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id, such as hydrochloric acid. Astatine is extracted from this acid using an organic solvent such as butyl or isopropyl ether, diisopropylether DIPE, or thiosemicarbazide. Using liquidliquid extraction, the astatine product can be repeatedly washed with an acid, such as HCl, and extracted into the organic solvent layer. A separation yield of 93 using nitric acid has been reported, falling to 72 by the time purification procedures were completed distillation of nitric acid, purging residual nitrogen oxides, and redissolving bismuth nitrate to enable liquidliquid extraction. Wet methods involve "multiple radioactivity handling steps" and have not been considered well suited for isolating larger quantities of astatine. However, wet extraction methods are being examined for use in production of larger quantities of astatine211, as it is thought that wet extraction methods can provide more consistency. They can enable the production of astatine in a specific oxidation state and may have greater applicability in ex
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perimental radiochemistry.
Uses and precautions
class"wikitable"
Several 211Atcontaining molecules and their experimental uses
! Agent
! Applications
211Atastatinetellurium colloids
Compartmental tumors
6211Atastato2methyl1,4naphtaquinol diphosphate
Adenocarcinomas
211Atlabeled methylene blue
Melanomas
Meta211Atastatobenzyl guanidine
Neuroendocrine tumors
5211Atastato2'deoxyuridine
Various
211Atlabeled biotin conjugates
Various pretargeting
211Atlabeled octreotide
Somatostatin receptor
211Atlabeled monoclonal antibodies and fragments
Various
211Atlabeled bisphosphonates
Bone metastases
Newly formed astatine211 is the subject of ongoing research in nuclear medicine. It must be used quickly as it decays with a halflife of 7.2 hours; this is long enough to permit multistep labeling strategies. Astatine211 has potential for targeted alphaparticle therapy, since it decays either via emission of an alpha particle to bismuth207, or via electron capture to an extremely shortlived nucli
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de, polonium211, which undergoes further alpha decay, very quickly reaching its stable granddaughter lead207. Polonium Xrays emitted as a result of the electron capture branch, in the range of 7792 keV, enable the tracking of astatine in animals and patients. Although astatine210 has a slightly longer halflife, it is wholly unsuitable because it usually undergoes beta plus decay to the extremely toxic polonium210.
The principal medicinal difference between astatine211 and iodine131 a radioactive iodine isotope also used in medicine is that iodine131 emits highenergy beta particles, and astatine does not. Beta particles have much greater penetrating power through tissues than do the much heavier alpha particles. An average alpha particle released by astatine211 can travel up to 70 m through surrounding tissues; an averageenergy beta particle emitted by iodine131 can travel nearly 30 times as far, to about 2 mm. The short halflife and limited penetrating power of alpha radiation through tissues offers advantag
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es in situations where the "tumor burden is low andor malignant cell populations are located in close proximity to essential normal tissues." Significant morbidity in cell culture models of human cancers has been achieved with from one to ten astatine211 atoms bound per cell.
Several obstacles have been encountered in the development of astatinebased radiopharmaceuticals for cancer treatment. World War II delayed research for close to a decade. Results of early experiments indicated that a cancerselective carrier would need to be developed and it was not until the 1970s that monoclonal antibodies became available for this purpose. Unlike iodine, astatine shows a tendency to dehalogenate from molecular carriers such as these, particularly at sp3 carbon sites less so from sp2 sites. Given the toxicity of astatine accumulated and retained in the body, this emphasized the need to ensure it remained attached to its host molecule. While astatine carriers that are slowly metabolized can be assessed for their effica
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cy, more rapidly metabolized carriers remain a significant obstacle to the evaluation of astatine in nuclear medicine. Mitigating the effects of astatineinduced radiolysis of labeling chemistry and carrier molecules is another area requiring further development. A practical application for astatine as a cancer treatment would potentially be suitable for a "staggering" number of patients; production of astatine in the quantities that would be required remains an issue.
Animal studies show that astatine, similarly to iodine although to a lesser extent, perhaps because of its slightly more metallic nature is preferentially and dangerously concentrated in the thyroid gland. Unlike iodine, astatine also shows a tendency to be taken up by the lungs and spleen, possibly because of inbody oxidation of At to At. If administered in the form of a radiocolloid it tends to concentrate in the liver. Experiments in rats and monkeys suggest that astatine211 causes much greater damage to the thyroid gland than does iodine
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131, with repetitive injection of the nuclide resulting in necrosis and cell dysplasia within the gland. Early research suggested that injection of astatine into female rodents caused morphological changes in breast tissue; this conclusion remained controversial for many years. General agreement was later reached that this was likely caused by the effect of breast tissue irradiation combined with hormonal changes due to irradiation of the ovaries. Trace amounts of astatine can be handled safely in fume hoods if they are wellaerated; biological uptake of the element must be avoided.
See also
Radiation protection
Notes
References
Bibliography
External links
Astatine at The Periodic Table of Videos University of Nottingham
Astatine Halogen or Metal?
Halogens
Metalloids
Chemical elements
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An atom is the smallest unit of ordinary matter that forms a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms are extremely small, typically around 100 picometers across. They are so small that accurately predicting their behavior using classical physicsas if they were tennis balls, for exampleis not possible due to quantum effects.
Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and a number of neutrons. Only the most common variety of hydrogen has no neutrons. More than 99.94 of an atom's mass is in the nucleus. The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. If the number of protons and electrons are equal, then the atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively such atoms are called ions.
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The electrons of an atom are attracted to the protons in an atomic nucleus by the electromagnetic force. The protons and neutrons in the nucleus are attracted to each other by the nuclear force. This force is usually stronger than the electromagnetic force that repels the positively charged protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus splits and leaves behind different elements. This is a form of nuclear decay.
The number of protons in the nucleus is the atomic number and it defines to which chemical element the atom belongs. For example, any atom that contains 29 protons is copper. The number of neutrons defines the isotope of the element. Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to associate and dissociate is responsible for most of the physical changes observed in nature. Chemistry is the discipl
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ine that studies these changes.
History of atomic theory
In philosophy
The basic idea that matter is made up of tiny, indivisible particles appears in many ancient cultures such as those of Greece and India. The word atom is derived from the ancient Greek word atomos a combination of the negative term "a" and "," the term for "cut" that means "uncuttable". This ancient idea was based in philosophical reasoning rather than scientific reasoning; modern atomic theory is not based on these old concepts. Nonetheless, the term "atom" was used throughout the ages by thinkers who suspected that matter was ultimately granular in nature. It has since been discovered that "atoms" can be split, but the misnomer is still used.
Dalton's law of multiple proportions
In the early 1800s, the English chemist John Dalton compiled experimental data gathered by himself and other scientists and discovered a pattern now known as the "law of multiple proportions". He noticed that in chemical compounds which contain a particular
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chemical element, the content of that element in these compounds will differ by ratios of small whole numbers. This pattern suggested to Dalton that each chemical element combines with other elements by some basic and consistent unit of mass.
For example, there are two types of tin oxide one is a black powder that is 88.1 tin and 11.9 oxygen, and the other is a white powder that is 78.7 tin and 21.3 oxygen. Adjusting these figures, in the black oxide there is about 13.5 g of oxygen for every 100 g of tin, and in the white oxide there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 12. In these oxides, for every tin atom there are one or two oxygen atoms respectively SnO and SnO2.
As a second example, Dalton considered two iron oxides a black powder which is 78.1 iron and 21.9 oxygen, and a red powder which is 70.4 iron and 29.6 oxygen. Adjusting these figures, in the black oxide there is about 28 g of oxygen for every 100 g of iron, and in the red oxide there is about 42 g of oxy
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gen for every 100 g of iron. 28 and 42 form a ratio of 23. In these respective oxides, for every two atoms of iron, there are two or three atoms of oxygen Fe2O2 and Fe2O3.
As a final example nitrous oxide is 63.3 nitrogen and 36.7 oxygen, nitric oxide is 44.05 nitrogen and 55.95 oxygen, and nitrogen dioxide is 29.5 nitrogen and 70.5 oxygen. Adjusting these figures, in nitrous oxide there is 80 g of oxygen for every 140 g of nitrogen, in nitric oxide there is about 160 g of oxygen for every 140 g of nitrogen, and in nitrogen dioxide there is 320 g of oxygen for every 140 g of nitrogen. 80, 160, and 320 form a ratio of 124. The respective formulas for these oxides are N2O, NO, and NO2.
Kinetic theory of gases
In the late 18th century, a number of scientists found that they could better explain the behavior of gases by describing them as collections of submicroscopic particles and modelling their behavior using statistics and probability. Unlike Dalton's atomic theory, the kinetic theory of gases describes no
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t how gases react chemically with each other to form compounds, but how they behave physically diffusion, viscosity, conductivity, pressure, etc.
