half-life
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Half-life is the time required for a quantity to reduce to half its initial value. The term is used in nuclear physics to describe how quickly unstable atoms undergo. The term is used more generally to characterize any type of exponential or non-exponential decay. For example, the medical sciences refer to the biological half-life of drugs, the converse of half-life is doubling time. The original term, half-life period, dating to Ernest Rutherfords discovery of the principle in 1907, was shortened to half-life in the early 1950s. Rutherford applied the principle of a radioactive elements half-life to studies of age determination of rocks by measuring the period of radium to lead-206. Half-life is constant over the lifetime of an exponentially decaying quantity, the accompanying table shows the reduction of a quantity as a function of the number of half-lives elapsed. A half-life usually describes the decay of discrete entities, such as radioactive atoms, in that case, it does not work to use the definition that states half-life is the time required for exactly half of the entities to decay. For example, if there are 3 radioactive atoms with a half-life of one second, instead, the half-life is defined in terms of probability, Half-life is the time required for exactly half of the entities to decay on average. In other words, the probability of a radioactive atom decaying within its half-life is 50%, for example, the image on the right is a simulation of many identical atoms undergoing radioactive decay. Note that after one half-life there are not exactly one-half of the remaining, only approximately. Nevertheless, when there are many identical atoms decaying, the law of large numbers suggests that it is a good approximation to say that half of the atoms remain after one half-life. There are various simple exercises that demonstrate probabilistic decay, for example involving flipping coins or running a computer program. The three parameters t1⁄2, τ, and λ are all related in the following way. Amount approaches zero as t approaches infinity as expected, some quantities decay by two exponential-decay processes simultaneously. There is a half-life describing any exponential-decay process, for example, The current flowing through an RC circuit or RL circuit decays with a half-life of RCln or lnL/R, respectively. For this example, the half time might be used instead of half life. In a first-order chemical reaction, the half-life of the reactant is ln/λ, in radioactive decay, the half-life is the length of time after which there is a 50% chance that an atom will have undergone nuclear decay
21.
Tooltip
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The tooltip or infotip or a hint is a common graphical user interface element. It is used in conjunction with a cursor, usually a pointer, the user hovers the pointer over an item, without clicking it, and a tooltip may appear—a small hover box with information about the item being hovered over. Tooltips do not usually appear on operating systems, because there is no cursor. A common variant, especially in software, is displaying a description of the tool in a status bar. Another system, in the classic Mac OS, that aims to solve the same problem, microsoft invented another term, ScreenTip, and uses it in its end-user documentation. Demonstrations of tooltip usage are prevalent on web pages, some browsers, notably Microsofts Internet Explorer, will also display the alt attribute of an image as a tooltip in the same manner if an alt attribute is specified and a title attribute is not. This is a behaviour of the alt attribute according to the W3C specification. If a title attribute is specified, it will override the alt attribute for tooltip content. CSS, HTML, and JavaScript allow web designers to create customized tooltip, more recently, these tooltips are used in various parts of an interface, not only on toolbars. Some software and applications, such as GIMP, provide an option for users to turn off some or all tooltips, however, such options are left to the discretion of the developer, and are often not implemented
22.
Isotopes of hydrogen
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Hydrogen has three naturally occurring isotopes, sometimes denoted 1H, 2H, and 3H. The first two of these are stable while 3H has a half-life of 12.32 years, all heavier isotopes are synthetic and have a half-life less than one zeptosecond. Of these, 5H is the most stable, and 7H is the least, Hydrogen is the only element whose isotopes have different names that are in common use today. The 2H isotope is usually called deuterium, while the 3H isotope is usually called tritium, the symbols D and T are sometimes used for deuterium and tritium. The IUPAC states in the 2005 Red Book that while the use of D and T is common, the ordinary isotope of hydrogen, with no neutrons, is sometimes called protium. 1H is the most common hydrogen isotope with an abundance of more than 99. 98%, because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium. The proton has never observed to decay and hydrogen-1 is therefore considered a stable isotope. Some grand unified theories proposed in the 1970s predict that proton decay can occur with a half-life between 1031 and 1036 years, if this prediction is found to be true, then hydrogen-1 are only observationally stable. To date however, experiments have shown that the minimum proton half life is in excess of 1034 years, 2H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. The nucleus of deuterium is called a deuteron, deuterium comprises 0.0026 –0. 0184% of hydrogen samples on Earth, with the lower number tending to be found in samples of hydrogen gas and the higher enrichment typical of ocean water. Deuterium on Earth has been enriched with respect to its concentration in the Big Bang. Deuterium is not radioactive, and does not represent a significant toxicity hazard, water enriched in molecules that include deuterium instead of protium is called heavy water. Deuterium and its compounds are used as a label in chemical experiments. Heavy water is used as a moderator and coolant for nuclear reactors. Deuterium is also a fuel for commercial nuclear fusion. 3H is known as tritium and contains one proton and two neutrons in its nucleus and it is radioactive, decaying into helium-3 through β− decay with a half-life of 12.32 years. Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases, tritium has also been released during nuclear weapons tests. It is used in thermonuclear weapons, as a tracer in isotope geochemistry
23.
