A proton is a subatomic particle, symbol p or p+, with a positive electric charge of +1e elementary charge and a mass less than that of a neutron. Protons and neutrons, each with masses of one atomic mass unit, are collectively referred to as "nucleons". One or more protons are present in the nucleus of every atom; the number of protons in the nucleus is the defining property of an element, is referred to as the atomic number. Since each element has a unique number of protons, each element has its own unique atomic number; the word proton is Greek for "first", this name was given to the hydrogen nucleus by Ernest Rutherford in 1920. In previous years, Rutherford had discovered that the hydrogen nucleus could be extracted from the nuclei of nitrogen by atomic collisions. Protons were therefore a candidate to be a fundamental particle, hence a building block of nitrogen and all other heavier atomic nuclei. In the modern Standard Model of particle physics, protons are hadrons, like neutrons, the other nucleon, are composed of three quarks.
Although protons were considered fundamental or elementary particles, they are now known to be composed of three valence quarks: two up quarks of charge +2/3e and one down quark of charge –1/3e. The rest masses of quarks contribute only about 1% of a proton's mass, however; the remainder of a proton's mass is due to quantum chromodynamics binding energy, which includes the kinetic energy of the quarks and the energy of the gluon fields that bind the quarks together. Because protons are not fundamental particles, they possess a physical size, though not a definite one. At sufficiently low temperatures, free protons will bind to electrons. However, the character of such bound protons does not change, they remain protons. A fast proton moving through matter will slow by interactions with electrons and nuclei, until it is captured by the electron cloud of an atom; the result is a protonated atom, a chemical compound of hydrogen. In vacuum, when free electrons are present, a sufficiently slow proton may pick up a single free electron, becoming a neutral hydrogen atom, chemically a free radical.
Such "free hydrogen atoms" tend to react chemically with many other types of atoms at sufficiently low energies. When free hydrogen atoms react with each other, they form neutral hydrogen molecules, which are the most common molecular component of molecular clouds in interstellar space. Protons are composed of three valence quarks, making them baryons; the two up quarks and one down quark of a proton are held together by the strong force, mediated by gluons. A modern perspective has a proton composed of the valence quarks, the gluons, transitory pairs of sea quarks. Protons have a positive charge distribution which decays exponentially, with a mean square radius of about 0.8 fm. Protons and neutrons are both nucleons, which may be bound together by the nuclear force to form atomic nuclei; the nucleus of the most common isotope of the hydrogen atom is a lone proton. The nuclei of the heavy hydrogen isotopes deuterium and tritium contain one proton bound to one and two neutrons, respectively. All other types of atomic nuclei are composed of two or more protons and various numbers of neutrons.
The concept of a hydrogen-like particle as a constituent of other atoms was developed over a long period. As early as 1815, William Prout proposed that all atoms are composed of hydrogen atoms, based on a simplistic interpretation of early values of atomic weights, disproved when more accurate values were measured. In 1886, Eugen Goldstein discovered canal rays and showed that they were positively charged particles produced from gases. However, since particles from different gases had different values of charge-to-mass ratio, they could not be identified with a single particle, unlike the negative electrons discovered by J. J. Thomson. Wilhelm Wien in 1898 identified the hydrogen ion as particle with highest charge-to-mass ratio in ionized gases. Following the discovery of the atomic nucleus by Ernest Rutherford in 1911, Antonius van den Broek proposed that the place of each element in the periodic table is equal to its nuclear charge; this was confirmed experimentally by Henry Moseley in 1913 using X-ray spectra.
In 1917, Rutherford proved that the hydrogen nucleus is present in other nuclei, a result described as the discovery of protons. Rutherford had earlier learned to produce hydrogen nuclei as a type of radiation produced as a product of the impact of alpha particles on nitrogen gas, recognize them by their unique penetration signature in air and their appearance in scintillation detectors; these experiments were begun when Rutherford had noticed that, when alpha particles were shot into air, his scintillation detectors showed the signatures of typical hydrogen nuclei as a product. After experimentation Rutherford traced the reaction to the nitrogen in air, found that when alphas were produced into pure nitrogen gas, the effect was larger. Rutherford determined that this hydrogen could have come only from the nitrogen, therefore nitrogen must contain hydrogen nuclei. One hydrogen nucleus was being knocked off by the impact of the alpha particle, producing oxygen-17 in the process; this was 14N + α → 17O + p.
