Supercritical liquid–gas boundaries
Supercritical liquid–gas boundaries are lines in the p–T diagram that delimit more liquid-like and more gas-like states of a supercritical fluid. They comprise the Fisher–Widom line, the Widom line, the Frenkel line. According to textbook knowledge, it is possible to transform a liquid continuously into a gas, without undergoing a phase transition, by heating and compressing enough to go around the critical point. However, different criteria still allow to distinguish liquid-like and more gas-like states of a supercritical fluid; these criteria result in different boundaries in the pT plane. These lines emanate either from the critical point, or from the liquid–vapor boundary somewhat below the critical point, they do not correspond to weaker singularities. The Fisher–Widom line is the boundary between monotonic and oscillating asymptotics of the pair correlation function G; the Widom line is a generalization thereof so named by H. Eugene Stanley. However, it was first measured experimentally in 1956 by Jones and Walker, subsequently named the'hypercritical line' by Bernal in 1964, who suggested a structural interpretation.
The Frenkel line is a boundary between "rigid" and "non-rigid" fluids characterized by the onset of transverse sound modes. One of the above mentioned criteria is based on the velocity autocorrelation function: below the Frenkel line the vacf demonstrates oscillatory behaviour, while above it the vacf monotonically decays to zero; the second criterion is based on the fact that at moderate temperatures liquids can sustain transverse excitations, which disappear upon heating. One further criterion is based on isochoric heat capacity measurements; the isochoric heat capacity per particle of a monatomic liquid near to the melting line is close to 3 k B. The contribution to the heat capacity due to the potential part of transverse excitations is 1 k B; therefore at the Frenkel line, where transverse excitations vanish, the isochoric heat capacity per particle should be c V = 2 k B, a direct prediction from the phonon theory of liquid thermodynamics. Anisimov et al. without referring to Frenkel, Fisher or Widom, reviewed thermodynamic derivatives and transport coefficients in supercritical water, found pronounced extrema as function of pressure up to 100 K above Tc. 8.
G. G. Simeoni, T. Bryk, F. A. Gorelli, M. Krisch, G. Ruocco, M. Santoro & T. Scopigno, The Widom line as the crossover between liquid-like and gas-like behaviour in supercritical fluids, Nature Physics 6, 503–507 http://www.nature.com/nphys/journal/v6/n7/full/nphys1683.html?foxtrotcallback=true
Xenon is a chemical element with symbol Xe and atomic number 54. It is a colorless, odorless noble gas found in the Earth's atmosphere in trace amounts. Although unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized. Xenon is used in flash lamps and arc lamps, as a general anesthetic; the first excimer laser design used a xenon dimer molecule as the lasing medium, the earliest laser designs used xenon flash lamps as pumps. Xenon is used to search for hypothetical weakly interacting massive particles and as the propellant for ion thrusters in spacecraft. Occurring xenon consists of eight stable isotopes. More than 40 unstable xenon isotopes undergo radioactive decay, the isotope ratios of xenon are an important tool for studying the early history of the Solar System. Radioactive xenon-135 is produced by beta decay from iodine-135, is the most significant neutron absorber in nuclear reactors. Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers in September 1898, shortly after their discovery of the elements krypton and neon.
They found xenon in the residue left over from evaporating components of liquid air. Ramsay suggested the name xenon for this gas from the Greek word ξένον, neuter singular form of ξένος, meaning'foreign','strange', or'guest'. In 1902, Ramsay estimated the proportion of xenon in the Earth's atmosphere to be one part in 20 million. During the 1930s, American engineer Harold Edgerton began exploring strobe light technology for high speed photography; this led him to the invention of the xenon flash lamp in which light is generated by passing brief electric current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method. In 1939, American physician Albert R. Behnke Jr. began exploring the causes of "drunkenness" in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, discovered that this caused the divers to perceive a change in depth. From his results, he deduced. Although Russian toxicologist Nikolay V. Lazarev studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice.
Xenon was first used as a surgical anesthetic in 1951 by American anesthesiologist Stuart C. Cullen, who used it with two patients. Xenon and the other noble gases were for a long time considered to be chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride was a powerful oxidizing agent that could oxidize oxygen gas to form dioxygenyl hexafluoroplatinate. Since O2 and xenon have the same first ionization potential, Bartlett realized that platinum hexafluoride might be able to oxidize xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate. Bartlett thought its composition to be Xe+−, but work revealed that it was a mixture of various xenon-containing salts. Since many other xenon compounds have been discovered, in addition to some compounds of the noble gases argon and radon, including argon fluorohydride, krypton difluoride, radon fluoride.
