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Isotopes

Isotopes

Isotopes are forms of an element whose nuclei have the same atomic number–-the number of protons in the nucleus--but different atomic masses because they contain different numbers of neutrons. For example, the isotope Carbon-14 may, via radioactive decay, become Carbon-13, but both are Carbon isotopes. The word isotope, meaning at the same place, comes from the fact that all isotopes of an element are located at the same place on the periodic table. Collectively, the isotopes of the elements form the set of nuclides. A nuclide is a particular type of atomic nucleus, or more generally an agglomeration of protons and neutrons. Strictly speaking, it is more correct to say that an element such as fluorine consists of one stable nuclide rather than that it has one stable isotope. In scientific nomenclature, isotopes (nuclides) are specified by the name of the particular element by a hyphen and the number of nucleons (protons and neutrons) in the atomic nucleus (e.g., helium-3, carbon-12, carbon-14, iron-57, uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g., ³He, ¹²C, ¹⁴C, ⁵⁷Fe, ²³⁸U).

Variation in properties between isotopes

In a neutral atom, the number of electrons equals the number of protons. Thus, different isotopes of a given element also have the same number of electrons and the same electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, isotopes exhibit nearly identical chemical behavior. The primary exception is that, due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. (This phenomenon is termed the kinetic isotope effect). This "mass effect" is most pronounced for protium (¹H) vis-à-vis deuterium (²H), because deuterium has twice the mass of protium. For heavier elements the relative mass difference between isotopes is much less, and the mass effect is usually negligible. Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have nearly identical electronic structure, and therefore have similar physical and chemical properties. The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range. Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Atomic nuclei consist of protons and neutrons bound together by the strong nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion and stabilizing the nucleus. For this reason neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus; for example, although the neutron/proton ratio of ³He is 1/2, the neutron/proton ratio of ²³⁸U is greater than 3/2. If too many neutrons or too few neutrons are present, the nucleus becomes unstable and subject to nuclear decay.

Occurrence in nature

Several isotopes of each element can be found in nature. The relative abundance of an isotope is strongly correlated with its tendency toward nuclear decay; short-lived nuclides quickly decay away, while their long-lived counterparts endure. However, this does not mean that short-lived species disappear entirely; many are continually produced through the decay of longer-lived nuclides. The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. According to generally accepted cosmology, virtually all nuclides other than isotopes of hydrogen and helium were built in stars and supernovae. Their respective abundances here result from the quantities formed by these processes, their spread through the galaxy, and their rates of decay. After the initial coalescence of the solar system, isotopes were redistributed according to mass (see also Origin of the solar system). The isotopic composition of elements is different on different planets, making it possible to determine the origin of meteorites.

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element.

Use of chemical properties

- One of the most common applications is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy (see "Properties"). If radioactive isotopes are used, they can be detected by the radiation they emit (this is radioisotopic labelling).
- A technique similar to radioisotopic labelling is radiometric dating (most famously radiocarbon dating). It can be used to study chemical processes that the experimenter does not witness, by using naturally-occurring isotopic tracers.
- Isotopic substitution can be used to determine the mechanism of a reaction via the kinetic isotope effect.

Use of nuclear properties

- Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are ¹H, ²D, ¹³C, and ³¹P.
- Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as ⁵⁷Fe.
- Radionuclides also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes. The process of isotope separation represents a significant technological challenge.

External links

- [http://physics.nist.gov/cgi-bin/Compositions/stand_alone.pl?ele=&ascii=html&isotype=some Atomic weights of all isotopes]
- [http://atom.kaeri.re.kr/ Atomgewichte, Zerfallsenergieen und Halbwärtszeiten aller Isotope]
- [http://ie.lbl.gov/education/isotopes.htm Exploring the Table of the Isotopes] at the LBNL
- Category:Nuclear chemistry Category:Nuclear physics ko:동위원소 ja:同位体 simple:Isotope th:ไอโซโทป

Chemical element

A chemical element, often called simply element, is a chemical substance that canot be divided or changed into other chemical substances by any ordinary chemical technique. The smallest unit of this kind of chemical substances is an atom. An element is a class of substances that contain the same number of protons in all its atoms.

Chemistry terminology

Earlier an element or pure element was defined as a substance which "cannot be further broken down into another compound with different chemical properties" -- which should be taken to mean it consists of atoms of one element. However, due to allotropy, the isotope effect, and the confusion with the more useful term referring to the general class of atoms (irrespective of what compound it may be in), this usage is in disfavor amongst contemporary chemists, and sees restricted, mostly historical, use. This definition was motivated by the observation that these elements could not be dissociated by chemical means into other compounds. For example, water could be converted into hydrogen and oxygen, but hydrogen and oxygen could not be further decomposed, thus "elemental". There are also many counterexamples (for example "elemental oxygen" (O₂) can be decomposed by solely chemical means into oxygen ions and atoms which have drastically different chemical properties). The remainder of this article will concern itself with the first definition.

Description

The atomic number of an element, Z, is equal to the number of protons which defines the element. For example, all carbon atoms contain 6 protons in their nucleus, so for carbon Z=6. These atoms may have different amounts of neutrons, and are known as isotopes of the element. The atomic mass of an element, A, is measured in unified atomic mass units (u) is the average mass of all the atoms of the element in an environment of interest (usually the earth's crust and atmosphere). Since electrons are light, and neutrons are barely more than the mass of the proton, this usually corresponds to the sum of the protons and neutrons in the nucleus of the most abundant isotope, though this is not always the case (notably chlorine, which is about three-quarters ³⁵Cl and a quarter ³⁷Cl). Some isotopes are radioactive and decay into other elements upon radiating an alpha or beta particle. Some elements have no nonradioactive isotopes, in particular all elements with Z >= 84. The lightest elements are hydrogen and helium. Hydrogen is thought to be the first element to appear after the Big Bang. All the heavier elements, are made naturally and artificially through various methods of nucleosynthesis. As of 2005, there are 116 known elements: 93 occur naturally on earth (including technetium and plutonium), and 94 (including promethium) have been detected so far in the universe. The 23 elements not found on earth are derived artificially; the first purportedly synthesized element was technetium, in 1937, although the trace amounts of naturally occurring technetium were not known then. All artificially derived elements are radioactive with short half-lives so that any such atoms that were present at the formation of Earth are extremely likely to have already decayed. Lists of the elements by name, by symbol, by atomic number, by density, by melting point and by boiling point are available. The most convenient presentation of the elements is in the periodic table, which groups elements with similar chemical properties together.

