# Atom

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Helium atom
An illustration of the helium atom, depicting the nucleus (pink) and the electron cloud distribution (black). The nucleus (upper right) in helium-4 is in reality spherically symmetric and closely resembles the electron cloud, although for more complicated nuclei this is not always the case. The black bar is one angstrom (10−10 m or 100  pm ).
Classification
Smallest recognized division of a chemical element
Properties
Mass range 1.67×10−27 to 4.52×10−25 kg
Electric charge zero (neutral), or ion charge
Diameter range62 pm (He) to 520 pm (Cs) (data page)
Components Electrons and a compact nucleus of protons and neutrons

An atom is the smallest constituent unit of ordinary matter that constitutes a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms are extremely small; typical sizes are around 100  picometers (1×10−10 m, a ten-millionth of a millimeter, or 1/254,000,000 of an inch). They are so small that accurately predicting their behavior using classical physics  – as if they were billiard balls, for example – is not possible. This is due to quantum effects. Current atomic models now use quantum principles to better explain and predict this behavior.

In classical physics and general chemistry, matter is any substance that has mass and takes up space by having volume. All everyday objects that can be touched are ultimately composed of atoms, which are made up of interacting subatomic particles, and in everyday as well as scientific usage, "matter" generally includes atoms and anything made up of them, and any particles that act as if they have both rest mass and volume. However it does not include massless particles such as photons, or other energy phenomena or waves such as light or sound. Matter exists in various states. These include classical everyday phases such as solid, liquid, and gas – for example water exists as ice, liquid water, and gaseous steam – but other states are possible, including plasma, Bose–Einstein condensates, fermionic condensates, and quark–gluon plasma.

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.

Solid is one of the four fundamental states of matter. In solids, particles are closely packed. It is characterized by structural rigidity and resistance to changes of shape or volume. Unlike a liquid, a solid object does not flow to take on the shape of its container, nor does it expand to fill the entire volume available to it like a gas does. The atoms in a solid are tightly bound to each other, either in a regular geometric lattice or irregularly, and are typically low in energy. Solids cannot be compressed with little pressure whereas gases can be compressed with little pressure because in gases molecules are loosely packed.

## Contents

Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and a number of neutrons. Only the most common variety of hydrogen has no neutrons. Protons and neutrons are called nucleons. More than 99.94% of an atom's mass is in the nucleus. The protons have a positive electric charge whereas the electrons have a negative electric charge. The neutrons have no electric charge. If the number of protons and electrons are equal, then the atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively. These atoms are called ions.

The atomic nucleus is the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford based on the 1909 Geiger–Marsden gold foil experiment. After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. An atom is composed of a positively-charged nucleus, with a cloud of negatively-charged electrons surrounding it, bound together by electrostatic force. Almost all of the mass of an atom is located in the nucleus, with a very small contribution from the electron cloud. Protons and neutrons are bound together to form a nucleus by the nuclear force.

The electron is a subatomic particle, symbol
e
or
β
, whose electric charge is negative one elementary charge. Electrons belong to the first generation of the lepton particle family, and are generally thought to be elementary particles because they have no known components or substructure. The electron has a mass that is approximately 1/1836 that of the proton. Quantum mechanical properties of the electron include an intrinsic angular momentum (spin) of a half-integer value, expressed in units of the reduced Planck constant, ħ. Being fermions, no two electrons can occupy the same quantum state, in accordance with the Pauli exclusion principle. Like all elementary particles, electrons exhibit properties of both particles and waves: they can collide with other particles and can be diffracted like light. The wave properties of electrons are easier to observe with experiments than those of other particles like neutrons and protons because electrons have a lower mass and hence a longer de Broglie wavelength for a given energy.

A proton is a subatomic particle, symbol
p
or
p+
, with a positive electric charge of +1e elementary charge and a mass slightly less than that of a neutron. Protons and neutrons, each with masses of approximately one atomic mass unit, are collectively referred to as "nucleons".

The electrons of an atom are attracted to the protons in an atomic nucleus by the electromagnetic force. The protons and neutrons in the nucleus are attracted to each other by the nuclear force. This force is usually stronger than the electromagnetic force that repels the positively charged protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus shatters and leaves behind different elements. This is a kind of nuclear decay. All electrons, nucleons, and nuclei alike are subatomic particles. The behavior of electrons in atoms is closer to a wave than a particle.

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 almost 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.

Nuclear transmutation is the conversion of one chemical element or an isotope into another chemical element. Because any element is defined by its number of protons in its atoms, i.e. in the atomic nucleus, nuclear transmutation occurs in any process where the number of protons or neutrons in the nucleus is changed.

In the physical sciences, subatomic particles are particles much smaller than atoms. The two types of subatomic particles are: elementary particles, which according to current theories are not made of other particles; and composite particles. Particle physics and nuclear physics study these particles and how they interact. The idea of a particle underwent serious rethinking when experiments showed that light could behave like a stream of particles as well as exhibiting wave-like properties. This led to the new concept of wave–particle duality to reflect that quantum-scale "particles" behave like both particles and waves. Another new concept, the uncertainty principle, states that some of their properties taken together, such as their simultaneous position and momentum, cannot be measured exactly. In more recent times, wave–particle duality has been shown to apply not only to photons but to increasingly massive particles as well.

The number of protons in the nucleus, called the atomic number , defines to which chemical element the atom belongs. For example, each copper atom contains 29 protons. The number of neutrons defines the isotope of the element. Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to associate and dissociate is responsible for most of the physical changes observed in nature. Chemistry is the discipline that studies these changes.

The atomic number or proton number of a chemical element is the number of protons found in the nucleus of every atom of that element. The atomic number uniquely identifies a chemical element. It is identical to the charge number of the nucleus. In an uncharged atom, the atomic number is also equal to the number of electrons.

Copper is a chemical element with the symbol Cu and atomic number 29. It is a soft, malleable, and ductile metal with very high thermal and electrical conductivity. A freshly exposed surface of pure copper has a pinkish-orange color. Copper is used as a conductor of heat and electricity, as a building material, and as a constituent of various metal alloys, such as sterling silver used in jewelry, cupronickel used to make marine hardware and coins, and constantan used in strain gauges and thermocouples for temperature measurement.

Isotopes are variants of a particular chemical element which differ in neutron number, and consequently in nucleon number. All isotopes of a given element have the same number of protons but different numbers of neutrons in each atom.

## History of atomic theory

### Atoms in philosophy

The idea that matter is made up of discrete units is a very old idea, appearing in many ancient cultures such as Greece and India. The word atomos, meaning "uncuttable", was coined by the ancient Greek philosophers Leucippus and his pupil Democritus (c. 460 – c. 370 BC). [1] [2] [3] [4] Democritus taught that atoms were infinite in number, uncreated, and eternal, and that the qualities of an object result from the kind of atoms that compose it. [2] [3] [4] Democritus's atomism was refined and elaborated by the later philosopher Epicurus (341–270 BC). [3] [4] During the Early Middle Ages, atomism was mostly forgotten in western Europe, but survived among some groups of Islamic philosophers. [3] During the twelfth century, atomism became known again in western Europe through references to it in the newly-rediscovered writings of Aristotle. [3]

Ancient Greek philosophy arose in the 6th century BC and continued throughout the Hellenistic period and the period in which Greece and most Greek-inhabited lands were part of the Roman Empire. Philosophy was used to make sense out of the world in a non-religious way. It dealt with a wide variety of subjects, including astronomy, mathematics, political philosophy, ethics, metaphysics, ontology, logic, biology, rhetoric and aesthetics.

Leucippus is reported in some ancient sources to have been a philosopher who was the earliest Greek to develop the theory of atomism—the idea that everything is composed entirely of various imperishable, indivisible elements called atoms. Leucippus often appears as the master to his pupil Democritus, a philosopher also touted as the originator of the atomic theory.

Democritus was an Ancient Greek pre-Socratic philosopher primarily remembered today for his formulation of an atomic theory of the universe.

In the fourteenth century, the rediscovery of major works describing atomist teachings, including Lucretius's De rerum natura and Diogenes Laërtius's Lives and Opinions of Eminent Philosophers , led to increased scholarly attention on the subject. [3] Nonetheless, because atomism was associated with the philosophy of Epicureanism, which contradicted orthodox Christian teachings, belief in atoms was not considered acceptable. [3] The French Catholic priest Pierre Gassendi (1592–1655) revived Epicurean atomism with modifications, arguing that atoms were created by God and, though extremely numerous, are not infinite. [3] [4] Gassendi's modified theory of atoms was popularized in France by the physician François Bernier (1620–1688) and in England by the natural philosopher Walter Charleton (1619–1707). [3] The chemist Robert Boyle (1627–1691) and the physicist Isaac Newton (1642–1727) both defended atomism and, by the end of the seventeenth century, it had become accepted by portions of the scientific community. [3]

Titus Lucretius Carus was a Roman poet and philosopher. His only known work is the philosophical poem De rerum natura, a didactic work about the tenets and philosophy of Epicureanism, and which is usually translated into English as On the Nature of Things. Lucretius has been credited with originating the concept of the three-age system which was formalised in 1836 by C. J. Thomsen.