By 1971, more than 80 xenon compounds were known. In November 1989, IBM scientists demonstrated a technology capable of manipulating individual atoms; the program, called IBM in atoms, used a scanning tunneling microscope to arrange 35 individual xenon atoms on a substrate of chilled crystal of nickel to spell out the three letter company initialism. It was the first time atoms had been positioned on a flat surface. Xenon has atomic number 54. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3. As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point. Liquid xenon has a high polarizability due to its large atomic volume, thus is an excellent solvent, it can dissolve hydrocarbons, biological molecules, water. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is greater than the average density of granite, 2.75 g/cm3.
Under gigapascals of pressure, xenon forms a metallic phase. Solid xenon changes from face-centered cubic to hexagonal close packed crystal phase under pressure and begins to turn metallic at about 140 GPa, with no noticeable volume change in the hcp phase, it is metallic at 155 GPa. When metallized, xenon appears sky blue because it absorbs red light and transmits other visible frequencies; such behavior is unusual for a metal and is explained by the small width of the electron bands in that state. Liquid or solid xenon nanoparticles can be formed at room temperature by implanting Xe+ ions into a solid matrix. Many solids have lattice constants smaller than solid Xe; this results in compression of the implanted Xe to pressures that may be sufficient for its liquefaction or solidification. Xenon is a member of the zero-valence elements that are called inert gases, it is inert to most common chemical reactions because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are bound.
In a gas-filled tube, xenon em
Sulfuric acid known as vitriol, is a mineral acid composed of the elements sulfur and hydrogen, with molecular formula H2SO4. It is a colorless and syrupy liquid, soluble in water, in a reaction, exothermic, its corrosiveness can be ascribed to its strong acidic nature, and, if at a high concentration, its dehydrating and oxidizing properties. It is hygroscopic absorbing water vapor from the air. Upon contact, sulfuric acid can cause severe chemical burns and secondary thermal burns. Sulfuric acid is a important commodity chemical, a nation's sulfuric acid production is a good indicator of its industrial strength, it is produced with different methods, such as contact process, wet sulfuric acid process, lead chamber process and some other methods. Sulfuric acid is a key substance in the chemical industry, it is most used in fertilizer manufacture, but is important in mineral processing, oil refining, wastewater processing, chemical synthesis. It has a wide range of end applications including in domestic acidic drain cleaners, as an electrolyte in lead-acid batteries, in various cleaning agents.
Although nearly 100% sulfuric acid can be made, the subsequent loss of SO3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes; some common concentrations are: "Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in the lead chamber itself and tower acid being the acid recovered from the bottom of the Glover tower. They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid is prepared by adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C or higher. Sulfuric acid reacts with its anhydride, SO3, to form H2S2O7, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less Nordhausen acid.
Concentrations of oleum are either expressed in terms of % SO3 or as % H2SO4. Pure H2S2O7 is a solid with melting point of 36 °C. Pure sulfuric acid has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C. Pure sulfuric acid is a viscous clear liquid, like oil, this explains the old name of the acid. Commercial sulfuric acid is sold in several different purity grades. Technical grade H2SO4 is impure and colored, but is suitable for making fertilizer. Pure grades, such as United States Pharmacopeia grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are available. Nine hydrates are known, but four of them were confirmed to be tetrahydrate and octahydrate. Anhydrous H2SO4 is a polar liquid, having a dielectric constant of around 100, it has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis. 2 H2SO4 ⇌ H3SO+4 + HSO−4The equilibrium constant for the autoprotolysis is Kap = = 2.7×10−4The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 smaller.
In spite of the viscosity of the acid, the effective conductivities of the H3SO+4 and HSO−4 ions are high due to an intramolecular proton-switch mechanism, making sulfuric acid a good conductor of electricity. It is an excellent solvent for many reactions; because the hydration reaction of sulfuric acid is exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent; this reaction is best thought of as the formation of hydronium ions: H2SO4 + H2O → H3O+ + HSO−4 Ka1 = 2.4×106 HSO−4 + H2O → H3O+ + SO2−4 Ka2 = 1.0×10−2 HSO−4 is the bisulfate anion and SO2−4 is the sulfate anion. Ka1 and Ka2 are the acid dissociation constants; because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent.