(This reaction wo
Boron is a chemical element with symbol B and atomic number 5. Produced by cosmic ray spallation and supernovae and not by stellar nucleosynthesis, it is a low-abundance element in the Solar system and in the Earth's crust. Boron is concentrated on Earth by the water-solubility of its more common occurring compounds, the borate minerals; these are mined industrially as evaporites, such as kernite. The largest known boron deposits are in the largest producer of boron minerals. Elemental boron is a metalloid, found in small amounts in meteoroids but chemically uncombined boron is not otherwise found on Earth. Industrially pure boron is produced with difficulty because of refractory contamination by carbon or other elements. Several allotropes of boron exist: amorphous boron is a brown powder; the primary use of elemental boron is as boron filaments with applications similar to carbon fibers in some high-strength materials. Boron is used in chemical compounds. About half of all boron consumed globally is an additive in fiberglass for insulation and structural materials.
The next leading use is in polymers and ceramics in high-strength, lightweight structural and refractory materials. Borosilicate glass is desired for its greater strength and thermal shock resistance than ordinary soda lime glass. Boron as sodium perborate is used as a bleach. A small amount of boron is used as a dopant in semiconductors, reagent intermediates in the synthesis of organic fine chemicals. A few boron-containing organic pharmaceuticals are in study. Natural boron is composed of two stable isotopes, one of which has a number of uses as a neutron-capturing agent. In biology, borates have low toxicity in mammals, but are more toxic to arthropods and are used as insecticides. Boric acid is mildly antimicrobial, several natural boron-containing organic antibiotics are known. Boron is an essential plant nutrient and boron compounds such as borax and boric acid are used as fertilizers in agriculture, although it's only required in small amounts, with excess being toxic. Boron compounds play a strengthening role in the cell walls of all plants.
There is no consensus on whether boron is an essential nutrient for mammals, including humans, although there is some evidence it supports bone health. The word boron was coined from borax, the mineral from which it was isolated, by analogy with carbon, which boron resembles chemically. Borax, its mineral form known as tincal, glazes were used in China from AD 300, some crude borax reached the West, where the Perso-Arab alchemist Jābir ibn Hayyān mentioned it in AD 700. Marco Polo brought some glazes back to Italy in the 13th century. Agricola, around 1600, reports the use of borax as a flux in metallurgy. In 1777, boric acid was recognized in the hot springs near Florence and became known as sal sedativum, with medical uses; the rare mineral is called sassolite, found at Sasso, Italy. Sasso was the main source of European borax from 1827 to 1872. Boron compounds were rarely used until the late 1800s when Francis Marion Smith's Pacific Coast Borax Company first popularized and produced them in volume at low cost.
Boron was not recognized as an element until it was isolated by Sir Humphry Davy and by Joseph Louis Gay-Lussac and Louis Jacques Thénard. In 1808 Davy observed that electric current sent through a solution of borates produced a brown precipitate on one of the electrodes. In his subsequent experiments, he used potassium to reduce boric acid instead of electrolysis, he named the element boracium. Gay-Lussac and Thénard used iron to reduce boric acid at high temperatures. By oxidizing boron with air, they showed. Jöns Jakob Berzelius identified boron as an element in 1824. Pure boron was arguably first produced by the American chemist Ezekiel Weintraub in 1909; the earliest routes to elemental boron involved the reduction of boric oxide with metals such as magnesium or aluminium. However, the product is always contaminated with borides of those metals. Pure boron can be prepared by reducing volatile boron halides with hydrogen at high temperatures. Ultrapure boron for use in the semiconductor industry is produced by the decomposition of diborane at high temperatures and further purified by the zone melting or Czochralski processes.
