A crystal or crystalline solid is a solid material whose constituents are arranged in a ordered microscopic structure, forming a crystal lattice that extends in all directions. In addition, macroscopic single crystals are identifiable by their geometrical shape, consisting of flat faces with specific, characteristic orientations; the scientific study of crystals and crystal formation is known as crystallography. The process of crystal formation via mechanisms of crystal growth is called crystallization or solidification; the word crystal derives from the Ancient Greek word κρύσταλλος, meaning both "ice" and "rock crystal", from κρύος, "icy cold, frost". Examples of large crystals include snowflakes and table salt. Most inorganic solids are not crystals but polycrystals, i.e. many microscopic crystals fused together into a single solid. Examples of polycrystals include most metals, rocks and ice. A third category of solids is amorphous solids, where the atoms have no periodic structure whatsoever.
Examples of amorphous solids include glass and many plastics. Despite the name, lead crystal, crystal glass, related products are not crystals, but rather types of glass, i.e. amorphous solids. Crystals are used in pseudoscientific practices such as crystal therapy, along with gemstones, are sometimes associated with spellwork in Wiccan beliefs and related religious movements; the scientific definition of a "crystal" is based on the microscopic arrangement of atoms inside it, called the crystal structure. A crystal is a solid where the atoms form a periodic arrangement.. Not all solids are crystals. For example, when liquid water starts freezing, the phase change begins with small ice crystals that grow until they fuse, forming a polycrystalline structure. In the final block of ice, each of the small crystals is a true crystal with a periodic arrangement of atoms, but the whole polycrystal does not have a periodic arrangement of atoms, because the periodic pattern is broken at the grain boundaries.
Most macroscopic inorganic solids are polycrystalline, including all metals, ice, etc. Solids that are neither crystalline nor polycrystalline, such as glass, are called amorphous solids called glassy, vitreous, or noncrystalline; these have no periodic order microscopically. There are distinct differences between crystalline solids and amorphous solids: most notably, the process of forming a glass does not release the latent heat of fusion, but forming a crystal does. A crystal structure is characterized by its unit cell, a small imaginary box containing one or more atoms in a specific spatial arrangement; the unit cells are stacked in three-dimensional space to form the crystal. The symmetry of a crystal is constrained by the requirement that the unit cells stack with no gaps. There are 219 possible crystal symmetries, called crystallographic space groups; these are grouped into 7 crystal systems, such as hexagonal crystal system. Crystals are recognized by their shape, consisting of flat faces with sharp angles.
These shape characteristics are not necessary for a crystal—a crystal is scientifically defined by its microscopic atomic arrangement, not its macroscopic shape—but the characteristic macroscopic shape is present and easy to see. Euhedral crystals are those with well-formed flat faces. Anhedral crystals do not because the crystal is one grain in a polycrystalline solid; the flat faces of a euhedral crystal are oriented in a specific way relative to the underlying atomic arrangement of the crystal: they are planes of low Miller index. This occurs; as a crystal grows, new atoms attach to the rougher and less stable parts of the surface, but less to the flat, stable surfaces. Therefore, the flat surfaces tend to grow larger and smoother, until the whole crystal surface consists of these plane surfaces. One of the oldest techniques in the science of crystallography consists of measuring the three-dimensional orientations of the faces of a crystal, using them to infer the underlying crystal symmetry.
A crystal's habit is its visible external shape. This is determined by the crystal structure, the specific crystal chemistry and bonding, the conditions under which the crystal formed. By volume and weight, the largest concentrations of crystals in the Earth are part of its solid bedrock. Crystals found in rocks range in size from a fraction of a millimetre to several centimetres across, although exceptionally large crystals are found; as of 1999, the world's largest known occurring crystal is a crystal of beryl from Malakialina, Madagascar, 18 m long and 3.5 m in diameter, weighing 380,000 kg. Some crystals have formed by magmatic and metamorphic processes, giving origin to large masses of crystalline rock; the vast majority of igneous rocks are formed from molten magma and the degree of crystallization depends on the conditions under which they solidified. Such rocks as granite, which have cooled slowly and under great pressures, have crystallized.