Nomenclature

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When it was learned, existing names (e.g., gold, mercury, iron) were kept in most countries, and national differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For example, the Germans use "Wasserstoff" for "hydrogen" and "Sauerstoff" for "oxygen," while some romance languages use "natrium" for "sodium" and "kalium" for "potassium," and the French prefer the obsolete but historic term "azote" for "nitrogen." But for international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that "aluminium" and "caesium" take the place of the US spellings "aluminum" and "cesium," while the US "sulfur" takes the place of the British "sulphur." But chemicals which are practicable to be sold in bulk within many countries, however, still have national names, and those which do not use the Latin alphabet cannot be expected to use the IUPAC name. According to IUPAC, the full name of an element is not capitalized, even if it is derived from a proper noun (unless it would be capitalized by some other rule, for instance if it begins a sentence). And in the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have too quick a decay rate to ever be sold in bulk. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question which delayed the naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy). Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century (e.g., as "lutetium" refers to Paris, France, the Germans were reticent about relinquishing naming rights to the French, often calling it "cassiopeium"). And notably, the British discoverer of "niobium" originally named it "columbium," after the New World, though this did not catch on in Europe. The Americans had to accept the international name just when it was becoming an economically important material late in the twentieth century.

Chemical symbols

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of one atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules. These were superseded by the current typographical system in which chemical symbols are not used as mere abbreviations though each consists letters of the Latin alphabet - they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully international, for they were based on the Latin abbreviations of the names of metals: Fe comes from Ferrum; Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Besides a name, later chemical elements are also given a unique chemical symbol, based on the name of the element, not necessarily derived from the colloquial English name. (e.g., sodium has chemical symbol 'Na' after the Latin natrium). The same applies to "W" (wolframium) for Tungsten , "Hg" (Hydrargyrum) for mercury and "K" for potassium. Stricly taken, a symbol like Tu for tungsten or M or Me for mercury seems to be more logical. Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not be confused with a roman numeral. The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always minuscule (small letters).

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical (not to be confused with radical_(chemistry), meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of Yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

Nonelement symbols

Nonelements, especially in organic and organometallic chemistry, often acquire symbols which are inspired by the elemental symbols. A few examples: Cy - cyclohexyl; Ph - phenyl; Bz - benzoyl; Bn - benzyl; Cp - Cyclopentadiene; Pr - propyl; Me - methyl; Et - ethyl; Tf - triflate; Ts - tosyl.

External links

- [http://www.vanderkrogt.net/elements/ Elementymology & Elements Multidict] word history and language dictionary

Chemical information

- [http://www.webelements.com/ WebElements]
- [http://www.vcs.ethz.ch/chemglobe/ptoe/ ChemGlobe]
- [http://pearl1.lanl.gov/periodic/default.htm Los Alamos National Laboratory]
- [http://www.chemicalelements.com/ ChemicalElements] ko:화학 원소 ms:Unsur kimia ja:元素 simple:Element th:ธาตุเคมี

Nucleus

Nucleus usually refers to the center of something, but can mean:
- atomic nucleus, the collection of protons and neutrons in the center of an atom that carries the bulk of the atom's mass and positive charge
- cell nucleus, the membrane-bound subcellular organelle found in eukaryotes, visible via microscopy, which contains, primarily, the cell's chromosomes
- nucleus (neuroanatomy), a central nervous system structure composed mainly of gray matter that mediates electrical signaling within a particular subsystem
- comet nucleus, the solid core of a comet
- galaxy nucleus, the central region of a galaxy
- ice nucleus, the center of an ice crystal
- cloud condensation nuclei, the basis for the development of a cloud droplet
- syllable nucleus, the central part of a syllable
- sentence nucleus, the syllable which receives the greatest stress in a word
- Nucleus CMS, an open-source weblog system
- Nucleus RTOS, a brand of operating system
- Nucleus (band), a British jazz-rock band led by Ian Carr

Etymology

"Nucleus" is New Latin, the diminutive of the Latin nux (nut). ko:핵 th:นิวเคลียส

Proton

:For alternative meanings see proton (disambiguation). In physics, the proton (Greek proton = first) is a subatomic particle with an electric charge of one positive fundamental unit (1.602 × 10⁻¹⁹ coulomb) and a mass of 938.3 MeV/c² (1.6726 × 10⁻²⁷ kg), or about 1836 times the mass of an electron. The proton is observed to be stable, with a lower limit on its half-life of about 10³⁵ years, although some theories predict that the proton may decay. The proton and neutron are both nucleons. The nucleus of the most common isotope (called protium) of the hydrogen atom is a single proton. The nuclei of other atoms are composed of protons and neutrons held together by the strong nuclear force. The number of protons in the nucleus determines the chemical properties of the atom and which chemical element it is. Protons are classified as baryons and are composed of two up quarks and one down quark, which are also held together by the strong nuclear force, mediated by gluons. The proton's antimatter equivalent is the antiproton, which has the same magnitude charge as the proton but the opposite sign. In chemistry and biochemistry, the term proton may refer to the hydrogen ion, H⁺. In this context, a proton donor is an acid and a proton acceptor a base (see acid-base reaction theories).

History

Ernest Rutherford is generally credited with the discovery of the proton. In 1918 Rutherford noticed that when alpha particles were shot into nitrogen gas, his scintillation detectors showed the signatures of hydrogen nuclei. Rutherford determined that the only place this hydrogen could have come from was the nitrogen, and therefore nitrogen must contain hydrogen nuclei. He thus suggested that the hydrogen nucleus, which was known to have an atomic number of 1, was an elementary particle. Prior to Rutherford, Eugene Goldstein had observed canal rays, which were composed of positively charged ions.

Technological applications

Protons can exist in spin states. This property is exploited by nuclear magnetic resonance spectroscopy. In NMR spectroscopy, a magnetic field is applied to a substance in order to detect the shielding around the protons in the nuclei of that substance, which is provided by the surrounding electron clouds. Scientists can use this information to then construct the molecular structure of the molecule under study.

Antiproton

The antiproton is the antiparticle of the proton. It was discovered in the year 1955 by Emilio Segre and Owen Chamberlain, for which they were awarded a 1959 Nobel Prize in Physics. CPT-symmetry puts strong constraints on the relative properties of particles and antiparticles and, therefore, is open to stringent tests. For example, the charges of the proton and antiproton must sum to exactly zero. This equality has been tested to one part in 10^-8. The equality of their masses is also tested to better than one part in 10^-8. By holding antiprotons in a Penning trap, the equality of the charge to mass ratio of the proton and the antiproton has been tested to 1 part in 9×10^-11. The magnetic moment of the antiproton has been found with error of 8×10^-3 nuclear Bohr magnetons, and is found to be equal and opposite to that of the proton.