Brownian motion
In 1827, botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as "Brownian motion". This was thought to be caused by water molecules knocking the grains about. In 1905, Albert Einstein proved the reality of these molecules and their motions by producing the first statistical physics analysis of Brownian motion. French physicist Jean Perrin used Einstein's work to experimentally determine the mass and dimensions of molecules, thereby providing physical evidence for the particle nature of matter.
Discovery of the electron
In 1897, J. J. Thomson discovered that cathode rays are not electromagnetic waves but made of particles that are 1,800 times lighter than hydrogen the lightest atom. Thomson concluded that these particles came from the atoms
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within the cathode they were subatomic particles. He called these new particles corpuscles but they were later renamed electrons. Thomson also showed that electrons were identical to particles given off by photoelectric and radioactive materials. It was quickly recognized that electrons are the particles that carry electric currents in metal wires. Thomson concluded that these electrons emerged from the very atoms of the cathode in his instruments, which meant that atoms are not indivisible as the name atomos suggests.
Discovery of the nucleus
J. J. Thomson thought that the negativelycharged electrons were distributed throughout the atom in a sea of positive charge that was distributed across the whole volume of the atom. This model is sometimes known as the plum pudding model.
Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden came to have doubts about the Thomson model after they encountered difficulties when they tried to build an instrument to measure the chargetomass ratio of alpha
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particles these are positivelycharged particles emitted by certain radioactive substances such as radium. The alpha particles were being scattered by the air in the detection chamber, which made the measurements unreliable. Thomson had encountered a similar problem in his work on cathode rays, which he solved by creating a nearperfect vacuum in his instruments. Rutherford didn't think he'd run into this same problem because alpha particles are much heavier than electrons. According to Thomson's model of the atom, the positive charge in the atom is not concentrated enough to produce an electric field strong enough to deflect an alpha particle, and the electrons are so lightweight they should be pushed aside effortlessly by the much heavier alpha particles. Yet there was scattering, so Rutherford and his colleagues decided to investigate this scattering carefully.
Between 1908 and 1913, Rutheford and his colleagues performed a series of experiments in which they bombarded thin foils of metal with alpha partic
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les. They spotted alpha particles being deflected by angles greater than 90. To explain this, Rutherford proposed that the positive charge of the atom is not distributed throughout the atom's volume as Thomson believed, but is concentrated in a tiny nucleus at the center. Only such an intense concentration of charge could produce an electric field strong enough to deflect the alpha particles as observed.
Discovery of isotopes
While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table. The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J. J. Thomson created a technique for isotope separation through his work on ionized gases, which subsequently led to the discovery of stable isotopes.
Bohr model
In 1913, the physicist Niels Bohr proposed a model in which the electrons of an atom were assumed to orbit the
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nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. This quantization was used to explain why the electrons' orbits are stable given that normally, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation, see synchrotron radiation and why elements absorb and emit electromagnetic radiation in discrete spectra.
Later in the same year Henry Moseley provided additional experimental evidence in favor of Niels Bohr's theory. These results refined Ernest Rutherford's and Antonius van den Broek's model, which proposed that the atom contains in its nucleus a number of positive nuclear charges that is equal to its atomic number in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today
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.
Chemical bonds between atoms were explained by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons. As the chemical properties of the elements were known to largely repeat themselves according to the periodic law, in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus.
The Bohr model of the atom was the first complete physical model of the atom. It described the overall structure of the atom, how atoms bond to each other, and predicted the spectral lines of hydrogen. Bohr's model was not perfect and was soon superseded by the more accurate Schrdinger model, but it was sufficient to evaporate any remaining doubts that matter is composed of atoms. For chemists, the idea of the atom had been a useful heuristic tool, but physicists had doubts as to whether matter really is made up of atoms as
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nobody had yet developed a complete physical model of the atom.
The Schrdinger model
The SternGerlach experiment of 1922 provided further evidence of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom's angular momentum, or spin. As this spin direction is initially random, the beam would be expected to deflect in a random direction. Instead, the beam was split into two directional components, corresponding to the atomic spin being oriented up or down with respect to the magnetic field.
In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics matrix mechanics. One year earlier, Louis de Broglie had proposed the de Broglie hypothesis that all particles behave like waves to some extent, and in 1926 Erwin Schrdinger used this idea to develop the Schrdinger equation, a mathematical model of the atom wave mechanics that describ
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ed the electrons as threedimensional waveforms rather than point particles.
A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa.
This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed.
Discovery of the neutron
The development of the mass spectrometer allowed the mass of atoms to be measured with increased accuracy. The device uses a magnet to bend the
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trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom's mass to its charge. The chemist Francis William Aston used this instrument to show that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole number rule. The explanation for these different isotopes awaited the discovery of the neutron, an uncharged particle with a mass similar to the proton, by the physicist James Chadwick in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus.
Fission, highenergy physics and condensed matter
In 1938, the German chemist Otto Hahn, a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. Instead, his chemical experiments showed barium as a product. A year later, Lise Meitner and her nephew Otto Frisch verified that Hahn's result were the first experimental nuclear fission. In 1944, Hahn received the
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Nobel Prize in Chemistry. Despite Hahn's efforts, the contributions of Meitner and Frisch were not recognized.
In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks. The standard model of particle physics was developed that so far has successfully explained the properties of the nucleus in terms of these subatomic particles and the forces that govern their interactions.
Structure
Subatomic particles
Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton and the neutron.
The electron is by far the least massive of these particles at , with a negative electrical charge and a size that is too small to be measured using availa
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ble techniques. It was the lightest particle with a positive rest mass measured, until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details.
Protons have a positive charge and a mass 1,836 times that of the electron, at . The number of protons in an atom is called its atomic number. Ernest Rutherford 1919 observed that nitrogen under alphaparticle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton.
Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the
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electron, or . Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons collectively known as nucleons have comparable dimensionson the order of although the 'surface' of these particles is not sharply defined. The neutron was discovered in 1932 by the English physicist James Chadwick.
In the Standard Model of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks each with charge and one down quark with a charge of . Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles.
The quarks are held together by the strong interaction or strong force, which is mediated by gluons. The protons and neutro
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ns, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different rangeproperties see the article on the nuclear force for more. The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.
Nucleus
All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to femtometres, where is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a shortranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other.
Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons m
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ay vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay.
The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud.
A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mut
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ual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus.
The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3 to 10 keV to overcome their mutual repulsionthe coulomb barrierand fuse together into a single nucleus. Nuclear fission is the opposite process, causing a nucleus to split into two smaller nucleiusually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element.
If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, t
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hen the difference between these two values can be emitted as a type of usable energy such as a gamma ray, or the kinetic energy of a beta particle, as described by Albert Einstein's massenergy equivalence formula, , where is the mass loss and is the speed of light. This deficit is part of the binding energy of the new nucleus, and it is the nonrecoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate.
The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickela total nucleon number of about 60is usually an exothermic process that releases more energy than is required to bring them together. It is this energyreleasing process that makes nuclear fusion in stars a selfsustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about
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60, is an endothermic process. These more massive nuclei can not undergo an energyproducing fusion reaction that can sustain the hydrostatic equilibrium of a star.
Electron cloud
The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations.
Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of threedimensional standing wavea wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the pr
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obability that an electron appears to be at a particular location when its position is measured. Only a discrete or quantized set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation.
Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines.
The amount of energy needed to remove or add an electronthe electron binding energyis far less than the binding energy of nucleons. For exa
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mple, it requires only 13.6 eV to strip a groundstate electron from a hydrogen atom, compared to 2.23 million eV for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals.
Properties
Nuclear properties
By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons hydrogen1, by far the most common form, also called protium, one neutron deuterium, two neutrons tritium a
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nd more than two neutrons. The known elements form a set of atomic numbers, from the singleproton element hydrogen up to the 118proton element oganesson. All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 bismuth is so slight as to be practically negligible.
About 339 nuclides occur naturally on Earth, of which 252 about 74 have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stable theoretically, while another 162 bringing the total to 252 have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 34 radioactive nuclides have halflives longer than 100 million years, and are longlived enough to have been present since the birth of the Solar System. This collection of 286 nuclides are known as primordial nuclides. Finally, an additional 53 shortlived nuclides are known to occur naturally, as daughter product
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s of primordial nuclide decay such as radium from uranium, or as products of natural energetic processes on Earth, such as cosmic ray bombardment for example, carbon14.
For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.2 stable isotopes per element. Twentysix elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes.
Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 252 known stab
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le nuclides, only four have both an odd number of protons and odd number of neutrons hydrogen2 deuterium, lithium6, boron10 and nitrogen14. Also, only four naturally occurring, radioactive oddodd nuclides have a halflife over a billion years potassium40, vanadium50, lanthanum138 and tantalum180m. Most oddodd nuclei are highly unstable with respect to beta decay, because the decay products are eveneven, and are therefore more strongly bound, due to nuclear pairing effects.
