Isotopes of boron
Boron occurs as isotopes 10B and 11B, the latter of which makes up about 80% of natural boron. There are 13 radioisotopes that have been discovered, with mass numbers from 7 to 21, all with short half-lives, the longest being that of 8B, with a half-life of only 770 milliseconds and 12B with a half-life of 20.2 ms. All other isotopes have half-lives shorter than 17.35 ms. Those isotopes with mass below 10 decay into helium while those with mass above 11 become carbon; the precision of the isotope abundances and atomic mass is limited through variations. The given ranges should be applicable to any normal terrestrial material. Commercially available materials may have been subjected to an undisclosed or inadvertent isotopic fractionation. Substantial deviations from the given mass and composition can occur. Values marked # are not purely derived from experimental data, but at least from systematic trends. Spins with weak assignment arguments are enclosed in parentheses. Uncertainties are given in concise form in parentheses after the corresponding last digits.
Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC, which use folical uncertainties. Nuclide masses are given by IUPAP Commission on Symbols, Nomenclature, Atomic Masses and Fundamental Constants. Isotope abundances are given by IUPAC Commission on Atomic Weights. Neutrinos from boron-8 beta decays within the sun are an important background to dark matter direct detection experiments, they are the first component of the neutrino floor that dark matter direct detection experiments are expected to encounter. Boron-10 is used in boron neutron capture therapy as an experimental treatment of some brain cancers. Isotope masses from: Audi, Georges. P.. "Atomic weights of the elements. Review 2000". Pure and Applied Chemistry. 75: 683–800. Doi:10.1351/pac200375060683. Wieser, Michael E.. "Atomic weights of the elements 2005". Pure and Applied Chemistry. 78: 2051–2066. Doi:10.1351/pac200678112051. Lay summary. Half-life and isomer data selected from the following sources.
See editing notes on this article's talk page. Audi, Georges. "NuDat 2.x database". Brookhaven National Laboratory. Holden, Norman E.. "11. Table of the Isotopes". In Lide, David R. CRC Handbook of Physics. Boca Raton, Florida: CRC Press. ISBN 978-0-8493-0485-9
In quantum mechanics and particle physics, spin is an intrinsic form of angular momentum carried by elementary particles, composite particles, atomic nuclei. Spin is one of two types of angular momentum in quantum mechanics, the other being orbital angular momentum; the orbital angular momentum operator is the quantum-mechanical counterpart to the classical angular momentum of orbital revolution and appears when there is periodic structure to its wavefunction as the angle varies. The existence of spin angular momentum is inferred from experiments, such as the Stern–Gerlach experiment, in which silver atoms were observed to possess two possible discrete angular momenta despite having no orbital angular momentum. In some ways, spin is like a vector quantity. All elementary particles of a given kind have the same magnitude of spin angular momentum, indicated by assigning the particle a spin quantum number; the SI unit of spin is the or, just as with classical angular momentum. In practice, spin is given as a dimensionless spin quantum number by dividing the spin angular momentum by the reduced Planck constant ħ, which has the same units of angular momentum, although this is not the full computation of this value.
The "spin quantum number" is called "spin", leaving its meaning as the unitless "spin quantum number" to be inferred from context. When combined with the spin-statistics theorem, the spin of electrons results in the Pauli exclusion principle, which in turn underlies the periodic table of chemical elements. Wolfgang Pauli in 1924 was the first to propose a doubling of electron states due to a two-valued non-classical "hidden rotation". In 1925, George Uhlenbeck and Samuel Goudsmit at Leiden University suggested the simple physical interpretation of a particle spinning around its own axis, in the spirit of the old quantum theory of Bohr and Sommerfeld. Ralph Kronig anticipated the Uhlenbeck-Goudsmit model in discussion with Hendrik Kramers several months earlier in Copenhagen, but did not publish; the mathematical theory was worked out in depth by Pauli in 1927. When Paul Dirac derived his relativistic quantum mechanics in 1928, electron spin was an essential part of it; as the name suggests, spin was conceived as the rotation of a particle around some axis.