External links

- [http://pdg.lbl.gov/ Particle Data Group] Category:Nucleon ko:양성자 ms:Proton ja:陽子 simple:Proton th:โปรตอน

Neutron

Classification

Neutron

Properties

Mass:	1.674 927 16(13) × 10⁻²⁷ kg
	939.565 530(38) MeV/c²
Electric charge:	0 C
Spin:	½
Magnetic dipole moment:	-1.91304 μ_N
Quark composition:	2 Down, 1 Up

In physics, the neutron is a subatomic particle with no net electric charge and a mass of 939.573 MeV/c² ( kg, slightly more than a proton). Its spin is ½. Its antiparticle is called the antineutron. The neutron and proton are instances of a nucleon. The nucleus of most atoms (all except the most common isotope of hydrogen, which consists of a single proton only) consists of protons and neutrons.

Properties

Outside the nucleus, neutrons are unstable and have a mean lifetime of 886 seconds (about 15 minutes, uncertainty about 2 s [http://www.ill.fr/pages/menu_g/docs/universe2003.pdf]), decaying by emitting an electron and antineutrino to become a proton. Neutrons in this unstable form are known as free neutrons. The same decay method (beta decay) occurs in some nuclei. Particles inside the nucleus are typically resonances between neutrons and protons, which transform into one another by the emission and absorption of pions. A neutron is classified as a baryon, and consists of two down quarks and one up quark. The neutron's antimatter equivalent is the antineutron. The number of neutrons determines the isotope of an element. (For example, the carbon-12 isotope has 6 protons and 6 neutrons, while the carbon-14 isotope has 6 protons and 8 neutrons.) Isotopes are atoms of the same element that have the same atomic number but different masses due to a different number of neutrons.

Neutron Interactions

The neutron interacts through all four of the common classifications of physical interaction. These four are the electromagnetic, weak nuclear, strong nuclear and gravitational interactions. Although it is true that the neutron has zero net charge, it is nonetheless composed of electrically charged quarks, in the same way that a neutral atom is nonetheless composed of protons and electrons. As such, the neutron experiences the electromagnetic interaction. The net charge is zero, so if you are far enough away from the neutron that it appears to occupy no volume, then the total effect of the electric force will add up to zero. The movement of the charges inside the neutrons do not cancel however, and this is what gives the neutron its nonzero magnetic moment. Gravity is often not discussed when talking about neutrons. This is because neutrons are usually studied in terms of subatomic interactions. In the subatomic world, gravity is undetectable relative to the other forces which are much stronger. This having been said, a neutron accelerates at the same rate in the earth's gravitational field as a lead brick. Charged particles (such as protons, electrons, or alpha particles) and electromagnetic radiation (such as gamma rays) lose energy in passing through matter. They exert electric forces which ionize atoms of the material through which they pass. The energy taken up in ionization equals the energy lost by the charged particle, which slows down, or by the gamma ray, which is absorbed or scattered (see compton scattering). The neutron, in contrast, is seen by atoms it passes as containing no electric charge, and so does not create any ionization. As far as the nuclear forces are concerned, it is a different story. Nuclear forces play the leading role when neutrons pass through regular matter. Consequently, a free neutron goes on its way unchecked until it makes a "head-on" collision with an atomic nucleus. When this happens, the neutrons and target nuclei can be scattered (deflected or slowed down), absorbed, or transformed into something different. In the case of the reaction n + ³He → ¹H + ³H (n:neutron; ³He: nucleus consisting of two protons and one neutron; ¹H: nucleus consisting of a only proton; ³H: nucleus consisting of one proton and two neutrons) for example, the proton and the neutron appear to have exchanged places, and kinetic energy is released. In many cases, secondary particles are created and energy can be used up or released. Neutrons, like other particles, can undergo elastic collisions. A collision is elastic under the special case where kinetic energy is conserved. Billiard balls for example typically undergo elastic collisions. The law of conservation of momentum also applies as it does for any collision. If the nucleus that is struck in an elastic collision is heavy, it acquires relatively little speed, but if it is a proton, which is approximately equal in mass to the neutron, it is projected forward with a large fraction of the original speed of the neutron, which is itself correspondingly slowed.

Neutron Detection

The common means of detecting a charged particle by looking for a track of ionization does not work for neutrons directly. Neutrons that elastically scatter off of another atom can create an ionization track that is detectable, but the experiments are not as simple to carry out and other means for detecting neutrons, consisting of allowing them to interact with atomic nuclei, are more commonly used. A common method for detecting neutrons involves converting the energy released from such reactions into electrical signals. The nuclides ³He, ⁶Li, ¹⁰B, ²³³U, ²³⁵U, ²³⁷Np and ²³⁹Pu are useful for this purpose. A good discussion on neutron detection is found in chapter 14 of the book Radiation Detection and Measurement by Glenn F. Knoll (John Wiley & Sons, 1979).

Neutron Uses

The neutron plays an important role in many nuclear reactions. For example, neutron capture often results in neutron activation, inducing radioactivity. In particular, knowledge of neutrons and their behavior has been important in the development of nuclear reactors and nuclear weapons. The development of [http://www.nature.com/nature/journal/v357/n6377/abs/357390a0.html "neutron lenses"] based on total internal reflection within hollow glass capillary tubes or by reflection from dimpled aluminum plates has driven ongoing research into [http://www.physorg.com/news599.html neutron microscopy] and [http://www.nasa.gov/vision/earth/technologies/nuggets.html neutron/gamma ray tomography]. One use of neutron emitters is the detection of light nuclei, particularly the hydrogen found in water molecules. When a fast neutron collides with a light nucleus, it loses a large fraction of its energy. By measuring the rate at which slow neutrons return to the probe after reflecting off of hydrogen nuclei, a neutron probe may determine the water content in soil.

Neutron Sources

Due to the fact that free neutrons are unstable, they (neutron radiation) can be obtained only from nuclear disintegrations, nuclear reactions, and high-energy reactions (such as in cosmic radiation showers or accelerator collisions). Free neutron beams are obtained from neutron sources by neutron transport. For access to intense neutron sources, researchers must go to specialist facilties, such as the ISIS facility in the UK, which is currently the world's most intense pulsed neutron and muon source. Neutrons' lack of total electric charge prevents engineers or experimentalists from being able to steer or accelerate them. Charged particles can be accelerated, decelerated, or deflected by electric or magnetic fields. However, these methods have almost no effect on neutrons (there is a small effect of a magnetic field on the free neutron because of its magnetic moment).