Isotopes of helium
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Although there are nine known isotopes of helium, only helium-3 and helium-4 are stable. All radioisotopes are short-lived, the longest-lived being 6He with a half-life of 806.7 milliseconds, the least stable is 5He, with a half-life of 7. 6×10−22 s, although it is possible that 2He has an even shorter half-life. In the Earths atmosphere, there is one 3He atom for every million 4He atoms, however, helium is unusual in that its isotopic abundance varies greatly depending on its origin. In the interstellar medium, the proportion of 3He is around a hundred times higher. Rocks from the Earths crust have isotope ratios varying by as much as a factor of ten, this is used in geology to investigate the origin of rocks, the different formation processes of the two stable isotopes of helium produce the differing isotope abundances. Equal mixtures of liquid 3He and 4He below 0.8 K will separate into two phases due to their dissimilarity. Dilution refrigerators take advantage of the immiscibility of these two isotopes to achieve temperatures of a few millikelvins, helium-2 or 2He, also known as a diproton, is an extremely unstable isotope of helium that consists of two protons with no neutrons. According to theoretical calculations it would have much more stable if the strong force had been 2% greater. There may have been observations of 2He, in 2000, physicists first observed a new type of radioactive decay in which a nucleus emits two protons at once—perhaps a 2He nucleus. Galindo-Uribarri and co-workers chose an isotope of neon with a structure that prevents it from emitting protons one at a time. This means that the two protons are ejected simultaneously, the team fired a beam of fluorine ions at a proton-rich target to produce 18Ne, which then decayed into oxygen and two protons. Any protons ejected from the target itself were identified by their characteristic energies, there are two ways in which the two-proton emission may proceed. The neon nucleus might eject a diproton—a pair of protons bound together as a 2He nucleus—which then decays into separate protons, alternatively, the protons may be emitted separately but simultaneously—so-called democratic decay. The experiment was not sensitive enough to establish which of two processes was taking place. More evidence of 2He was found in 2008 at the Istituto Nazionale di Fisica Nucleare, a beam of 20Ne ions was directed at a target of beryllium foil. This collision converted some of the heavier neon nuclei in the beam into 18Ne nuclei and these nuclei then collided with a foil of lead. The second collision had the effect of exciting the 18Ne nucleus into an unstable condition. As in the experiment at Oak Ridge, the 18Ne nucleus decayed into an 16O nucleus, plus two protons detected exiting from the same direction
24.
Hydrogen
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Hydrogen is a chemical element with chemical symbol H and atomic number 1. With a standard weight of circa 1.008, hydrogen is the lightest element on the periodic table. Its monatomic form is the most abundant chemical substance in the Universe, non-remnant stars are mainly composed of hydrogen in the plasma state. The most common isotope of hydrogen, termed protium, has one proton, the universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, odorless, tasteless, non-toxic, nonmetallic, since hydrogen readily forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays an important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a charge when it is known as a hydride. The hydrogen cation is written as though composed of a bare proton, Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. Industrial production is mainly from steam reforming natural gas, and less often from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production, mostly for the fertilizer market, Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks. Hydrogen gas is flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol,2 H2 + O2 →2 H2O +572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74%, the explosive reactions may be triggered by spark, heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C, the detection of a burning hydrogen leak may require a flame detector, such leaks can be very dangerous. Hydrogen flames in other conditions are blue, resembling blue natural gas flames, the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a mixture of hydrogen to oxygen combined with carbon compounds from the airship skin. H2 reacts with every oxidizing element, the ground state energy level of the electron in a hydrogen atom is −13.6 eV, which is equivalent to an ultraviolet photon of roughly 91 nm wavelength. The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, however, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity. The most complicated treatments allow for the effects of special relativity
25.
Helium
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Helium is a chemical element with symbol He and atomic number 2. It is a colorless, odorless, tasteless, non-toxic, inert, monatomic gas and its boiling point is the lowest among all the elements. Its abundance is similar to figure in the Sun and in Jupiter. This is due to the high nuclear binding energy of helium-4 with respect to the next three elements after helium. This helium-4 binding energy also accounts for why it is a product of nuclear fusion and radioactive decay. Most helium in the universe is helium-4, and is believed to have formed during the Big Bang. Large amounts of new helium are being created by fusion of hydrogen in stars. Helium is named for the Greek god of the Sun, Helios and it was first detected as an unknown yellow spectral line signature in sunlight during a solar eclipse in 1868 by French astronomer Jules Janssen. Janssen is jointly credited with detecting the element along with Norman Lockyer, Janssen observed during the solar eclipse of 1868 while Lockyer observed from Britain. Lockyer was the first to propose that the line was due to a new element, the formal discovery of the element was made in 1895 by two Swedish chemists, Per Teodor Cleve and Nils Abraham Langlet, who found helium emanating from the uranium ore cleveite. In 1903, large reserves of helium were found in gas fields in parts of the United States. Liquid helium is used in cryogenics, particularly in the cooling of superconducting magnets, a well-known but minor use is as a lifting gas in balloons and airships. As with any gas whose density differs from that of air, inhaling a small volume of helium temporarily changes the timbre, on Earth it is relatively rare—5.2 ppm by volume in the atmosphere. Most terrestrial helium present today is created by the radioactive decay of heavy radioactive elements. Previously, terrestrial helium—a non-renewable resource, because once released into the atmosphere it readily escapes into space—was thought to be in short supply. The first evidence of helium was observed on August 18,1868, the line was detected by French astronomer Jules Janssen during a total solar eclipse in Guntur, India. This line was assumed to be sodium. He concluded that it was caused by an element in the Sun unknown on Earth, Lockyer and English chemist Edward Frankland named the element with the Greek word for the Sun, ἥλιος
26.