The production of boron compounds does not involve the formation of elemental boron, but exploits the convenient availability of borates. Boron is similar to carbon in its capability to form stable covalently bonded molecular networks. Nominally disordered boron contains regular boron icosahedra which are, bonded randomly to each other without long-range order. Crystalline boron is a hard, black material with a melting point of above 2000 °C, it forms four major polymorphs: β-rhombohedral, γ and β-tetragonal. Most of the phases are based on B12 icosahedra, but the γ-phase can be described as a rocksalt-type arrangement of the icosahedra and B2 atomic pairs, it can be produced by compressing other boron phases to 12–20 GPa and heating to 1500–1800 °C. The T phase is produced at similar pressures, but higher temperatures of 1800–2200 °C; as to the α and β phases, they might both coexist at ambient conditions with the β phase being more stable
Plasma is one of the four fundamental states of matter, was first described by chemist Irving Langmuir in the 1920s. Plasma can be artificially generated by heating or subjecting a neutral gas to a strong electromagnetic field to the point where an ionized gaseous substance becomes electrically conductive, long-range electromagnetic fields dominate the behaviour of the matter. Plasma and ionized gases have properties and display behaviours unlike those of the other states, the transition between them is a matter of nomenclature and subject to interpretation. Based on the surrounding environmental temperature and density ionized or ionized forms of plasma may be produced. Neon signs and lightning are examples of ionized plasma; the Earth's ionosphere is a plasma and the magnetosphere contains plasma in the Earth's surrounding space environment. The interior of the Sun is an example of ionized plasma, along with the solar corona and stars. Positive charges in ions are achieved by stripping away electrons orbiting the atomic nuclei, where the total number of electrons removed is related to either increasing temperature or the local density of other ionized matter.
This can be accompanied by the dissociation of molecular bonds, though this process is distinctly different from chemical processes of ion interactions in liquids or the behaviour of shared ions in metals. The response of plasma to electromagnetic fields is used in many modern technological devices, such as plasma televisions or plasma etching. Plasma may be the most abundant form of ordinary matter in the universe, although this hypothesis is tentative based on the existence and unknown properties of dark matter. Plasma is associated with stars, extending to the rarefied intracluster medium and the intergalactic regions; the word plasma comes from Ancient Greek πλάσμα, meaning'moldable substance' or'jelly', describes the behaviour of the ionized atomic nuclei and the electrons within the surrounding region of the plasma. Each of these nuclei are suspended in a movable sea of electrons. Plasma was first identified in a Crookes tube, so described by Sir William Crookes in 1879; the nature of this "cathode ray" matter was subsequently identified by British physicist Sir J.
J. Thomson in 1897; the term "plasma" was coined by Irving Langmuir in 1928. Lewi Tonks and Harold Mott-Smith, both of whom worked with Irving Langmuir in the 1920s, recall that Langmuir first used the word "plasma" in analogy with blood. Mott-Smith recalls, in particular, that the transport of electrons from thermionic filaments reminded Langmuir of "the way blood plasma carries red and white corpuscles and germs."Langmuir described the plasma he observed as follows: "Except near the electrodes, where there are sheaths containing few electrons, the ionized gas contains ions and electrons in about equal numbers so that the resultant space charge is small. We shall use the name plasma to describe this region containing balanced charges of ions and electrons." Plasma is a state of matter in which an ionized gaseous substance becomes electrically conductive to the point that long-range electric and magnetic fields dominate the behaviour of the matter. The plasma state can be contrasted with the other states: solid and gas.
Plasma is an electrically neutral medium of unbound negative particles. Although these particles are unbound, they are not "free" in the sense of not experiencing forces. Moving charged particles generate an electric current within a magnetic field, any movement of a charged plasma particle affects and is affected by the fields created by the other charges. In turn this governs collective behaviour with many degrees of variation. Three factors define a plasma: The plasma approximation: The plasma approximation applies when the plasma parameter, Λ, representing the number of charge carriers within a sphere surrounding a given charged particle, is sufficiently high as to shield the electrostatic influence of the particle outside of the sphere. Bulk interactions: The Debye screening length is short compared to the physical size of the plasma; this criterion means that interactions in the bulk of the plasma are more important than those at its edges, where boundary effects may take place. When this criterion is satisfied, the plasma is quasineutral.