De rerum natura is a first-century BC didactic poem by the Roman poet and philosopher Lucretius with the goal of explaining Epicurean philosophy to a Roman audience. The poem, written in some 7,400 dactylic hexameters, is divided into six untitled books, and explores Epicurean physics through poetic language and metaphors. Namely, Lucretius explores the principles of atomism; the nature of the mind and soul; explanations of sensation and thought; the development of the world and its phenomena; and explains a variety of celestial and terrestrial phenomena. The universe described in the poem operates according to these physical principles, guided by fortuna ("chance"), and not the divine intervention of the traditional Roman deities.

Diogenes Laërtius was a biographer of the Greek philosophers. Nothing is definitively known about his life, but his surviving Lives and Opinions of Eminent Philosophers is a principal source for the history of ancient Greek philosophy. His reputation is controversial among scholars because he often repeats information from his sources without critically evaluating it. He also frequently focuses on trivial or insignificant details of his subjects' lives while ignoring important details of their philosophical teachings and he sometimes fails to distinguish between earlier and later teachings of specific philosophical schools. However, unlike many other ancient secondary sources, Diogenes Laërtius generally reports philosophical teachings without attempting to reinterpret or expand on them, which means his accounts are often closer to the primary sources. Due to the loss of so many of the primary sources on which Diogenes relied, his work has become the foremost surviving source on the history of Greek philosophy.

### First evidence-based theory

In the early 1800s, John Dalton used the concept of atoms to explain why elements always react in ratios of small whole numbers (the law of multiple proportions). For instance, there are two types of tin oxide: one is 88.1% tin and 11.9% oxygen and the other is 78.7% tin and 21.3% oxygen (tin(II) oxide and tin dioxide respectively). This means that 100g of tin will combine either with 13.5g or 27g of oxygen. 13.5 and 27 form a ratio of 1:2, a ratio of small whole numbers. This common pattern in chemistry suggested to Dalton that elements react in multiples of discrete units — in other words, atoms. In the case of tin oxides, one tin atom will combine with either one or two oxygen atoms. [5]

Dalton also believed atomic theory could explain why water absorbs different gases in different proportions. For example, he found that water absorbs carbon dioxide far better than it absorbs nitrogen. [6] Dalton hypothesized this was due to the differences between the masses and configurations of the gases' respective particles, and carbon dioxide molecules (CO2) are heavier and larger than nitrogen molecules (N2).

### Brownian motion

In 1827, botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as "Brownian motion". This was thought to be caused by water molecules knocking the grains about. In 1905, Albert Einstein proved the reality of these molecules and their motions by producing the first statistical physics analysis of Brownian motion. [7] [8] [9] French physicist Jean Perrin used Einstein's work to experimentally determine the mass and dimensions of atoms, thereby conclusively verifying Dalton's atomic theory. [10]

### Discovery of the electron

The physicist J.J. Thomson measured the mass of cathode rays, showing they were made of particles, but were around 1800 times lighter than the lightest atom, hydrogen. Therefore, they were not atoms, but a new particle, the first subatomic particle to be discovered, which he originally called " corpuscle " but was later named electron, after particles postulated by George Johnstone Stoney in 1874. He also showed they were identical to particles given off by photoelectric and radioactive materials. [11] It was quickly recognized that they are the particles that carry electric currents in metal wires, and carry the negative electric charge within atoms. Thomson was given the 1906 Nobel Prize in Physics for this work. Thus he overturned the belief that atoms are the indivisible, ultimate particles of matter. [12] Thomson also incorrectly postulated that the low mass, negatively charged electrons were distributed throughout the atom in a uniform sea of positive charge. This became known as the plum pudding model.

### Discovery of the nucleus

In 1909, Hans Geiger and Ernest Marsden, under the direction of Ernest Rutherford, bombarded a metal foil with alpha particles to observe how they scattered. They expected all the alpha particles to pass straight through with little deflection, because Thomson's model said that the charges in the atom are so diffuse that their electric fields could not affect the alpha particles much. However, Geiger and Marsden spotted alpha particles being deflected by angles greater than 90°, which was supposed to be impossible according to Thomson's model. To explain this, Rutherford proposed that the positive charge of the atom is concentrated in a tiny nucleus at the center of the atom. [13]

### Discovery of isotopes

While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table. [14] The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J.J. Thomson created a technique for isotope separation through his work on ionized gases, which subsequently led to the discovery of stable isotopes. [15]

### Bohr model

In 1913 the physicist Niels Bohr proposed a model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. [16] This quantization was used to explain why the electrons orbits are stable (given that normally, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation, see synchrotron radiation ) and why elements absorb and emit electromagnetic radiation in discrete spectra. [17]

Later in the same year Henry Moseley provided additional experimental evidence in favor of Niels Bohr's theory. These results refined Ernest Rutherford's and Antonius Van den Broek's model, which proposed that the atom contains in its nucleus a number of positive nuclear charges that is equal to its (atomic) number in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today. [18]

### Chemical bonding explained

Chemical bonds between atoms were now explained, by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons. [19] As the chemical properties of the elements were known to largely repeat themselves according to the periodic law, [20] in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus. [21]

### Further developments in quantum physics

The Stern–Gerlach experiment of 1922 provided further evidence of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom's angular momentum, or spin. As this spin direction is initially random, the beam would be expected to deflect in a random direction. Instead, the beam was split into two directional components, corresponding to the atomic spin being oriented up or down with respect to the magnetic field. [22]

In 1925 Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (Matrix Mechanics). [18] One year earlier, in 1924, Louis de Broglie had proposed that all particles behave to an extent like waves and, in 1926, Erwin Schrödinger used this idea to develop a mathematical model of the atom (Wave Mechanics) that described the electrons as three-dimensional waveforms rather than point particles.

A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. [18] In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. [23] This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed. [24] [25]

### Discovery of the neutron

The development of the mass spectrometer allowed the mass of atoms to be measured with increased accuracy. The device uses a magnet to bend the trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom's mass to its charge. The chemist Francis William Aston used this instrument to show that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole number rule. [26] The explanation for these different isotopes awaited the discovery of the neutron, an uncharged particle with a mass similar to the proton, by the physicist James Chadwick in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus. [27]

### Fission, high-energy physics and condensed matter

In 1938, the German chemist Otto Hahn, a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. Instead, his chemical experiments showed barium as a product. [28] [29] A year later, Lise Meitner and her nephew Otto Frisch verified that Hahn's result were the first experimental nuclear fission. [30] [31] In 1944, Hahn received the Nobel prize in chemistry. Despite Hahn's efforts, the contributions of Meitner and Frisch were not recognized. [32]

In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. [33] Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks. The standard model of particle physics was developed that so far has successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions. [34]

## Structure

### Subatomic particles

Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton and the neutron; all three are fermions. However, the hydrogen-1 atom has no neutrons and the hydron ion has no electrons.

The electron is by far the least massive of these particles at 9.11×10−31 kg, with a negative electrical charge and a size that is too small to be measured using available techniques. [35] It was the lightest particle with a positive rest mass measured, until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details.

Protons have a positive charge and a mass 1,836 times that of the electron, at 1.6726×10−27 kg. The number of protons in an atom is called its atomic number. Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton.

Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the electron, or 1.6749×10−27 kg. [36] [37] Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of 2.5×10−15 m—although the 'surface' of these particles is not sharply defined. [38] The neutron was discovered in 1932 by the English physicist James Chadwick.

In the Standard Model of physics, electrons are truly elementary particles with no internal structure. However, both protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks (each with charge +2/3) and one down quark (with a charge of −1/3). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. [39] [40]

The quarks are held together by the strong interaction (or strong force), which is mediated by gluons. The protons and neutrons, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces. [39] [40]

### Nucleus

All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to 1.07 3A  fm, where A is the total number of nucleons. [41] This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other. [42]

Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay. [43]

The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud. [44]

A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay. However, with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus, which slightly modifies this trend of equal numbers of protons to neutrons. [44]

The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3–10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus. [45] Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element. [46] [47]

If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass–energy equivalence formula, ${\displaystyle E=mc^{2}}$, where ${\displaystyle m}$ is the mass loss and ${\displaystyle c}$ is the speed of light. This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate. [48]

The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together. [49] It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about 60, is an endothermic process. These more massive nuclei can not undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star. [44]

### Electron cloud

The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations.

Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured. [50] Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. [51] Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation. [52]

Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines. [51]

The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, [53] compared to 2.23 million eV for splitting a deuterium nucleus. [54] Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. [55]

## Properties

### Nuclear properties

By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form, [56] also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single proton element hydrogen up to the 118-proton element oganesson. [57] All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible. [58] [59]

About 339 nuclides occur naturally on Earth, [60] of which 254 (about 75%) have not been observed to decay, and are referred to as "stable isotopes". However, only 90 of these nuclides are stable to all decay, even in theory. Another 164 (bringing the total to 254) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 34 radioactive nuclides have half-lives longer than 80 million years, and are long-lived enough to be present from the birth of the solar system. This collection of 288 nuclides are known as primordial nuclides. Finally, an additional 51 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or else as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14). [61] [note 1]

For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.2 stable isotopes per element. Twenty-six elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes. [62] [ page needed ]

Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 254 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10 and nitrogen-14. Also, only four naturally occurring, radioactive odd–odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138 and tantalum-180m. Most odd–odd nuclei are highly unstable with respect to beta decay, because the decay products are even–even, and are therefore more strongly bound, due to nuclear pairing effects. [62] [ page needed ]

### Mass

The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles (called "nucleons") in a given atom is called the mass number. It is a positive integer and dimensionless (instead of having dimension of mass), because it expresses a count. An example of use of a mass number is "carbon-12," which has 12 nucleons (six protons and six neutrons).