Isotopes of lithium
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Naturally occurring lithium is composed of two stable isotopes, lithium-6 and lithium-7, with the latter being far more abundant, about 92.5 percent of the atoms. Both of the isotopes have an unexpectedly low nuclear binding energy per nucleon when compared with the adjacent lighter and heavier elements. The longest-lived radioisotope of lithium is lithium-8, which has a half-life of just 838 milliseconds, lithium-9 has a half-life of 178 milliseconds, and lithium-11 has a half-life of about 8.6 milliseconds. All of the isotopes of lithium have half-lives that are shorter than 10 nanoseconds. Lithium-7 and lithium-6 are two of the primordial nuclides that were produced in the Big Bang, with lithium-7 to be 10−9 of all primordial nuclides, a small percentage of lithium-6 is also known to be produced by nuclear reactions in certain stars. The isotopes of lithium separate somewhat during a variety of geological processes, lithium ions replace magnesium or iron in certain octahedral locations in clays, and lithium-6 is sometimes preferred over lithium-7. This results in some enrichment of lithium-7 in geological processes, lithium-6 is an important isotope in nuclear physics because when it is bombarded with neutrons, tritium is produced. Lithium-6 has a greater affinity than lithium-7 for the element mercury, when an amalgam of lithium and mercury is added to solutions containing lithium hydroxide, the lithium-6 becomes more concentrated in the amalgam and the lithium-7 more in the hydroxide solution. The colex separation method makes use of this by passing a counter-flow of amalgam, the fraction of lithium-6 is preferentially drained by the mercury, but the lithium-7 flows mostly with the hydroxide. At the bottom of the column, the lithium is separated from the amalgam, at the top, the lithium hydroxide solution is electrolyzed to liberate the lithium-7 fraction. The enrichment obtained with this method varies with the column length, lithium is heated to a temperature of about 550 °C in a vacuum. Lithium atoms evaporate from the surface and are collected on a cold surface positioned a few centimetres above the liquid surface. Since lithium-6 atoms have a mean free path, they are collected preferentially. The theoretical separation efficiency is about 8.0 percent, a multistage process may be used to obtain higher degrees of separation. Lithium-4 contains three protons and one neutron and it is the shortest-lived known isotope of lithium, with a half-life of about 9. 1×10−23 seconds and decays by proton emission to helium-3. Lithium-4 can be formed as an intermediate in some nuclear fusion reactions, lithium-6 is valuable as the source material for the production of tritium and as an absorber of neutrons in nuclear fusion reactions. Natural lithium contains about 7.5 percent lithium-6, with the rest being lithium-7, large amounts of lithium-6 have been separated out for placing into hydrogen bombs. The separation of lithium-6 has by now ceased in the large thermonuclear powers, lithium-6 is one of only three stable isotopes with a spin of 1 and has the smallest nonzero nuclear electric quadrupole moment of any stable nucleus
27.