Plasma frequency: The electron plasma frequency is large compared to the electron-neutral collision frequency. When this condition is valid, electrostatic interactions dominate over the processes of ordinary gas kinetics. Plasma temperature is measured in kelvin or electronvolts and is, informally, a measure of the thermal kinetic energy per particle. High temperatures are needed to sustain ionisation, a defining feature of a plasma; the degree of plasma ionisation is determined by the electron temperature relative to the ionization energy, in a relationship called the Saha equation. At low temperatures and electrons tend to recombine into bound states—atoms—and the plasma will become a gas. In most cases the electrons are close enough to thermal equilibrium that their temperature is well-defined; because of the large difference in ma
The octet rule is a chemical rule of thumb that reflects observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas. The rule is applicable to carbon, nitrogen and the halogens, but to metals such as sodium or magnesium; the valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide. The electrons shared by the two atoms in a covalent bond are counted once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, two from the oxygen itself and two from the carbon. All four of these electrons are counted in both the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity and the second a nonmetal of high electronegativity. A chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively.
The first electron affinity of chlorine is -328.8 kJ per mole of chlorine atoms. Adding a second electron to chlorine requires energy, energy that cannot be recovered by the formation of a chemical bond; the result is that chlorine will often form a compound in which it has eight electrons in its outer shell. A sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy, +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the second electron resides in the deeper second electron shell, the second ionization energy required for its removal is much larger: +4562.4 kJ per mole. Thus sodium will, in most cases, form a compound in which it has lost a single electron and have a full outer shell of eight electrons, or octet; the energy required to transfer an electron from a sodium atom to a chlorine atom is small: +495.8 − 328.8 = +167 kJ mol−1.
This energy is offset by the lattice energy of sodium chloride: −787.3 kJ mol−1. This completes the explanation of the octet rule in this case. In the late 19th century it was known that coordination compounds were formed by the combination of atoms or molecules in such a manner that the valencies of the atoms involved became satisfied. In 1893, Alfred Werner showed that the number of atoms or groups associated with a central atom is 4 or 6. In 1904 Richard Abegg was one of the first to extend the concept of coordination number to a concept of valence in which he distinguished atoms as electron donors or acceptors, leading to positive and negative valence states that resemble the modern concept of oxidation states. Abegg noted that the difference between the maximum positive and negative valences of an element under his model is eight. In 1916, Gilbert N. Lewis referred to this insight as Abegg's rule and used it to help formulate his cubical atom model and the "rule of eight", which began to distinguish between valence and valence electrons.
In 1919 Irving Langmuir refined these concepts further and renamed them the "cubical octet atom" and "octet theory". The "octet theory" evolved into what is now known as the "octet rule". Walther Kossel and Gilbert N. Lewis saw that noble gases did not have the tendency of taking part in chemical reactions under ordinary conditions. On the basis of this observation they concluded that atoms of noble gases are stable and on the basis of this conclusion they proposed a theory of valency known as "Electronic Theory of valency" in 1916: During the formation of a chemical bond, atoms combine together by gaining, losing or sharing electrons in such a way that they acquire nearest noble gas configuration; the quantum theory of the atom explains the eight electrons as a closed shell with an s2p6 electron configuration. A closed-shell configuration is one in which low-lying energy levels are full and higher energy levels are empty. For example, the neon atom ground state has an empty n = 3 shell. According to the octet rule, the atoms before and after neon in the periodic table, tend to attain a similar configuration by gaining, losing, or sharing electrons.