The actual mass of an atom at rest is often expressed in daltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately 1.66×10−27 kg. [63] Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da. [64] The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 Da). However, this number will not be exactly an integer except in the case of carbon-12 (see below). [65] The heaviest stable atom is lead-208, [58] with a mass of 207.9766521 Da. [66]

As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms (about 6.022×1023). This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 Da, and so a mole of carbon-12 atoms weighs exactly 0.012 kg. [63]

### Shape and size

Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. [67] However, this assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. [68] On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right). [69] Consequently, the smallest atom is helium with a radius of 32  pm, while one of the largest is caesium at 225 pm. [70]

When subjected to external forces, like electrical fields, the shape of an atom may deviate from spherical symmetry. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by group-theoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystal-electrical fields may occur at low-symmetry lattice sites. [71] [72] Significant ellipsoidal deformations have been shown to occur for sulfur ions [73] and chalcogen ions [74] in pyrite-type compounds.

Atomic dimensions are thousands of times smaller than the wavelengths of light (400–700  nm) so they cannot be viewed using an optical microscope. However, individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width. [75] A single drop of water contains about 2  sextillion (2×1021) atoms of oxygen, and twice the number of hydrogen atoms. [76] A single carat diamond with a mass of 2×10−4 kg contains about 10 sextillion (1022) atoms of carbon. [note 2] If an apple were magnified to the size of the Earth, then the atoms in the apple would be approximately the size of the original apple. [77]

Every element has one or more isotopes that have unstable nuclei that are subject to radioactive decay, causing the nucleus to emit particles or electromagnetic radiation. Radioactivity can occur when the radius of a nucleus is large compared with the radius of the strong force, which only acts over distances on the order of 1 fm. [78]

The most common forms of radioactive decay are: [79] [80]

• Alpha decay: this process is caused when the nucleus emits an alpha particle, which is a helium nucleus consisting of two protons and two neutrons. The result of the emission is a new element with a lower atomic number.
• Beta decay (and electron capture): these processes are regulated by the weak force, and result from a transformation of a neutron into a proton, or a proton into a neutron. The neutron to proton transition is accompanied by the emission of an electron and an antineutrino, while proton to neutron transition (except in electron capture) causes the emission of a positron and a neutrino. The electron or positron emissions are called beta particles. Beta decay either increases or decreases the atomic number of the nucleus by one. Electron capture is more common than positron emission, because it requires less energy. In this type of decay, an electron is absorbed by the nucleus, rather than a positron emitted from the nucleus. A neutrino is still emitted in this process, and a proton changes to a neutron.
• Gamma decay: this process results from a change in the energy level of the nucleus to a lower state, resulting in the emission of electromagnetic radiation. The excited state of a nucleus which results in gamma emission usually occurs following the emission of an alpha or a beta particle. Thus, gamma decay usually follows alpha or beta decay.

Other more rare types of radioactive decay include ejection of neutrons or protons or clusters of nucleons from a nucleus, or more than one beta particle. An analog of gamma emission which allows excited nuclei to lose energy in a different way, is internal conversion—a process that produces high-speed electrons that are not beta rays, followed by production of high-energy photons that are not gamma rays. A few large nuclei explode into two or more charged fragments of varying masses plus several neutrons, in a decay called spontaneous nuclear fission.

Each radioactive isotope has a characteristic decay time period—the half-life—that is determined by the amount of time needed for half of a sample to decay. This is an exponential decay process that steadily decreases the proportion of the remaining isotope by 50% every half-life. Hence after two half-lives have passed only 25% of the isotope is present, and so forth. [78]

### Magnetic moment

Elementary particles possess an intrinsic quantum mechanical property known as spin. This is analogous to the angular momentum of an object that is spinning around its center of mass, although strictly speaking these particles are believed to be point-like and cannot be said to be rotating. Spin is measured in units of the reduced Planck constant (ħ), with electrons, protons and neutrons all having spin ½ ħ, or "spin-½". In an atom, electrons in motion around the nucleus possess orbital angular momentum in addition to their spin, while the nucleus itself possesses angular momentum due to its nuclear spin. [81]

The magnetic field produced by an atom—its magnetic moment—is determined by these various forms of angular momentum, just as a rotating charged object classically produces a magnetic field. However, the most dominant contribution comes from electron spin. Due to the nature of electrons to obey the Pauli exclusion principle, in which no two electrons may be found in the same quantum state, bound electrons pair up with each other, with one member of each pair in a spin up state and the other in the opposite, spin down state. Thus these spins cancel each other out, reducing the total magnetic dipole moment to zero in some atoms with even number of electrons. [82]

In ferromagnetic elements such as iron, cobalt and nickel, an odd number of electrons leads to an unpaired electron and a net overall magnetic moment. The orbitals of neighboring atoms overlap and a lower energy state is achieved when the spins of unpaired electrons are aligned with each other, a spontaneous process known as an exchange interaction. When the magnetic moments of ferromagnetic atoms are lined up, the material can produce a measurable macroscopic field. Paramagnetic materials have atoms with magnetic moments that line up in random directions when no magnetic field is present, but the magnetic moments of the individual atoms line up in the presence of a field. [82] [83]

The nucleus of an atom will have no spin when it has even numbers of both neutrons and protons, but for other cases of odd numbers, the nucleus may have a spin. Normally nuclei with spin are aligned in random directions because of thermal equilibrium. However, for certain elements (such as xenon-129) it is possible to polarize a significant proportion of the nuclear spin states so that they are aligned in the same direction—a condition called hyperpolarization. This has important applications in magnetic resonance imaging. [84] [85]

### Energy levels

The potential energy of an electron in an atom is negative, its dependence of its position reaches the minimum (the most absolute value) inside the nucleus, and vanishes when the distance from the nucleus goes to infinity, roughly in an inverse proportion to the distance. In the quantum-mechanical model, a bound electron can only occupy a set of states centered on the nucleus, and each state corresponds to a specific energy level; see time-independent Schrödinger equation for theoretical explanation. An energy level can be measured by the amount of energy needed to unbind the electron from the atom, and is usually given in units of electronvolts (eV). The lowest energy state of a bound electron is called the ground state, i.e. stationary state, while an electron transition to a higher level results in an excited state. [86] The electron's energy raises when n increases because the (average) distance to the nucleus increases. Dependence of the energy on is caused not by electrostatic potential of the nucleus, but by interaction between electrons.

For an electron to transition between two different states, e.g. ground state to first excited state, it must absorb or emit a photon at an energy matching the difference in the potential energy of those levels, according to the Niels Bohr model, what can be precisely calculated by the Schrödinger equation. Electrons jump between orbitals in a particle-like fashion. For example, if a single photon strikes the electrons, only a single electron changes states in response to the photon; see Electron properties.

The energy of an emitted photon is proportional to its frequency, so these specific energy levels appear as distinct bands in the electromagnetic spectrum. [87] Each element has a characteristic spectrum that can depend on the nuclear charge, subshells filled by electrons, the electromagnetic interactions between the electrons and other factors. [88]

When a continuous spectrum of energy is passed through a gas or plasma, some of the photons are absorbed by atoms, causing electrons to change their energy level. Those excited electrons that remain bound to their atom spontaneously emit this energy as a photon, traveling in a random direction, and so drop back to lower energy levels. Thus the atoms behave like a filter that forms a series of dark absorption bands in the energy output. (An observer viewing the atoms from a view that does not include the continuous spectrum in the background, instead sees a series of emission lines from the photons emitted by the atoms.) Spectroscopic measurements of the strength and width of atomic spectral lines allow the composition and physical properties of a substance to be determined. [89]

Close examination of the spectral lines reveals that some display a fine structure splitting. This occurs because of spin–orbit coupling, which is an interaction between the spin and motion of the outermost electron. [90] When an atom is in an external magnetic field, spectral lines become split into three or more components; a phenomenon called the Zeeman effect. This is caused by the interaction of the magnetic field with the magnetic moment of the atom and its electrons. Some atoms can have multiple electron configurations with the same energy level, which thus appear as a single spectral line. The interaction of the magnetic field with the atom shifts these electron configurations to slightly different energy levels, resulting in multiple spectral lines. [91] The presence of an external electric field can cause a comparable splitting and shifting of spectral lines by modifying the electron energy levels, a phenomenon called the Stark effect. [92]

If a bound electron is in an excited state, an interacting photon with the proper energy can cause stimulated emission of a photon with a matching energy level. For this to occur, the electron must drop to a lower energy state that has an energy difference matching the energy of the interacting photon. The emitted photon and the interacting photon then move off in parallel and with matching phases. That is, the wave patterns of the two photons are synchronized. This physical property is used to make lasers, which can emit a coherent beam of light energy in a narrow frequency band. [93]