Isotopes of beryllium
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Beryllium has 12 known isotopes, but only one of these isotopes is stable and a primordial nuclide. As such, beryllium is considered a monoisotopic element and it is also a mononuclidic element, because its other isotopes have such short half-lives that none are primordial and their abundance is very low. Beryllium is unique as being the only element with both an even number of protons and an odd number of neutrons. There are 25 other monoisotopic elements but all have odd atomic numbers, of the 11 radioisotopes of beryllium, the most stable are 10Be with a half-life of 1.39 million years and 7Be with a half-life of 53.22 days. All other radioisotopes have half-lives under 13.85 seconds, most under 0.03 seconds, the least stable isotope is 6Be, with a half-life measured as 5.03 × 10−21 seconds. The half-life for the decay of 8Be is only 6. 7×10−17 seconds, beryllium is prevented from having a stable isotope with 4 protons and 6 neutrons by the very large mismatch in proton/neutron ratio for such a light element. Nevertheless, this isotope, 10Be, has a half-life of 1.39 million years, still other possible beryllium isotopes have even more severe mismatches in neutron and proton number, and thus are even less stable. Most 9Be in the universe is thought to be formed by cosmic ray nucleosynthesis from cosmic ray spallation in the period between the Big Bang and the formation of the solar system. The isotopes 7Be, with a half-life of 53.22 days and these two radioisotopes of beryllium in the atmosphere track the sun spot cycle and solar activity, since this affects the magnetic field that shields the Earth from cosmic rays. The rate at which the short-lived 7Be is transferred from the air to the ground is controlled in part by the weather, 7Be decay in the sun is one of the sources of solar neutrinos, and the first type ever detected using the Homestake experiment. Presence of 7Be in sediments is often used to establish that they are fresh, i. e. less than about 3–4 months in age, values marked # are not purely derived from experimental data, but at least partly from systematic trends. Spins with weak assignment arguments are enclosed in parentheses, most isotopes of beryllium within the proton/neutron drip lines decay via beta decay and/or a combination of beta decay and alpha decay or neutron emission. However, 7Be decays only via electron capture, a phenomenon to which its unusually long half-life may be attributed, also anomalous is 8Be, which decays via alpha decay to 4He. This alpha decay is often considered fission, which would be able to account for its extremely short half-life, the NUBASE evaluation of nuclear and decay properties. Isotopic compositions and standard atomic masses from, J. R. de Laeter, J. K. Böhlke, P. De Bièvre, H. Hidaka, H. S. Peiser, K. J. R. Rosman, atomic weights of the elements 2005. Half-life, spin, and isomer data selected from the following sources, see editing notes on this articles talk page. G. Audi, A. H. Wapstra, C, the NUBASE evaluation of nuclear and decay properties. National Nuclear Data Center, Brookhaven National Laboratory, information extracted from the NuDat 2.1 database
28.
Isotopes of boron
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Boron naturally occurs as isotopes 10B and 11B, the latter of which makes up about 80% of natural boron.2 ms. All other isotopes have half-lives shorter than 17.35 ms, with the least stable isotope being 7B and those isotopes with mass below 10 decay into helium while those with mass above 11 mostly become carbon. The precision of the isotope abundances and atomic mass is limited through variations, the given ranges should be applicable to any normal terrestrial material. Commercially available materials may have been subjected to an undisclosed or inadvertent isotopic fractionation, substantial deviations from the given mass and composition can occur. Values marked # are not purely derived from data. Spins with weak assignment arguments are enclosed in parentheses, uncertainties are given in concise form in parentheses after the corresponding last digits. Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC, nuclide masses are given by IUPAP Commission on Symbols, Units, Nomenclature, Atomic Masses and Fundamental Constants. Isotope abundances are given by IUPAC Commission on Isotopic Abundances and Atomic Weights, boron-10 is used in boron neutron capture therapy as an experimental treatment of some brain cancers. Isotope masses from, G. Audi, A. H. Wapstra, the NUBASE evaluation of nuclear and decay properties. Isotopic compositions and standard atomic masses from, J. R. de Laeter, J. K. Böhlke, P. De Bièvre, H. Hidaka, H. S. Peiser, K. J. R. Rosman, Atomic weights of the elements 2005. Half-life, spin, and isomer data selected from the following sources, see editing notes on this articles talk page. G. Audi, A. H. Wapstra, C, the NUBASE evaluation of nuclear and decay properties. CRC Handbook of Chemistry and Physics
29.
Lithium
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Lithium is a chemical element with the symbol Li and atomic number 3. It is a soft, silver-white metal belonging to the metal group of chemical elements. Under standard conditions, it is the lightest metal and the least dense solid element, like all alkali metals, lithium is highly reactive and flammable. For this reason, it is stored in mineral oil. When cut open, it exhibits a metallic luster, but contact with moist air corrodes the surface quickly to a silvery gray. Because of its reactivity, lithium never occurs freely in nature, and instead, appears only in compounds. Lithium occurs in a number of minerals, but due to its solubility as an ion, is present in ocean water and is commonly obtained from brines. On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride, the nucleus of the lithium atom verges on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative instability, lithium is less common in the solar system than 25 of the first 32 chemical elements even though the nuclei are very light in atomic weight. For related reasons, lithium has important links to nuclear physics, the transmutation of lithium atoms to helium in 1932 was the first fully man-made nuclear reaction, and lithium-6 deuteride serves as a fusion fuel in staged thermonuclear weapons. These uses consume more than three quarters of lithium production, Lithium is found in variable amounts in foods, primary food sources are grains and vegetables, in some areas, the drinking water also provides significant amounts of the element. Human dietary lithium intakes depend on location and the type of foods consumed, traces of lithium were detected in human organs and fetal tissues already in the late 19th century, leading to early suggestions as to possible specific functions in the organism. However, it took another century until evidence for the essentiality of lithium became available, in studies conducted from the 1970s to the 1990s, rats and goats maintained on low-lithium rations were shown to exhibit higher mortalities as well as reproductive and behavioral abnormalities. Lithium appears to play an important role during the early fetal development as evidenced by the high lithium contents of the embryo during the early gestational period. The available experimental evidence now appears to be sufficient to accept lithium as essential, the lithium ion Li+ administered as any of several lithium salts has proven to be useful as a mood-stabilizing drug in the treatment of bipolar disorder in humans. Like the other metals, lithium has a single valence electron that is easily given up to form a cation. Because of this, lithium is a conductor of heat and electricity as well as a highly reactive element. Lithiums low reactivity is due to the proximity of its electron to its nucleus
30.