The argon atom has an analogous 3s2 3p6 configuration. There is an empty 3d level, but it is at higher energy than 3s and 3p, so that 3s2 3p6 is still considered a closed shell for chemical purposes; the atoms before and after argon tend to attain this configuration in compounds. There are, some hypervalent molecules in which the 3d level may play a part in the bonding, although this is controversial. For helium there is no 1p level according to the quantum theory, so that 1s2 is a closed shell with no p electrons; the atoms before and after helium follow a duet rule and tend to have the same 1s2 configuration as helium. The octet rule is only applicable to main group elements and there are many molecules that do not obey the octet rule; these molecules can be divided into two types: unstable intermediates that react so as to attain stability, stable molecules that follow other electron counting rules. Although stable odd-electron molecules and hypervalent molecules are taught as v
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
In nuclear chemistry, nuclear fusion is a reaction in which two or more atomic nuclei are combined to form one or more different atomic nuclei and subatomic particles. The difference in mass between the reactants and products is manifested as either the release or absorption of energy; this difference in mass arises due to the difference in atomic "binding energy" between the atomic nuclei before and after the reaction. Fusion is other high magnitude stars. A fusion process that produces a nucleus lighter than iron-56 or nickel-62 will yield a net energy release; these elements have the smallest mass per nucleon and the largest binding energy per nucleon, respectively. Fusion of light elements toward these releases energy, while a fusion producing nuclei heavier than these elements will result in energy retained by the resulting nucleons, the resulting reaction is endothermic; the opposite is true for nuclear fission. This means that the lighter elements, such as helium, are in general more fusible.
The extreme astrophysical event of a supernova can produce enough energy to fuse nuclei into elements heavier than iron. In 1920, Arthur Eddington suggested hydrogen-helium fusion could be the primary source of stellar energy. Quantum tunneling was discovered by Friedrich Hund in 1929, shortly afterwards Robert Atkinson and Fritz Houtermans used the measured masses of light elements to show that large amounts of energy could be released by fusing small nuclei. Building on the early experiments in nuclear transmutation by Ernest Rutherford, laboratory fusion of hydrogen isotopes was accomplished by Mark Oliphant in 1932. In the remainder of that decade, the theory of the main cycle of nuclear fusion in stars were worked out by Hans Bethe. Research into fusion for military purposes began in the early 1940s as part of the Manhattan Project. Fusion was accomplished in 1951 with the Greenhouse Item nuclear test. Nuclear fusion on a large scale in an explosion was first carried out on 1 November 1952, in the Ivy Mike hydrogen bomb test.
Research into developing controlled thermonuclear fusion for civil purposes began in earnest in the 1940s, it continues to this day. The release of energy with the fusion of light elements is due to the interplay of two opposing forces: the nuclear force, which combines together protons and neutrons, the Coulomb force, which causes protons to repel each other. Protons are positively charged and repel each other by the Coulomb force, but they can nonetheless stick together, demonstrating the existence of another, short-range, force referred to as nuclear attraction. Light nuclei are sufficiently small and proton-poor allowing the nuclear force to overcome repulsion; this is because the nucleus is sufficiently small that all nucleons feel the short-range attractive force at least as as they feel the infinite-range Coulomb repulsion. Building up nuclei from lighter nuclei by fusion releases the extra energy from the net attraction of particles. For larger nuclei, however, no energy is released, since the nuclear force is short-range and cannot continue to act across longer atomic length scales.
Thus, energy is not released with the fusion of such nuclei. Fusion powers stars and produces all elements in a process called nucleosynthesis; the Sun is a main-sequence star, and, as such, generates its energy by nuclear fusion of hydrogen nuclei into helium. In its core, the Sun fuses 620 million metric tons of hydrogen and makes 606 million metric tons of helium each second; the fusion of lighter elements in stars releases the mass that always accompanies it. For example, in the fusion of two hydrogen nuclei to form helium, 0.7% of the mass is carried away in the form of kinetic energy of an alpha particle or other forms of energy, such as electromagnetic radiation. It takes considerable energy to force nuclei to fuse those of the lightest element, hydrogen; when accelerated to high enough speeds, nuclei can overcome this electrostatic repulsion and brought close enough such that the attractive nuclear force is greater than the repulsive Coulomb force. The strong force grows once the nuclei are close enough, the fusing nucleons can "fall" into each other and the result is fusion and net energy produced.