Mass
The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles called "nucleons" in a given atom is called the mass number. It is a positive integer and dimensionless instead of having dimension of mass, because it expresses a count. An example of use of a mass number is "carbon12," which has 12 nucleons six protons and six neutrons.
The actual mass of an atom at rest is often expressed in daltons Da, also called the unified atomic mass unit u. This unit is defined as a t
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welfth of the mass of a free neutral atom of carbon12, which is approximately . Hydrogen1 the lightest isotope of hydrogen which is also the nuclide with the lowest mass has an atomic weight of 1.007825 Da. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal within 1 to its mass number times the atomic mass unit for example the mass of a nitrogen14 is roughly 14 Da, but this number will not be exactly an integer except by definition in the case of carbon12. The heaviest stable atom is lead208, with a mass of .
As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms about . This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon12 atom has an atomic mass of exactly 12 Da, and so a mole of
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carbon12 atoms weighs exactly 0.012 kg.
Shape and size
Atoms lack a welldefined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. This assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms coordination number and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows left to right. Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm.
When subjected to external forces, like electrical fields, the shape of an atom may
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deviate from spherical symmetry. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by grouptheoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystalelectrical fields may occur at lowsymmetry lattice sites. Significant ellipsoidal deformations have been shown to occur for sulfur ions and chalcogen ions in pyritetype compounds.
Atomic dimensions are thousands of times smaller than the wavelengths of light 400700 nm so they cannot be viewed using an optical microscope, although individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width. A single drop of water contains about 2 sextillion atoms of oxygen, and twice the number of hydrogen atoms. A single carat diamond with a mass of contains about 10 sextillion 1022 atoms of carbon. If an apple were magnified to the size of th
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e Earth, then the atoms in the apple would be approximately the size of the original apple.
Radioactive decay
Every element has one or more isotopes that have unstable nuclei that are subject to radioactive decay, causing the nucleus to emit particles or electromagnetic radiation. Radioactivity can occur when the radius of a nucleus is large compared with the radius of the strong force, which only acts over distances on the order of 1 fm.
The most common forms of radioactive decay are
Alpha decay this process is caused when the nucleus emits an alpha particle, which is a helium nucleus consisting of two protons and two neutrons. The result of the emission is a new element with a lower atomic number.
Beta decay and electron capture these processes are regulated by the weak force, and result from a transformation of a neutron into a proton, or a proton into a neutron. The neutron to proton transition is accompanied by the emission of an electron and an antineutrino, while proton to neutron transition excep
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t in electron capture causes the emission of a positron and a neutrino. The electron or positron emissions are called beta particles. Beta decay either increases or decreases the atomic number of the nucleus by one. Electron capture is more common than positron emission, because it requires less energy. In this type of decay, an electron is absorbed by the nucleus, rather than a positron emitted from the nucleus. A neutrino is still emitted in this process, and a proton changes to a neutron.
Gamma decay this process results from a change in the energy level of the nucleus to a lower state, resulting in the emission of electromagnetic radiation. The excited state of a nucleus which results in gamma emission usually occurs following the emission of an alpha or a beta particle. Thus, gamma decay usually follows alpha or beta decay.
Other more rare types of radioactive decay include ejection of neutrons or protons or clusters of nucleons from a nucleus, or more than one beta particle. An analog of gamma emissio
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n which allows excited nuclei to lose energy in a different way, is internal conversiona process that produces highspeed electrons that are not beta rays, followed by production of highenergy photons that are not gamma rays. A few large nuclei explode into two or more charged fragments of varying masses plus several neutrons, in a decay called spontaneous nuclear fission.
Each radioactive isotope has a characteristic decay time periodthe halflifethat is determined by the amount of time needed for half of a sample to decay. This is an exponential decay process that steadily decreases the proportion of the remaining isotope by 50 every halflife. Hence after two halflives have passed only 25 of the isotope is present, and so forth.
Magnetic moment
Elementary particles possess an intrinsic quantum mechanical property known as spin. This is analogous to the angular momentum of an object that is spinning around its center of mass, although strictly speaking these particles are believed to be pointlike and cannot
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be said to be rotating. Spin is measured in units of the reduced Planck constant , with electrons, protons and neutrons all having spin , or "spin". In an atom, electrons in motion around the nucleus possess orbital angular momentum in addition to their spin, while the nucleus itself possesses angular momentum due to its nuclear spin.
The magnetic field produced by an atomits magnetic momentis determined by these various forms of angular momentum, just as a rotating charged object classically produces a magnetic field, but the most dominant contribution comes from electron spin. Due to the nature of electrons to obey the Pauli exclusion principle, in which no two electrons may be found in the same quantum state, bound electrons pair up with each other, with one member of each pair in a spin up state and the other in the opposite, spin down state. Thus these spins cancel each other out, reducing the total magnetic dipole moment to zero in some atoms with even number of electrons.
In ferromagnetic elements
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such as iron, cobalt and nickel, an odd number of electrons leads to an unpaired electron and a net overall magnetic moment. The orbitals of neighboring atoms overlap and a lower energy state is achieved when the spins of unpaired electrons are aligned with each other, a spontaneous process known as an exchange interaction. When the magnetic moments of ferromagnetic atoms are lined up, the material can produce a measurable macroscopic field. Paramagnetic materials have atoms with magnetic moments that line up in random directions when no magnetic field is present, but the magnetic moments of the individual atoms line up in the presence of a field.
The nucleus of an atom will have no spin when it has even numbers of both neutrons and protons, but for other cases of odd numbers, the nucleus may have a spin. Normally nuclei with spin are aligned in random directions because of thermal equilibrium, but for certain elements such as xenon129 it is possible to polarize a significant proportion of the nuclear spin s
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tates so that they are aligned in the same directiona condition called hyperpolarization. This has important applications in magnetic resonance imaging.
Energy levels
The potential energy of an electron in an atom is negative relative to when the distance from the nucleus goes to infinity; its dependence on the electron's position reaches the minimum inside the nucleus, roughly in inverse proportion to the distance. In the quantummechanical model, a bound electron can occupy only a set of states centered on the nucleus, and each state corresponds to a specific energy level; see timeindependent Schrdinger equation for a theoretical explanation. An energy level can be measured by the amount of energy needed to unbind the electron from the atom, and is usually given in units of electronvolts eV. The lowest energy state of a bound electron is called the ground state, i.e. stationary state, while an electron transition to a higher level results in an excited state. The electron's energy increases along with n be
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cause the average distance to the nucleus increases. Dependence of the energy on is caused not by the electrostatic potential of the nucleus, but by interaction between electrons.
For an electron to transition between two different states, e.g. ground state to first excited state, it must absorb or emit a photon at an energy matching the difference in the potential energy of those levels, according to the Niels Bohr model, what can be precisely calculated by the Schrdinger equation.
Electrons jump between orbitals in a particlelike fashion. For example, if a single photon strikes the electrons, only a single electron changes states in response to the photon; see Electron properties.
The energy of an emitted photon is proportional to its frequency, so these specific energy levels appear as distinct bands in the electromagnetic spectrum. Each element has a characteristic spectrum that can depend on the nuclear charge, subshells filled by electrons, the electromagnetic interactions between the electrons and o
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ther factors.
When a continuous spectrum of energy is passed through a gas or plasma, some of the photons are absorbed by atoms, causing electrons to change their energy level. Those excited electrons that remain bound to their atom spontaneously emit this energy as a photon, traveling in a random direction, and so drop back to lower energy levels. Thus the atoms behave like a filter that forms a series of dark absorption bands in the energy output. An observer viewing the atoms from a view that does not include the continuous spectrum in the background, instead sees a series of emission lines from the photons emitted by the atoms. Spectroscopic measurements of the strength and width of atomic spectral lines allow the composition and physical properties of a substance to be determined.
Close examination of the spectral lines reveals that some display a fine structure splitting. This occurs because of spinorbit coupling, which is an interaction between the spin and motion of the outermost electron. When an a
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tom is in an external magnetic field, spectral lines become split into three or more components; a phenomenon called the Zeeman effect. This is caused by the interaction of the magnetic field with the magnetic moment of the atom and its electrons. Some atoms can have multiple electron configurations with the same energy level, which thus appear as a single spectral line. The interaction of the magnetic field with the atom shifts these electron configurations to slightly different energy levels, resulting in multiple spectral lines. The presence of an external electric field can cause a comparable splitting and shifting of spectral lines by modifying the electron energy levels, a phenomenon called the Stark effect.
If a bound electron is in an excited state, an interacting photon with the proper energy can cause stimulated emission of a photon with a matching energy level. For this to occur, the electron must drop to a lower energy state that has an energy difference matching the energy of the interacting pho
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ton. The emitted photon and the interacting photon then move off in parallel and with matching phases. That is, the wave patterns of the two photons are synchronized. This physical property is used to make lasers, which can emit a coherent beam of light energy in a narrow frequency band.