Beryllium
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Beryllium is a chemical element with symbol Be and atomic number 4. It is a rare element in the universe, usually occurring as a product of the spallation of larger atomic nuclei that have collided with cosmic rays. Within the cores of stars beryllium is depleted as it is fused and it is a divalent element which occurs naturally only in combination with other elements in minerals. Notable gemstones which contain beryllium include beryl and chrysoberyl, as a free element it is a steel-gray, strong, lightweight and brittle alkaline earth metal. Beryllium improves many physical properties when added as an element to aluminium, copper, iron. Beryllium does not form oxides until it reaches high temperatures. Tools made of copper alloys are strong and hard and do not create sparks when they strike a steel surface. The high thermal conductivities of beryllium and beryllium oxide have led to their use in thermal management applications, Beryllium is a steel gray and hard metalloid that is brittle at room temperature and has a close-packed hexagonal crystal structure. It has exceptional stiffness and a high melting point. The modulus of elasticity of beryllium is approximately 50% greater than that of steel, the combination of this modulus and a relatively low density results in an unusually fast sound conduction speed in beryllium – about 12.9 km/s at ambient conditions. Other significant properties are high specific heat and thermal conductivity, which make beryllium the metal with the best heat dissipation characteristics per unit weight, in combination with the relatively low coefficient of linear thermal expansion, these characteristics result in a unique stability under conditions of thermal loading. Naturally occurring beryllium, save for slight contamination by cosmogenic radioisotopes, is essentially pure beryllium-9, Beryllium has a large scattering cross section for high-energy neutrons, about 6 barns for energies above approximately 10 keV. The single primordial beryllium isotope 9Be also undergoes a neutron reaction with neutron energies over about 1.9 MeV, to produce 8Be, thus, for high-energy neutrons, beryllium is a neutron multiplier, releasing more neutrons than it absorbs. Beryllium also releases neutrons under bombardment by gamma rays.8 seconds, β− is an electron, tritium is a radioisotope of concern in nuclear reactor waste streams. As a metal, beryllium is transparent to most wavelengths of X-rays and gamma rays, making it useful for the windows of X-ray tubes. Both stable and unstable isotopes of beryllium are created in stars, primordial beryllium contains only one stable isotope, 9Be, and therefore beryllium is a monoisotopic element. Radioactive cosmogenic 10Be is produced in the atmosphere of the Earth by the cosmic ray spallation of oxygen, 10Be accumulates at the soil surface, where its relatively long half-life permits a long residence time before decaying to boron-10. The production of 10Be is inversely proportional to activity, because increased solar wind during periods of high solar activity decreases the flux of galactic cosmic rays that reach the Earth
31.
Boron
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Boron is a chemical element with symbol B and atomic number 5. Produced entirely by cosmic ray spallation and supernovae and not by stellar nucleosynthesis, it is an element in the Solar system. Boron is concentrated on Earth by the water-solubility of its more common naturally occurring compounds and these are mined industrially as evaporites, such as borax and kernite. The largest known deposits are in Turkey, the largest producer of boron minerals. Elemental boron is a metalloid that is found in small amounts in meteoroids, industrially, very pure boron is produced with difficulty because of refractory contamination by carbon or other elements. Several allotropes of boron exist, amorphous boron is a powder, crystalline boron is silvery to black, extremely hard. The primary use of boron is as boron filaments with applications similar to carbon fibers in some high-strength materials. Boron is primarily used in chemical compounds, about half of all consumption globally, boron is used as an additive in glass fibers of boron-containing fiberglass for insulation and structural materials. The next leading use is in polymers and ceramics in high-strength, lightweight structural, borosilicate glass is desired for its greater strength and thermal shock resistance than ordinary soda lime glass. Boron compounds are used as fertilizers in agriculture and in sodium perborate bleaches, a small amount of boron is used as a dopant in semiconductors, and reagent intermediates in the synthesis of organic fine chemicals. A few boron-containing organic pharmaceuticals are used or are in study, natural boron is composed of two stable isotopes, one of which has a number of uses as a neutron-capturing agent. In biology, borates have low toxicity in mammals, but are toxic to arthropods and are used as insecticides. Boric acid is mildly antimicrobial, and several natural boron-containing organic antibiotics are known, small amounts of boron compounds play a strengthening role in the cell walls of all plants, making boron a necessary plant nutrient. Boron is involved in the metabolism of calcium in both plants and animals and it is considered an essential nutrient for humans, and boron deficiency is implicated in osteoporosis. The word boron was coined from borax, the mineral from which it was isolated, by analogy with carbon, marco Polo brought some glazes back to Italy in the 13th century. Agricola, around 1600, reports the use of borax as a flux in metallurgy, in 1777, boric acid was recognized in the hot springs near Florence, Italy, and became known as sal sedativum, with primarily medical uses. The rare mineral is called sassolite, which is found at Sasso, Sasso was the main source of European borax from 1827 to 1872, when American sources replaced it. Boron compounds were relatively rarely used until the late 1800s when Francis Marion Smiths Pacific Coast Borax Company first popularized and produced them in volume at low cost
32.