### Valence and bonding behavior

Valency is the combining power of an element. It is equal to number of hydrogen atoms that atom can combine or displace in forming compounds. [94] The outermost electron shell of an atom in its uncombined state is known as the valence shell, and the electrons in that shell are called valence electrons. The number of valence electrons determines the bonding behavior with other atoms. Atoms tend to chemically react with each other in a manner that fills (or empties) their outer valence shells. [95] For example, a transfer of a single electron between atoms is a useful approximation for bonds that form between atoms with one-electron more than a filled shell, and others that are one-electron short of a full shell, such as occurs in the compound sodium chloride and other chemical ionic salts. However, many elements display multiple valences, or tendencies to share differing numbers of electrons in different compounds. Thus, chemical bonding between these elements takes many forms of electron-sharing that are more than simple electron transfers. Examples include the element carbon and the organic compounds. [96]

The chemical elements are often displayed in a periodic table that is laid out to display recurring chemical properties, and elements with the same number of valence electrons form a group that is aligned in the same column of the table. (The horizontal rows correspond to the filling of a quantum shell of electrons.) The elements at the far right of the table have their outer shell completely filled with electrons, which results in chemically inert elements known as the noble gases. [97] [98]

### States

Quantities of atoms are found in different states of matter that depend on the physical conditions, such as temperature and pressure. By varying the conditions, materials can transition between solids, liquids, gases and plasmas. [99] Within a state, a material can also exist in different allotropes. An example of this is solid carbon, which can exist as graphite or diamond. [100] Gaseous allotropes exist as well, such as dioxygen and ozone.

At temperatures close to absolute zero, atoms can form a Bose–Einstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale. [101] [102] This super-cooled collection of atoms then behaves as a single super atom, which may allow fundamental checks of quantum mechanical behavior. [103]

## Identification

The scanning tunneling microscope is a device for viewing surfaces at the atomic level. It uses the quantum tunneling phenomenon, which allows particles to pass through a barrier that would normally be insurmountable. Electrons tunnel through the vacuum between two planar metal electrodes, on each of which is an adsorbed atom, providing a tunneling-current density that can be measured. Scanning one atom (taken as the tip) as it moves past the other (the sample) permits plotting of tip displacement versus lateral separation for a constant current. The calculation shows the extent to which scanning-tunneling-microscope images of an individual atom are visible. It confirms that for low bias, the microscope images the space-averaged dimensions of the electron orbitals across closely packed energy levels—the Fermi level local density of states. [104] [105]

An atom can be ionized by removing one of its electrons. The electric charge causes the trajectory of an atom to bend when it passes through a magnetic field. The radius by which the trajectory of a moving ion is turned by the magnetic field is determined by the mass of the atom. The mass spectrometer uses this principle to measure the mass-to-charge ratio of ions. If a sample contains multiple isotopes, the mass spectrometer can determine the proportion of each isotope in the sample by measuring the intensity of the different beams of ions. Techniques to vaporize atoms include inductively coupled plasma atomic emission spectroscopy and inductively coupled plasma mass spectrometry, both of which use a plasma to vaporize samples for analysis. [106]

A more area-selective method is electron energy loss spectroscopy, which measures the energy loss of an electron beam within a transmission electron microscope when it interacts with a portion of a sample. The atom-probe tomograph has sub-nanometer resolution in 3-D and can chemically identify individual atoms using time-of-flight mass spectrometry. [107]

Spectra of excited states can be used to analyze the atomic composition of distant stars. Specific light wavelengths contained in the observed light from stars can be separated out and related to the quantized transitions in free gas atoms. These colors can be replicated using a gas-discharge lamp containing the same element. [108] Helium was discovered in this way in the spectrum of the Sun 23 years before it was found on Earth. [109]

## Origin and current state

Baryonic matter forms about 4% of the total energy density of the observable Universe, with an average density of about 0.25 particles/m3 (mostly protons and electrons). [110] Within a galaxy such as the Milky Way, particles have a much higher concentration, with the density of matter in the interstellar medium (ISM) ranging from 105 to 109 atoms/m3. [111] The Sun is believed to be inside the Local Bubble, so the density in the solar neighborhood is only about 103 atoms/m3. [112] Stars form from dense clouds in the ISM, and the evolutionary processes of stars result in the steady enrichment of the ISM with elements more massive than hydrogen and helium.

Up to 95% of the Milky Way's baryonic matter are concentrated inside stars, where conditions are unfavorable for atomic matter. The total baryonic mass is about 10% of the mass of the galaxy; [113] the remainder of the mass is an unknown dark matter. [114] High temperature inside stars makes most "atoms" fully ionized, that is, separates all electrons from the nuclei. In stellar remnants—with exception of their surface layers—an immense pressure make electron shells impossible.

### Formation

Electrons are thought to exist in the Universe since early stages of the Big Bang. Atomic nuclei forms in nucleosynthesis reactions. In about three minutes Big Bang nucleosynthesis produced most of the helium, lithium, and deuterium in the Universe, and perhaps some of the beryllium and boron. [115] [116] [117]

Ubiquitousness and stability of atoms relies on their binding energy, which means that an atom has a lower energy than an unbound system of the nucleus and electrons. Where the temperature is much higher than ionization potential, the matter exists in the form of plasma—a gas of positively charged ions (possibly, bare nuclei) and electrons. When the temperature drops below the ionization potential, atoms become statistically favorable. Atoms (complete with bound electrons) became to dominate over charged particles 380,000 years after the Big Bang—an epoch called recombination, when the expanding Universe cooled enough to allow electrons to become attached to nuclei. [118]

Since the Big Bang, which produced no carbon or heavier elements, atomic nuclei have been combined in stars through the process of nuclear fusion to produce more of the element helium, and (via the triple alpha process) the sequence of elements from carbon up to iron; [119] see stellar nucleosynthesis for details.

Isotopes such as lithium-6, as well as some beryllium and boron are generated in space through cosmic ray spallation. [120] This occurs when a high-energy proton strikes an atomic nucleus, causing large numbers of nucleons to be ejected.

Elements heavier than iron were produced in supernovae through the r-process and in AGB stars through the s-process, both of which involve the capture of neutrons by atomic nuclei. [121] Elements such as lead formed largely through the radioactive decay of heavier elements. [122]

### Earth

Most of the atoms that make up the Earth and its inhabitants were present in their current form in the nebula that collapsed out of a molecular cloud to form the Solar System. The rest are the result of radioactive decay, and their relative proportion can be used to determine the age of the Earth through radiometric dating. [123] [124] Most of the helium in the crust of the Earth (about 99% of the helium from gas wells, as shown by its lower abundance of helium-3) is a product of alpha decay. [125]

There are a few trace atoms on Earth that were not present at the beginning (i.e., not "primordial"), nor are results of radioactive decay. Carbon-14 is continuously generated by cosmic rays in the atmosphere. [126] Some atoms on Earth have been artificially generated either deliberately or as by-products of nuclear reactors or explosions. [127] [128] Of the transuranic elements—those with atomic numbers greater than 92—only plutonium and neptunium occur naturally on Earth. [129] [130] Transuranic elements have radioactive lifetimes shorter than the current age of the Earth [131] and thus identifiable quantities of these elements have long since decayed, with the exception of traces of plutonium-244 possibly deposited by cosmic dust. [123] Natural deposits of plutonium and neptunium are produced by neutron capture in uranium ore. [132]

The Earth contains approximately 1.33×1050 atoms. [133] Although small numbers of independent atoms of noble gases exist, such as argon, neon, and helium, 99% of the atmosphere is bound in the form of molecules, including carbon dioxide and diatomic oxygen and nitrogen. At the surface of the Earth, an overwhelming majority of atoms combine to form various compounds, including water, salt, silicates and oxides. Atoms can also combine to create materials that do not consist of discrete molecules, including crystals and liquid or solid metals. [134] [135] This atomic matter forms networked arrangements that lack the particular type of small-scale interrupted order associated with molecular matter. [136]

### Rare and theoretical forms

#### Superheavy elements

While isotopes with atomic numbers higher than lead (82) are known to be radioactive, an "island of stability" has been proposed for some elements with atomic numbers above 103. These superheavy elements may have a nucleus that is relatively stable against radioactive decay. [137] The most likely candidate for a stable superheavy atom, unbihexium, has 126 protons and 184 neutrons. [138]

#### Exotic matter

Each particle of matter has a corresponding antimatter particle with the opposite electrical charge. Thus, the positron is a positively charged antielectron and the antiproton is a negatively charged equivalent of a proton. When a matter and corresponding antimatter particle meet, they annihilate each other. Because of this, along with an imbalance between the number of matter and antimatter particles, the latter are rare in the universe. The first causes of this imbalance are not yet fully understood, although theories of baryogenesis may offer an explanation. As a result, no antimatter atoms have been discovered in nature. [139] [140] However, in 1996 the antimatter counterpart of the hydrogen atom (antihydrogen) was synthesized at the CERN laboratory in Geneva. [141] [142]

Other exotic atoms have been created by replacing one of the protons, neutrons or electrons with other particles that have the same charge. For example, an electron can be replaced by a more massive muon, forming a muonic atom. These types of atoms can be used to test the fundamental predictions of physics. [143] [144] [145]

## Notes

1. For more recent updates see Interactive Chart of Nuclides (Brookhaven National Laboratory) Archived 25 August 2009 at the Wayback Machine .
2. A carat is 200 milligrams. By definition, carbon-12 has 0.012 kg per mole. The Avogadro constant defines 6×1023 atoms per mole.