Concentrated sulfuric acid has a powerful dehydrating property, removing water from other chemical compounds including sugar and other carbohydrates and producing carbon and steam. In the laboratory, this is demonstrated by mixing table sugar into sulfuric acid; the sugar changes from white to dark brown and to black as carbon is formed. A rigid column of black, porous carbon will emerge as well; the carbon will smell of caramel due to the heat generated. C 12 H 22 O 11 ⏞ sucrose → H 2 SO 4 12 C + 11 H 2
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
Krypton is a chemical element with symbol Kr and atomic number 36. It is a member of group 18 elements. A colorless, tasteless noble gas, krypton occurs in trace amounts in the atmosphere and is used with other rare gases in fluorescent lamps. With rare exceptions, krypton is chemically inert. Krypton, like the other noble gases, is used in photography. Krypton light has many spectral lines, krypton plasma is useful in bright, high-powered gas lasers, each of which resonates and amplifies a single spectral line. Krypton fluoride makes a useful laser medium. From 1960 to 1983, the official length of a meter was defined by the 605 nm wavelength of the orange spectral line of krypton-86, because of the high power and relative ease of operation of krypton discharge tubes. Krypton was discovered in Britain in 1898 by Sir William Ramsay, a Scottish chemist, Morris Travers, an English chemist, in residue left from evaporating nearly all components of liquid air. Neon was discovered by a similar procedure by the same workers just a few weeks later.
William Ramsay was awarded the 1904 Nobel Prize in Chemistry for discovery of a series of noble gases, including krypton. In 1960, the International Conference on Weights and Measures defined the meter as 1,650,763.73 wavelengths of light emitted by the krypton-86 isotope. This agreement replaced the 1889 international prototype meter located in Paris, a metal bar made of a platinum-iridium alloy; this obsoleted the 1927 definition of the ångström based on the red cadmium spectral line, replacing it with 1 Å = 10−10 m. The krypton-86 definition lasted until the October 1983 conference, which redefined the meter as the distance that light travels in vacuum during 1/299,792,458 s. Krypton is characterized by several sharp emission lines the strongest being yellow. Krypton is one of the products of uranium fission. Solid krypton is white and has a face-centered cubic crystal structure, a common property of all noble gases. Occurring krypton in Earth's atmosphere is composed of five stable isotopes, plus one isotope with such a long half-life that it can be considered stable..
In addition, about thirty unstable isotopes and isomers are known. Traces of 81Kr, a cosmogenic nuclide produced by the cosmic ray irradiation of 80Kr occur in nature: this isotope is radioactive with a half-life of 230,000 years. Krypton is volatile and does not stay in solution in near-surface water, but 81Kr has been used for dating old groundwater.85Kr is an inert radioactive noble gas with a half-life of 10.76 years. It is produced by the fission of uranium and plutonium, such as in nuclear bomb testing and nuclear reactors. 85Kr is released during the reprocessing of fuel rods from nuclear reactors. Concentrations at the North Pole are 30% higher than at the South Pole due to convective mixing. Like the other noble gases, krypton is chemically unreactive; the rather restricted chemistry of krypton in its only known nonzero oxidation state of +2 parallels that of the neighboring element bromine in the +1 oxidation state. Before the 1960s, no noble gas compounds had been synthesized. However, following the first successful synthesis of xenon compounds in 1962, synthesis of krypton difluoride was reported in 1963.
In the same year, KrF4 was reported by Grosse, et al. but was subsequently shown to be a mistaken identification. Under extreme conditions, krypton reacts with fluorine to form KrF2 according to the following equation: Kr + F2 → KrF2Compounds with krypton bonded to atoms other than fluorine have been discovered. There are unverified reports of a barium salt of a krypton oxoacid. ArKr+ and KrH+ polyatomic ions have been investigated and there is evidence for KrXe or KrXe+; the reaction of KrF2 with B3 produces an unstable compound, Kr2, that contains a krypton-oxygen bond. A krypton-nitrogen bond is found in the cation +, produced by the reaction of KrF2 with + below −50 °C. HKrCN and HKrC≡CH were reported to be stable up to 40 K. Krypton hydride crystals can be grown at pressures above 5 GPa, they have a face-centered cubic structure where krypton octahedra are surrounded by randomly oriented hydrogen molecules. Earth has retained all of the noble gases. Krypton's concentration in the atmosphere is about 1 ppm.