Deuterium
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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 hydrogen isotope, Deuterium has a natural abundance in Earths oceans of about one atom in 6420 of hydrogen. Thus deuterium accounts for approximately 0. 0156% of all the naturally occurring hydrogen in the oceans, the abundance of deuterium changes slightly from one kind of natural water to another. The deuterium isotopes name is formed from the Greek deuteros meaning second, Deuterium was discovered and named in 1931 by Harold Urey. When the neutron was discovered in 1932, this made the structure of deuterium obvious. Soon after deuteriums discovery, Urey and others produced samples of water in which the deuterium content had been highly concentrated. Deuterium is destroyed in the interiors of stars faster than it is produced, 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 and this is the ratio found in the gas giant planets, such as Jupiter. However, other bodies are found to have different ratios of deuterium to hydrogen-1. This is thought to be as a result of natural isotope separation processes that occur from solar heating of ices in comets, like the water-cycle in Earths weather, such heating processes may enrich deuterium with respect to protium. The analysis of ratios in comets found results very similar to the mean ratio in Earths oceans. This reinforces theories that much of Earths ocean water is of cometary origin, 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 frequently represented by the chemical symbol D, since it is an isotope of hydrogen with mass number 2, it is also represented by 2H. IUPAC allows both D and 2H, although 2H is preferred, a distinct chemical symbol is used for convenience because of the isotopes common use in various scientific processes. In quantum mechanics the energy levels of electrons in atoms depend on the mass of the system of electron. For hydrogen, this amount is about 1837/1836, or 1.000545, the energies of spectroscopic lines for deuterium and light-hydrogen therefore differ by the ratios of these two numbers, which is 1.000272. The wavelengths of all deuterium spectroscopic lines are shorter than the lines of light hydrogen
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Helium-3
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Helium-3 is a light, non-radioactive isotope of helium with two protons and one neutron. Its hypothetical existence was first proposed in 1934 by the Australian nuclear physicist Mark Oliphant while he was working at the University of Cambridge Cavendish Laboratory, Oliphant had performed experiments in which fast deuterons collided with deuteron targets. Helium-3 was thought to be a radioactive isotope until it was found in samples of natural helium. Other than 1H, helium-3 is the stable isotope of any element with more protons than neutrons. Helium-3 occurs as a primordial nuclide, escaping from the Earths crust into the atmosphere, some of the helium-3 found in the terrestrial atmosphere is also a relic of atmospheric and underwater nuclear weapons testing. Most of this comes from the decay of tritium, which decays into helium-3 with a life of 12.3 years. Furthermore, some nuclear reactors periodically release some helium-3 and tritium into the atmosphere, the nuclear reactor disaster at Chernobyl released a huge amount of radioactive tritium into the atmosphere, and smaller accidents have caused smaller releases. Furthermore, significant amounts of tritium and helium-3 have been produced in national arsenal nuclear reactors by the irradiation of lithium-6. The tritium is used to boost nuclear weapons, and some of this inevitably escapes during its production, transportation, hence, helium-3 enters the atmosphere both through its direct release and through the radioactive decay of tritium. Because of its atomic mass of 3.02 atomic mass units, helium-3 has some physical properties different from those of helium-4. Because of the weak, induced dipole–dipole interaction between atoms, their macroscopic physical properties are mainly determined by their zero-point energy. Also, the properties of helium-3 cause it to have a higher zero-point energy than helium-4. This implies that helium-3 can overcome dipole–dipole interactions with less energy than helium-4 can. Its latent heat of vaporization is also considerably lower at 0.026 kilojoule per mole compared with the 0.0829 kilojoule per mole of helium-4, the appeal of helium-3 fusion stems from the aneutronic nature of its reaction products. The lone high-energy by-product, the proton, can be contained using electric and magnetic fields, the momentum energy of this proton will interact with the containing electromagnetic field, resulting in direct net electricity generation. Because of the higher Coulomb barrier, the temperatures required for 21H + 32He fusion are much higher than those of conventional D-T fusion. Moreover, since both reactants need to be mixed together to fuse, reactions between nuclei of the same reactant will occur, and the D-D reaction does produce a neutron. Reaction rates vary with temperature, but the D-3He reaction rate is never greater than 3.56 times the D-D reaction rate
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Carbon
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Carbon is a chemical element with symbol C and atomic number 6. It is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds, three isotopes occur naturally, 12C and 13C being stable, while 14C is a radioactive isotope, decaying with a half-life of about 5,730 years. Carbon is one of the few elements known since antiquity, Carbon is the 15th most abundant element in the Earths crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is the second most abundant element in the body by mass after oxygen. The atoms of carbon can bond together in different ways, termed allotropes of carbon, the best known are graphite, diamond, and amorphous carbon. The physical properties of carbon vary widely with the allotropic form, for example, graphite is opaque and black while diamond is highly transparent. Graphite is soft enough to form a streak on paper, while diamond is the hardest naturally occurring material known, graphite is a good electrical conductor while diamond has a low electrical conductivity. Under normal conditions, diamond, carbon