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
Island of stability
In nuclear physics, the island of stability is the prediction that a set of superheavy nuclides with magic numbers of protons and neutrons will temporarily reverse the trend of decreasing stability in elements heavier than uranium. Various predictions have been made regarding the exact location of the island of stability, though it is thought to center near copernicium and flerovium isotopes approaching the predicted closed shell at neutron number N = 184, it is thought that the closed shell will confer additional stability towards fission, while leading to longer half-lives towards alpha decay. While these effects are expected to be greatest near atomic number Z = 114 and N = 184, the region of increased stability is expected to encompass several neighboring elements, there may be additional islands of stability around heavier nuclei that are doubly magic. Estimates of the stability of the elements on the island are around a half-life of minutes or days. Although the nuclear shell model predicting magic numbers has existed since the 1940s, the existence of long-lived superheavy nuclides has not been definitively demonstrated.
Like the rest of the superheavy elements, the nuclides on the island of stability have never been found in nature. Scientists have not found a way to carry out such a reaction; the successful synthesis of superheavy elements up to oganesson in recent years demonstrates a slight stabilizing effect around elements 110–114 that may continue in unknown isotopes, supporting the existence of the island of stability. The composition of an atomic nucleus is determined by the number of protons Z and the number of neutrons N; the atomic number Z determines the position of an element in the periodic table, but the more than 3000 nuclides are represented in a chart with Z and N for its axes and the half-life for radioactive decay indicated for each unstable nuclide. 253 of the nuclides are thought to be stable, these follow a general trend in which the number of neutrons rises more than the number of protons. The last element in the periodic table that has a stable isotope is lead, with stability decreasing in heavier elements, the half-lives of nuclei decrease when there is a lopsided neutron-proton ratio.
The stability of a nucleus is determined by its binding energy, with higher binding energy conferring greater stability. The binding energy per nucleon increases with atomic number to a broad plateau around A = 60 declines. If a nucleus can be split into two parts that have a lower total energy, it is unstable; the nucleus can hold together for a finite time because there is a potential barrier opposing the split, but this barrier can be crossed by quantum tunnelling. The lower the barrier and the masses of the constituents, the greater the probability per unit time of a split. Protons in a nucleus are bound together by the strong force, which counterbalances the Coulomb repulsion between positively charged protons. In heavier nuclei, larger numbers of neutrons are needed to reduce repulsion and confer additional stability. So, as physicists started to synthesize elements that are not found in nature, they found the stability decreased as the nuclei became heavier. Thus, they speculated; the discoverers of plutonium considered thinking it was the last.
And it seemed that element 108 might be the limit. The possible existence of superheavy elements with atomic numbers well beyond that of uranium had been suggested as far back as 1955 by John Archibald Wheeler, but the idea did not attract wide interest until a decade after improvements in the nuclear shell model. In this model, the atomic nucleus is built up in "shells", analogous to electron shells in atoms. Independently of each other and protons have energy levels that are close together, but after a given shell is filled, it takes more energy to start filling the next. Thus, the binding energy per nucleon reaches a local maximum and nuclei with filled shells are more stable than those without; the numbers of nucleons for which shells are filled are called magic numbers, magic numbers of 2, 8, 20, 28, 50, 82 and 126 have been observed for neutrons, the next number is predicted to be 184. Protons share the first six of these magic numbers, 126 had been predicted since the 1940s. Nuclides with a magic number of each are referred to as "doubly magic" and are more stable than nearby nuclides as a result of greater binding energies.
In the late 1960s, more sophisticated shell models by William Myers and Władysław Świątecki, by H. Meldner, taking into account Coulomb repulsion, changed the prediction for the next proton magic number from 126 to 114; some Russian physicists argued for the existence of the doubly magic nuclide 298Fl, rather than 310Ubh, predicted to be doubly magic as early as 1957. Myers and Świątecki appear to have coined the term "island of stability", Glenn Seaborg a discoverer of many of the superheavy elements adopted the term and promoted it. Subsequently, estimates of the proton magic number have ranged from 114 to 126, there is still no consensus. Interest in island of stability grew over the next few years, as some calculations suggested that it
Nuclear physics is the field of physics that studies atomic nuclei and their constituents and interactions. Other forms of nuclear matter are studied. Nuclear physics should not be confused with atomic physics, which studies the atom as a whole, including its electrons. Discoveries in nuclear physics have led to applications in many fields; this includes nuclear power, nuclear weapons, nuclear medicine and magnetic resonance imaging and agricultural isotopes, ion implantation in materials engineering, radiocarbon dating in geology and archaeology. Such applications are studied in the field of nuclear engineering. Particle physics evolved out of nuclear physics and the two fields are taught in close association. Nuclear astrophysics, the application of nuclear physics to astrophysics, is crucial in explaining the inner workings of stars and the origin of the chemical elements; the history of nuclear physics as a discipline distinct from atomic physics starts with the discovery of radioactivity by Henri Becquerel in 1896, while investigating phosphorescence in uranium salts.
