Isotopes are variants of a particular chemical element which differ in neutron number, in nucleon number. All isotopes of a given element have the same number of protons but different numbers of neutrons in each atom; the term isotope is formed from the Greek roots isos and topos, meaning "the same place". It was coined by a Scottish doctor and writer Margaret Todd in 1913 in a suggestion to chemist Frederick Soddy; the number of protons within the atom's nucleus is called atomic number and is equal to the number of electrons in the neutral atom. Each atomic number identifies a specific element, but not the isotope; the number of nucleons in the nucleus is the atom's mass number, each isotope of a given element has a different mass number. For example, carbon-12, carbon-13, carbon-14 are three isotopes of the element carbon with mass numbers 12, 13, 14, respectively; the atomic number of carbon is 6, which means that every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7, 8 respectively.
A nuclide is a species of an atom with a specific number of protons and neutrons in the nucleus, for example carbon-13 with 6 protons and 7 neutrons. The nuclide concept emphasizes nuclear properties over chemical properties, whereas the isotope concept emphasizes chemical over nuclear; the neutron number has large effects on nuclear properties, but its effect on chemical properties is negligible for most elements. In the case of the lightest elements where the ratio of neutron number to atomic number varies the most between isotopes it has only a small effect, although it does matter in some circumstances; the term isotopes is intended to imply comparison, for example: the nuclides 126C, 136C, 146C are isotopes, but 4018Ar, 4019K, 4020Ca are isobars. However, because isotope is the older term, it is better known than nuclide, is still sometimes used in contexts where nuclide might be more appropriate, such as nuclear technology and nuclear medicine. An isotope and/or nuclide is specified by the name of the particular element followed by a hyphen and the mass number.
When a chemical symbol is used, e.g. "C" for carbon, standard notation is to indicate the mass number with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left. Because the atomic number is given by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript; the letter m is sometimes appended after the mass number to indicate a nuclear isomer, a metastable or energetically-excited nuclear state, for example 180m73Ta. The common pronunciation of the AZE notation is different from how it is written: 42He is pronounced as helium-four instead of four-two-helium, 23592U as uranium two-thirty-five or uranium-two-three-five instead of 235-92-uranium; some isotopes/nuclides are radioactive, are therefore referred to as radioisotopes or radionuclides, whereas others have never been observed to decay radioactively and are referred to as stable isotopes or stable nuclides.
For example, 14C is a radioactive form of carbon, whereas 12C and 13C are stable isotopes. There are about 339 occurring nuclides on Earth, of which 286 are primordial nuclides, meaning that they have existed since the Solar System's formation. Primordial nuclides include 32 nuclides with long half-lives and 253 that are formally considered as "stable nuclides", because they have not been observed to decay. In most cases, for obvious reasons, if an element has stable isotopes, those isotopes predominate in the elemental abundance found on Earth and in the Solar System. However, in the cases of three elements the most abundant isotope found in nature is one long-lived radioisotope of the element, despite these elements having one or more stable isotopes. Theory predicts that many "stable" isotopes/nuclides are radioactive, with long half-lives; some stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay products have yet been observed, so these isotopes are said to be "observationally stable".
The predicted half-lives for these nuclides greatly exceed the estimated age of the universe, in fact there are 27 known radionuclides with half-lives longer than the age of the universe. Adding in the radioactive nuclides that have been created artificially, there are 3,339 known nuclides; these include 905 nuclides that are either stable or have half-lives
In nuclear physics, beta decay is a type of radioactive decay in which a beta ray is emitted from an atomic nucleus. For example, beta decay of a neutron transforms it into a proton by the emission of an electron accompanied by an antineutrino, or conversely a proton is converted into a neutron by the emission of a positron with a neutrino, thus changing the nuclide type. Neither the beta particle nor its associated neutrino exist within the nucleus prior to beta decay, but are created in the decay process. By this process, unstable atoms obtain a more stable ratio of protons to neutrons; the probability of a nuclide decaying due to beta and other forms of decay is determined by its nuclear binding energy. The binding energies of all existing nuclides form what is called the nuclear band or valley of stability. For either electron or positron emission to be energetically possible, the energy release or Q value must be positive. Beta decay is a consequence of the weak force, characterized by lengthy decay times.