Discovery

In 1930 Walther Bothe and H. Becker in Germany found that if the very energetic alpha particles emitted from polonium fell on certain of the light elements, specifically beryllium, boron, or lithium, an unusually penetrating radiation was produced. At first this radiation was thought to be gamma radiation although it was more penetrating than any gamma rays known, and the details of experimental results were very difficult to interpret on this basis. The next important contribution was reported in 1932 by Irène Joliot-Curie and Frédéric Joliot in Paris. They showed that if this unknown radiation fell on paraffin or any other hydrogen-containing compound it ejected protons of very high energy. This was not in itself inconsistent with the assumed gamma ray nature of the new radiation, but detailed quantitative analysis of the data became increasingly difficult to reconcile with such a hypothesis. Finally (later in 1932) the physicist James Chadwick in England performed a series of experiments showing that the gamma ray hypothesis was untenable. He suggested that in fact the new radiation consisted of uncharged particles of approximately the mass of the proton, and he performed a series of experiments verifying his suggestion. Such uncharged particles were eventually called neutrons, apparently from the Latin root for neutral and the Greek ending -on (by imitation of electron and proton).

Current developments

The existence of stable clusters of four neutrons, or tetraneutrons, has been hypothesised by a team led by Francisco-Miguel Marqués at the CNRS Laboratory for Nuclear Physics based on observations of the disintegration of beryllium-14 nuclei. This is particularly interesting, because current theory suggests that these clusters should not be stable, and therefore should not exist.

Antineutron

The antineutron is the antiparticle of the neutron. It was discovered by Bruce Cork in the year 1956, a year after the antiproton was discovered. CPT-symmetry puts strong constraints on the relative properties of particles and antiparticles and, therefore, is open to stringent tests. The masses of the neutron and antineutron are equal to one part in (9±5)×10^-5.

Fluorine

Fluorine (from L. fluere, meaning "to flow"), is the chemical element in the periodic table that has the symbol F and atomic number 9. It is a poisonous pale yellow-green, univalent gaseous halogen that is the most chemically reactive and electronegative of all the elements. In its pure form, it is highly dangerous, causing severe chemical burns on contact with skin.

Notable characteristics

Pure fluorine is a corrosive pale yellow gas that is a powerful oxidizing agent. It is the most reactive and electronegative of all the elements, and readily forms compounds with most other elements. Fluorine even combines with the noble gases krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. It is so reactive that, glass, metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. It is far too reactive to be found in elemental form and has such an affinity for most elements, including silicon, that it can neither be prepared nor should be kept in glass vessels. In moist air it reacts with water to form the equally dangerous hydrofluoric acid. In aqueous solution, fluorine commonly occurs as the fluoride ion F^-. Other forms are fluoro-complexes (such as [FeF₄]^-) or H₂F⁺. Fluorides are compounds that combine fluoride with some positively charged counterpart. They often consist of ions. Fluorine compounds with metals are among the most stable of salts.

Applications

Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMs fabrication. Other uses:
- Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
- Fluorine is indirectly used in the production of low friction plastics such as Teflon, and in halons such as Freon
- Along with some of its compounds, fluorine is used in the production of uranium (from the hexafluoride) and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, textiles, etc.
- Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration. Chlorofluorocarbons have been banned for these applications because they contribute to the ozone hole.
- Sulfur hexafluoride is an extremely inert and nontoxic gas. These classes of compounds are potent greenhouse gases
- Many important agents for general anaesthesia are fluorohydrocarbon derivatives, e.g. sevoflurane, desflurane, and isoflurane.
- Potassium hexafluoroaluminate, the so-called cryolite, is used in electrolysis of aluminium.
- Sodium fluoride has been used as an insecticide, especially against cockroaches.
- Some other fluorides are often added to toothpaste and, somewhat controversially, to municipal water supplies to prevent dental cavities.
- It has been used in the past to help molten metal flow, hence the name.
- Fluorine-18, a radioactive isotope that emits positrons, is often used in positron emission tomography because of its half-life of 110 minutes. Some researchers - including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. Experiments failed since fluorine was so hard to handle.

History

Fluorine in the form of fluorspar (calcium fluoride) was described in 1529 by Georgius Agricola for its use as a flux, which is a substance that is used to promote the fusion of metals or minerals. In 1670 Schwandhard found that glass was etched when it was exposed to fluorspar that was treated with acid. Karl Scheele and many later researchers, including Humphry Davy, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating calcium fluoride (fluorspar) with concentrated sulfuric acid. It was eventually realized that hydrofluoric acid contained a previously unknown element. This element was not isolated for many years after this due to its extreme reactivity - it is separated from its compounds only with difficulty and then it immediately attacks the remaining materials of the compound. Finally in 1886 fluorine was isolated by Henri Moissan after almost 74 years of continuous effort. It was an effort which cost several researchers their health or even their lives, and for Moissan, it earned him the 1906 Nobel Prize in chemistry. The first large scale production of fluorine was needed for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF₆) was used to separate the U-235 and U-238 isotopes of uranium. Today both the gaseous diffusion process and the gas centrifuge process use gaseous (UF₆) to produce enriched uranium for nuclear power applications. The derevation of elemental flourine from hydroflouric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These men came to be referred to as "Flourine Martyrs".

Precautions

Both fluorine and HF must be handled with great care and any contact with skin and eyes should be strictly avoided. All equipment must be passivated before exposure to fluorine. Both elemental fluorine and fluoride ions are highly toxic. When it is a free element, fluorine has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. It is recommended that the maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 µL/L (part per million by volume) (lower than, for example, hydrogen cyanide). Fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite. However, safe handling procedures enable the transport of liquid fluorine by the ton.

Preparation

Elemental fluorine is prepared industrially by Moissan's original process: electrolysis of anhydrous HF in which KHF₂ has been dissolved to provide enough ions for conduction to take place. In 1986, preparing for a conference to celebrate the 100th aniversary of the discovery of fluorine, Karl Christe discovered a purely-chemical preparation by reacting together at 150C solutions in anhydrous HF of K₂MnF₆ and of SbF₅. This is not a practical synthesis, but demonstrates that electrolysis is not essential.