## Related Research Articles

Alpha decay or α-decay is a type of radioactive decay in which an atomic nucleus emits an alpha particle and thereby transforms or 'decays' into a different atomic nucleus, with a mass number that is reduced by four and an atomic number that is reduced by two. An alpha particle is identical to the nucleus of a helium-4 atom, which consists of two protons and two neutrons. It has a charge of +2 e and a mass of 4 u. For example, uranium-238 decays to form thorium-234. Alpha particles have a charge +2 e, but as a nuclear equation describes a nuclear reaction without considering the electrons – a convention that does not imply that the nuclei necessarily occur in neutral atoms – the charge is not usually shown.

In nuclear physics, beta decay (β-decay) is a type of radioactive decay in which a beta particle is emitted from an atomic nucleus, transforming the original nuclide to its isobar. 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 in so-called positron emission. Neither the beta particle nor its associated (anti-)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.

The neutron is a subatomic particle, symbol
n
or
n0
, with no net electric charge and a mass slightly greater than that of a proton. Protons and neutrons constitute the nuclei of atoms. Since protons and neutrons behave similarly within the nucleus, and each has a mass of approximately one atomic mass unit, they are both referred to as nucleons. Their properties and interactions are described by nuclear physics.

Nuclear physics is the field of physics that studies atomic nuclei and their constituents and interactions. Other forms of nuclear matter are also studied. Nuclear physics should not be confused with atomic physics, which studies the atom as a whole, including its electrons.

Neutronium is a hypothetical substance composed purely of neutrons. The word was coined by scientist Andreas von Antropoff in 1926 for the conjectured "element of atomic number zero" that he placed at the head of the periodic table. However, the meaning of the term has changed over time, and from the last half of the 20th century onward it has been also used to refer to extremely dense substances resembling the neutron-degenerate matter theorized to exist in the cores of neutron stars; hereinafter "degenerate neutronium" will refer to this. Science fiction and popular literature frequently use the term "neutronium" to refer to a highly dense phase of matter composed primarily of neutrons.

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 2 or more smaller, lighter nuclei. The fission process often produces gamma photons, and releases a very large amount of energy even by the energetic standards of radioactive decay.

Electron capture is a process in which the proton-rich nucleus of an electrically neutral atom absorbs an inner atomic electron, usually from the K or L electron shell. This process thereby changes a nuclear proton to a neutron and simultaneously causes the emission of an electron neutrino.

A nuclide is an atomic species characterized by the specific constitution of its nucleus, i.e., by its number of protons Z, its number of neutrons N, and its nuclear energy state.

In nuclear physics and nuclear chemistry, a nuclear reaction is semantically considered to be the process in which two nuclei, or else a nucleus of an atom and a subatomic particle from outside the atom, collide to produce one or more nuclides that are different from the nuclide(s) that began the process. Thus, a nuclear reaction must cause a transformation of at least one nuclide to another. If a nucleus interacts with another nucleus or particle and they then separate without changing the nature of any nuclide, the process is simply referred to as a type of nuclear scattering, rather than a nuclear reaction.

The mass number, also called atomic mass number or nucleon number, is the total number of protons and neutrons in an atomic nucleus. It is approximately equal to the atomic mass of the atom expressed in atomic mass units. Because protons and neutrons both are baryons, the mass number A is identical with the baryon number B as of the nucleus as of the whole atom or ion. The mass number is different for each different isotope of a chemical element. Hence, the difference between the mass number and the atomic number Z gives the number of neutrons (N) in a given nucleus: N = A - Z.

Helium-4 is a stable isotope of the element helium. It is by far the more abundant of the two naturally occurring isotopes of helium, making up about 99.99986% of the helium on Earth. Its nucleus is identical to an alpha particle, and consists of two protons and two neutrons.

Double electron capture is a decay mode of an atomic nucleus. For a nuclide with number of nucleons A and atomic number Z, double electron capture is only possible if the mass of the nuclide of is lower.

Although there are nine known isotopes of helium (2He), only helium-3 and helium-4 are stable. All radioisotopes are short-lived, the longest-lived being 6
He
with a half-life of 806.7 milliseconds. The least stable is 5
He
, with a half-life of 7.6×10−22 s, although it is possible that 2
He
has an even shorter half-life.

Nuclear binding energy is the minimum energy that would be required to disassemble the nucleus of an atom into its component parts. These component parts are neutrons and protons, which are collectively called nucleons. The binding energy is always a positive number, as we need to spend energy in moving these nucleons, attracted to each other by the strong nuclear force, away from each other. The mass of an atomic nucleus is less than the sum of the individual masses of the free constituent protons and neutrons, according to Einstein's equation E=mc2. This 'missing mass' is known as the mass defect, and represents the energy that was released when the nucleus was formed.

In nuclear physics, the valley of stability is a characterization of the stability of nuclides to radioactivity based on their binding energy. Nuclides are composed of protons and neutrons. The shape of the valley refers to the profile of binding energy as a function of the numbers of neutrons and protons, with the lowest part of the valley corresponding to the region of most stable nuclei. The line of stable nuclides down the center of the valley of stability is known as the line of beta stability. The sides of the valley correspond to increasing instability to beta decay. The decay of a nuclide becomes more energetically favorable the further it is from the line of beta stability. The boundaries of the valley correspond to the nuclear drip lines, where nuclides become so unstable they emit single protons or single neutrons. Regions of instability within the valley at high atomic number also include radioactive decay by alpha radiation or spontaneous fission. The shape of the valley is roughly an elongated paraboloid corresponding to the nuclide binding energies as a function of neutron and atomic numbers.

The discovery of the neutron and its properties was central to the extraordinary developments in atomic physics in the first half of the 20th century. Early in the century, Ernest Rutherford developed a crude model of the atom, based on the gold foil experiment of Hans Geiger and Ernest Marsden. In this model, atoms had their mass and positive electric charge concentrated in a very small nucleus. By 1920 chemical isotopes had been discovered, the atomic masses had been determined to be (approximately) integer multiples of the mass of the hydrogen atom, and the atomic number had been identified as the charge on the nucleus. Throughout the 1920s, the nucleus was viewed as composed of combinations of protons and electrons, the two elementary particles known at the time, but that model presented several experimental and theoretical contradictions.