This picture is correct so far as spin obeys the same mathematical laws as quantized angular momenta do. On the other hand, spin has some peculiar properties that distinguish it from orbital angular momenta: Spin quantum numbers may take half-integer values. Although the direction of its spin can be changed, an elementary particle cannot be made to spin faster or slower; the spin of a charged particle is associated with a magnetic dipole moment with a g-factor differing from 1. This could only occur classically if the internal charge of the particle were distributed differently from its mass; the conventional definition of the spin quantum number, s, is s = n/2, where n can be any non-negative integer. Hence the allowed values of s are 1/2, 1, 3/2, 2, etc.. The value of s for an elementary particle depends only on the type of particle, cannot be altered in any known way; the spin angular momentum, S, of any physical system is quantized. The allowed values of S are S = ℏ s = h 4 π n, where h is the Planck constant and ℏ = h/2π is the reduced Planck constant.
In contrast, orbital angular momentum can only take on integer values of s. Those particles with half-integer spins, such as 1/2, 3/2, 5/2, are known as fermions, while those particles with integer spins, such as 0, 1, 2, are known as bosons; the two families of particles obey different rules and broadly have different roles in the world around us. A key distinction between the two families is. In contrast, bosons obey the rules of Bose–Einstein statistics and have no such restriction, so they may "bunch together" if in identical states. Composite particles can have spins different from their component particles. For example, a helium atom in the ground state has spin 0 and behaves like a boson though the quarks and electrons which make it up are all fermions; this has some profound consequences: Quarks and leptons, which make up what is classically known as matter, are all fermions with spin 1/2. The common idea that "matter takes up space" comes from the Pauli exclusion principle acting on these particles to prevent the fermions that make up matter from being in the same quantum state.
Further compaction would require electrons to occupy the same energy states, therefore a kind of pressure acts to resist the fermions being overly close. Elementary fermions with other spins are not known to exist. Elementary particles which are thought of as carrying forces are all bosons with spin 1, they include the photon which carries the electromagnetic force, the gluon, the W and Z bosons. The ability of bosons to occupy the same quantu
In physics, the electronvolt is a unit of energy equal to 1.6×10−19 joules in SI units. The electronvolt was devised as a standard unit of measure through its usefulness in electrostatic particle accelerator sciences, because a particle with electric charge q has an energy E = qV after passing through the potential V. Like the elementary charge on which it is based, it is not an independent quantity but is equal to 1 J/C √2hα / μ0c0, it is a common unit of energy within physics used in solid state, atomic and particle physics. It is used with the metric prefixes milli-, kilo-, mega-, giga-, tera-, peta- or exa-. In some older documents, in the name Bevatron, the symbol BeV is used, which stands for billion electronvolts. An electronvolt is the amount of kinetic energy gained or lost by a single electron accelerating from rest through an electric potential difference of one volt in vacuum. Hence, it has a value of one volt, 1 J/C, multiplied by the electron's elementary charge e, 1.6021766208×10−19 C.
Therefore, one electronvolt is equal to 1.6021766208×10−19 J. The electronvolt, as opposed to volt, is not an SI unit, its derivation is empirical, which means its value in SI units must be obtained by experiment and is therefore not known unlike the litre, the light-year and such other non-SI units. Electronvolt is a unit of energy; the SI unit for energy is joule. 1 eV is equal to 1.6021766208×10−19 J. By mass–energy equivalence, the electronvolt is a unit of mass, it is common in particle physics, where units of mass and energy are interchanged, to express mass in units of eV/c2, where c is the speed of light in vacuum. It is common to express mass in terms of "eV" as a unit of mass using a system of natural units with c set to 1; the mass equivalent of 1 eV/c2 is 1 eV / c 2 = ⋅ 1 V 2 = 1.783 × 10 − 36 kg. For example, an electron and a positron, each with a mass of 0.511 MeV/c2, can annihilate to yield 1.022 MeV of energy. The proton has a mass of 0.938 GeV/c2. In general, the masses of all hadrons are of the order of 1 GeV/c2, which makes the GeV a convenient unit of mass for particle physics: 1 GeV/c2 = 1.783×10−27 kg.