The discovery of the electron by J. J. Thomson a year was an indication that the atom had internal structure. At the beginning of the 20th century the accepted model of the atom was J. J. Thomson's "plum pudding" model in which the atom was a positively charged ball with smaller negatively charged electrons embedded inside it. In the years that followed, radioactivity was extensively investigated, notably by Marie and Pierre Curie as well as by Ernest Rutherford and his collaborators. By the turn of the century physicists had discovered three types of radiation emanating from atoms, which they named alpha and gamma radiation. Experiments by Otto Hahn in 1911 and by James Chadwick in 1914 discovered that the beta decay spectrum was continuous rather than discrete; that is, electrons were ejected from the atom with a continuous range of energies, rather than the discrete amounts of energy that were observed in gamma and alpha decays. This was a problem for nuclear physics at the time, because it seemed to indicate that energy was not conserved in these decays.
The 1903 Nobel Prize in Physics was awarded jointly to Becquerel for his discovery and to Marie and Pierre Curie for their subsequent research into radioactivity. Rutherford was awarded the Nobel Prize in Chemistry in 1908 for his "investigations into the disintegration of the elements and the chemistry of radioactive substances". In 1905 Albert Einstein formulated the idea of mass–energy equivalence. While the work on radioactivity by Becquerel and Marie Curie predates this, an explanation of the source of the energy of radioactivity would have to wait for the discovery that the nucleus itself was composed of smaller constituents, the nucleons. In 1906 Ernest Rutherford published "Retardation of the α Particle from Radium in passing through matter." Hans Geiger expanded on this work in a communication to the Royal Society with experiments he and Rutherford had done, passing alpha particles through air, aluminum foil and gold leaf. More work was published in 1909 by Geiger and Ernest Marsden, further expanded work was published in 1910 by Geiger.
In 1911–1912 Rutherford went before the Royal Society to explain the experiments and propound the new theory of the atomic nucleus as we now understand it. The key experiment behind this announcement was performed in 1910 at the University of Manchester: Ernest Rutherford's team performed a remarkable experiment in which Geiger and Marsden under Rutherford's supervision fired alpha particles at a thin film of gold foil; the plum pudding model had predicted that the alpha particles should come out of the foil with their trajectories being at most bent. But Rutherford instructed his team to look for something that shocked him to observe: a few particles were scattered through large angles completely backwards in some cases, he likened it to firing a bullet at tissue paper and having it bounce off. The discovery, with Rutherford's analysis of the data in 1911, led to the Rutherford model of the atom, in which the atom had a small dense nucleus containing most of its mass, consisting of heavy positively charged particles with embedded electrons in order to balance out the charge.
As an example, in this model nitrogen-14 consisted of a nucleus with 14 protons and 7 electrons and the nucleus was surrounded by 7 more orbiting electrons. Around 1920, Arthur Eddington anticipated the discovery and mechanism of nuclear fusion processes in stars, in his paper The Internal Constitution of the Stars. At that time, the source of stellar energy was a complete mystery; this was a remarkable development since at that time fusion and thermonuclear energy, that stars are composed of hydrogen, had not yet been discovered. The Rutherford model worked quite well until studies of nuclear spin were carried out by Franco Rasetti at the California Institute of Technology in 1929. By 1925 it was known that protons and electrons each had a spin of +/-1⁄2. In the Rutherford model of nitrogen-14, 20 of the total 21 nuclear particles should have paired up to cancel each other's spin, the final odd particle should have left the nucleus with a net spin of 1⁄2. Rasetti discovered, that nitrogen-14 had a spin of 1.