Nucleons are composed of up quarks and down quarks, the weak force allows a quark to change type by the exchange of a W boson and the creation of an electron/antineutrino or positron/neutrino pair. For example, a neutron, composed of two down quarks and an up quark, decays to a proton composed of a down quark and two up quarks. Decay times for many nuclides that are subject to beta decay can be thousands of years. Electron capture is sometimes included as a type of beta decay, because the basic nuclear process, mediated by the weak force, is the same. In electron capture, an inner atomic electron is captured by a proton in the nucleus, transforming it into a neutron, an electron neutrino is released; the two types of beta decay are known as beta beta plus. In beta minus decay, a neutron is converted to a proton, the process creates an electron and an electron antineutrino. Β+ decay is known as positron emission. Beta decay conserves a quantum number known as the lepton number, or the number of electrons and their associated neutrinos.
These particles have lepton number +1, while their antiparticles have lepton number −1. Since a proton or neutron has lepton number zero, β+ decay must be accompanied with an electron neutrino, while β− decay must be accompanied by an electron antineutrino. An example of electron emission is the decay of carbon-14 into nitrogen-14 with a half-life of about 5,730 years: 146C → 147N + e− + νeIn this form of decay, the original element becomes a new chemical element in a process known as nuclear transmutation; this new element has an unchanged mass number A, but an atomic number Z, increased by one. As in all nuclear decays, the decaying element is known as the parent nuclide while the resulting element is known as the daughter nuclide. Another example is the decay of hydrogen-3 into helium-3 with a half-life of about 12.3 years: 31H → 32He + e− + νeAn example of positron emission is the decay of magnesium-23 into sodium-23 with a half-life of about 11.3 s: 2312Mg → 2311Na + e+ + νeβ+ decay results in nuclear transmutation, with the resulting element having an atomic number, decreased by one.
The beta spectrum, or distribution of energy values for the beta particles, is continuous. The total energy of the decay process is divided between the electron, the antineutrino, the recoiling nuclide. In the figure to the right, an example of an electron with 0.40 MeV energy from the beta decay of 210Bi is shown. In this example, the total decay energy is 1.16 MeV, so the antineutrino has the remaining energy: 1.16-0.40=0.76 MeV. An electron at the far right of the curve would have the maximum possible kinetic energy, leaving the energy of the neutrino to be only its small rest mass. Radioactivity was discovered in 1896 by Henri Becquerel in uranium, subsequently observed by Marie and Pierre Curie in thorium and in the new elements polonium and radium. In 1899, Ernest Rutherford separated radioactive emissions into two types: alpha and beta, based on penetration of objects and ability to cause ionization. Alpha rays could be stopped by thin sheets of paper or aluminium, whereas beta rays could penetrate several millimetres of aluminium.
In 1900, Paul Villard identified a still more penetrating type of radiation, which Rutherford identified as a fundamentally new type in 1903 and termed gamma rays. Alpha and gamma are the first three letters of the Greek alphabet. In 1900, Becquerel measured the mass-to-charge ratio for beta particles by the method of J. J. Thomson used to identify the electron, he found that m/e for a beta particle is the same as for Thomson's electron, therefore suggested that the beta particle is in fact an electron. In 1901, Rutherford and Frederick Soddy showed that alpha and beta radioactivity involves the transmutation of atoms into atoms of other chemical elements. In 1913, after the products of more radioactive decays were known and Kazimierz Fajans independently proposed their radioactive displacement law, which states that beta emission from one element produces another element one place to the right in the periodic table, while alpha emission produces an element two places to the left; the study of beta decay provided the first physical evidence for the existence of the neutrino.