The unified atomic mass unit, 1 gram divided by Avogadro's number, is the mass of a hydrogen atom, the mass of the proton. To convert to megaelectronvolts, use the formula: 1 u = 931.4941 MeV/c2 = 0.9314941 GeV/c2. In high-energy physics, the electronvolt is used as a unit of momentum. A potential difference of 1 volt causes an electron to gain an amount of energy; this gives rise to usage of eV as units of momentum, for the energy supplied results in acceleration of the particle. The dimensions of momentum units are LMT−1; the dimensions of energy units are L2MT−2. Dividing the units of energy by a fundamental constant that has units of velocity, facilitates the required conversion of using energy units to describe momentum. In the field of high-energy particle physics, the fundamental velocity unit is the speed of light in vacuum c. By dividing energy in eV by the speed of light, one can describe the momentum of an electron in units of eV/c; the fundamental velocity constant c is dropped from the units of momentum by way of defining units of length such that the value of c is unity.
For example, if the momentum p of an electron is said to be 1 GeV the conversion to MKS can be achieved by: p = 1 GeV / c = ⋅ ⋅ = 5.344286 × 10 − 19 kg ⋅ m / s. In particle physics, a system of "natural units" in which the speed of light in vacuum c and the reduced Planck constant ħ are dimensionless and equal to unity is used: c = ħ = 1. In these units, both distances and times are expressed in inverse energy units (while energy and mass are expressed in the same units, see mas
International Standard Serial Number
An International Standard Serial Number is an eight-digit serial number used to uniquely identify a serial publication, such as a magazine. The ISSN is helpful in distinguishing between serials with the same title. ISSN are used in ordering, interlibrary loans, other practices in connection with serial literature; the ISSN system was first drafted as an International Organization for Standardization international standard in 1971 and published as ISO 3297 in 1975. ISO subcommittee TC 46/SC 9 is responsible for maintaining the standard; when a serial with the same content is published in more than one media type, a different ISSN is assigned to each media type. For example, many serials are published both in electronic media; the ISSN system refers to these types as electronic ISSN, respectively. Conversely, as defined in ISO 3297:2007, every serial in the ISSN system is assigned a linking ISSN the same as the ISSN assigned to the serial in its first published medium, which links together all ISSNs assigned to the serial in every medium.
The format of the ISSN is an eight digit code, divided by a hyphen into two four-digit numbers. As an integer number, it can be represented by the first seven digits; the last code digit, which may be 0-9 or an X, is a check digit. Formally, the general form of the ISSN code can be expressed as follows: NNNN-NNNC where N is in the set, a digit character, C is in; the ISSN of the journal Hearing Research, for example, is 0378-5955, where the final 5 is the check digit, C=5. To calculate the check digit, the following algorithm may be used: Calculate the sum of the first seven digits of the ISSN multiplied by its position in the number, counting from the right—that is, 8, 7, 6, 5, 4, 3, 2, respectively: 0 ⋅ 8 + 3 ⋅ 7 + 7 ⋅ 6 + 8 ⋅ 5 + 5 ⋅ 4 + 9 ⋅ 3 + 5 ⋅ 2 = 0 + 21 + 42 + 40 + 20 + 27 + 10 = 160 The modulus 11 of this sum is calculated. For calculations, an upper case X in the check digit position indicates a check digit of 10. To confirm the check digit, calculate the sum of all eight digits of the ISSN multiplied by its position in the number, counting from the right.
The modulus 11 of the sum must be 0. There is an online ISSN checker. ISSN codes are assigned by a network of ISSN National Centres located at national libraries and coordinated by the ISSN International Centre based in Paris; the International Centre is an intergovernmental organization created in 1974 through an agreement between UNESCO and the French government. The International Centre maintains a database of all ISSNs assigned worldwide, the ISDS Register otherwise known as the ISSN Register. At the end of 2016, the ISSN Register contained records for 1,943,572 items. ISSN and ISBN codes are similar in concept. An ISBN might be assigned for particular issues of a serial, in addition to the ISSN code for the serial as a whole. An ISSN, unlike the ISBN code, is an anonymous identifier associated with a serial title, containing no information as to the publisher or its location. For this reason a new ISSN is assigned to a serial each time it undergoes a major title change. Since the ISSN applies to an entire serial a new identifier, the Serial Item and Contribution Identifier, was built on top of it to allow references to specific volumes, articles, or other identifiable components.