Bromine is a chemical element with symbol Br and atomic number 35. It is the third-lightest halogen, is a fuming red-brown liquid at room temperature that evaporates to form a coloured gas, its properties are thus intermediate between those of iodine. Isolated independently by two chemists, Carl Jacob Löwig and Antoine Jérôme Balard, its name was derived from the Ancient Greek βρῶμος, referencing its sharp and disagreeable smell. Elemental bromine is reactive and thus does not occur free in nature, but in colourless soluble crystalline mineral halide salts, analogous to table salt. While it is rather rare in the Earth's crust, the high solubility of the bromide ion has caused its accumulation in the oceans. Commercially the element is extracted from brine pools in the United States and China; the mass of bromine in the oceans is about one three-hundredth. At high temperatures, organobromine compounds dissociate to yield free bromine atoms, a process that stops free radical chemical chain reactions.
This effect makes organobromine compounds useful as fire retardants, more than half the bromine produced worldwide each year is put to this purpose. The same property causes ultraviolet sunlight to dissociate volatile organobromine compounds in the atmosphere to yield free bromine atoms, causing ozone depletion; as a result, many organobromide compounds—such as the pesticide methyl bromide—are no longer used. Bromine compounds are still used in well drilling fluids, in photographic film, as an intermediate in the manufacture of organic chemicals. Large amounts of bromide salts are toxic from the action of soluble bromide ion, causing bromism. However, a clear biological role for bromide ion and hypobromous acid has been elucidated, it now appears that bromine is an essential trace element in humans; the role of biological organobromine compounds in sea life such as algae has been known for much longer. As a pharmaceutical, the simple bromide ion has inhibitory effects on the central nervous system, bromide salts were once a major medical sedative, before replacement by shorter-acting drugs.
They retain niche uses as antiepileptics. Bromine was discovered independently by two chemists, Carl Jacob Löwig and Antoine Balard, in 1825 and 1826, respectively. Löwig isolated bromine from a mineral water spring from his hometown Bad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with diethyl ether. After evaporation of the ether a brown liquid remained. With this liquid as a sample of his work he applied for a position in the laboratory of Leopold Gmelin in Heidelberg; the publication of the results was delayed and Balard published his results first. Balard found bromine chemicals in the ash of seaweed from the salt marshes of Montpellier; the seaweed was used to produce iodine, but contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine; the properties of the resulting substance were intermediate between those of iodine. After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique.
In his publication, Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Brôme derives from the Greek βρωμος. Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors. Bromine was not produced in large quantities until 1858, when the discovery of salt deposits in Stassfurt enabled its production as a by-product of potash. Apart from some minor medical applications, the first commercial use was the daguerreotype. In 1840, bromine was discovered to have some advantages over the used iodine vapor to create the light sensitive silver halide layer in daguerreotypy. Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, but were superseded by chloral hydrate and by the barbiturates. In the early years of the First World War, bromine compounds such as xylyl bromide were used as poison gas. Bromine is the third halogen.
Its properties are thus similar to those of fluorine and iodine, tend to be intermediate between those of the two neighbouring halogens and iodine. Bromine has the electron configuration 3d104s24p5, with the seven electrons in the fourth and outermost shell acting as its valence electrons. Like all halogens, it is thus one electron short of a full octet, is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell. Corresponding to periodic trends, it is intermediate in electronegativity between chlorine and iodine, is less reactive than chlorine and more reactive than iodine, it is a weaker oxidising agent than chlorine, but a stronger one than iodine. Conversely, the bromide ion is a weaker reducing agent than iodide, but a stronger one than chloride; these similarities led to chlorine and iodine together being classified as one of the original triads of Johann Wolfgang Döbereiner, whose work foreshadowed the periodic law for chemical elements. It is intermediate in atomic radius between chlorine and iodi
Rubidium is a chemical element with symbol Rb and atomic number 37. Rubidium is a soft, silvery-white metallic element of the alkali metal group, with a standard atomic weight of 85.4678. Elemental rubidium is reactive, with properties similar to those of other alkali metals, including rapid oxidation in air. On Earth, natural rubidium comprises two isotopes: 72% is the stable isotope, 85Rb. German chemists Robert Bunsen and Gustav Kirchhoff discovered rubidium in 1861 by the newly developed technique, flame spectroscopy; the name comes from the Latin word rubidus, meaning the color of its emission spectrum. Rubidium's compounds have various chemical and electronic applications. Rubidium metal is vaporized and has a convenient spectral absorption range, making it a frequent target for laser manipulation of atoms. Rubidium is not a known nutrient for any living organisms. However, rubidium ions have the same charge as potassium ions, are taken up and treated by animal cells in similar ways. Rubidium is a soft, silvery-white metal.