## References

1. Pullman, Bernard (1998). The Atom in the History of Human Thought. Oxford, England: Oxford University Press. pp. 31–33. ISBN   978-0-19-515040-7.
2. Kenny, Anthony (2004). Ancient Philosophy. A New History of Western Philosophy. 1. Oxford, England: Oxford University Press. pp. 26–28. ISBN   978-0-19-875273-8.
3. Pyle, Andrew (2010). "Atoms and Atomism". In Grafton, Anthony; Most, Glenn W.; Settis, Salvatore (eds.). The Classical Tradition. Cambridge, Massachusetts and London: The Belknap Press of Harvard University Press. pp. 103–104. ISBN   978-0-674-03572-0.
4. Cohen, Henri; Lefebvre, Claire, eds. (2017). Handbook of Categorization in Cognitive Science (Second ed.). Amsterdam, The Netherlands: Elsevier. p. 427. ISBN   978-0-08-101107-2.
5. Andrew G. van Melsen (1952). From Atomos to Atom. Mineola, NY: Dover Publications. ISBN   978-0-486-49584-2.
6. Dalton, John. "On the Absorption of Gases by Water and Other Liquids Archived 11 June 2011 at the Wayback Machine ", in Memoirs of the Literary and Philosophical Society of Manchester. 1803. Retrieved on August 29, 2007.
7. Einstein, Albert (1905). "Über die von der molekularkinetischen Theorie der Wärme geforderte Bewegung von in ruhenden Flüssigkeiten suspendierten Teilchen" (PDF). Annalen der Physik (in German). 322 (8): 549–560. Bibcode:1905AnP...322..549E. doi:10.1002/andp.19053220806. Archived (PDF) from the original on 18 July 2007. Retrieved 4 February 2007.
8. Mazo, Robert M. (2002). Brownian Motion: Fluctuations, Dynamics, and Applications. Oxford University Press. pp. 1–7. ISBN   978-0-19-851567-8. OCLC   48753074.
9. Lee, Y.K.; Hoon, K. (1995). "Brownian Motion". Imperial College. Archived from the original on 18 December 2007. Retrieved 18 December 2007.
10. Patterson, G. (2007). "Jean Perrin and the triumph of the atomic doctrine". Endeavour . 31 (2): 50–53. doi:10.1016/j.endeavour.2007.05.003. PMID   17602746.
11. Thomson, J.J. (August 1901). "On bodies smaller than atoms". The Popular Science Monthly: 323–335. Retrieved 21 June 2009.
12. "J.J. Thomson". Nobel Foundation. 1906. Archived from the original on 12 May 2013. Retrieved 20 December 2007.
13. Rutherford, E. (1911). "The Scattering of α and β Particles by Matter and the Structure of the Atom" (PDF). Philosophical Magazine . 21 (125): 669–688. doi:10.1080/14786440508637080. Archived (PDF) from the original on 31 May 2016. Retrieved 29 April 2016.
14. "Frederick Soddy, The Nobel Prize in Chemistry 1921". Nobel Foundation. Archived from the original on 9 April 2008. Retrieved 18 January 2008.
15. Thomson, Joseph John (1913). "Rays of positive electricity". Proceedings of the Royal Society . A. 89 (607): 1–20. Bibcode:1913RSPSA..89....1T. doi:. Archived from the original on 4 November 2016. Retrieved 12 February 2008.
16. Stern, David P. (16 May 2005). "The Atomic Nucleus and Bohr's Early Model of the Atom". NASA/Goddard Space Flight Center. Archived from the original on 20 August 2007. Retrieved 20 December 2007.
17. Bohr, Niels (11 December 1922). "Niels Bohr, The Nobel Prize in Physics 1922, Nobel Lecture". Nobel Foundation. Archived from the original on 15 April 2008. Retrieved 16 February 2008.
18. Pais, Abraham (1986). Inward Bound: Of Matter and Forces in the Physical World. New York: Oxford University Press. pp. 228–230. ISBN   978-0-19-851971-3.
19. Lewis, Gilbert N. (1916). "The Atom and the Molecule" (PDF). Journal of the American Chemical Society . 38 (4): 762–786. doi:10.1021/ja02261a002. Archived (PDF) from the original on 25 August 2019. Retrieved 25 August 2019.
20. Scerri, Eric R. (2007). The periodic table: its story and its significance. Oxford University Press US. pp. 205–226. ISBN   978-0-19-530573-9.
21. Langmuir, Irving (1919). "The Arrangement of Electrons in Atoms and Molecules". Journal of the American Chemical Society . 41 (6): 868–934. doi:10.1021/ja02227a002. Archived from the original on 21 June 2019. Retrieved 27 June 2019.
22. Scully, Marlan O.; Lamb, Willis E.; Barut, Asim (1987). "On the theory of the Stern-Gerlach apparatus". Foundations of Physics . 17 (6): 575–583. Bibcode:1987FoPh...17..575S. doi:10.1007/BF01882788.
23. Chad Orzel (16 September 2014). "What is the Heisenberg Uncertainty Principle?". TED-Ed. Archived from the original on 13 September 2015. Retrieved 26 October 2015 via YouTube.
24. Brown, Kevin (2007). "The Hydrogen Atom". MathPages. Archived from the original on 13 May 2008. Retrieved 21 December 2007.
25. Harrison, David M. (2000). "The Development of Quantum Mechanics". University of Toronto. Archived from the original on 25 December 2007. Retrieved 21 December 2007.
26. Aston, Francis W. (1920). "The constitution of atmospheric neon". Philosophical Magazine . 39 (6): 449–455. doi:10.1080/14786440408636058.
27. Chadwick, James (12 December 1935). "Nobel Lecture: The Neutron and Its Properties". Nobel Foundation. Archived from the original on 12 October 2007. Retrieved 21 December 2007.
28. Bowden, Mary Ellen (1997). "Otto Hahn, Lise Meitner, and Fritz Strassmann". Chemical achievers : the human face of the chemical sciences. Philadelphia, PA: Chemical Heritage Foundation. pp. 76–80, 125. ISBN   978-0-941901-12-3.
29. "Otto Hahn, Lise Meitner, and Fritz Strassmann". Science History Institute. June 2016. Archived from the original on 21 March 2018. Retrieved 20 March 2018.
30. Meitner, Lise; Frisch, Otto Robert (1939). "Disintegration of uranium by neutrons: a new type of nuclear reaction". Nature . 143 (3615): 239–240. Bibcode:1939Natur.143..239M. doi:10.1038/143239a0.
31. Schroeder, M. "Lise Meitner – Zur 125. Wiederkehr Ihres Geburtstages" (in German). Archived from the original on 19 July 2011. Retrieved 4 June 2009.
32. Crawford, E.; Sime, Ruth Lewin; Walker, Mark (1997). "A Nobel tale of postwar injustice". Physics Today . 50 (9): 26–32. Bibcode:1997PhT....50i..26C. doi:10.1063/1.881933.
33. Kullander, Sven (28 August 2001). "Accelerators and Nobel Laureates". Nobel Foundation. Archived from the original on 13 April 2008. Retrieved 31 January 2008.
34. "The Nobel Prize in Physics 1990". Nobel Foundation. 17 October 1990. Archived from the original on 14 May 2008. Retrieved 31 January 2008.
35. Demtröder, Wolfgang (2002). Atoms, Molecules and Photons: An Introduction to Atomic- Molecular- and Quantum Physics (1st ed.). Springer. pp. 39–42. ISBN   978-3-540-20631-6. OCLC   181435713.
36. Woan, Graham (2000). The Cambridge Handbook of Physics. Cambridge University Press. p. 8. ISBN   978-0-521-57507-2. OCLC   224032426.
37. Mohr, P.J.; Taylor, B.N. and Newell, D.B. (2014), "The 2014 CODATA Recommended Values of the Fundamental Physical Constants" Archived 21 February 2012 at WebCite (Web Version 7.0). The database was developed by J. Baker, M. Douma, and S. Kotochigova. (2014). National Institute of Standards and Technology, Gaithersburg, Maryland 20899.
38. MacGregor, Malcolm H. (1992). The Enigmatic Electron. Oxford University Press. pp. 33–37. ISBN   978-0-19-521833-6. OCLC   223372888.
39. Particle Data Group (2002). "The Particle Adventure". Lawrence Berkeley Laboratory. Archived from the original on 4 January 2007. Retrieved 3 January 2007.
40. Schombert, James (18 April 2006). "Elementary Particles". University of Oregon. Archived from the original on 21 August 2011. Retrieved 3 January 2007.
41. Jevremovic, Tatjana (2005). Nuclear Principles in Engineering. Springer. p. 63. ISBN   978-0-387-23284-3. OCLC   228384008.
42. Pfeffer, Jeremy I.; Nir, Shlomo (2000). Modern Physics: An Introductory Text. Imperial College Press. pp. 330–336. ISBN   978-1-86094-250-1. OCLC   45900880.
43. Wenner, Jennifer M. (10 October 2007). "How Does Radioactive Decay Work?". Carleton College. Archived from the original on 11 May 2008. Retrieved 9 January 2008.
44. Raymond, David (7 April 2006). "Nuclear Binding Energies". New Mexico Tech. Archived from the original on 1 December 2002. Retrieved 3 January 2007.
45. Mihos, Chris (23 July 2002). "Overcoming the Coulomb Barrier". Case Western Reserve University. Archived from the original on 12 September 2006. Retrieved 13 February 2008.
46. Staff (30 March 2007). "ABC's of Nuclear Science". Lawrence Berkeley National Laboratory. Archived from the original on 5 December 2006. Retrieved 3 January 2007.
47. Makhijani, Arjun; Saleska, Scott (2 March 2001). "Basics of Nuclear Physics and Fission". Institute for Energy and Environmental Research. Archived from the original on 16 January 2007. Retrieved 3 January 2007.
48. Shultis, J. Kenneth; Faw, Richard E. (2002). Fundamentals of Nuclear Science and Engineering. CRC Press. pp. 10–17. ISBN   978-0-8247-0834-4. OCLC   123346507.