Compounds

Fluorine can often be substituted for hydrogen when it occurs in organic compounds. Through this mechanism, fluorine can have a very large number of compounds. Fluorine compounds involving noble gases were first synthesised by Neil Bartlett in 1962 - xenon hexafluoroplatinate, XePtF₆, being the first. Fluorides of krypton and radon have also been prepared. Also Argon Fluorohydride has been prepared, althought it is only stable at cryogenic temperatures. This element is recovered from fluorite, cryolite, and fluorapatite. See also: Fluorocarbon
- Ammonium fluoride (NH₄F)
- Antimony pentafluoride (SbF₅)
- Boron trifluoride (BF₃)
- Bromine pentafluoride (BrF₅)
- Bromine trifluoride (BrF₃)
- Caesium fluoride (CsF)
- Calcium fluoride (CaF₂)
- Chlorine pentafluoride (ClF₅)
- Fluorosulfuric acid (FSO₃(H)
- Hydrofluoric Acid (HF)
- Iodine pentafluoride (IF₅)
- Iodine heptafluoride (IF₇)
- Lithium fluoride (LiF)
- Nitrogen trifluoride (NF₃)
- Nitrosyl fluoride (NOF)
- Nitryl fluoride (NO₂F)
- Phosphorus trifluoride (PF₃)
- Phosphorus pentafluoride (PF₅)
- Potassium fluoride (KF)
- Radon difluoride (RnF₂)
- Silver(I) fluoride (AgF)
- Sulfur hexafluoride (SF₆)
- Thionyl fluoride (SOF₂)
- Tungsten(VI) fluoride (WF₆)
- Uranium hexafluoride (UF₆)
- Xenon hexafluoroplatinate (XePtF₆)
- Xenon tetrafluoride (XeF₄)

References

- [http://periodic.lanl.gov/elements/9.html Los Alamos National Laboratory – Fluorine]

External links

- [http://www.webelements.com/webelements/elements/text/F/index.html WebElements.com – Fluorine]
- [http://education.jlab.org/itselemental/ele009.html It's Elemental – Fluorine]
- [http://www.chemie-master.de/pse/pse.php?modul=F Picture of liquid fluorine – chemie-master.de]
- [http://www.chemsoc.org/viselements/pages/fluorine.html Chemsoc.org]
- [http://nautilus.fis.uc.pt/st2.5/index-en.html Periodic Table of Elements] Category:Chemical elements Category:Halogens ko:플루오르 ja:フッ素 th:ฟลูออรีน

Carbon-12

Carbon-12 is the more abundant (98.89%) of the two stable isotopes of the element carbon. It contains 6 protons, 6 neutrons and 6 electrons. Carbon-12 is of particular importance as it is used as the standard from which all other isotopes' atomic weight is measured and thus the measurement of Avogadro's number.

History

Prior to 1959 both the IUPAP and IUPAC tended to use used oxygen to define the mole, the chemists defining the mole as the number of atoms of oxygen which had mass 16 g, the physicists using a similar definition but with the oxygen-16 isotope only. The two organizations agreed in 1959/1960 to define the mole as: "The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon 12; its symbol is "mol."" This was adopted by the CIPM (International Committee for Weights and Measures) in 1967, and in 1971 it was adopted by the 14th CGPM (General Conference on Weights and Measures). In 1980 the CIPM clarified the above definition, defining that the carbon-12 atoms are unbound and in their ground state.

Carbon-14

Carbon-14, or ¹⁴C, is a radioactive isotope of carbon discovered February 27, 1940, by Martin Kamen and Sam Ruben. Its nucleus contains 6 protons and 8 neutrons. Its presence in organic materials is used in radiocarbon dating. It occurs naturally and has a relative abundance up to 0.00000000012%. The half-life of carbon-14 is 5730 years. It decays into nitrogen-14 through beta-decay. Carbon-14 is produced in the upper layers of the troposphere and the stratosphere by thermal neutrons absorbed by nitrogen atoms. When cosmic rays enter the atmosphere, they undergo various transformations, including the production of neutrons. The resulting neutrons participate in the following reaction: :n + ¹⁴N → ¹⁴C + ¹H This reaction is relatively common, as nitrogen constitutes nearly 80% of Earth's atmosphere. The highest rate of carbon-14 production takes place at altitudes of 30,000 to 50,000 feet, and at higher geomagnetic latitudes, but the carbon-14 spreads evenly throughout the atmosphere and reacts with oxygen to form carbon dioxide. Carbon dioxide also permeates the oceans, dissolving in the water. Carbon-14 can also be produced in ice by fast neutrons causing spallation reactions in oxygen. Most of man-made chemicals are made of fossil fuels, like petroleum or coal, where the carbon-14 decayed. Presence of carbon-14 in the isotopic signature of a sample of material therefore indicates its possible biogenic origin.

References

Kamen, Martin D. Radiant Science, Dark Politics: A Memoir of the Nuclear Age, Forward by Edwin M. McMillan, Berkeley: University of California Press, 1985. Category:Isotopes ja:炭素14

Chemical symbol

:See also chemical formula. A chemical symbol is an abbreviation or short representation of the name of a chemical element. Natural elements all have symbols of one or two letters; some man-made elements have temporary symbols of three letters. Chemical symbols are listed in the periodic table and are used as shorthand and in chemical equations, e.g., :

2H_2 + O_2\ \rightarrow\ 2H_2O

. Because chemical symbols are often derived from the Latin or Greek name of the element, they may not bear much similar to the common English name, e.g., Na for sodium (Latin natrium) and Au for gold (Latin aurum). In China, each chemical element is assigned an ideograph as its symbol; most of them have been explicitly created for this purpose (see Chinese characters for chemical elements). Chemical symbols may also be changed to show if one particular isotope of an atom that is specified, as well as to show other attributes such as ionization and oxidation state of a chemical compound. For complete listings of the chemical elements and their symbols, see:
- List of elements by symbol
- List of elements by name
- List of elements by number
- Periodic table of the elements
- Category:Symbols ko:원소 기호 ja:元素記号

Electron

The electron is a fundamental subatomic particle which carries a negative electric charge.

Overview

Within an atom the electrons surround the nucleus of protons and neutrons in an electron configuration. The word electron was coined in 1894 and is derived from the term electric, whose ultimate origin is the Greek word 'ηλεκτρον, meaning amber. Electrons in motion constitute electric current which may be used by scientists and engineers to measure many physical properties. Electric current existing for a finite time gives rise to a movement of charge (electricity) that may be harnessed as a practical means to perform work. The variations in electric field generated by differing numbers of electrons and their configurations in atoms determine the chemical properties of the elements. These fields play a fundamental role in chemical bonds and chemistry.

Electrons in practice

Classification of electrons

The electron is one of a class of subatomic particles called leptons which are believed to be fundamental particles (that is, they cannot be broken down into smaller constituent parts). The word "particle" is somewhat misleading however, because quantum mechanics shows that electrons also behave like a wave, e.g. in the double-slit experiment; this is called wave-particle duality. The antiparticle of an electron is the positron, which has the same mass but positive rather than negative charge. The term negatron is sometimes used to refer to standard electrons so that the term electron may be used to describe both positrons and negatrons, as proposed by Carl D. Anderson. Under ordinary circumstances, however, electron refers to the negatively charged particle alone.