It can be extracted from liquid air by fractional distillation. The amount of krypton in space is uncertain, because measurement is derived from meteoric activity and solar winds; the first measurements suggest an abundance of krypton in space. Krypton's multiple emission lines make ionized krypton gas discharges appear whitish, which in turn makes krypton-based bulbs useful in photography as a brilliant white light source. Krypton is used in some photographic flashes for high speed photography. Krypton gas is combined with other gases to make luminous signs that glow with a bright greenish-yellow light. Krypton is mixed with argon in energy efficient fluorescent lamps, reducing the power consumption, but reducing the light output and raising the c
A supercritical fluid is any substance at a temperature and pressure above its critical point, where distinct liquid and gas phases do not exist. It can effuse through solids like a gas, dissolve materials like a liquid. In addition, close to the critical point, small changes in pressure or temperature result in large changes in density, allowing many properties of a supercritical fluid to be "fine-tuned". Supercritical fluids occur in the atmospheres of the gas giants Jupiter and Saturn, in those of the ice giants Uranus and Neptune. In a range of industrial and laboratory processes, they are used as a substitute for organic solvents. Carbon dioxide and water are the most used supercritical fluids, being used for decaffeination and power generation, respectively. In general terms, supercritical fluids have properties between those of a liquid. In Table 1, the critical properties are shown for some substances that are used as supercritical fluids. Table 2 shows density and viscosity for typical liquids and supercritical fluids.
In addition, there is no surface tension in a supercritical fluid, as there is no liquid/gas phase boundary. By changing the pressure and temperature of the fluid, the properties can be "tuned" to be more liquid-like or more gas-like. One of the most important properties is the solubility of material in the fluid. Solubility in a supercritical fluid tends to increase with density of the fluid. Since density increases with pressure, solubility tends to increase with pressure; the relationship with temperature is a little more complicated. At constant density, solubility will increase with temperature. However, close to the critical point, the density can drop with a slight increase in temperature. Therefore, close to the critical temperature, solubility drops with increasing temperature rises again. All supercritical fluids are miscible with each other so for a mixture a single phase can be guaranteed if the critical point of the mixture is exceeded; the critical point of a binary mixture can be estimated as the arithmetic mean of the critical temperatures and pressures of the two components, Tc = × TcA + × TcB.
For greater accuracy, the critical point can be calculated using equations of state, such as the Peng Robinson, or group contribution methods. Other properties, such as density, can be calculated using equations of state. Figures 1 and 2 show two-dimensional projections of a phase diagram. In the pressure-temperature phase diagram the boiling separates the gas and liquid region and ends in the critical point, where the liquid and gas phases disappear to become a single supercritical phase; the appearance of a single phase can be observed in the density-pressure phase diagram for carbon dioxide. At well below the critical temperature, e.g. 280 K, as the pressure increases, the gas compresses and condenses into a much denser liquid, resulting in the discontinuity in the line. The system consists of 2 phases in a dense liquid and a low density gas; as the critical temperature is approached, the density of the gas at equilibrium becomes higher, that of the liquid lower. At the critical point, there is no difference in density, the 2 phases become one fluid phase.
Thus, above the critical temperature a gas cannot be liquefied by pressure. At above the critical temperature, in the vicinity of the critical pressure, the line is vertical. A small increase in pressure causes a large increase in the density of the supercritical phase. Many other physical properties show large gradients with pressure near the critical point, e.g. viscosity, the relative permittivity and the solvent strength, which are all related to the density. At higher temperatures, the fluid starts to behave like a gas, as can be seen in Figure 2. For carbon dioxide at 400 K, the density increases linearly with pressure. Many pressurized gases are supercritical fluids. For example, nitrogen has a critical point of 3.4 MPa. Therefore, nitrogen in a gas cylinder above this pressure is a supercritical fluid; these are more known as permanent gases. At room temperature, they are well above their critical temperature, therefore behave as a gas, similar to CO2 at 400 K above. However, they can not be liquified by pressure.
In recent years, a significant effort has been devoted to investigation of various properties of supercritical fluids. This has been an exciting field with a long history since 1822 when Baron Charles Cagniard de la Tour discovered supercritical fluids while conducting experiments involving the discontinuities of the sound in a sealed cannon barrel filled with various fluids at high temperature. More supercritical fluids have found application in a variety of fields, ranging from the extraction of floral fragrance from flowers to applications in food science such as creating decaffeinated coffee, functional food ingredients, cosmetics, powders, bio- and functional materials, nano-systems, natural products, biotechnology and bio-fuels, microelectronics and environment. Much of the excitement and interest of the past decade is due to the enormous progress made in increasing the power of relevant experimental tools; the development of new experimental methods and improvement of existing ones continues to play an important role in this field, with recent research focusing on dynamic properties of fluids.