It is the second most electropositive of the stable alkali metals and melts at a temperature of 39.3 °C. Like other alkali metals, rubidium metal reacts violently with water; as with potassium and caesium, this reaction is vigorous enough to ignite the hydrogen gas it produces. Rubidium has been reported to ignite spontaneously in air, it forms amalgams with mercury and alloys with gold, caesium and potassium, but not lithium. Rubidium has a low ionization energy of only 406 kJ/mol. Rubidium and potassium show a similar purple color in the flame test, distinguishing the two elements requires more sophisticated analysis, such as spectroscopy. Rubidium chloride is the most used rubidium compound: among several other chlorides, it is used to induce living cells to take up DNA. Other common rubidium compounds are the corrosive rubidium hydroxide, the starting material for most rubidium-based chemical processes. Rubidium silver iodide has the highest room temperature conductivity of any known ionic crystal, a property exploited in thin film batteries and other applications.
Rubidium forms a number of oxides when exposed to air, including rubidium monoxide, Rb6O, Rb9O2. Rubidium forms salts with halides, producing rubidium fluoride, rubidium chloride, rubidium bromide, rubidium iodide. Although rubidium is monoisotopic, rubidium in the Earth's crust is composed of two isotopes: the stable 85Rb and the radioactive 87Rb. Natural rubidium is radioactive, with specific activity of about 670 Bq/g, enough to expose a photographic film in 110 days. Twenty four additional rubidium isotopes have been synthesized with half-lives of less than 3 months. Rubidium-87 has a half-life of 48.8×109 years, more than three times the age of the universe of ×109 years, making it a primordial nuclide. It substitutes for potassium in minerals, is therefore widespread. Rb has been used extensively in dating rocks. During fractional crystallization, Sr tends to concentrate in plagioclase, leaving Rb in the liquid phase. Hence, the Rb/Sr ratio in residual magma may increase over time, the progressing differentiation results in rocks with elevated Rb/Sr ratios.
The highest ratios occur in pegmatites. If the initial amount of Sr is known or can be extrapolated the age can be determined by measurement of the Rb and Sr concentrations and of the 87Sr/86Sr ratio; the dates indicate the true age of the minerals only if the rocks have not been subsequently altered. Rubidium-82, one of the element's non-natural isotopes, is produced by electron-capture decay of strontium-82 with a half-life of 25.36 days. With a half-life of 76 seconds, rubidium-82 decays by positron emission to stable krypton-82. Rubidium is the twenty-third most abundant element in the Earth's crust as abundant as zinc and rather more common than copper, it occurs in the minerals leucite, pollucite and zinnwaldite, which contain as much as 1% rubidium oxide. Lepidolite contains between 0.3% and 3.5% rubidium, is the commercial source of the element. Some potassium minerals and potassium chlorides contain the element in commercially significant quantities. Seawater contains an average of 125 µg/L of rubidium compared to the much higher value for potassium of 408 mg/L and the much lower value of 0.3 µg/L for caesium.
Because of its large ionic radius, rubidium is one of the "incompatible elements." During magma crystallization, rubidium is concentrated together with its heavier analogue caesium in the liquid phase and crystallizes last. Therefore, the largest deposits of rubidium and caesium are zone pegmatite ore bodies formed by this enrichment process; because rubidium substitutes for potassium in the crystallization of magma, the enrichment is far less effective than that of caesium. Zone pegmatite ore bodies containing mineable quantities of caesium as pollucite or the lithium minerals lepidolite are a source for rubidium as a by-product. Two notable sources of rubidium are th
Chlorine is a chemical element with symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are intermediate between them. Chlorine is a yellow-green gas at room temperature, it is an reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the Pauling scale, behind only oxygen and fluorine. The most common compound of chlorine, sodium chloride, has been known since ancient times. Around 1630, chlorine gas was first synthesised in a chemical reaction, but not recognised as a fundamentally important substance. Carl Wilhelm Scheele wrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρός, translit. Khlôros, lit.'pale green' based on its colour.