Properties and behavior of electrons

Electrons have a negative electric charge of −1.6 × 10⁻¹⁹ coulombs, and a mass of about 9.11 × 10⁻³¹ kg (0.51 MeV/c²), which is approximately ¹⁄₁₈₃₆ of the mass of the proton. These are commonly represented as e⁻. According to quantum mechanics, electrons can be represented by wavefunctions, from which the electron density can be determined. The exact momentum and position of an electron cannot be simultaneously determined. This is a limitation described by the Heisenberg uncertainty principle, which, in this instance, simply states that the more accurately we know a particle's position, the less accurately we can know its momentum and vice versa. The electron has spin ½, which implies it is a fermion, i.e., it follows the Fermi-Dirac statistics. While most electrons are found in atoms, others move independently in matter, or together as an electron beam in a vacuum. In some superconductors, electrons move in Cooper pairs, in which their motion is coupled to nearby matter via lattice vibrations called phonons. When electrons move, free of the nuclei of atoms, and there is a net flow, this flow is called electricity, or an electric current. A body has a static charge when the body has more or fewer electrons than are required to balance the positive charge of the nuclei. When there is an excess of electrons, the object is said to be negatively charged. When there are fewer electrons than protons, the object is said to be positively charged. When the number of electrons and the number of protons are equal, their charges cancel out and the object is said to be electrically neutral. A macroscopic body can acquire charge through rubbing, i.e. the phenomena of triboelectricity. Electrons and positrons can annihilate each other and produce a pair of photons. Conversely, high-energy photons may transform into an electron and a positron by a process called pair production. The electron is an elementary particle — that means that it has no substructure (at least, experiments have not found any so far, and there is good reason to believe that there is not any). Hence, it is usually described as point-like, i.e. with no spatial extension. However, if one gets very near an electron, one notices that its properties (charge and mass) seem to change. This is an effect common to all elementary particles: the particle influences the vacuum fluctuations in its vicinity, so that the properties one observes from far away are the sum of the bare properties and the vacuum effects (see renormalization). There is a physical constant called the classical electron radius, with a value of 2.8179 × 10⁻¹⁵ m. Note that this is the radius that one could infer from its charge if the physics were only described by the classical theory of electrodynamics and there were no quantum mechanics (hence, it is an outdated concept that nevertheless sometimes still proves useful in calculations). The speed of an electron in a vacuum can approach, but never reach c, the speed of light in a vacuum. This is due to an effect of special relativity. The effects of special relativity are based on a quantity known as gamma or the Lorentz factor. Gamma is a function of v, the velocity of the particle, and c. The following is the formula for gamma: :

\gamma = 1 / \sqrt

The energy necessary to accelerate a particle is gamma minus one times the rest mass. For example, the linear accelerator at Stanford can [http://www2.slac.stanford.edu/vvc/theory/relativity.html accelerate] an electron to roughly 51 GeV. This gives you a gamma of 100,000 given that the rest mass of an electron is 0.51 MeV/c² (the relativistic mass of this fast electron is 100 000 times its rest mass). Solving the equation above for the speed of the electron gives a speed of: :

(1-\frac   \gamma ^)c

= 0.999 999 999 95 c. (The formula applies for large γ.)

Electrons in the universe

It is believed that the number of electrons existing in the known universe is at least 10⁷⁹. This number amounts to a density of about one electron per cubic metre of space. Based on the classical electron radius and assuming a dense sphere packing, it can be calculated that the number of electrons that would fit in the observable universe is on the order of 10¹³⁰. Of course, this number is even less meaningful than the classical electron radius itself.

Electrons in industry

Electron beams are used in welding as well as lithography.

Electrons in the laboratory

Early experiments

The quantum or discrete nature of electron's charge was observed by Robert Millikan in the Oil-drop experiment of 1909.

Use of electrons in the laboratory

Electron microscopes are used to magnify details up to 500,000 times. Quantum effects of electrons are used in Scanning tunneling microscope to study features at the atomic scale.

Electrons in theory

In relativistic quantum mechanics, the electron is described by the Dirac Equation. Quantum electrodynamics (QED) models an electron as a charged particle surrounded a sea of interacting virtual particles, modifying the sea of virtual particles which makes up a vacuum. Although this theory involves difficult theoretical problems where calculations produce infinite terms, a practical (although mathematically dubious) method called renormalization was discovered whereby infinite terms can be cancelled to produce finite predictions about the electron. The correction of just over 0.1% to the predicted value of the electron's gyromagnetic ratio from exactly 2 (as predicted by Dirac's single particle model), and its extraordinarily precise agreement with the experimentally determined value, is viewed as one of the pinnacles of modern physics. There are now indications that string theory and its descendants may provide a model of the electron and other fundamental particles where the infinities in calculations do not appear, because the electron is no longer seen as a dimensionless point. At present, string theory is very much a 'work in progress' and lacks predictions analogous to those made by QED that can be experimentally verified. In the Standard Model of particle physics, it forms a doublet in SU(2) with the electron neutrino, as they interact through the weak interaction. The electron has two more massive partners, with the same charge but different masses: the muon and the tau lepton. The antimatter counterpart of the electron is its antiparticle, the positron. The positron has the same amount of electrical charge as the electron, except that the charge is positive. It has the same mass and spin as the electron. When an electron and a positron meet, they may annihilate each other, giving rise to two gamma-ray photons, each having an energy of 0.511 MeV (511 keV). See also Electron-positron annihilation. Electrons are also a key element in electromagnetism, an approximate theory that is adequate for macroscopic systems, and for classical modelling of microscopic systems.

History

The electron as a unit of charge in electrochemistry had been posited by G. Johnstone Stoney in 1874. In 1894, he also invented the word itself. The discovery that the electron was a subatomic particle was made in 1897 by J.J. Thomson at the Cavendish Laboratory at Cambridge University, while he was studying "cathode rays". Influenced by the work of James Clerk Maxwell, and the discovery of the X-ray, he deduced that cathode rays existed and were negatively charged "particles", which he called "corpuscles". He published his discovery in 1897. The periodic law states that the chemical properties of elements largely repeat themselves periodically and is the foundation of the periodic table of elements. The law itself was initially explained by the atomic mass of the elements. However, as there were anomalies in the periodic table, efforts were made to find a better explanation for it. In 1913, Henry Moseley introduced the concept of the atomic number and explained the periodic law with the number of protons each element has. In the same year, Niels Bohr showed that electrons are the actual foundation of the table. In 1916, Gilbert Newton Lewis and Irving Langmuir explained the chemical bonding of elements by electronic interactions.