Valence and bonding behavior
Valency is the combining power of an element. It is determined by the number of bonds it can form to other atoms or groups. The outermost electron shell of an atom in its uncombined state is known as the valence shell, and the electrons in
that shell are called valence electrons. The number of valence electrons determines the bonding
behavior with other atoms. Atoms tend to chemically react with each other in a manner that fills or empties their outer valence shells. For example, a transfer of a single electron between atoms is a useful approximation for bonds that form between atoms with oneelectron more than a filled shell, and others that are oneelectron short of a full shell, such as occurs
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in the compound sodium chloride and other chemical ionic salts. Many elements display multiple valences, or tendencies to share differing numbers of electrons in different compounds. Thus, chemical bonding between these elements takes many forms of electronsharing that are more than simple electron transfers. Examples include the element carbon and the organic compounds.
The chemical elements are often displayed in a periodic table that is laid out to display recurring chemical properties, and elements with the same number of valence electrons form a group that is aligned in the same column of the table. The horizontal rows correspond to the filling of a quantum shell of electrons. The elements at the far right of the table have their outer shell completely filled with electrons, which results in chemically inert elements known as the noble gases.
States
Quantities of atoms are found in different states of matter that depend on the physical conditions, such as temperature and pressure. By varying the condi
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tions, materials can transition between solids, liquids, gases and plasmas. Within a state, a material can also exist in different allotropes. An example of this is solid carbon, which can exist as graphite or diamond. Gaseous allotropes exist as well, such as dioxygen and ozone.
At temperatures close to absolute zero, atoms can form a BoseEinstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale. This supercooled collection of atoms
then behaves as a single super atom, which may allow fundamental checks of quantum mechanical behavior.
Identification
While atoms are too small to be seen, devices such as the scanning tunneling microscope STM enable their visualization at the surfaces of solids. The microscope uses the quantum tunneling phenomenon, which allows particles to pass through a barrier that would be insurmountable in the classical perspective. Electrons tunnel through the vacuum between two biased el
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ectrodes, providing a tunneling current that is exponentially dependent on their separation. One electrode is a sharp tip ideally ending with a single atom. At each point of the scan of the surface the tip's height is adjusted so as to keep the tunneling current at a set value. How much the tip moves to and away from the surface is interpreted as the height profile. For low bias, the microscope images the averaged electron orbitals across closely packed energy levelsthe local density of the electronic states near the Fermi level. Because of the distances involved, both electrodes need to be extremely stable; only then periodicities can be observed that correspond to individual atoms. The method alone is not chemically specific, and cannot identify the atomic species present at the surface.
Atoms can be easily identified by their mass. If an atom is ionized by removing one of its electrons, its trajectory when it passes through a magnetic field will bend. The radius by which the trajectory of a moving ion is
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turned by the magnetic field is determined by the mass of the atom. The mass spectrometer uses this principle to measure the masstocharge ratio of ions. If a sample contains multiple isotopes, the mass spectrometer can determine the proportion of each isotope in the sample by measuring the intensity of the different beams of ions. Techniques to vaporize atoms include inductively coupled plasma atomic emission spectroscopy and inductively coupled plasma mass spectrometry, both of which use a plasma to vaporize samples for analysis.
The atomprobe tomograph has subnanometer resolution in 3D and can chemically identify individual atoms using timeofflight mass spectrometry.
Electron emission techniques such as Xray photoelectron spectroscopy XPS and Auger electron spectroscopy AES, which measure the binding energies of the core electrons, are used to identify the atomic species present in a sample in a nondestructive way. With proper focusing both can be made areaspecific. Another such method is electron energy
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loss spectroscopy EELS, which measures the energy loss of an electron beam within a transmission electron microscope when it interacts with a portion of a sample.
Spectra of excited states can be used to analyze the atomic composition of distant stars. Specific light wavelengths contained in the observed light from stars can be separated out and related to the quantized transitions in free gas atoms. These colors can be replicated using a gasdischarge lamp containing the same element. Helium was discovered in this way in the spectrum of the Sun 23 years before it was found on Earth.
Origin and current state
Baryonic matter forms about 4 of the total energy density of the observable Universe, with an average density of about 0.25 particlesm3 mostly protons and electrons. Within a galaxy such as the Milky Way, particles have a much higher concentration, with the density of matter in the interstellar medium ISM ranging from 105 to 109 atomsm3. The Sun is believed to be inside the Local Bubble, so the density i
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n the solar neighborhood is only about 103 atomsm3. Stars form from dense clouds in the ISM, and the evolutionary processes of stars result in the steady enrichment of the ISM with elements more massive than hydrogen and helium.
Up to 95 of the Milky Way's baryonic matter are concentrated inside stars, where conditions are unfavorable for atomic matter. The total baryonic mass is about 10 of the mass of the galaxy; the remainder of the mass is an unknown dark matter. High temperature inside stars makes most "atoms" fully ionized, that is, separates all electrons from the nuclei. In stellar remnantswith exception of their surface layersan immense pressure make electron shells impossible.
Formation
Electrons are thought to exist in the Universe since early stages of the Big Bang. Atomic nuclei forms in nucleosynthesis reactions. In about three minutes Big Bang nucleosynthesis produced most of the helium, lithium, and deuterium in the Universe, and perhaps some of the beryllium and boron.
Ubiquitousness and
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stability of atoms relies on their binding energy, which means that an atom has a lower energy than an unbound system of the nucleus and electrons. Where the temperature is much higher than ionization potential, the matter exists in the form of plasmaa gas of positively charged ions possibly, bare nuclei and electrons. When the temperature drops below the ionization potential, atoms become statistically favorable. Atoms complete with bound electrons became to dominate over charged particles 380,000 years after the Big Bangan epoch called recombination, when the expanding Universe cooled enough to allow electrons to become attached to nuclei.
Since the Big Bang, which produced no carbon or heavier elements, atomic nuclei have been combined in stars through the process of nuclear fusion to produce more of the element helium, and via the triple alpha process the sequence of elements from carbon up to iron; see stellar nucleosynthesis for details.
Isotopes such as lithium6, as well as some beryllium and boron a
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re generated in space through cosmic ray spallation. This occurs when a highenergy proton strikes an atomic nucleus, causing large numbers of nucleons to be ejected.
Elements heavier than iron were produced in supernovae and colliding neutron stars through the rprocess, and in AGB stars through the sprocess, both of which involve the capture of neutrons by atomic nuclei. Elements such as lead formed largely through the radioactive decay of heavier elements.
Earth
Most of the atoms that make up the Earth and its inhabitants were present in their current form in the nebula that collapsed out of a molecular cloud to form the Solar System. The rest are the result of radioactive decay, and their relative proportion can be used to determine the age of the Earth through radiometric dating. Most of the helium in the crust of the Earth about 99 of the helium from gas wells, as shown by its lower abundance of helium3 is a product of alpha decay.
There are a few trace atoms on Earth that were not present at the begin
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ning i.e., not "primordial", nor are results of radioactive decay. Carbon14 is continuously generated by cosmic rays in the atmosphere. Some atoms on Earth have been artificially generated either deliberately or as byproducts of nuclear reactors or explosions. Of the transuranic elementsthose with atomic numbers greater than 92only plutonium and neptunium occur naturally on Earth. Transuranic elements have radioactive lifetimes shorter than the current age of the Earth and thus identifiable quantities of these elements have long since decayed, with the exception of traces of plutonium244 possibly deposited by cosmic dust. Natural deposits of plutonium and neptunium are produced by neutron capture in uranium ore.
The Earth contains approximately atoms. Although small numbers of independent atoms of noble gases exist, such as argon, neon, and helium, 99 of the atmosphere is bound in the form of molecules, including carbon dioxide and diatomic oxygen and nitrogen. At the surface of the Earth, an overwhelming m
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ajority of atoms combine to form various compounds, including water, salt, silicates and oxides. Atoms can also combine to create materials that do not consist of discrete molecules, including crystals and liquid or solid metals. This atomic matter forms networked arrangements that lack the particular type of smallscale interrupted order associated with molecular matter.
Rare and theoretical forms
Superheavy elements
All nuclides with atomic numbers higher than 82 lead are known to be radioactive. No nuclide with an atomic number exceeding 92 uranium exists on Earth as a primordial nuclide, and heavier elements generally have shorter halflives. Nevertheless, an "island of stability" encompassing relatively longlived isotopes of superheavy elements with atomic numbers 110 to 114 might exist. Predictions for the halflife of the most stable nuclide on the island range from a few minutes to millions of years. In any case, superheavy elements with Z 104 would not exist due to increasing Coulomb repulsion which
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results in spontaneous fission with increasingly short halflives in the absence of any stabilizing effects.