Because of its great reactivity, all chlorine in the Earth's crust is in the form of ionic chloride compounds, which includes table salt. It is the second-most abundant halogen and twenty-first most abundant chemical element in Earth's crust; these crustal deposits are dwarfed by the huge reserves of chloride in seawater. Elemental chlorine is commercially produced from brine by electrolysis; the high oxidising potential of elemental chlorine led to the development of commercial bleaches and disinfectants, a reagent for many processes in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, many intermediates for the production of plastics and other end products which do not contain the element; as a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary. Elemental chlorine at high concentrations is dangerous and poisonous for all living organisms, was used in World War I as the first gaseous chemical warfare agent.
In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils as part of the immune response against bacteria; the most common compound of chlorine, sodium chloride, has been known since ancient times. Its importance in food was well known in classical antiquity and was sometimes used as payment for services for Roman generals and military tribunes. Elemental chlorine was first isolated around 1200 with the discovery of aqua regia and its ability to dissolve gold, since chlorine gas is one of the products of this reaction: it was however not recognised as a new substance. Around 1630, chlorine was recognized as a gas by the Flemish chemist and physician Jan Baptist van Helmont.
The element was first studied in detail in 1774 by Swedish chemist Carl Wilhelm Scheele, he is credited with the discovery. Scheele produced chlorine by reacting MnO2 with HCl: 4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow-green color, the smell similar to aqua regia, he called it "dephlogisticated muriatic acid air" since it is a gas and it came from hydrochloric acid. He failed to establish chlorine as an element. Common chemical theory at that time held that an acid is a compound that contains oxygen, so a number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and the yet undiscovered element, muriaticum. In 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid air by reacting it with charcoal to release the free element muriaticum, they did not succeed and published a report in which they considered the possibility that dephlogisticated muriatic acid air is an element, but were not convinced.
In 1810, Sir Humphry Davy tried the same experiment again, concluded that the substance was an element, not a compound. He announced his results to the Royal Society on 15 November that year. At that time, he named this new element "chlorine", from the Greek word χλωρος, meaning green-yellow; the name "halogen", meaning "salt producer", was used for chlorine in 1811 by Johann Salomo Christoph Schweigger. This term was used as a generic term to describe all the elements in the chlorine family, after a suggestion by Jöns Jakob Berzelius in 1826. In 1823, Michael Faraday liquefied chlorine for the first time, demonstrated that what was known as "solid chlorine" had a structure of chlorine hydrate. Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785. Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel, by passing chlorine gas through a solution of sodium carbonate; the resulting liqu
Infrared radiation, sometimes called infrared light, is electromagnetic radiation with longer wavelengths than those of visible light, is therefore invisible to the human eye, although IR at wavelengths up to 1050 nanometers s from specially pulsed lasers can be seen by humans under certain conditions. IR wavelengths extend from the nominal red edge of the visible spectrum at 700 nanometers, to 1 millimeter. Most of the thermal radiation emitted by objects near room temperature is infrared; as with all EMR, IR carries radiant energy and behaves both like a wave and like its quantum particle, the photon. Infrared radiation was discovered in 1800 by astronomer Sir William Herschel, who discovered a type of invisible radiation in the spectrum lower in energy than red light, by means of its effect on a thermometer. More than half of the total energy from the Sun was found to arrive on Earth in the form of infrared; the balance between absorbed and emitted infrared radiation has a critical effect on Earth's climate.
Infrared radiation is emitted or absorbed by molecules when they change their rotational-vibrational movements. It excites vibrational modes in a molecule through a change in the dipole moment, making it a useful frequency range for study of these energy states for molecules of the proper symmetry. Infrared spectroscopy examines transmission of photons in the infrared range. Infrared radiation is used in industrial, military, law enforcement, medical applications. Night-vision devices using active near-infrared illumination allow people or animals to be observed without the observer being detected. Infrared astronomy uses sensor-equipped telescopes to penetrate dusty regions of space such as molecular clouds, detect objects such as planets, to view red-shifted objects from the early days of the universe. Infrared thermal-imaging cameras are used to detect heat loss in insulated systems, to observe changing blood flow in the skin, to detect overheating of electrical apparatus. Extensive uses for military and civilian applications include target acquisition, night vision and tracking.