nanotubes, and graphene have the highest thermal conductivities of all known materials, all carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form. They are chemically resistant and require high temperature to react even with oxygen, the most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and transition metal carbonyl complexes. The largest sources of carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil. For this reason, carbon has often referred to as the king of the elements. The allotropes of carbon graphite, one of the softest known substances, and diamond. It bonds readily with other small atoms including other carbon atoms, Carbon is known to form almost ten million different compounds, a large majority of all chemical compounds. Carbon also has the highest sublimation point of all elements, although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature. Carbon is the element, with a ground-state electron configuration of 1s22s22p2. Its first four ionisation energies,1086.5,2352.6,4620.5 and 6222.7 kJ/mol, are higher than those of the heavier group 14 elements. Carbons covalent radii are normally taken as 77.2 pm,66.7 pm and 60.3 pm, although these may vary depending on coordination number, in general, covalent radius decreases with lower coordination number and higher bond order. Carbon compounds form the basis of all life on Earth
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Tritium
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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, naturally occurring tritium is extremely 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, the name of this isotope is derived from the Greek word τρίτος meaning third. While tritium has several different experimentally determined values of its half-life and 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 electrons kinetic energy varies, with an average of 5.7 keV, beta particles from tritium can penetrate only about 6.0 mm of air, and they are incapable of passing through the dead outermost layer of human skin. The unusually low energy released in the beta decay makes the decay appropriate for absolute neutrino mass measurements in the laboratory. The low energy of tritiums 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, and is an exothermic reaction yielding 4.8 MeV, in comparison, the fusion of deuterium with tritium releases about 17.6 MeV of energy. High-energy neutrons can produce tritium from lithium-7 in an endothermic reaction. This was discovered when the 1954 Castle Bravo nuclear test produced a high yield. High-energy neutrons irradiating boron-10 will also produce tritium, A more common result of boron-10 neutron capture is 7Li. The reactions requiring high neutron energies are not attractive production methods for peaceful applications, Tritium is also produced in heavy water-moderated reactors whenever a deuterium nucleus captures a neutron. This reaction has a quite small absorption cross section, making water a good neutron moderator. Even so, cleaning tritium from the moderator may be desirable after several years to reduce the risk of its escaping to the environment. Ontario Power Generations Tritium Removal Facility processes up to 2,500 tonnes of water a year. Deuteriums 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 a product of the nuclear fission of uranium-235, plutonium-239. The release or recovery of tritium needs to be considered in the operation of reactors, especially in the reprocessing of nuclear fuels
36.
Helium-4
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Helium-4 is a non-radioactive isotope of the element helium. It is by far the most abundant of the two naturally occurring isotopes of helium, making up about 99. 99986% of the helium on Earth and its nucleus is identical to an alpha particle, and consists of two protons and two neutrons. Alpha decay of elements in the Earths crust is the source of most naturally occurring helium-4 on Earth. While it is produced by nuclear fusion in stars, most helium-4 in the Sun and in the universe is thought to have been produced by the Big Bang. However, primordial helium-4 is largely absent from the Earth, having escaped during the phase of Earths formation. Radioactive decay from other elements is the source of most of the helium-4 found on Earth, helium-4 makes up about one quarter of the ordinary matter in the universe by mass, with almost all of the rest being hydrogen. When liquid helium-4 is cooled to below 2.17 kelvins it becomes a superfluid, for example, if superfluid helium-4 is kept in an open vessel, a thin film will climb up the sides of the vessel and overflow. In this state and situation, it is called a Rollin film, the total spin of the helium-4 nucleus is an integer, and therefore it is a boson. The superfluid behavior is now understood to be a manifestation of Bose–Einstein condensation and it is theorized that, at 0.2 K and 50 Atm, solid helium-4 may be a superglass. Helium-4 also exists on the Moon and—as on Earth—it is the most abundant helium isotope, the helium atom is the second simplest atom, but the extra electron introduces a third body, so the solution to its wave equation becomes a three-body problem, which has no analytic solution. However, numerical approximations of the equations of mechanics have given a good estimate of the key atomic properties of helium-4, such as its size. The nucleus of the atom is identical with an alpha particle. High-energy electron-scattering experiments show its charge to decrease exponentially from a maximum at a central point, adding another of any of these particles would require angular momentum, and would release substantially less energy. This arrangement is thus energetically extremely stable for all particles. For example, the stability and low energy of the cloud of helium causes heliums chemical inertness. Some stable helium-3 is produced in fusion reactions from hydrogen, but it is a small fraction. The stability of helium-4 is the reason that hydrogen is converted to helium-4, the unusual stability of the helium-4 nucleus is also important cosmologically. The energy of nuclear binding per nucleon is stronger than in any of these elements, and thus no energetic drive was available to make elements 3,4
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Nitrogen