49. Fewell, M.P. (1995). "The atomic nuclide with the highest mean binding energy". American Journal of Physics . 63 (7): 653–658. Bibcode:1995AmJPh..63..653F. doi:10.1119/1.17828.
50. Mulliken, Robert S. (1967). "Spectroscopy, Molecular Orbitals, and Chemical Bonding". Science . 157 (3784): 13–24. Bibcode:1967Sci...157...13M. doi:10.1126/science.157.3784.13. PMID   5338306.
51. Brucat, Philip J. (2008). "The Quantum Atom". University of Florida. Archived from the original on 7 December 2006. Retrieved 4 January 2007.
52. Manthey, David (2001). "Atomic Orbitals". Orbital Central. Archived from the original on 10 January 2008. Retrieved 21 January 2008.
53. Herter, Terry (2006). "Lecture 8: The Hydrogen Atom". Cornell University. Archived from the original on 22 February 2012. Retrieved 14 February 2008.
54. Bell, R.E.; Elliott, L.G. (1950). "Gamma-Rays from the Reaction H1(n,γ)D2 and the Binding Energy of the Deuteron". Physical Review . 79 (2): 282–285. Bibcode:1950PhRv...79..282B. doi:10.1103/PhysRev.79.282.
55. Smirnov, Boris M. (2003). Physics of Atoms and Ions. Springer. pp. 249–272. ISBN   978-0-387-95550-6.
56. Matis, Howard S. (9 August 2000). "The Isotopes of Hydrogen". Guide to the Nuclear Wall Chart. Lawrence Berkeley National Lab. Archived from the original on 18 December 2007. Retrieved 21 December 2007.
57. Weiss, Rick (17 October 2006). "Scientists Announce Creation of Atomic Element, the Heaviest Yet". Washington Post. Archived from the original on 21 August 2011. Retrieved 21 December 2007.
58. Sills, Alan D. (2003). Earth Science the Easy Way. Barron's Educational Series. pp. 131–134. ISBN   978-0-7641-2146-3. OCLC   51543743.
59. Dumé, Belle (23 April 2003). "Bismuth breaks half-life record for alpha decay". Physics World. Archived from the original on 14 December 2007. Retrieved 21 December 2007.
60. Lindsay, Don (30 July 2000). "Radioactives Missing From The Earth". Don Lindsay Archive. Archived from the original on 28 April 2007. Retrieved 23 May 2007.
61. Tuli, Jagdish K. (April 2005). "Nuclear Wallet Cards". National Nuclear Data Center, Brookhaven National Laboratory. Archived from the original on 3 October 2011. Retrieved 16 April 2011.
62. CRC Handbook (2002).
63. Mills, Ian; Cvitaš, Tomislav; Homann, Klaus; Kallay, Nikola; Kuchitsu, Kozo (1993). Quantities, Units and Symbols in Physical Chemistry (PDF) (2nd ed.). Oxford: International Union of Pure and Applied Chemistry, Commission on Physiochemical Symbols Terminology and Units, Blackwell Scientific Publications. p. 70. ISBN   978-0-632-03583-0. OCLC   27011505. Archived (PDF) from the original on 10 November 2011. Retrieved 10 December 2011.
64. Chieh, Chung (22 January 2001). "Nuclide Stability". University of Waterloo. Archived from the original on 30 August 2007. Retrieved 4 January 2007.
65. "Atomic Weights and Isotopic Compositions for All Elements". National Institute of Standards and Technology. Archived from the original on 31 December 2006. Retrieved 4 January 2007.
66. Audi, G.; Wapstra, A.H.; Thibault, C. (2003). "The Ame2003 atomic mass evaluation (II)" (PDF). Nuclear Physics A . 729 (1): 337–676. Bibcode:2003NuPhA.729..337A. doi:10.1016/j.nuclphysa.2003.11.003. Archived (PDF) from the original on 16 October 2005. Retrieved 1 May 2015.
67. Ghosh, D.C.; Biswas, R. (2002). "Theoretical calculation of Absolute Radii of Atoms and Ions. Part 1. The Atomic Radii". Int. J. Mol. Sci. 3 (11): 87–113. doi:.
68. Shannon, R.D. (1976). "Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides" (PDF). Acta Crystallographica A . 32 (5): 751–767. Bibcode:1976AcCrA..32..751S. doi:10.1107/S0567739476001551.
69. Dong, Judy (1998). "Diameter of an Atom". The Physics Factbook. Archived from the original on 4 November 2007. Retrieved 19 November 2007.
70. Zumdahl, Steven S. (2002). Introductory Chemistry: A Foundation (5th ed.). Houghton Mifflin. ISBN   978-0-618-34342-3. OCLC   173081482. Archived from the original on 4 March 2008. Retrieved 5 February 2008.
71. Bethe, Hans (1929). "Termaufspaltung in Kristallen". Annalen der Physik. 3 (2): 133–208. Bibcode:1929AnP...395..133B. doi:10.1002/andp.19293950202.
72. Birkholz, Mario (1995). "Crystal-field induced dipoles in heteropolar crystals – I. concept". Z. Phys. B. 96 (3): 325–332. Bibcode:1995ZPhyB..96..325B. CiteSeerX  . doi:10.1007/BF01313054.
73. Birkholz, M.; Rudert, R. (2008). "Interatomic distances in pyrite-structure disulfides – a case for ellipsoidal modeling of sulfur ions]". Physica Status Solidi B. 245 (9): 1858–1864. Bibcode:2008PSSBR.245.1858B. doi:10.1002/pssb.200879532.
74. Birkholz, M. (2014). "Modeling the Shape of Ions in Pyrite-Type Crystals". Crystals. 4 (3): 390–403. doi:.
75. Staff (2007). "Small Miracles: Harnessing nanotechnology". Oregon State University. Archived from the original on 21 May 2011. Retrieved 7 January 2007. – describes the width of a human hair as 105 nm and 10 carbon atoms as spanning 1 nm.
76. Padilla, Michael J.; Miaoulis, Ioannis; Cyr, Martha (2002). Prentice Hall Science Explorer: Chemical Building Blocks. Upper Saddle River, New Jersey: Prentice-Hall, Inc. p. 32. ISBN   978-0-13-054091-1. OCLC   47925884. There are 2,000,000,000,000,000,000,000 (that's 2 sextillion) atoms of oxygen in one drop of water—and twice as many atoms of hydrogen.
77. Feynman, Richard (1995). Six Easy Pieces. The Penguin Group. p. 5. ISBN   978-0-14-027666-4. OCLC   40499574.
78. "Radioactivity". Splung.com. Archived from the original on 4 December 2007. Retrieved 19 December 2007.
79. L'Annunziata, Michael F. (2003). Handbook of Radioactivity Analysis. Academic Press. pp. 3–56. ISBN   978-0-12-436603-9. OCLC   16212955.
80. Firestone, Richard B. (22 May 2000). "Radioactive Decay Modes". Berkeley Laboratory. Archived from the original on 29 September 2006. Retrieved 7 January 2007.
81. Hornak, J.P. (2006). "Chapter 3: Spin Physics". The Basics of NMR. Rochester Institute of Technology. Archived from the original on 3 February 2007. Retrieved 7 January 2007.
82. Schroeder, Paul A. (25 February 2000). "Magnetic Properties". University of Georgia. Archived from the original on 29 April 2007. Retrieved 7 January 2007.
83. Goebel, Greg (1 September 2007). "[4.3] Magnetic Properties of the Atom". Elementary Quantum Physics. In The Public Domain website. Archived from the original on 29 June 2011. Retrieved 7 January 2007.
84. Yarris, Lynn (Spring 1997). "Talking Pictures". Berkeley Lab Research Review. Archived from the original on 13 January 2008. Retrieved 9 January 2008.
85. Liang, Z.-P.; Haacke, E.M. (1999). Webster, J.G. (ed.). Encyclopedia of Electrical and Electronics Engineering: Magnetic Resonance Imaging. vol. 2. John Wiley & Sons. pp. 412–426. ISBN   978-0-471-13946-1.
86. Zeghbroeck, Bart J. Van (1998). "Energy levels". Shippensburg University. Archived from the original on 15 January 2005. Retrieved 23 December 2007.
87. Fowles, Grant R. (1989). Introduction to Modern Optics. Courier Dover Publications. pp. 227–233. ISBN   978-0-486-65957-2. OCLC   18834711.
88. Martin, W.C.; Wiese, W.L. (May 2007). "Atomic Spectroscopy: A Compendium of Basic Ideas, Notation, Data, and Formulas". National Institute of Standards and Technology. Archived from the original on 8 February 2007. Retrieved 8 January 2007.
89. "Atomic Emission Spectra – Origin of Spectral Lines". Avogadro Web Site. Archived from the original on 28 February 2006. Retrieved 10 August 2006.
90. Fitzpatrick, Richard (16 February 2007). "Fine structure". University of Texas at Austin. Archived from the original on 21 August 2011. Retrieved 14 February 2008.
91. Weiss, Michael (2001). "The Zeeman Effect". University of California-Riverside. Archived from the original on 2 February 2008. Retrieved 6 February 2008.
92. Beyer, H.F.; Shevelko, V.P. (2003). Introduction to the Physics of Highly Charged Ions. CRC Press. pp. 232–236. ISBN   978-0-7503-0481-8. OCLC   47150433.
93. Watkins, Thayer. "Coherence in Stimulated Emission". San José State University. Archived from the original on 12 January 2008. Retrieved 23 December 2007.
94. oxford dictionary – valency
95. Reusch, William (16 July 2007). "Virtual Textbook of Organic Chemistry". Michigan State University. Archived from the original on 29 October 2007. Retrieved 11 January 2008.
96. "Covalent bonding – Single bonds". chemguide. 2000. Archived from the original on 1 November 2008. Retrieved 20 November 2008.
97. Husted, Robert; et al. (11 December 2003). "Periodic Table of the Elements". Los Alamos National Laboratory. Archived from the original on 10 January 2008. Retrieved 11 January 2008.