External links

- [http://www.aip.org/history/electron/ The Discovery of the Electron] from the American Institute of Physics History Center
- [http://pdg.lbl.gov/ Particle Data Group]
- Stoney, G. Johnstone, "[http://dbhs.wvusd.k12.ca.us/webdocs/Chem-History/Stoney-1894.html Of the 'Electron,' or Atom of Electricity]". Philosophical Magazine. Series 5, Volume 38, p. 418-420 October 1894.
- Eric Weisstein's World of Physics: [http://scienceworld.wolfram.com/physics/Electron.html Electron]

References

-
-
- Brumfiel, G. (6 January 2005). Can electrons do the splits? In Nature, 433, 11. ko:전자 ja:電子 simple:Electron th:อิเล็กตรอน

Protium

Protium can be several things:
- In chemistry, protium is the most common isotope of the element hydrogen; that has one proton and no neutrons. :See Hydrogen atom.
- In botany, Protium is a genus of chiefly tropical American trees in the family Burseraceae , having fragrant wood and yielding gum "elemi".
- In computer language, "Protium" is a universal, symbolic programming language system, based on a systematic a priori analysis of the tasks required for computation. It is polymorphic in type with considerable character flexibility. It is pasigraphic after the fashion of John Wilkins and Charles K. Bliss, and does not favour one host natural language over another. :See [http://hopl.murdoch.edu.au The Encyclopedia of Programming Languages]
- Protium is also the name of an IRC network. :See [http://www.protium.org The Protium IRC Network] ms:Protium

Molecules

A molecule is the smallest particle of a pure chemical substance that still retains its chemical composition and properties. The science of molecules is called molecular chemistry or molecular physics, depending on the focus. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, however, this distinction is vague. According to the strict definition, molecules can consist of one atom (as in noble gases) or more atoms bonded together. The concept of monatomic (single-atom) molecule is used almost exclusively in the kinetic theory of gases. In molecular sciences, a molecule consists of a stable system (bound state) comprising two or more atoms. The term unstable molecule is used for very reactive species, i.e., short-lived assemblies (resonances) of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, Van der Waals complexes, or systems of colliding atoms as in Bose-Einstein condensates. A peculiar use of the term molecular is as a synonym to covalent, which arises from the fact that, unlike molecular covalent compounds, ionic compounds do not yield well-defined smallest particles that would be consistent with the definition above. No typical "smallest particle" can be defined for covalent crystals, or network solids, which are composed of repeating unit cells that extend indefinitely either in a plane (such as in graphite) or three-dimensionally (such as in diamond). Although the concept of molecules was first introduced in 1811 by Avogadro, and was accepted by many chemists as a result of Dalton's laws of Definite and Multiple Proportions (1803-1808), with notable exceptions (Boltzmann, Maxwell, Gibbs), the existence of molecules as anything other than convenient mathematical constructs was still an open debate in the physics community until the work of Perrin (1911), and was strenuously resisted by early positvists such as Mach. The modern theory of molecules makes great use of the many numerical techniques offered by computational chemistry. Dozens of molecules have now been identified in interstellar space by microwave spectroscopy. microwave spectroscopy (right) representations of the terpenoid, atisane. In the 3D model on the left, carbon atoms are represented by gray spheres; white spheres represent the hydrogen atoms and the cylinders represent the bonds. The model is enveloped in a "mesh" representation of the molecular surface, colored by areas of positive (red) and negative (blue) electric charge. In the 3D model (center), the light-blue spheres represent carbon atoms, the white spheres are hydrogen atoms, and the cylinders in between the atoms correspond to single bonds.]]

Chemical bond

:See main article chemical bond In a molecule, the atoms are joined by shared pairs of electrons in a chemical bond. It may consist of atoms of the same chemical element, as with oxygen (O₂), or of different elements, as with water (H₂O).

Size

Most molecules are much too small to be seen with the naked eye, but there are exceptions. DNA, a macromolecule, can reach macroscopic sizes. The smallest molecule is the hydrogen molecule. The interatomic distance is 0.15 nanometres (1.5 Å). But the size of its electron cloud is difficult to define precisely. Under standard conditions molecules have a dimension of a few to a few dozen Å.

Empirical formula

:See main article empirical formula The empirical formula of a molecule is the simplest integer ratio of the chemical elements that constitute the compound. For example, in their pure forms, water is always composed of a 2:1 ratio of hydrogen to oxygen, and ethyl alcohol or ethanol is always composed of carbon, hydrogen, and oxygen in a 2:6:1 ratio. However, this does not determine the kind of molecule uniquely - dimethyl ether has the same ratio as ethanol, for instance. Molecules with the same atoms in different arrangements are called isomers. The empirical formula is often the same as the molecular formula but not always. For example the molecule acetylene has molecular formula C2H2, but the simplest integer ratio of elements is CH.

Chemical formula

:See main article chemical formula The chemical formula reflects the exact number of atoms that compose a molecule. The molecular mass can be calculated from the chemical formula and is expressed in conventional units equal to 1/12 from the mass of a ¹²C isotope atom. For network solids, the term formula unit is used in stoichiometric calculations.

Molecular geometry

:See main article molecular geometry Molecules have fixed equilibrium geometries—bond lengths and angles—. A pure substance is composed of molecules with the same geometrical structure. The chemical formula and the structure of a molecule are the two important factors that determine its properties, particularly its reactivity. Isomers share a chemical formula but normally have very different properties because of their different structures. Stereoisomers, a particular type of isomers, may have very similar physico-chemical properties and at the same time very different biochemical activities.

Molecular spectroscopy

:See main article spectroscopy Molecular spectroscopy is the study of the response (spectrum) of a molecule to a signal of known energy (or frequency, according to Planck's formula). This signal is usually an electromagnetic wave or a beam of electrons, but new molecular spectroscopies, such as the positron spectroscopy, are under development. The molecular response can be signal absorption (absorption spectroscopy), emission of another signal (emission spectroscopy), fragmentation, or a change in its chemical nature. Spectroscopy is recognized as the most powerful tool in the investigation of the microscopic properties of molecules, and, in particular, their energy levels. Nowadays, in order to extract the maximum microscopic information from the experimental results, spectroscopical studies are very often coupled with computational chemical investigations. The theoretical background of spectroscopy is the scattering theory.

Isotopologue

Isotopologues (not to be confused with isotopomers) are chemical species that differ only in the isotopic composition of their molecules or ions. An example is water, where three of its hydrogen-related isotopologues are: HOH, HOD and DOD. Simply, the isotopologue of a chemical species has at least one atom with a different number of neutrons. Category:Chemical substances

Photon

In physics, the photon (from Greek φως "phos", meaning light) is a quantum of the electromagnetic field, for instance light. The term photon was coined by Gilbert Lewis in 1926. Gilbert Lewis The photon is one of the elementary particles. Its interactions with electrons and atomic nuclei account for a great many of the features of matter, such as the existence and stability of atoms, molecules, and solids. These interactions are studied in quantum electrodynamics (QED), which is the oldest part of the Standard Model of particle physics. In some respects a photon acts as a particle, for instance when registered by the light sensitive device in a camera. In other respects, a photon acts like a wave, as when passing through the optics in a camera. According to the so-called wave-particle duality in quantum physics, it is natural for the photon to display either aspect of its nature, according to the circumstances. Normally, light is formed from a large number of photons, with the intensity related to the number of them. At low intensity, it requires very sensitive instruments, used in astronomy or spectroscopy, for instance, to detect the individual photons.