In both alpha and gamma decay, the resulting particle has a narrow energy distribution, since the particle carries the energy from the diffe
Internal conversion is a radioactive decay process wherein an excited nucleus interacts electromagnetically with one of the orbital electrons of the atom. This causes the electron to be emitted from the atom. Thus, in an internal conversion process, a high-energy electron is emitted from the radioactive atom, but not from the nucleus. For this reason, the high-speed electrons resulting from internal conversion are not called beta particles, since the latter come from beta decay, where they are newly created in the nuclear decay process. Internal conversion is possible whenever gamma decay is possible, except in the case where the atom is ionised. During internal conversion, the atomic number does not change, thus no transmutation of one element to another takes place. Since an electron is lost from the atom, a hole appears in an electron shell, subsequently filled by other electrons that descend to that empty, lower energy level, in the process emit characteristic X-ray, Auger electron, or both.
The atom thus emits high-energy electrons and X-ray photons, none of which originate in that nucleus. The atom supplied the energy needed to eject the electron, which in turn caused the latter events and the other emissions. Since primary electrons from internal conversion carry a fixed part of the characteristic decay energy, they have a discrete energy spectrum, rather than the spread spectrum characteristic of beta particles. Whereas the energy spectrum of beta particles plots as a broad hump, the energy spectrum of internally converted electrons plots as a single sharp peak. In the quantum mechanical model of the electron, there is a non-zero probability of finding the electron within the nucleus. During the internal conversion process, the wavefunction of an inner shell electron is said to penetrate the volume of the atomic nucleus; when this happens, the electron may couple to an excited energy state of the nucleus and take the energy of the nuclear transition directly, without an intermediate gamma ray being first produced.
The kinetic energy of the emitted electron is equal to the transition energy in the nucleus, minus the binding energy of the electron to the atom. Most internal conversion electrons come from the K shell, as these two electrons have the highest probability of being within the nucleus. However, the s states in the L, M, N shells are able to couple to the nuclear fields and cause IC electron ejections from those shells. Ratios of K-shell to other L, M, or N shell internal conversion probabilities for various nuclides have been prepared. An amount of energy exceeding the atomic binding energy of the s electron must be supplied to that electron in order to eject it from the atom to result in IC. There are a few radionuclides in which the decay energy is not sufficient to convert a 1s electron, these nuclides, to decay by internal conversion, must decay by ejecting electrons from the L or M or N shells as these binding energies are lower. Although s electrons are more for IC processes due to their superior nuclear penetration compared to electrons with orbital angular momentum, spectral studies show that p electrons are ejected in the IC process.
After the IC electron has been emitted, the atom is left with a vacancy in one of its electron shells an inner one. This hole will be filled with an electron from one of the higher shells, which causes another outer electron to fill its place in turn, causing a cascade. One or more characteristic X-rays or Auger electrons will be emitted as the remaining electrons in the atom cascade down to fill the vacancies; the decay scheme on the left shows that 203Hg produces a continuous beta spectrum with maximum energy 214 keV, that leads to an excited state of the daughter nucleus 203Tl. This state decays quickly to the ground state of 203Tl, emitting a gamma quantum of 279 keV; the figure on the right shows the electron spectrum of 203Hg, measured by means of a magnetic spectrometer. It includes the continuous beta spectrum and K-, L-, M-lines due to internal conversion. Since the binding energy of the K electrons in 203Tl amounts to 85 keV, the K line has an energy of 279 - 85 = 194 keV; because of lesser binding energies, the L- and M-lines have higher energies.
Because of the finite energy resolution of the spectrometer, the "lines" have a Gaussian shape of finite width. Internal conversion is favoured whenever the energy available for a gamma transition is small, it is the primary mode of de-excitation for 0+→0+ transitions; the 0+→0+ transitions occur where an excited nucleus has zero-spin and positive parity, decays to a ground state which has zero-spin and positive parity. In such cases, de-excitation cannot take place by way of emission of a gamma ray, since this would violate conservation of angular momentum, hence other mechanisms like IC predominate; this shows that internal conversion is not a two-step process where a gamma ray would be first emitted and converted. The competition between internal conversion and gamma decay is quantified in the form of the internal conversion coefficient, defined as α = e / γ where e is the rate of conve
In nuclear physics and nuclear chemistry, nuclear fission is a nuclear reaction or a radioactive decay process in which the nucleus of an atom splits into smaller, lighter nuclei. The fission process produces free neutrons and gamma photons, releases a large amount of energy by the energetic standards of radioactive decay. Nuclear fission of heavy elements was discovered on December 17, 1938 by German Otto Hahn and his assistant Fritz Strassmann, explained theoretically in January 1939 by Lise Meitner and her nephew Otto Robert Frisch. Frisch named the process by analogy with biological fission of living cells. For heavy nuclides, it is an exothermic reaction which can release large amounts of energy both as electromagnetic radiation and as kinetic energy of the fragments. In order for fission to produce energy, the total binding energy of the resulting elements must be more negative than that of the starting element. Fission is a form of nuclear transmutation because the resulting fragments are not the same element as the original atom.