98. Baum, Rudy (2003). "It's Elemental: The Periodic Table". Chemical & Engineering News. Archived from the original on 21 August 2011. Retrieved 11 January 2008.
99. Goodstein, David L. (2002). States of Matter. Courier Dover Publications. pp. 436–438. ISBN   978-0-13-843557-8.
100. Brazhkin, Vadim V. (2006). "Metastable phases, phase transformations, and phase diagrams in physics and chemistry". Physics-Uspekhi. 49 (7): 719–724. Bibcode:2006PhyU...49..719B. doi:10.1070/PU2006v049n07ABEH006013.
101. Myers, Richard (2003). The Basics of Chemistry. Greenwood Press. p. 85. ISBN   978-0-313-31664-7. OCLC   50164580.
102. Staff (9 October 2001). "Bose-Einstein Condensate: A New Form of Matter". National Institute of Standards and Technology. Archived from the original on 3 January 2008. Retrieved 16 January 2008.
103. Colton, Imogen; Fyffe, Jeanette (3 February 1999). "Super Atoms from Bose-Einstein Condensation". The University of Melbourne. Archived from the original on 29 August 2007. Retrieved 6 February 2008.
104. Jacox, Marilyn; Gadzuk, J. William (November 1997). "Scanning Tunneling Microscope". National Institute of Standards and Technology. Archived from the original on 7 January 2008. Retrieved 11 January 2008.
105. "The Nobel Prize in Physics 1986". The Nobel Foundation. Archived from the original on 17 September 2008. Retrieved 11 January 2008. – in particular, see the Nobel lecture by G. Binnig and H. Rohrer.
106. Jakubowski, N.; Moens, Luc; Vanhaecke, Frank (1998). "Sector field mass spectrometers in ICP-MS". Spectrochimica Acta Part B: Atomic Spectroscopy. 53 (13): 1739–1763. Bibcode:1998AcSpe..53.1739J. doi:10.1016/S0584-8547(98)00222-5.
107. Müller, Erwin W.; Panitz, John A.; McLane, S. Brooks (1968). "The Atom-Probe Field Ion Microscope". Review of Scientific Instruments . 39 (1): 83–86. Bibcode:1968RScI...39...83M. doi:10.1063/1.1683116.
108. Lochner, Jim; Gibb, Meredith; Newman, Phil (30 April 2007). "What Do Spectra Tell Us?". NASA/Goddard Space Flight Center. Archived from the original on 16 January 2008. Retrieved 3 January 2008.
109. Winter, Mark (2007). "Helium". WebElements. Archived from the original on 30 December 2007. Retrieved 3 January 2008.
110. Hinshaw, Gary (10 February 2006). "What is the Universe Made Of?". NASA/WMAP. Archived from the original on 31 December 2007. Retrieved 7 January 2008.
111. Choppin, Gregory R.; Liljenzin, Jan-Olov; Rydberg, Jan (2001). Radiochemistry and Nuclear Chemistry. Elsevier. p. 441. ISBN   978-0-7506-7463-8. OCLC   162592180.
112. Davidsen, Arthur F. (1993). "Far-Ultraviolet Astronomy on the Astro-1 Space Shuttle Mission". Science . 259 (5093): 327–334. Bibcode:1993Sci...259..327D. doi:10.1126/science.259.5093.327. PMID   17832344.
113. Lequeux, James (2005). The Interstellar Medium. Springer. p. 4. ISBN   978-3-540-21326-0. OCLC   133157789.
114. Smith, Nigel (6 January 2000). "The search for dark matter". Physics World. Archived from the original on 16 February 2008. Retrieved 14 February 2008.
115. Croswell, Ken (1991). "Boron, bumps and the Big Bang: Was matter spread evenly when the Universe began? Perhaps not; the clues lie in the creation of the lighter elements such as boron and beryllium". New Scientist (1794): 42. Archived from the original on 7 February 2008. Retrieved 14 January 2008.
116. Copi, Craig J.; Schramm, DN; Turner, MS (1995). "Big-Bang Nucleosynthesis and the Baryon Density of the Universe". Science (Submitted manuscript). 267 (5195): 192–199. arXiv:. Bibcode:1995Sci...267..192C. doi:10.1126/science.7809624. PMID   7809624. Archived from the original on 14 August 2019. Retrieved 27 July 2018.
117. Hinshaw, Gary (15 December 2005). "Tests of the Big Bang: The Light Elements". NASA/WMAP. Archived from the original on 17 January 2008. Retrieved 13 January 2008.
118. Abbott, Brian (30 May 2007). "Microwave (WMAP) All-Sky Survey". Hayden Planetarium. Archived from the original on 13 February 2013. Retrieved 13 January 2008.
119. Hoyle, F. (1946). "The synthesis of the elements from hydrogen". Monthly Notices of the Royal Astronomical Society . 106 (5): 343–383. Bibcode:1946MNRAS.106..343H. doi:10.1093/mnras/106.5.343.
120. Knauth, D.C.; Knauth, D.C.; Lambert, David L.; Crane, P. (2000). "Newly synthesized lithium in the interstellar medium". Nature . 405 (6787): 656–658. Bibcode:2000Natur.405..656K. doi:10.1038/35015028. PMID   10864316.
121. Mashnik, Stepan G. (2000). "On Solar System and Cosmic Rays Nucleosynthesis and Spallation Processes". arXiv:.
122. Kansas Geological Survey (4 May 2005). "Age of the Earth". University of Kansas. Archived from the original on 5 July 2008. Retrieved 14 January 2008.
123. Manuel 2001, pp. 407–430, 511–519.
124. Dalrymple, G. Brent (2001). "The age of the Earth in the twentieth century: a problem (mostly) solved". Geological Society, London, Special Publications. 190 (1): 205–221. Bibcode:2001GSLSP.190..205D. doi:10.1144/GSL.SP.2001.190.01.14. Archived from the original on 11 November 2007. Retrieved 14 January 2008.
125. Anderson, Don L.; Foulger, G.R.; Meibom, Anders (2 September 2006). "Helium: Fundamental models". MantlePlumes.org. Archived from the original on 8 February 2007. Retrieved 14 January 2007.
126. Pennicott, Katie (10 May 2001). "Carbon clock could show the wrong time". PhysicsWeb. Archived from the original on 15 December 2007. Retrieved 14 January 2008.
127. Yarris, Lynn (27 July 2001). "New Superheavy Elements 118 and 116 Discovered at Berkeley Lab". Berkeley Lab. Archived from the original on 9 January 2008. Retrieved 14 January 2008.
128. Diamond, H; et al. (1960). "Heavy Isotope Abundances in Mike Thermonuclear Device". Physical Review . 119 (6): 2000–2004. Bibcode:1960PhRv..119.2000D. doi:10.1103/PhysRev.119.2000.
129. Poston Sr.; John W. (23 March 1998). "Do transuranic elements such as plutonium ever occur naturally?". Scientific American. Archived from the original on 27 March 2015. Retrieved 1 May 2015.
130. Keller, C. (1973). "Natural occurrence of lanthanides, actinides, and superheavy elements". Chemiker Zeitung. 97 (10): 522–530. OSTI   4353086.
131. Zaider, Marco; Rossi, Harald H. (2001). Radiation Science for Physicians and Public Health Workers. Springer. p. 17. ISBN   978-0-306-46403-4. OCLC   44110319.
132. "Oklo Fossil Reactors". Curtin University of Technology. Archived from the original on 18 December 2007. Retrieved 15 January 2008.
133. Weisenberger, Drew. "How many atoms are there in the world?". Jefferson Lab. Archived from the original on 22 October 2007. Retrieved 16 January 2008.
134. Pidwirny, Michael. "Fundamentals of Physical Geography". University of British Columbia Okanagan. Archived from the original on 21 January 2008. Retrieved 16 January 2008.
135. Anderson, Don L. (2002). "The inner inner core of Earth". Proceedings of the National Academy of Sciences . 99 (22): 13966–13968. Bibcode:2002PNAS...9913966A. doi:10.1073/pnas.232565899. PMC  . PMID   12391308.
136. Pauling, Linus (1960). The Nature of the Chemical Bond. Cornell University Press. pp. 5–10. ISBN   978-0-8014-0333-0. OCLC   17518275.
137. Anonymous (2 October 2001). "Second postcard from the island of stability". CERN Courier. Archived from the original on 3 February 2008. Retrieved 14 January 2008.
138. Jacoby, Mitch (2006). "As-yet-unsynthesized superheavy atom should form a stable diatomic molecule with fluorine". Chemical & Engineering News . 84 (10): 19. doi:10.1021/cen-v084n010.p019a.
139. Koppes, Steve (1 March 1999). "Fermilab Physicists Find New Matter-Antimatter Asymmetry". University of Chicago. Archived from the original on 19 July 2008. Retrieved 14 January 2008.
140. Cromie, William J. (16 August 2001). "A lifetime of trillionths of a second: Scientists explore antimatter". Harvard University Gazette. Archived from the original on 3 September 2006. Retrieved 14 January 2008.
141. Hijmans, Tom W. (2002). "Particle physics: Cold antihydrogen". Nature . 419 (6906): 439–440. Bibcode:2002Natur.419..439H. doi:. PMID   12368837.
142. Staff (30 October 2002). "Researchers 'look inside' antimatter". BBC News. Archived from the original on 22 February 2007. Retrieved 14 January 2008.
143. Barrett, Roger (1990). "The Strange World of the Exotic Atom". New Scientist (1728): 77–115. Archived from the original on 21 December 2007. Retrieved 4 January 2008.
144. Indelicato, Paul (2004). "Exotic Atoms". Physica Scripta . T112 (1): 20–26. arXiv:. Bibcode:2004PhST..112...20I. doi:10.1238/Physica.Topical.112a00020. Archived from the original on 4 November 2018. Retrieved 4 November 2018.
145. Ripin, Barrett H. (July 1998). "Recent Experiments on Exotic Atoms". American Physical Society. Archived from the original on 23 July 2012. Retrieved 15 February 2008.