Symbol

A photon is usually given the symbol

\gamma

(gamma), although in nuclear physics this symbol refers to a very high-energy photon (a gamma ray).

Properties

Photons are commonly associated with visible light, but this is actually only a very limited part of the electromagnetic spectrum. All electromagnetic radiation is quantized as photons: that is, the smallest amount of electromagnetic radiation that can exist is one photon, whatever its wavelength, frequency, energy, or momentum. Photons are fundamental particles. They can be created and destroyed when interacting with other particles, but are not known to decay on their own. Unlike most particles, photons have no detectable intrinsic mass, or "rest mass" (as opposed to relativistic mass). Photons are always moving at the speed of light with respect to all observers. Although they lack mass, photons have both energy and momentum proportional to their frequency (or inversely proportional to their wavelength). This momentum can be transferred when a photon collides with matter. The force due to a large number of photons falling on a surface is known as radiation pressure, which may be used for propulsion with a solar sail. Photons are deflected by a gravitational field twice as much as Newtonian mechanics predicts for a mass traveling at the speed of light with the same momentum as the photon. This observation is commonly cited as evidence supporting Einstein's theory of gravitation, general relativity. In general relativity, photons always travel in a "straight" line, after taking into account the curvature of spacetime. (In curved space, such lines are called geodesics).

Creation

Photons are produced by atoms when a bound electron moves from one orbital to another orbital with less (more negative) energy. Photons can also be emitted by an unstable nucleus when it undergoes some types of nuclear decay. Furthermore, photons are produced whenever charged particles are accelerated. Atoms continuously emit photons due to their collisions with each other. The wavelength distribution of these photons thus are related to their absolute temperature. The Maxwell-Boltzmann distribution provides the probability of a photon being a certain wavelength when emitted by a collection of atoms at a given temperature. The spectrum of such photons is normally peaked in the range between microwave and infrared, but hot objects (such as the surface of the Sun) will emit visible light as well. As temperature is further increased, some photons will reach even higher frequencies, such as ultraviolet and X-ray. Radio, television, radar and other types of transmitters used for telecommunication and remote sensing routinely create a wide variety of low-energy photons by the oscillation of electric fields in conductors. Magnetrons emit coherent photons used in household microwave ovens. Klystron tubes are used when microwave emissions must be more finely controlled. Masers and lasers create monochromatic photons by stimulated emission. More energetic photons can be created by nuclear transitions, particle-antiparticle annihilation, and in high-energy particle collisions.

Spin

Photons have spin 1, and they are therefore classified as bosons. Photons mediate the electromagnetic interaction; they are the gauge bosons of quantum electrodynamics (QED), which is a U(1) gauge theory. In general, a boson with spin 1 should have three distinct spin projections (−1, 0 and +1). However, the zero projection would require a frame where the photon is at rest. Since the photon's mass is zero, it always travels at the speed of light, and such a frame does not exist. Thus photons in empty space show only two spin projections, corresponding to the right- and left-handed circular polarizations of classical electromagnetic waves. The more familiar linear polarization is formed by a mixture of right- and left-circularly polarized photons.

Quantum state

Visible light from ordinary sources (like the Sun or a lamp) is a mixture of many photons of different wavelengths. One sees this in the frequency spectrum, for instance by passing the light through a prism. In so-called "mixed states", which these sources tend to produce, light can consist of photons in thermal equilibrium (so-called black-body radiation). Here they in many ways resemble a gas of particles. For example, they exert pressure, known as radiation pressure. On the other hand, an assembly of photons can also exist in much more well-organized coherent states, such as in the light emitted by an ideal laser. The high degree of precision obtained with laser instruments is due to this organization. The quantum state of a photon assembly, like that of other quantum particles, is the so-called Fock state denoted

|n\rangle

, meaning

n

photons in one of the distinct "modes" of the electromagnetic field. If the field is multimode (involves several different wavelength photons), its quantum state is a tensor product of photon states, for example: :

|n_\rangle\otimes|n_\rangle\otimes\dots\otimes|n_\rangle\dots

Here

k_i

denote the possible modes, and

n_

the number of photons in each mode

Molecular absorption

A typical molecule,

M

, has many different energy levels. When a molecule absorbs a photon, its energy is increased by an amount equal to the energy of the photon. The molecule then enters an excited state,

M^ -  \,

. :

M + \gamma \to M^ -  \,

Photons in vacuo

In empty space (vacuum) all photons move at the speed of light, c, defined as 299,792,458 meters per second, or approximately 3×10⁸ m s⁻¹. The meter is defined as the distance traveled by light in vacuum in 1/299,792,458 of a second, so the speed of light does not suffer any experimental uncertainty, unlike the meter or the second, which rely on the second being defined by means of a very accurate clock. According to one principle of Einstein's special relativity, all observations of the speed of light in vacuo are same in all directions to any observer in an inertial frame of reference. This principle is generally accepted in

Isotopes

Variation in properties between isotopes

Occurrence in nature

Applications of isotopes

Use of chemical properties

Use of nuclear properties

See also

External links

Chemical element

Chemistry terminology

Description

Nomenclature

Chemical symbols

Specific chemical elements

General chemical symbols

Nonelement symbols

See also

External links

Chemical information

Nucleus

Etymology

Proton

History

Technological applications

Antiproton

See also

External links

Neutron

Properties

Neutron Interactions

Neutron Detection

Neutron Uses

Neutron Sources

Discovery

Current developments

Antineutron

See also

Fields concerning neutrons

Types of neutrons

Objects containing neutrons

Neutron sources

Processes involving neutrons

Fluorine

Notable characteristics

Applications

History

Precautions

Preparation

Compounds

References

External links

Carbon-12

History

See also

Carbon-14

References

Chemical symbol

Electron

Overview

Electrons in practice

Classification of electrons

Properties and behavior of electrons

Electrons in the universe

Electrons in industry

Electrons in the laboratory

Early experiments

Use of electrons in the laboratory

Electrons in theory

History

See also

External links

References

Protium

Molecules

Chemical bond

Size

Empirical formula

Chemical formula

Molecular geometry

Molecular spectroscopy