The two nuclei produced are most of comparable but different sizes with a mass ratio of products of about 3 to 2, for common fissile isotopes. Most fissions are binary fissions, but three positively charged fragments are produced, in a ternary fission; the smallest of these fragments in ternary processes ranges in size from a proton to an argon nucleus. Apart from fission induced by a neutron and exploited by humans, a natural form of spontaneous radioactive decay is referred to as fission, occurs in high-mass-number isotopes. Spontaneous fission was discovered in 1940 by Flyorov and Kurchatov in Moscow, when they decided to confirm that, without bombardment by neutrons, the fission rate of uranium was indeed negligible, as predicted by Niels Bohr; the unpredictable composition of the products distinguishes fission from purely quantum-tunneling processes such as proton emission, alpha decay, cluster decay, which give the same products each time. Nuclear fission drives the explosion of nuclear weapons.
Both uses are possible because certain substances called nuclear fuels undergo fission when struck by fission neutrons, in turn emit neutrons when they break apart. This makes a self-sustaining nuclear chain reaction possible, releasing energy at a controlled rate in a nuclear reactor or at a rapid, uncontrolled rate in a nuclear weapon; the amount of free energy contained in nuclear fuel is millions of times the amount of free energy contained in a similar mass of chemical fuel such as gasoline, making nuclear fission a dense source of energy. The products of nuclear fission, are on average far more radioactive than the heavy elements which are fissioned as fuel, remain so for significant amounts of time, giving rise to a nuclear waste problem. Concerns over nuclear waste accumulation and over the destructive potential of nuclear weapons are a counterbalance to the peaceful desire to use fission as an energy source. Nuclear fission can occur without neutron bombardment as a type of radioactive decay.
This type of fission is rare except in a few heavy isotopes. In engineered nuclear devices all nuclear fission occurs as a "nuclear reaction" — a bombardment-driven process that results from the collision of two subatomic particles. In nuclear reactions, a subatomic particle causes changes to it. Nuclear reactions are thus driven by the mechanics of bombardment, not by the constant exponential decay and half-life characteristic of spontaneous radioactive processes. Many types of nuclear reactions are known. Nuclear fission differs from other types of nuclear reactions, in that it can be amplified and sometimes controlled via a nuclear chain reaction. In such a reaction, free neutrons released by each fission event can trigger yet more events, which in turn release more neutrons and cause more fission; the chemical element isotopes that can sustain a fission chain reaction are called nuclear fuels, are said to be fissile. The most common nuclear fuels are 239Pu; these fuels break apart into a bimodal range of chemical elements with atomic masses centering near 95 and 135 u.
Most nuclear fuels undergo spontaneous fission only slowly, decaying instead via an alpha-beta decay chain over periods of millennia to eons. In a nuclear reactor or nuclear weapon, the overwhelming majority of fission events are induced by bombardment with another particle, a neutron, itself produced by prior fission events. Nuclear fission in fissile fuels are the result of the nuclear excitation energy produced when a fissile nucleus captures a neutron; this energy, resulting from the neutron capture, is a result of the attractive nuclear force acting between the neutron and nucleus. It is enough to deform the nucleus into a double-lobed "drop", to the point that nuclear fragments exceed the distances at which the nuclear force can hold two groups of charged nucleons together and, when this happens, the two fragments complete their separation and are driven further apart by their mutually repulsive charges, in a process which becomes irreversible with greater and greater distance. A similar process occurs in fissionable iso
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
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
The neutron is a subatomic particle, symbol n or n0, with no net electric charge and a mass larger than that of a proton. Protons and neutrons constitute the nuclei of atoms. Since protons and neutrons behave within the nucleus, each has a mass of one atomic mass unit, they are both referred to as nucleons, their properties and interactions are described by nuclear physics. The chemical and nuclear properties of the nucleus are determined by the number of protons, called the atomic number, the number of neutrons, called the neutron number; the atomic mass number is the total number of nucleons. For example, carbon has atomic number 6, its abundant carbon-12 isotope has 6 neutrons, whereas its rare carbon-13 isotope has 7 neutrons; some elements occur in nature with only one stable isotope, such as fluorine. Other elements occur with many stable isotopes, such as tin with ten stable isotopes. Within the nucleus and neutrons are bound together through the nuclear force. Neutrons are required for the stability of nuclei, with the exception of the single-proton hydrogen atom.