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Nitrogen is a chemical element with symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772, although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. Nitrogen is the lightest member of group 15 of the periodic table, the name comes from the Greek πνίγειν to choke, directly referencing nitrogens asphyxiating properties. It is an element in the universe, estimated at about seventh in total abundance in the Milky Way. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2, dinitrogen forms about 78% of Earths atmosphere, making it the most abundant uncombined element. Nitrogen occurs in all organisms, primarily in amino acids, in the nucleic acids, the human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, many industrially important compounds, such as ammonia, nitric acid, organic nitrates, and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen, the second strongest bond in any diatomic molecule. Synthetically produced ammonia and nitrates are key industrial fertilisers, and fertiliser nitrates are key pollutants in the eutrophication of water systems. Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric, Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. Many notable nitrogen-containing drugs, such as the caffeine and morphine or the synthetic amphetamines. Nitrogen compounds have a long history, ammonium chloride having been known to Herodotus. They were well known by the Middle Ages, alchemists knew nitric acid as aqua fortis, as well as other nitrogen compounds such as ammonium salts and nitrate salts. The mixture of nitric and hydrochloric acids was known as aqua regia, celebrated for its ability to dissolve gold, the discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air. Though he did not recognise it as a different chemical substance, he clearly distinguished it from Joseph Blacks fixed air. The fact that there was a component of air that does not support combustion was clear to Rutherford, Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as air or azote, from the Greek word άζωτικός. In an atmosphere of nitrogen, animals died and flames were extinguished
38.
Isotopes of oxygen
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There are three stable isotopes of oxygen, 16O, 17O, and 18O. Radioactive isotopes with mass numbers from 12O to 24O have also characterized, all short-lived. The shortest-lived is 12O with a half-life of 580×10−24 second, naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant. Depending on the source, the standard atomic weight varies within the range of. Known oxygen isotopes range in number from 12 to 24. Most 16O is synthesized at the end of the fusion process in stars, the triple-alpha reaction creates 12C. The neon burning process creates additional 16O, both 17O and 18O are secondary isotopes, meaning that their nucleosynthesis requires seed nuclei. 17O is primarily made by the burning of hydrogen into helium during the CNO cycle, most 18O is produced when 14N captures a 4He nucleus, making 18O common in the helium-rich zones of stars. Approximately a billion degrees Celsius is required for two oxygen nuclei to undergo nuclear fusion to form the nucleus of sulfur. Measurements of the ratio of oxygen-18 to oxygen-16 are often used to interpret changes in paleoclimate, the isotopic composition of oxygen atoms in the Earths atmosphere is 99. 759% 16O,0. 037% 17O and 0. 204% 18O. This disparity allows analysis of temperature patterns via historic ice cores, an atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C. Since physicists referred to 16O only, while chemists meant the naturally-abundant mixture of isotopes, fourteen radioisotopes have been characterized, with the most stable being 15O with a half-life of 122.24 s and 14O with a half-life of 70.606 s. All of the radioactive isotopes have half-lives that are less than 27 s. For example, 24O has a half-life of 61 ms, the most common decay mode for isotopes lighter than the stable isotopes is β+ decay and the most common mode after is β− decay. Oxygen-13 is an isotope of oxygen. It consists of 8 protons and electrons, and 5 neutrons and it has a spin of 3/2-, and a half-life of 8.58 ms. Its atomic mass is 13.0248 Da and it decays to nitrogen-13 by electron capture, and has a decay energy of 17.765 MeV. Oxygen-15 is an isotope of oxygen, frequently used in positron emission tomography and it has 8 protons,7 neutrons, and 8 electrons
39.
Oxygen
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Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the group on the periodic table and is a highly reactive nonmetal. By mass, oxygen is the third-most abundant element in the universe, after hydrogen, at standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. This is an important part of the atmosphere and diatomic oxygen gas constitutes 20. 8% of the Earths atmosphere, additionally, as oxides the element makes up almost half of the Earths crust. Most of the mass of living organisms is oxygen as a component of water, conversely, oxygen is continuously replenished by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone, strongly absorbs ultraviolet UVB radiation, but ozone is a pollutant near the surface where it is a by-product of smog. At low earth orbit altitudes, sufficient atomic oxygen is present to cause corrosion of spacecraft, the name oxygen was coined in 1777 by Antoine Lavoisier, whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle, Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philos work by observing that a portion of air is consumed during combustion and respiration, Oxygen was discovered by the Polish alchemist Sendivogius, who considered it the philosophers stone. In the late 17th century, Robert Boyle proved that air is necessary for combustion, English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. From this he surmised that nitroaereus is consumed in both respiration and combustion, Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract De respiratione. Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element. This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, one part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. The fact that a substance like wood gains overall weight in burning was hidden by the buoyancy of the combustion products
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