Humans at normal body temperature radiate chiefly at wavelengths around 10 μm. Non-military uses include thermal efficiency analysis, environmental monitoring, industrial facility inspections, detection of grow-ops, remote temperature sensing, short-range wireless communication and weather forecasting. Infrared radiation extends from the nominal red edge of the visible spectrum at 700 nanometers to 1 millimeter; this range of wavelengths corresponds to a frequency range of 430 THz down to 300 GHz. Below infrared is the microwave portion of the electromagnetic spectrum. Sunlight, at an effective temperature of 5,780 kelvins, is composed of near-thermal-spectrum radiation, more than half infrared. At zenith, sunlight provides an irradiance of just over 1 kilowatt per square meter at sea level. Of this energy, 527 watts is infrared radiation, 445 watts is visible light, 32 watts is ultraviolet radiation. Nearly all the infrared radiation in sunlight is shorter than 4 micrometers. On the surface of Earth, at far lower temperatures than the surface of the Sun, some thermal radiation consists of infrared in the mid-infrared region, much longer than in sunlight.
However, black body or thermal radiation is continuous: it gives off radiation at all wavelengths. Of these natural thermal radiation processes, only lightning and natural fires are hot enough to produce much visible energy, fires produce far more infrared than visible-light energy. In general, objects emit infrared radiation across a spectrum of wavelengths, but sometimes only a limited region of the spectrum is of interest because sensors collect radiation only within a specific bandwidth. Thermal infrared radiation has a maximum emission wavelength, inversely proportional to the absolute temperature of object, in accordance with Wien's displacement law. Therefore, the infrared band is subdivided into smaller sections. A used sub-division scheme is: NIR and SWIR is sometimes called "reflected infrared", whereas MWIR and LWIR is sometimes referred to as "thermal infrared". Due to the nature of the blackbody radiation curves, typical "hot" objects, such as exhaust pipes appear brighter in the MW compared to the same object viewed in the LW.
The International Commission on Illumination recommended the division of infrared radiation into the following three bands: ISO 20473 specifies the following scheme: Astronomers divide the infrared spectrum as follows: These divisions are not precise and can vary depending on the publication. The three regions are used for observation of different temperature ranges, hence different environments in space; the most common photometric system used in astronomy allocates capital letters to different spectral regions according to filters used. These letters are understood in reference to atmospheric windows and appear, for instance, in the titles of many papers. A third scheme divides up the band based on the response of various detectors: Near-infrared: from 0.7 to 1.0 µm. Short-wave infrared: 1.0 to 3 µm. InGaAs covers to about 1.8 µm. Mid-wave infrared: 3 to 5 µm (defined by the atmospheric window and covered by indium antimonide and mercury cadmium telluride and by lead
Lithium is a chemical element with symbol Li and atomic number 3. It is a silvery-white alkali metal. Under standard conditions, it is the lightest solid element. Like all alkali metals, lithium is reactive and flammable, is stored in mineral oil; when cut, it exhibits a metallic luster, but moist air corrodes it to a dull silvery gray black tarnish. It never occurs in nature, but only in compounds, such as pegmatitic minerals, which were once the main source of lithium. Due to its solubility as an ion, it is present in ocean water and is obtained from brines. Lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride; the nucleus of the lithium atom verges on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative nuclear instability, lithium is less common in the solar system than 25 of the first 32 chemical elements though its nuclei are light: it is an exception to the trend that heavier nuclei are less common.
For related reasons, lithium has important uses in nuclear physics. The transmutation of lithium atoms to helium in 1932 was the first man-made nuclear reaction, lithium deuteride serves as a fusion fuel in staged thermonuclear weapons. Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, lithium grease lubricants, flux additives for iron and aluminium production, lithium batteries, lithium-ion batteries; these uses consume more than three quarters of lithium production. Lithium is present in biological systems in trace amounts. Lithium salts have proven to be useful as a mood-stabilizing drug in the treatment of bipolar disorder in humans. Like the other alkali metals, lithium has a single valence electron, given up to form a cation; because of this, lithium is a good conductor of heat and electricity as well as a reactive element, though it is the least reactive of the alkali metals. Lithium's low reactivity is due to the proximity of its valence electron to its nucleus.
However, molten lithium is more reactive than its solid form. Lithium metal is soft enough to be cut with a knife; when cut, it possesses a silvery-white color that changes to gray as it oxidizes to lithium oxide. While it has one of the lowest melting points among all metals, it has the highest melting and boiling points of the alkali metals. Lithium has a low density, comparable with pine wood, it is the least dense of all elements. Furthermore, apart from helium and hydrogen, it is less dense than any liquid element, being only two thirds as dense as liquid nitrogen. Lithium can float on the lightest hydrocarbon oils and is one of only three metals that can float on water, the other two being sodium and potassium. Lithium's coefficient of thermal expansion is twice that of aluminium and four times that of iron. Lithium is superconductive below 400 μK at standard pressure and at higher temperatures at high pressures. At temperatures below 70 K, like sodium, undergoes diffusionless phase change transformations.