Neutrons are produced copiously in nuclear fusion. They are a primary contributor to the nucleosynthesis of chemical elements within stars through fission and neutron capture processes; the neutron is essential to the production of nuclear power. In the decade after the neutron was discovered by James Chadwick in 1932, neutrons were used to induce many different types of nuclear transmutations. With the discovery of nuclear fission in 1938, it was realized that, if a fission event produced neutrons, each of these neutrons might cause further fission events, etc. in a cascade known as a nuclear chain reaction. These events and findings led to the first self-sustaining nuclear reactor and the first nuclear weapon. Free neutrons, while not directly ionizing atoms, cause ionizing radiation; as such they can be a biological hazard, depending upon dose. A small natural "neutron background" flux of free neutrons exists on Earth, caused by cosmic ray showers, by the natural radioactivity of spontaneously fissionable elements in the Earth's crust.
Dedicated neutron sources like neutron generators, research reactors and spallation sources produce free neutrons for use in irradiation and in neutron scattering experiments. An atomic nucleus is formed by a number of protons, Z, a number of neutrons, N, bound together by the nuclear force; the atomic number defines the chemical properties of the atom, the neutron number determines the isotope or nuclide. The terms isotope and nuclide are used synonymously, but they refer to chemical and nuclear properties, respectively. Speaking, isotopes are two or more nuclides with the same number of protons; the atomic mass number, symbol A, equals Z+N. Nuclides with the same atomic mass number are called isobars; 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 most common nuclide of the common chemical element lead, 208Pb, has 82 protons and 126 neutrons, for example. The table of nuclides comprises all the known nuclides. Though it is not a chemical element, the neutron is included in this table; the free neutron has 1.674927471 × 10 − 27 kg, or 1.00866491588 u. The neutron has a mean square radius of about 0.8×10−15 m, or 0.8 fm, it is a spin-½ fermion. The neutron has no measurable electric charge. With its positive electric charge, the proton is directly influenced by electric fields, whereas the neutron is unaffected by electric fields; the neutron has a magnetic moment, however. The neutron's magnetic moment has a negative value, because its orientation is opposite to the neutron's spin. A free neutron is unstable, decaying to a proton and antineutrino with a mean lifetime of just under 15 minutes; this radioactive decay, known as beta decay, is possible because the mass of the neutron is greater than the proton. The free proton is stable. Neutrons or protons bound in a nucleus can be stable or unstable, depending on the nuclide.
Beta decay, in which neutrons decay to protons, or vice versa, is governed by the weak force, it requires the emission or absorption of electrons and neutrinos, or their antiparticles. Protons and neutrons behave identically under the influence of the nuclear force within the nucleus; the concept of isospin, in which the proton and neutron are viewed as two quantum states of the same particle, is used to model the interactions of nucleons by the nuclear or weak forces. Because of the strength of the nuclear force at short distances, the binding energy of nucleons is more than seven orders of magnitude larger than the electromagnetic energy binding electrons in atoms. Nuclear reactions therefore have an energy density, more than ten million times that of chemical reactions; because of the mass–energy equivalence, nuclear binding energies reduce the mass of nuclei. The ability of the nuclear force to store energy arising from the electromagnetic repulsion of nuclear components is the basis for most of the energy that makes nuclear reactors or bombs possible.
In nuclear fission, the absorption of a neutron by a heavy nuclide causes the nuclide to become unstable and break into light nuclides and additional neu