Exotic matter
Each particle of matter has a corresponding antimatter particle with the opposite electrical charge. Thus, the positron is a positively charged antielectron and the antiproton is a negatively charged equivalent of a proton. When a matter and corresponding antimatter particle meet, they annihilate each other. Because of this, along with an imbalance between the number of matter and antimatter particles, the latter are rare in the universe. The first causes of this imbalance are not yet fully understood, although theories of baryogenesis may offer an explanation. As a result, no antimatter atoms have been discovered in nature. In 1996, the antimatter counterpart of the hydrogen atom antihydrogen was synthesized at the CERN laboratory in Geneva.
Other exotic atoms have been created by replacing one of the protons, neutrons or electrons with other particles that have the same charge. For e
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xample, an electron can be replaced by a more massive muon, forming a muonic atom. These types of atoms can be used to test fundamental predictions of physics.
See also
Notes
References
Bibliography
Further reading
External links
Chemistry
Articles containing video clips
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Arable land from the , "able to be ploughed" is any land capable of being ploughed and used to grow crops. Alternatively, for the purposes of agricultural statistics, the term often has a more precise definition
A more concise definition appearing in the Eurostat glossary similarly refers to actual rather than potential uses "land worked ploughed or tilled regularly, generally under a system of crop rotation".
Nonarable land can sometimes be converted to arable land through methods such as loosening and tilling breaking up of the soil, though in more extreme cases the degree of modification required to make certain types of land arable can become prohibitively expensive.
In Britain, arable land has traditionally been contrasted with pasturable land such as heaths, which could be used for sheeprearing but not as farmland.
Arable land area
According to the Food and Agriculture Organization of the United Nations, in 2013, the world's arable land amounted to 1.407 billion hectares, out of a total of 4.924 b
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illion hectares of land used for agriculture.
Arable land hectares per person
Nonarable land
Agricultural land that is not arable according to the FAO definition above includes
Meadows and pasturesland used as pasture and grazed range, and those natural grasslands and sedge meadows that are used for hay production in some regions.
Permanent cropland that produces crops from woody vegetation, e.g. orchard land, vineyards, coffee plantations, rubber plantations, and land producing nut trees;
Other nonarable land includes land that is not suitable for any agricultural use. Land that is not arable, in the sense of lacking capability or suitability for cultivation for crop production, has one or more limitationsa lack of sufficient freshwater for irrigation, stoniness, steepness, adverse climate, excessive wetness with the impracticality of drainage, excessive salts, or a combination of these, among others. Although such limitations may preclude cultivation, and some will in some cases preclude any agricultu
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ral use, large areas unsuitable for cultivation may still be agriculturally productive. For example, United States NRCS statistics indicate that about 59 percent of US nonfederal pasture and unforested rangeland is unsuitable for cultivation, yet such land has value for grazing of livestock. In British Columbia, Canada, 41 percent of the provincial Agricultural Land Reserve area is unsuitable for the production of cultivated crops, but is suitable for uncultivated production of forage usable by grazing livestock. Similar examples can be found in many rangeland areas elsewhere.
Land incapable of being cultivated for the production of crops can sometimes be converted to arable land. New arable land makes more food and can reduce starvation. This outcome also makes a country more selfsufficient and politically independent, because food importation is reduced. Making nonarable land arable often involves digging new irrigation canals and new wells, aqueducts, desalination plants, planting trees for shade in the d
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esert, hydroponics, fertilizer, nitrogen fertilizer, pesticides, reverse osmosis water processors, PET film insulation or other insulation against heat and cold, digging ditches and hills for protection against the wind, and installing greenhouses with internal light and heat for protection against the cold outside and to provide light in cloudy areas. Such modifications are often prohibitively expensive. An alternative is the seawater greenhouse, which desalinates water through evaporation and condensation using solar energy as the only energy input. This technology is optimized to grow crops on desert land close to the sea.
The use of artifices does not make the land arable. Rock still remains rock, and shallowless than turnable soil is still not considered toilable. The use of artifice is an openair none recycled water hydroponics relationship. The below described circumstances are not in perspective, have limited duration, and have a tendency to accumulate trace materials in soil that either there or els
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ewhere cause deoxygenation. The use of vast amounts of fertilizer may have unintended consequences for the environment by devastating rivers, waterways, and river endings through the accumulation of nondegradable toxins and nitrogenbearing molecules that remove oxygen and cause nonaerobic processes to form.
Examples of infertile nonarable land being turned into fertile arable land include
Aran Islands These islands off the west coast of Ireland not to be confused with the Isle of Arran in Scotland's Firth of Clyde were unsuitable for arable farming because they were too rocky. The people covered the islands with a shallow layer of seaweed and sand from the ocean. Today, crops are grown there, even though the islands are still considered nonarable.
Israel The construction of desalination plants along Israel's coast allowed agriculture in some areas that were formerly desert. The desalination plants, which remove the salt from ocean water, have produced a new source of water for farming, drinking, and washin
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g.
Slash and burn agriculture uses nutrients in wood ash, but these expire within a few years.
Terra preta, fertile tropical soils produced by adding charcoal.
Examples of fertile arable land being turned into infertile land include
Droughts such as the "Dust Bowl" of the Great Depression in the US turned farmland into desert.
Each year, arable land is lost due to desertification and humaninduced erosion. Improper irrigation of farmland can wick the sodium, calcium, and magnesium from the soil and water to the surface. This process steadily concentrates salt in the root zone, decreasing productivity for crops that are not salttolerant.
Rainforest deforestation The fertile tropical forests are converted into infertile desert land. For example, Madagascar's central highland plateau has become virtually totally barren about ten percent of the country as a result of slashandburn deforestation, an element of shifting cultivation practiced by many natives.
See also
Development easement
Land use statistics
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by country
List of environment topics
Soil fertility
References
External links
Article from Technorati on Shrinking Arable Farmland in the world
Surface area of the Earth
Agricultural land
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Aluminium or aluminum in American English and Canadian English is a chemical element with the symbol Al and atomic number 13. Aluminium has a density lower than those of other common metals, at approximately one third that of steel. It has a great affinity towards oxygen, and forms a protective layer of oxide on the surface when exposed to air. Aluminium visually resembles silver, both in its color and in its great ability to reflect light. It is soft, nonmagnetic and ductile. It has one stable isotope, 27Al; this isotope is very common, making aluminium the twelfth most common element in the Universe. The radioactivity of 26Al is used in radiodating.
Chemically, aluminium is a posttransition metal in the boron group; as is common for the group, aluminium forms compounds primarily in the 3 oxidation state. The aluminium cation Al3 is small and highly charged; as such, it is polarizing, and bonds aluminium forms tend towards covalency. The strong affinity towards oxygen leads to aluminium's common association
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with oxygen in nature in the form of oxides; for this reason, aluminium is found on Earth primarily in rocks in the crust, where it is the third most abundant element after oxygen and silicon, rather than in the mantle, and virtually never as the free metal.
The discovery of aluminium was announced in 1825 by Danish physicist Hans Christian rsted. The first industrial production of aluminium was initiated by French chemist Henri tienne SainteClaire Deville in 1856. Aluminium became much more available to the public with the HallHroult process developed independently by French engineer Paul Hroult and American engineer Charles Martin Hall in 1886, and the mass production of aluminium led to its extensive use in industry and everyday life. In World Wars I and II, aluminium was a crucial strategic resource for aviation. In 1954, aluminium became the most produced nonferrous metal, surpassing copper. In the 21st century, most aluminium was consumed in transportation, engineering, construction, and packaging in
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the United States, Western Europe, and Japan.
Despite its prevalence in the environment, no living organism is known to use aluminium salts metabolically, but aluminium is well tolerated by plants and animals. Because of the abundance of these salts, the potential for a biological role for them is of continuing interest, and studies continue.
Physical characteristics
Isotopes
Of aluminium isotopes, only is stable. This situation is common for elements with an odd atomic number. It is the only primordial aluminium isotope, i.e. the only one that has existed on Earth in its current form since the formation of the planet. Nearly all aluminium on Earth is present as this isotope, which makes it a mononuclidic element and means that its standard atomic weight is virtually the same as that of the isotope. This makes aluminium very useful in nuclear magnetic resonance NMR, as its single stable isotope has a high NMR sensitivity. The standard atomic weight of aluminium is low in comparison with many other metal
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s.
All other isotopes of aluminium are radioactive. The most stable of these is 26Al while it was present along with stable 27Al in the interstellar medium from which the Solar System formed, having been produced by stellar nucleosynthesis as well, its halflife is only 717,000 years and therefore a detectable amount has not survived since the formation of the planet. However, minute traces of 26Al are produced from argon in the atmosphere by spallation caused by cosmic ray protons. The ratio of 26Al to 10Be has been used for radiodating of geological processes over 105 to 106 year time scales, in particular transport, deposition, sediment storage, burial times, and erosion. Most meteorite scientists believe that the energy released by the decay of 26Al was responsible for the melting and differentiation of some asteroids after their formation 4.55 billion years ago.