In 1932 Chadwick realized that radiation, observed by Walther Bothe, Herbert Becker, Irène and Frédéric Joliot-Curie was due to a neutral particle of about the same mass as the proton, that he called the neutron (following a su
The nuclear force is a force that acts between the protons and neutrons of atoms. Neutrons and protons, both nucleons, are affected by the nuclear force identically. Since protons have charge +1 e, they experience an electric force that tends to push them apart, but at short range the attractive nuclear force is strong enough to overcome the electromagnetic force; the nuclear force binds nucleons into atomic nuclei. The nuclear force is powerfully attractive between nucleons at distances of about 1 femtometre, but it decreases to insignificance at distances beyond about 2.5 fm. At distances less than 0.7 fm, the nuclear force becomes repulsive. This repulsive component is responsible for the physical size of nuclei, since the nucleons can come no closer than the force allows. By comparison, the size of an atom, measured in angstroms, is five orders of magnitude larger; the nuclear force is not simple, since it depends on the nucleon spins, has a tensor component, may depend on the relative momentum of the nucleons.
The strong nuclear force is one of the fundamental forces of nature. The nuclear force plays an essential role in storing energy, used in nuclear power and nuclear weapons. Work is required to bring charged protons together against their electric repulsion; this energy is stored when the protons and neutrons are bound together by the nuclear force to form a nucleus. The mass of a nucleus is less than the sum total of the individual masses of the protons and neutrons; the difference in masses is known as the mass defect, which can be expressed as an energy equivalent. Energy is released; this energy is the electromagnetic potential energy, released when the nuclear force no longer holds the charged nuclear fragments together. A quantitative description of the nuclear force relies on equations that are empirical; these equations model the internucleon potential energies, or potentials. The constants for the equations are phenomenological, that is, determined by fitting the equations to experimental data.
The internucleon potentials attempt to describe the properties of nucleon–nucleon interaction. Once determined, any given potential can be used in, e.g. the Schrödinger equation to determine the quantum mechanical properties of the nucleon system. The discovery of the neutron in 1932 revealed that atomic nuclei were made of protons and neutrons, held together by an attractive force. By 1935 the nuclear force was conceived to be transmitted by particles called mesons; this theoretical development included a description of the Yukawa potential, an early example of a nuclear potential. Mesons, predicted by theory, were discovered experimentally in 1947. By the 1970s, the quark model had been developed, by which the mesons and nucleons were viewed as composed of quarks and gluons. By this new model, the nuclear force, resulting from the exchange of mesons between neighboring nucleons, is a residual effect of the strong force. While the nuclear force is associated with nucleons, more this force is felt between hadrons, or particles composed of quarks.
At small separations between nucleons the force becomes repulsive, which keeps the nucleons at a certain average separation if they are of different types. This repulsion arises from the Pauli exclusion force for identical nucleons. A Pauli exclusion force occurs between quarks of the same type within nucleons, when the nucleons are different. At distances larger than 0.7 fm the force becomes attractive between spin-aligned nucleons, becoming maximal at a center–center distance of about 0.9 fm. Beyond this distance the force drops exponentially, until beyond about 2.0 fm separation, the force is negligible. Nucleons have a radius of about 0.8 fm. At short distances, the attractive nuclear force is stronger than the repulsive Coulomb force between protons. However, the Coulomb force between protons has a much greater range as it varies as the inverse square of the charge separation, Coulomb repulsion thus becomes the only significant force between protons when their separation exceeds about 2 to 2.5 fm.
The nuclear force has a spin-dependent component. The force is stronger for particles with their spins aligned than for those with their spins anti-aligned. If two particles are the same, such as two neutrons or two protons, the force is not enough to bind the particles, since the spin vectors of two particles of the same type must point in opposite directions when the particles are near each other and are in the same quantum state; this requirement for fermions stems from the Pauli exclusion principle. For fermion particles of different types, such as a proton and neutron, particles may be close to each other and have aligned spins without violating the Pauli exclusion principle, the nuclear force may bind them, since the nuclear force is much stronger for spin-aligned particles, but if the particles' spins are anti-aligned the nuclear force is too weak to bind them if they are of different types. The nuclear force has a tensor component which depends on the interaction between the nucleon spins and the angular momentum of the nucleons, leading to deformation from a simple spherical shape
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