The fusion of lighter nuclei, which creates a heavier nucleus and a free neutron or proton releases more energy than it takes to force the nuclei together. Energy released in most nuclear reactions is much larger than in chemical reactions, because the binding energy that holds a nucleus together is greater than the energy that holds electrons to a nucleus. For example, the ionization energy gained by adding an electron to a hydrogen nucleus is 13.6 eV—less than one-millionth of the 17.6 MeV released in the deuterium–tritium reaction shown in the adjacent diagram. The complete conversion of one gram of matter would release 9×1013 joules of energy. Fusion reactions have an energy density many times greater than nuclear fission. Only direct conversion of mass into energy, such as that caused by the annihilatory collision of matter and antimatter, is more energetic per unit of mass than nuclear fusion. Research into using fusion for the p
A chemical element is a species of atom having the same number of protons in their atomic nuclei. For example, the atomic number of oxygen is 8, so the element oxygen consists of all atoms which have 8 protons. 118 elements have been identified, of which the first 94 occur on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have radionuclides, which decay over time into other elements. Iron is the most abundant element making up Earth, while oxygen is the most common element in the Earth's crust. Chemical elements constitute all of the ordinary matter of the universe; however astronomical observations suggest that ordinary observable matter makes up only about 15% of the matter in the universe: the remainder is dark matter. The two lightest elements and helium, were formed in the Big Bang and are the most common elements in the universe; the next three elements were formed by cosmic ray spallation, are thus rarer than heavier elements.
Formation of elements with from 6 to 26 protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis. The high abundance of oxygen and iron on Earth reflects their common production in such stars. Elements with greater than 26 protons are formed by supernova nucleosynthesis in supernovae, when they explode, blast these elements as supernova remnants far into space, where they may become incorporated into planets when they are formed; the term "element" is used for atoms with a given number of protons as well as for a pure chemical substance consisting of a single element. For the second meaning, the terms "elementary substance" and "simple substance" have been suggested, but they have not gained much acceptance in English chemical literature, whereas in some other languages their equivalent is used. A single element can form multiple substances differing in their structure; when different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds.
Only a minority of elements are found uncombined as pure minerals. Among the more common of such native elements are copper, gold and sulfur. All but a few of the most inert elements, such as noble gases and noble metals, are found on Earth in chemically combined form, as chemical compounds. While about 32 of the chemical elements occur on Earth in native uncombined forms, most of these occur as mixtures. For example, atmospheric air is a mixture of nitrogen and argon, native solid elements occur in alloys, such as that of iron and nickel; the history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur and gold. Civilizations extracted elemental copper, tin and iron from their ores by smelting, using charcoal. Alchemists and chemists subsequently identified many more; the properties of the chemical elements are summarized in the periodic table, which organizes the elements by increasing atomic number into rows in which the columns share recurring physical and chemical properties.
Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, most of them in low degrees of impurities. The lightest chemical elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe in a ratio of around 3:1 by mass, along with tiny traces of the next two elements and beryllium. All other elements found in nature were made by various natural methods of nucleosynthesis. On Earth, small amounts of new atoms are produced in nucleogenic reactions, or in cosmogenic processes, such as cosmic ray spallation. New atoms are naturally produced on Earth as radiogenic daughter isotopes of ongoing radioactive decay processes such as alpha decay, beta decay, spontaneous fission, cluster decay, other rarer modes of decay. Of the 94 occurring elements, those with atomic numbers 1 through 82 each have at least one stable isotope. Isotopes considered stable are those. Elements with atomic numbers 83 through 94 are unstable to the point that radioactive decay of all isotopes can be detected.
Some of these elements, notably bismuth and uranium, have one or more isotopes with half-lives long enough to survive as remnants of the explosive stellar nucleosynthesis that produced the heavy metals before the formation of our Solar System. At over 1.9×1019 years, over a billion times longer than the current estimated age of the universe, bismuth-209 has the longest known alpha decay half-life of any occurring element, is always considered on par with the 80 stable elements. The heaviest elements undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized; as of 2010, there are 118 known elements (in this context, "known" means observed well enough from just a few de