The remaining isotopes of aluminium, with mass numbers ranging from 22 to 43, all have halflives well under an hour. Three metastable states ar
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e known, all with halflives under a minute.
Electron shell
An aluminium atom has 13 electrons, arranged in an electron configuration of Ne 3s2 3p1, with three electrons beyond a stable noble gas configuration. Accordingly, the combined first three ionization energies of aluminium are far lower than the fourth ionization energy alone. Such an electron configuration is shared with the other wellcharacterized members of its group, boron, gallium, indium, and thallium; it is also expected for nihonium. Aluminium can relatively easily surrender its three outermost electrons in many chemical reactions see below. The electronegativity of aluminium is 1.61 Pauling scale.
A free aluminium atom has a radius of 143 pm. With the three outermost electrons removed, the radius shrinks to 39 pm for a 4coordinated atom or 53.5 pm for a 6coordinated atom. At standard temperature and pressure, aluminium atoms when not affected by atoms of other elements form a facecentered cubic crystal system bound by metallic bonding prov
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ided by atoms' outermost electrons; hence aluminium at these conditions is a metal. This crystal system is shared by many other metals, such as lead and copper; the size of a unit cell of aluminium is comparable to that of those other metals. The system, however, is not shared by the other members of its group; boron has ionization energies too high to allow metallization, thallium has a hexagonal closepacked structure, and gallium and indium have unusual structures that are not closepacked like those of aluminium and thallium. The few electrons that are available for metallic bonding in aluminium metal are a probable cause for it being soft with a low melting point and low electrical resistivity.
Bulk
Aluminium metal has an appearance ranging from silvery white to dull gray, depending on the surface roughness. A fresh film of aluminium serves as a good reflector approximately 92 of visible light and an excellent reflector as much as 98 of medium and far infrared radiation. Aluminium mirrors are the most r
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eflective of all metal mirrors for the near ultraviolet and far infrared light, and one of the most reflective in the visible spectrum, nearly on par with silver, and the two therefore look similar. Aluminium is also good at reflecting solar radiation, although prolonged exposure to sunlight in air adds wear to the surface of the metal; this may be prevented if aluminium is anodized, which adds a protective layer of oxide on the surface.
The density of aluminium is 2.70 gcm3, about 13 that of steel, much lower than other commonly encountered metals, making aluminium parts easily identifiable through their lightness. Aluminium's low density compared to most other metals arises from the fact that its nuclei are much lighter, while difference in the unit cell size does not compensate for this difference. The only lighter metals are the metals of groups 1 and 2, which apart from beryllium and magnesium are too reactive for structural use and beryllium is very toxic. Aluminium is not as strong or stiff as steel,
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but the low density makes up for this in the aerospace industry and for many other applications where light weight and relatively high strength are crucial.
Pure aluminium is quite soft and lacking in strength. In most applications various aluminium alloys are used instead because of their higher strength and hardness. The yield strength of pure aluminium is 711 MPa, while aluminium alloys have yield strengths ranging from 200 MPa to 600 MPa. Aluminium is ductile, with a percent elongation of 5070, and malleable allowing it to be easily drawn and extruded. It is also easily machined and cast.
Aluminium is an excellent thermal and electrical conductor, having around 60 the conductivity of copper, both thermal and electrical, while having only 30 of copper's density. Aluminium is capable of superconductivity, with a superconducting critical temperature of 1.2 kelvin and a critical magnetic field of about 100 gauss 10 milliteslas. It is paramagnetic and thus essentially unaffected by static magnetic fields. Th
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e high electrical conductivity, however, means that it is strongly affected by alternating magnetic fields through the induction of eddy currents.
Chemistry
Aluminium combines characteristics of pre and posttransition metals. Since it has few available electrons for metallic bonding, like its heavier group 13 congeners, it has the characteristic physical properties of a posttransition metal, with longerthanexpected interatomic distances. Furthermore, as Al3 is a small and highly charged cation, it is strongly polarizing and bonding in aluminium compounds tends towards covalency; this behavior is similar to that of beryllium Be2, and the two display an example of a diagonal relationship.
The underlying core under aluminium's valence shell is that of the preceding noble gas, whereas those of its heavier congeners gallium, indium, thallium, and nihonium also include a filled dsubshell and in some cases a filled fsubshell. Hence, the inner electrons of aluminium shield the valence electrons almost completely,
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unlike those of aluminium's heavier congeners. As such, aluminium is the most electropositive metal in its group, and its hydroxide is in fact more basic than that of gallium. Aluminium also bears minor similarities to the metalloid boron in the same group AlX3 compounds are valence isoelectronic to BX3 compounds they have the same valence electronic structure, and both behave as Lewis acids and readily form adducts. Additionally, one of the main motifs of boron chemistry is regular icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the AlZnMg class.
Aluminium has a high chemical affinity to oxygen, which renders it suitable for use as a reducing agent in the thermite reaction. A fine powder of aluminium metal reacts explosively on contact with liquid oxygen; under normal conditions, however, aluminium forms a thin oxide layer 5 nm at room temperature that protects the metal from further corrosion by oxygen, water, or dilute acid, a process terme
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d passivation. Because of its general resistance to corrosion, aluminium is one of the few metals that retains silvery reflectance in finely powdered form, making it an important component of silvercolored paints. Aluminium is not attacked by oxidizing acids because of its passivation. This allows aluminium to be used to store reagents such as nitric acid, concentrated sulfuric acid, and some organic acids.
In hot concentrated hydrochloric acid, aluminium reacts with water with evolution of hydrogen, and in aqueous sodium hydroxide or potassium hydroxide at room temperature to form aluminatesprotective passivation under these conditions is negligible. Aqua regia also dissolves aluminium. Aluminium is corroded by dissolved chlorides, such as common sodium chloride, which is why household plumbing is never made from aluminium. The oxide layer on aluminium is also destroyed by contact with mercury due to amalgamation or with salts of some electropositive metals. As such, the strongest aluminium alloys are less
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corrosionresistant due to galvanic reactions with alloyed copper, and aluminium's corrosion resistance is greatly reduced by aqueous salts, particularly in the presence of dissimilar metals.
Aluminium reacts with most nonmetals upon heating, forming compounds such as aluminium nitride AlN, aluminium sulfide Al2S3, and the aluminium halides AlX3. It also forms a wide range of intermetallic compounds involving metals from every group on the periodic table.
Inorganic compounds
The vast majority of compounds, including all aluminiumcontaining minerals and all commercially significant aluminium compounds, feature aluminium in the oxidation state 3. The coordination number of such compounds varies, but generally Al3 is either six or fourcoordinate. Almost all compounds of aluminiumIII are colorless.
In aqueous solution, Al3 exists as the hexaaqua cation AlH2O63, which has an approximate Ka of 105. Such solutions are acidic as this cation can act as a proton donor and progressively hydrolyze until a precipitate
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of aluminium hydroxide, AlOH3, forms. This is useful for clarification of water, as the precipitate nucleates on suspended particles in the water, hence removing them. Increasing the pH even further leads to the hydroxide dissolving again as aluminate, AlH2O2OH4, is formed.
Aluminium hydroxide forms both salts and aluminates and dissolves in acid and alkali, as well as on fusion with acidic and basic oxides. This behavior of AlOH3 is termed amphoterism and is characteristic of weakly basic cations that form insoluble hydroxides and whose hydrated species can also donate their protons. One effect of this is that aluminium salts with weak acids are hydrolyzed in water to the aquated hydroxide and the corresponding nonmetal hydride for example, aluminium sulfide yields hydrogen sulfide. However, some salts like aluminium carbonate exist in aqueous solution but are unstable as such; and only incomplete hydrolysis takes place for salts with strong acids, such as the halides, nitrate, and sulfate. For similar rea
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sons, anhydrous aluminium salts cannot be made by heating their "hydrates" hydrated aluminium chloride is in fact not AlCl36H2O but AlH2O6Cl3, and the AlO bonds are so strong that heating is not sufficient to break them and form AlCl bonds instead
2AlH2O6Cl3 Al2O3 6 HCl 9 H2O
All four trihalides are well known. Unlike the structures of the three heavier trihalides, aluminium fluoride AlF3 features sixcoordinate aluminium, which explains its involatility and insolubility as well as high heat of formation. Each aluminium atom is surrounded by six fluorine atoms in a distorted octahedral arrangement, with each fluorine atom being shared between the corners of two octahedra. Such AlF6 units also exist in complex fluorides such as cryolite, Na3AlF6. AlF3 melts at and is made by reaction of aluminium oxide with hydrogen fluoride gas at .