The Fisher-Widom line, the Widom line, or the Frenk
The Fahrenheit scale is a temperature scale based on one proposed in 1724 by Dutch–German–Polish physicist Daniel Gabriel Fahrenheit. It uses the degree Fahrenheit as the unit. Several accounts of how he defined his scale exist; the lower defining point, 0 °F, was established as the freezing temperature of a solution of brine made from equal parts of ice and salt. Further limits were established as the melting point of ice and his best estimate of the average human body temperature; the scale is now defined by two fixed points: the temperature at which water freezes into ice is defined as 32 °F, the boiling point of water is defined to be 212 °F, a 180 °F separation, as defined at sea level and standard atmospheric pressure. At the end of the 2010s, Fahrenheit was used as the official temperature scale only in the United States, its associated states in the Western Pacific, the Bahamas, the Cayman Islands and Liberia. Antigua and Barbuda and other islands which use the same meteorological service, such as Anguilla, the British Virgin Islands and Saint Kitts and Nevis, as well as Bermuda and the Turks and Caicos Islands, use Fahrenheit and Celsius.
All other countries in the world now use the Celsius scale, named after Swedish astronomer Anders Celsius. On the Fahrenheit scale, the freezing point of water is 32 degrees Fahrenheit and the boiling point is 212 °F; this puts the freezing points of water 180 degrees apart. Therefore, a degree on the Fahrenheit scale is 1⁄180 of the interval between the freezing point and the boiling point. On the Celsius scale, the freezing and boiling points of water are 100 degrees apart. A temperature interval of 1 °F is equal to an interval of 5⁄9 degrees Celsius; the Fahrenheit and Celsius scales intersect at −40°. Absolute zero is −273.15 °C or −459.67 °F. The Rankine temperature scale uses degree intervals of the same size as those of the Fahrenheit scale, except that absolute zero is 0 °R — the same way that the Kelvin temperature scale matches the Celsius scale, except that absolute zero is 0 K; the Fahrenheit scale uses the symbol ° to denote a point on the temperature scale and the letter F to indicate the use of the Fahrenheit scale, as well as to denote a difference between temperatures or an uncertainty in temperature.
For an exact conversion, the following formulas can be applied. Here, f is the value in Fahrenheit and c the value in Celsius: f °Fahrenheit to c °Celsius: °F × 5°C/9°F = /1.8 °C = c °C c °Celsius to f °Fahrenheit: + 32 °F = °F + 32 °F = f °FThis is an exact conversion making use of the identity −40 °F = −40 °C. Again, f is the value in Fahrenheit and c the value in Celsius: f °Fahrenheit to c °Celsius: − 40 = c. C °Celsius to f °Fahrenheit: − 40 = f. Fahrenheit proposed his temperature scale in 1724, basing it on two reference points of temperature. In his initial scale, the zero point was determined by placing the thermometer in a mixture "of ice, of water, of ammonium chloride or of sea salt"; this combination forms a eutectic system which stabilizes its temperature automatically: 0 °F was defined to be that stable temperature. The second point, 96 degrees, was the human body's temperature. According to a story in Germany, Fahrenheit chose the lowest air temperature measured in his hometown Danzig in winter 1708/09 as 0 °F, only had the need to be able to make this value reproducible using brine.
According to a letter Fahrenheit wrote to his friend Herman Boerhaave, his scale was built on the work of Ole Rømer, whom he had met earlier. In Rømer's scale, brine freezes at zero, water freezes and melts at 7.5 degrees, body temperature is 22.5, water boils at 60 degrees. Fahrenheit multiplied each value by four in order to eliminate fractions and make the scale more fine-grained, he re-calibrated his scale using the melting point of ice and normal human body temperature. Fahrenheit soon after observed; the use of the freezing and boiling points of water as thermometer fixed reference points became popular following the work of Anders Celsius and these fixed points were adopted by a committee of the Royal Society led by Henry Cavendish in 1776. Under this system, the Fahrenheit scale is redefined so that the freezing point of water is 32 °F, the boiling point is 212 °F or 180 degrees higher, it is for this reason that normal human body temperature is 98° on the revised scale. In the present-day Fahrenheit scale, 0 °F no longer corresponds to the eutectic temperature of ammonium chloride brine as described above.
Instead, that eutectic is at 4 °F on the final Fahrenheit scale. The Rankine temperature s