At 4.2 K it has a rhombohedral crystal system. At liquid-helium temperatures the rhombohedral structure is prevalent. Multiple allotropic forms have been identified for lithium at high pressures. Lithium has a mass specific heat capacity of 3.58 kilojoules per kilogram-kelvin, the highest of all solids. Because of this, lithium metal is used in coolants for heat transfer applications. Lithium reacts with water but with noticeably less vigor than other alkali metals; the reaction forms hydrogen lithium hydroxide in aqueous solution. Because of its reactivity with water, lithium is stored in a hydrocarbon sealant petroleum jelly. Though the heavier alkali metals can be stored in more dense substances, such as mineral oil, lithium is not dense enough to be submerged in these liquids. In moist air, lithium tarnishes to form a black coating of lithium hydroxide, lithium nitride and lithium carbonate; when placed over a flame, lithium compounds give off a striking crimson color, but when it burns the flame becomes a brilliant silver.
Lithium will burn in oxygen when exposed to water or water vapors. Lithium is flammable, it is explosive when exposed to air and to water, though less so than the other alkali metals; the lithium-water reaction at normal temperatures is brisk but nonviolent because the hydrogen produced does not ignite on its own. As with all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers. Lithium is one of the few metals. Lithium has a diagonal relationship with an element of similar atomic and ionic radius. Chemical resemblances between the two metals include the formation of a nitride by reaction with N2, the formation of an oxide and peroxide when burnt in O2, salts with similar solubilities, thermal instability of the carbonates and nitrides; the metal reacts with hy
Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a toxic pale yellow diatomic gas at standard conditions; as the most electronegative element, it is reactive, as it reacts with all other elements, except for helium and neon. Among the elements, fluorine ranks 24th in universal 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking.
The rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Molecules containing a Carbon–fluorine bond have high chemical and thermal stability. Pharmaceuticals such as atorvastatin and fluoxetine contain C-F bonds, the fluoride ion inhibits dental cavities, so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$15 billion a year. Fluorocarbon gases are greenhouse gases with global-warming potentials 100 to 20,000 times that of carbon dioxide. Organofluorine compounds persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals. Fluorine atoms have nine electrons, one fewer than neon, electron configuration 1s22s22p5: two electrons in a filled inner shell and seven in an outer shell requiring one more to be filled; the outer electrons are ineffective at nuclear shielding, experience a high effective nuclear charge of 9 − 2 = 7.
Fluorine's first ionization energy is third-highest among all elements, behind helium and neon, which complicates the removal of electrons from neutral fluorine atoms. It has a high electron affinity, second only to chlorine, tends to capture an electron to become isoelectronic with the noble gas neon. Fluorine atoms have a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen and neon; the bond energy of difluorine is much lower than that of either Cl2 or Br2 and similar to the cleaved peroxide bond. Conversely, bonds to other atoms are strong because of fluorine's high electronegativity. Unreactive substances like powdered steel, glass fragments, asbestos fibers react with cold fluorine gas. Reactions of elemental fluorine with metals require varying conditions. Alkali metals cause; some solid nonmetals react vigorously in liquid air temperature fluorine. Hydrogen sulfide and sulfur dioxide combine with fluorine, the latter sometimes explosively. Hydrogen, like some of the alkali metals, reacts explosively with fluorine.
Carbon, as lamp black, reacts at room temperature to yield fluoromethane. Graphite combines with fluorine above 400 °C to produce non-stoichiometric carbon monofluoride. Carbon dioxide and carbon monoxide react at or just above room temperature, whereas paraffins and other organic chemicals generate strong reactions: fully substituted haloalkanes such as carbon tetrachloride incombustible, may explode. Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the strong triple bond in elemental nitrogen. Oxygen does not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures. Heavier halogens react with fluorine as does the noble gas radon. At room temperature, fluorine is a gas of diatomic molecules, pale yellow, it has a characteristic halogen-like biting odor detectable at 20 ppb. Fluorine condenses into a bright yellow liquid at −188 °C, a transition temperature similar to those of oxygen and nitrogen.
Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at −220 °C and is transparent and sof