With heavier halides, the coordination numbers are lower. The other trihalides are dimeric or polymeric with tetrahedral fourcoordinate aluminium centers. Aluminium trichlor
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ide AlCl3 has a layered polymeric structure below its melting point of but transforms on melting to Al2Cl6 dimers. At higher temperatures those increasingly dissociate into trigonal planar AlCl3 monomers similar to the structure of BCl3. Aluminium tribromide and aluminium triiodide form Al2X6 dimers in all three phases and hence do not show such significant changes of properties upon phase change. These materials are prepared by treating aluminium metal with the halogen. The aluminium trihalides form many addition compounds or complexes; their Lewis acidic nature makes them useful as catalysts for the FriedelCrafts reactions. Aluminium trichloride has major industrial uses involving this reaction, such as in the manufacture of anthraquinones and styrene; it is also often used as the precursor for many other aluminium compounds and as a reagent for converting nonmetal fluorides into the corresponding chlorides a transhalogenation reaction.
Aluminium forms one stable oxide with the chemical formula Al2O3, com
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monly called alumina. It can be found in nature in the mineral corundum, alumina; there is also a alumina phase. Its crystalline form, corundum, is very hard Mohs hardness 9, has a high melting point of , has very low volatility, is chemically inert, and a good electrical insulator, it is often used in abrasives such as toothpaste, as a refractory material, and in ceramics, as well as being the starting material for the electrolytic production of aluminium metal. Sapphire and ruby are impure corundum contaminated with trace amounts of other metals. The two main oxidehydroxides, AlOOH, are boehmite and diaspore. There are three main trihydroxides bayerite, gibbsite, and nordstrandite, which differ in their crystalline structure polymorphs. Many other intermediate and related structures are also known. Most are produced from ores by a variety of wet processes using acid and base. Heating the hydroxides leads to formation of corundum. These materials are of central importance to the production of aluminium and a
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re themselves extremely useful. Some mixed oxide phases are also very useful, such as spinel MgAl2O4, Naalumina NaAl11O17, and tricalcium aluminate Ca3Al2O6, an important mineral phase in Portland cement.
The only stable chalcogenides under normal conditions are aluminium sulfide Al2S3, selenide Al2Se3, and telluride Al2Te3. All three are prepared by direct reaction of their elements at about and quickly hydrolyze completely in water to yield aluminium hydroxide and the respective hydrogen chalcogenide. As aluminium is a small atom relative to these chalcogens, these have fourcoordinate tetrahedral aluminium with various polymorphs having structures related to wurtzite, with twothirds of the possible metal sites occupied either in an orderly or random fashion; the sulfide also has a form related to alumina, and an unusual hightemperature hexagonal form where half the aluminium atoms have tetrahedral fourcoordination and the other half have trigonal bipyramidal fivecoordination.
Four pnictides aluminiu
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m nitride AlN, aluminium phosphide AlP, aluminium arsenide AlAs, and aluminium antimonide AlSb are known. They are all IIIV semiconductors isoelectronic to silicon and germanium, all of which but AlN have the zinc blende structure. All four can be made by hightemperature and possibly highpressure direct reaction of their component elements.
Aluminium alloys well with most other metals with the exception of most alkali metals and group 13 metals and over 150 intermetallics with other metals are known. Preparation involves heating fixed metals together in certain proportion, followed by gradual cooling and annealing. Bonding in them is predominantly metallic and the crystal structure primarily depends on efficiency of packing.
There are few compounds with lower oxidation states. A few aluminiumI compounds exist AlF, AlCl, AlBr, and AlI exist in the gaseous phase when the respective trihalide is heated with aluminium, and at cryogenic temperatures. A stable derivative of aluminium monoiodide is the cyclic add
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uct formed with triethylamine, Al4I4NEt34. Al2O and Al2S also exist but are very unstable. Very simple aluminiumII compounds are invoked or observed in the reactions of Al metal with oxidants. For example, aluminium monoxide, AlO, has been detected in the gas phase after explosion and in stellar absorption spectra. More thoroughly investigated are compounds of the formula R4Al2 which contain an AlAl bond and where R is a large organic ligand.
Organoaluminium compounds and related hydrides
A variety of compounds of empirical formula AlR3 and AlR1.5Cl1.5 exist. The aluminium trialkyls and triaryls are reactive, volatile, and colorless liquids or lowmelting solids. They catch fire spontaneously in air and react with water, thus necessitating precautions when handling them. They often form dimers, unlike their boron analogues, but this tendency diminishes for branchedchain alkyls e.g. Pri, Bui, Me3CCH2; for example, triisobutylaluminium exists as an equilibrium mixture of the monomer and dimer. These dimers, s
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uch as trimethylaluminium Al2Me6, usually feature tetrahedral Al centers formed by dimerization with some alkyl group bridging between both aluminium atoms. They are hard acids and react readily with ligands, forming adducts. In industry, they are mostly used in alkene insertion reactions, as discovered by Karl Ziegler, most importantly in "growth reactions" that form longchain unbranched primary alkenes and alcohols, and in the lowpressure polymerization of ethene and propene. There are also some heterocyclic and cluster organoaluminium compounds involving AlN bonds.
The industrially most important aluminium hydride is lithium aluminium hydride LiAlH4, which is used in as a reducing agent in organic chemistry. It can be produced from lithium hydride and aluminium trichloride. The simplest hydride, aluminium hydride or alane, is not as important. It is a polymer with the formula AlH3n, in contrast to the corresponding boron hydride that is a dimer with the formula BH32.
Natural occurrence
Space
Aluminium
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's perparticle abundance in the Solar System is 3.15 ppm parts per million. It is the twelfth most abundant of all elements and third most abundant among the elements that have odd atomic numbers, after hydrogen and nitrogen. The only stable isotope of aluminium, 27Al, is the eighteenth most abundant nucleus in the Universe. It is created almost entirely after fusion of carbon in massive stars that will later become Type II supernovas this fusion creates 26Mg, which, upon capturing free protons and neutrons becomes aluminium. Some smaller quantities of 27Al are created in hydrogen burning shells of evolved stars, where 26Mg can capture free protons. Essentially all aluminium now in existence is 27Al. 26Al was present in the early Solar System with abundance of 0.005 relative to 27Al but its halflife of 728,000 years is too short for any original nuclei to survive; 26Al is therefore extinct. Unlike for 27Al, hydrogen burning is the primary source of 26Al, with the nuclide emerging after a nucleus of 25Mg catch
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es a free proton. However, the trace quantities of 26Al that do exist are the most common gamma ray emitter in the interstellar gas; if the original 26Al were still present, gamma ray maps of the Milky Way would be brighter.
Earth
Overall, the Earth is about 1.59 aluminium by mass seventh in abundance by mass. Aluminium occurs in greater proportion in the Earth's crust than in the Universe at large, because aluminium easily forms the oxide and becomes bound into rocks and stays in the Earth's crust, while less reactive metals sink to the core. In the Earth's crust, aluminium is the most abundant metallic element 8.23 by mass and the third most abundant of all elements after oxygen and silicon. A large number of silicates in the Earth's crust contain aluminium. In contrast, the Earth's mantle is only 2.38 aluminium by mass. Aluminium also occurs in seawater at a concentration of 2 gkg.
Because of its strong affinity for oxygen, aluminium is almost never found in the elemental state; instead it is found in
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oxides or silicates. Feldspars, the most common group of minerals in the Earth's crust, are aluminosilicates. Aluminium also occurs in the minerals beryl, cryolite, garnet, spinel, and turquoise. Impurities in Al2O3, such as chromium and iron, yield the gemstones ruby and sapphire, respectively. Native aluminium metal is extremely rare and can only be found as a minor phase in low oxygen fugacity environments, such as the interiors of certain volcanoes. Native aluminium has been reported in cold seeps in the northeastern continental slope of the South China Sea. It is possible that these deposits resulted from bacterial reduction of tetrahydroxoaluminate AlOH4.
Although aluminium is a common and widespread element, not all aluminium minerals are economically viable sources of the metal. Almost all metallic aluminium is produced from the ore bauxite AlOxOH32x. Bauxite occurs as a weathering product of low iron and silica bedrock in tropical climatic conditions. In 2017, most bauxite was mined in Australia, Ch
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ina, Guinea, and India.
History
The history of aluminium has been shaped by usage of alum. The first written record of alum, made by Greek historian Herodotus, dates back to the 5th century BCE. The ancients are known to have used alum as a dyeing mordant and for city defense. After the Crusades, alum, an indispensable good in the European fabric industry, was a subject of international commerce; it was imported to Europe from the eastern Mediterranean until the mid15th century.
The nature of alum remained unknown. Around 1530, Swiss physician Paracelsus suggested alum was a salt of an earth of alum. In 1595, German doctor and chemist Andreas Libavius experimentally confirmed this. In 1722, German chemist Friedrich Hoffmann announced his belief that the base of alum was a distinct earth. In 1754, German chemist Andreas Sigismund Marggraf synthesized alumina by boiling clay in sulfuric acid and subsequently adding potash.