Separate ISSNs are needed for serials in different media. Thus, the print and electronic media versions of a serial need separate ISSNs. A CD-ROM version and a web version of a serial require different ISSNs since two different media are involved. However, the same ISSN can be used for different file formats of the same online serial; this "media-oriented identification" of serials made sense in the 1970s. In the 1990s and onward, with personal computers, better screens, the Web, it makes sense to consider only content, independent of media; this "content-oriented identification" of serials was a repressed demand during a decade, but no ISSN update or initiative occurred. A natural extension for ISSN, the unique-identification of the articles in the serials, was the main demand application. An alternative serials' contents model arrived with the indecs Content Model and its application, the digital object identifier, as ISSN-independent initiative, consolidated in the 2000s. Only in 2007, ISSN-L was defined in the
The periodic table known as the periodic table of elements, is a tabular display of the chemical elements, which are arranged by atomic number, electron configuration, recurring chemical properties. The structure of the table shows periodic trends; the seven rows of the table, called periods have metals on the left and non-metals on the right. The columns, called groups, contain elements with similar chemical behaviours. Six groups have accepted names as well as assigned numbers: for example, group 17 elements are the halogens. Displayed are four simple rectangular areas or blocks associated with the filling of different atomic orbitals; the organization of the periodic table can be used to derive relationships between the various element properties, to predict chemical properties and behaviours of undiscovered or newly synthesized elements. Russian chemist Dmitri Mendeleev published the first recognizable periodic table in 1869, developed to illustrate periodic trends of the then-known elements.
He predicted some properties of unidentified elements that were expected to fill gaps within the table. Most of his forecasts proved to be correct. Mendeleev's idea has been expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour; the modern periodic table now provides a useful framework for analyzing chemical reactions, continues to be used in chemistry, nuclear physics and other sciences. The elements from atomic numbers 1 through 118 have been discovered or synthesized, completing seven full rows of the periodic table; the first 94 elements all occur though some are found only in trace amounts and a few were discovered in nature only after having first been synthesized. Elements 95 to 118 have only been synthesized in nuclear reactors; the synthesis of elements having higher atomic numbers is being pursued: these elements would begin an eighth row, theoretical work has been done to suggest possible candidates for this extension.
Numerous synthetic radionuclides of occurring elements have been produced in laboratories. Each chemical element has a unique atomic number representing the number of protons in its nucleus. Most elements have differing numbers of neutrons among different atoms, with these variants being referred to as isotopes. For example, carbon has three occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, a small fraction have eight neutrons. Isotopes are never separated in the periodic table. Elements with no stable isotopes have the atomic masses of their most stable isotopes, where such masses are shown, listed in parentheses. In the standard periodic table, the elements are listed in order of increasing atomic number Z. A new row is started. Columns are determined by the electron configuration of the atom. Elements with similar chemical properties fall into the same group in the periodic table, although in the f-block, to some respect in the d-block, the elements in the same period tend to have similar properties, as well.
Thus, it is easy to predict the chemical properties of an element if one knows the properties of the elements around it. Since 2016, the periodic table has 118 confirmed elements, from element 1 to 118. Elements 113, 115, 117 and 118, the most recent discoveries, were confirmed by the International Union of Pure and Applied Chemistry in December 2015, their proposed names, moscovium and oganesson were announced by the IUPAC in June 2016 and made official in November 2016. The first 94 elements occur naturally. Of the 94 occurring elements, 83 are primordial and 11 occur only in decay chains of primordial elements. No element heavier than einsteinium has been observed in macroscopic quantities in its pure form, nor has astatine. A group or family is a vertical column in the periodic table. Groups have more significant periodic trends than periods and blocks, explained below. Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group have the same electron configurations in their valence shell.
Elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number. In some parts of the periodic table, such as the d-block and the f-block, horizontal similarities can be as important as, or more pronounced than, vertical similarities. Under an international naming convention, the groups are numbered numerically from 1 to 18 from the leftmost column to the rightmost column, they were known by roman numerals. In America, the roman numerals were followed by either an "A" if the group was in the s- or p-block, or a "B" if the group was in the d-block; the roman numerals used correspond to the last digit of today's naming convention (e.g. the
The atomic number or proton number of a chemical element is the number of protons found in the nucleus of an atom. It is identical to the charge number of the nucleus; the atomic number uniquely identifies a chemical element. In an uncharged atom, the atomic number is equal to the number of electrons; the sum of the atomic number Z and the number of neutrons, N, gives the mass number A of an atom. Since protons and neutrons have the same mass and the mass defect of nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in unified atomic mass units, is within 1% of the whole number A. Atoms with the same atomic number Z but different neutron numbers N, hence different atomic masses, are known as isotopes. A little more than three-quarters of occurring elements exist as a mixture of isotopes, the average isotopic mass of an isotopic mixture for an element in a defined environment on Earth, determines the element's standard atomic weight, it was these atomic weights of elements that were the quantities measurable by chemists in the 19th century.
The conventional symbol Z comes from the German word Zahl meaning number, before the modern synthesis of ideas from chemistry and physics denoted an element's numerical place in the periodic table, whose order is but not consistent with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence that this Z number was the nuclear charge and a physical characteristic of atoms, did the word Atomzahl come into common use in this context. Loosely speaking, the existence or construction of a periodic table of elements creates an ordering of the elements, so they can be numbered in order. Dmitri Mendeleev claimed. However, in consideration of the elements' observed chemical properties, he changed the order and placed tellurium ahead of iodine; this placement is consistent with the modern practice of ordering the elements by proton number, Z, but that number was not known or suspected at the time. A simple numbering based on periodic table position was never satisfactory, however.
Besides the case of iodine and tellurium several other pairs of elements were known to have nearly identical or reversed atomic weights, thus requiring their placement in the periodic table to be determined by their chemical properties. However the gradual identification of more and more chemically similar lanthanide elements, whose atomic number was not obvious, led to inconsistency and uncertainty in the periodic numbering of elements at least from lutetium onward. In 1911, Ernest Rutherford gave a model of the atom in which a central core held most of the atom's mass and a positive charge which, in units of the electron's charge, was to be equal to half of the atom's atomic weight, expressed in numbers of hydrogen atoms; this central charge would thus be half the atomic weight. In spite of Rutherford's estimation that gold had a central charge of about 100, a month after Rutherford's paper appeared, Antonius van den Broek first formally suggested that the central charge and number of electrons in an atom was equal to its place in the periodic table.
This proved to be the case. The experimental position improved after research by Henry Moseley in 1913. Moseley, after discussions with Bohr, at the same lab, decided to test Van den Broek's and Bohr's hypothesis directly, by seeing if spectral lines emitted from excited atoms fitted the Bohr theory's postulation that the frequency of the spectral lines be proportional to the square of Z. To do this, Moseley measured the wavelengths of the innermost photon transitions produced by the elements from aluminum to gold used as a series of movable anodic targets inside an x-ray tube; the square root of the frequency of these photons increased from one target to the next in an arithmetic progression. This led to the conclusion that the atomic number does correspond to the calculated electric charge of the nucleus, i.e. the element number Z. Among other things, Moseley demonstrated that the lanthanide series must have 15 members—no fewer and no more—which was far from obvious from the chemistry at that time.
After Moseley's death in 1915, the atomic numbers of all known elements from hydrogen to uranium were examined by his method. There were seven elements which were not found and therefore identified as still undiscovered, corresponding to atomic numbers 43, 61, 72, 75, 85, 87 and 91. From 1918 to 1947, all seven of these missing elements were discovered. By this time the first four transuranium elements had been discovered, so that the periodic table was complete with no gaps as far as curium. In 1915 the rea
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, phlogiston theory stated that all combustible materials were made of two parts.
One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.