Attempts to produce aluminium metal date back to 1760. The first successful attempt,
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however, was completed in 1824 by Danish physicist and chemist Hans Christian rsted. He reacted anhydrous aluminium chloride with potassium amalgam, yielding a lump of metal looking similar to tin. He presented his results and demonstrated a sample of the new metal in 1825. In 1827, German chemist Friedrich Whler repeated rsted's experiments but did not identify any aluminium. The reason for this inconsistency was only discovered in 1921. He conducted a similar experiment in the same year by mixing anhydrous aluminium chloride with potassium and produced a powder of aluminium. In 1845, he was able to produce small pieces of the metal and described some physical properties of this metal. For many years thereafter, Whler was credited as the discoverer of aluminium.
As Whler's method could not yield great quantities of aluminium, the metal remained rare; its cost exceeded that of gold. The first industrial production of aluminium was established in 1856 by French chemist Henri Etienne SainteClaire Deville and
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companions. Deville had discovered that aluminium trichloride could be reduced by sodium, which was more convenient and less expensive than potassium, which Whler had used. Even then, aluminium was still not of great purity and produced aluminium differed in properties by sample.
The first industrial largescale production method was independently developed in 1886 by French engineer Paul Hroult and American engineer Charles Martin Hall; it is now known as the HallHroult process. The HallHroult process converts alumina into metal. Austrian chemist Carl Joseph Bayer discovered a way of purifying bauxite to yield alumina, now known as the Bayer process, in 1889. Modern production of the aluminium metal is based on the Bayer and HallHroult processes.
Prices of aluminium dropped and aluminium became widely used in jewelry, everyday items, eyeglass frames, optical instruments, tableware, and foil in the 1890s and early 20th century. Aluminium's ability to form hard yet light alloys with other metals provided the
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metal with many uses at the time. During World War I, major governments demanded large shipments of aluminium for light strong airframes; during World War II, demand by major governments for aviation was even higher.
By the mid20th century, aluminium had become a part of everyday life and an essential component of housewares. In 1954, production of aluminium surpassed that of copper, historically second in production only to iron, making it the most produced nonferrous metal. During the mid20th century, aluminium emerged as a civil engineering material, with building applications in both basic construction and interior finish work, and increasingly being used in military engineering, for both airplanes and land armor vehicle engines. Earth's first artificial satellite, launched in 1957, consisted of two separate aluminium semispheres joined and all subsequent space vehicles have used aluminium to some extent. The aluminium can was invented in 1956 and employed as a storage for drinks in 1958.
Throughout the
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20th century, the production of aluminium rose rapidly while the world production of aluminium in 1900 was 6,800 metric tons, the annual production first exceeded 100,000 metric tons in 1916; 1,000,000 tons in 1941; 10,000,000 tons in 1971. In the 1970s, the increased demand for aluminium made it an exchange commodity; it entered the London Metal Exchange, the oldest industrial metal exchange in the world, in 1978. The output continued to grow the annual production of aluminium exceeded 50,000,000 metric tons in 2013.
The real price for aluminium declined from 14,000 per metric ton in 1900 to 2,340 in 1948 in 1998 United States dollars. Extraction and processing costs were lowered over technological progress and the scale of the economies. However, the need to exploit lowergrade poorer quality deposits and the use of fast increasing input costs above all, energy increased the net cost of aluminium; the real price began to grow in the 1970s with the rise of energy cost. Production moved from the industrializ
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ed countries to countries where production was cheaper. Production costs in the late 20th century changed because of advances in technology, lower energy prices, exchange rates of the United States dollar, and alumina prices. The BRIC countries' combined share in primary production and primary consumption grew substantially in the first decade of the 21st century. China is accumulating an especially large share of the world's production thanks to an abundance of resources, cheap energy, and governmental stimuli; it also increased its consumption share from 2 in 1972 to 40 in 2010. In the United States, Western Europe, and Japan, most aluminium was consumed in transportation, engineering, construction, and packaging. In 2021, prices for industrial metals such as aluminium have soared to nearrecord levels as energy shortages in China drive up costs for electricity.
Etymology
The names aluminium and aluminum are derived from the word alumine, an obsolete term for alumina, a naturally occurring oxide of alumini
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um. Alumine was borrowed from French, which in turn derived it from alumen, the classical Latin name for alum, the mineral from which it was collected. The Latin word alumen stems from the ProtoIndoEuropean root alu meaning "bitter" or "beer".
Coinage
British chemist Humphry Davy, who performed a number of experiments aimed to isolate the metal, is credited as the person who named the element. The first name proposed for the metal to be isolated from alum was alumium, which Davy suggested in an 1808 article on his electrochemical research, published in Philosophical Transactions of the Royal Society. It appeared that the name was coined from the English word alum and the Latin suffix ium; however, it was customary at the time that the elements should have names originating in the Latin language, and as such, this name was not adopted universally. This name was criticized by contemporary chemists from France, Germany, and Sweden, who insisted the metal should be named for the oxide, alumina, from which it wo
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uld be isolated. The English word name alum does not directly reference the Latin language, whereas aluminealumina easily references the Latin word alumen upon declension, alumen changes to alumin.
One example was a writing in French by Swedish chemist Jns Jacob Berzelius titled Essai sur la Nomenclature chimique, published in July 1811; in this essay, among other things, Berzelius used the name aluminium for the element that would be synthesized from alum. Another article in the same journal issue also refers to the metal whose oxide forms the basis of sapphire as to aluminium. A January 1811 summary of one of Davy's lectures at the Royal Society mentioned the name aluminium as a possibility. The following year, Davy published a chemistry textbook in which he used the spelling aluminum. Both spellings have coexisted since; however, their usage has split by region aluminum is the primary spelling in the United States and Canada while aluminium is in the rest of the Englishspeaking world.
Spelling
In 1812,
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British scientist Thomas Young wrote an anonymous review of Davy's book, in which he proposed the name aluminium instead of aluminum, which he felt had a "less classical sound". This name did catch on while the spelling was occasionally used in Britain, the American scientific language used from the start. Most scientists used throughout the world in the 19th century, and it was entrenched in many other European languages, such as French, German, or Dutch. In 1828, American lexicographer Noah Webster used exclusively the aluminum spelling in his American Dictionary of the English Language. In the 1830s, the spelling started to gain usage in the United States; by the 1860s, it had become the more common spelling there outside science. In 1892, Hall used the spelling in his advertising handbill for his new electrolytic method of producing the metal, despite his constant use of the spelling in all the patents he filed between 1886 and 1903. It remains unknown whether this spelling was introduced by mistake
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or intentionally; however, Hall preferred aluminum since its introduction because it resembled platinum, the name of a prestigious metal. By 1890, both spellings had been common in the U.S. overall, the spelling being slightly more common; by 1895, the situation had reversed; by 1900, aluminum had become twice as common as aluminium; during the following decade, the spelling dominated American usage. In 1925, the American Chemical Society adopted this spelling.
The International Union of Pure and Applied Chemistry IUPAC adopted aluminium as the standard international name for the element in 1990. In 1993, they recognized aluminum as an acceptable variant; the most recent 2005 edition of the IUPAC nomenclature of inorganic chemistry acknowledges this spelling as well. IUPAC official publications use the spelling as primary but list both where appropriate.
Production and refinement
The production of aluminium starts with the extraction of bauxite rock from the ground. The bauxite is processed and transf
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ormed using the Bayer process into alumina, which is then processed using the HallHroult process, resulting in the final aluminium metal.
Aluminium production is highly energyconsuming, and so the producers tend to locate smelters in places where electric power is both plentiful and inexpensive. As of 2019, the world's largest smelters of aluminium are located in China, India, Russia, Canada, and the United Arab Emirates, while China is by far the top producer of aluminium with a world share of fiftyfive percent.
According to the International Resource Panel's Metal Stocks in Society report, the global per capita stock of aluminium in use in society i.e. in cars, buildings, electronics, etc. is . Much of this is in moredeveloped countries per capita rather than lessdeveloped countries per capita.
Bayer process
Bauxite is converted to alumina by the Bayer process. Bauxite is blended for uniform composition and then is ground. The resulting slurry is mixed with a hot solution of sodium hydroxide; the mix
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ture is then treated in a digester vessel at a pressure well above atmospheric, dissolving the aluminium hydroxide in bauxite while converting impurities into relatively insoluble compounds
After this reaction, the slurry is at a temperature above its atmospheric boiling point. It is cooled by removing steam as pressure is reduced. The bauxite residue is separated from the solution and discarded. The solution, free of solids, is seeded with small crystals of aluminium hydroxide; this causes decomposition of the AlOH4 ions to aluminium hydroxide. After about half of aluminium has precipitated, the mixture is sent to classifiers. Small crystals of aluminium hydroxide are collected to serve as seeding agents; coarse particles are converted to alumina by heating; the excess solution is removed by evaporation, if needed purified, and recycled.
HallHroult process
The conversion of alumina to aluminium metal is achieved by the HallHroult process. In this energyintensive process, a solution of alumina in a molten
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