The melting point of a substance is the temperature at which it changes state from solid to liquid. At the melting point the solid and liquid phase exist in equilibrium; the melting point of a substance depends on pressure and is specified at a standard pressure such as 1 atmosphere or 100 kPa. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point; because of the ability of some substances to supercool, the freezing point is not considered as a characteristic property of a substance. When the "characteristic freezing point" of a substance is determined, in fact the actual methodology is always "the principle of observing the disappearance rather than the formation of ice", that is, the melting point. For most substances and freezing points are equal. For example, the melting point and freezing point of mercury is 234.32 kelvins. However, certain substances possess differing solid-liquid transition temperatures.
For example, agar melts at 85 °C and solidifies from 31 °C. The melting point of ice at 1 atmosphere of pressure is close to 0 °C. In the presence of nucleating substances, the freezing point of water is not always the same as the melting point. In the absence of nucleators water can exist as a supercooled liquid down to −48.3 °C before freezing. The chemical element with the highest melting point is tungsten, at 3,414 °C; the often-cited carbon does not melt at ambient pressure but sublimes at about 3,726.85 °C. Tantalum hafnium carbide is a refractory compound with a high melting point of 4215 K. At the other end of the scale, helium does not freeze at all at normal pressure at temperatures arbitrarily close to absolute zero. Many laboratory techniques exist for the determination of melting points. A Kofler bench is a metal strip with a temperature gradient. Any substance can be placed on a section of the strip, revealing its thermal behaviour at the temperature at that point. Differential scanning calorimetry gives information on melting point together with its enthalpy of fusion.
A basic melting point apparatus for the analysis of crystalline solids consists of an oil bath with a transparent window and a simple magnifier. The several grains of a solid are placed in a thin glass tube and immersed in the oil bath; the oil bath is heated and with the aid of the magnifier melting of the individual crystals at a certain temperature can be observed. In large/small devices, the sample is placed in a heating block, optical detection is automated; the measurement can be made continuously with an operating process. For instance, oil refineries measure the freeze point of diesel fuel online, meaning that the sample is taken from the process and measured automatically; this allows for more frequent measurements as the sample does not have to be manually collected and taken to a remote laboratory. For refractory materials the high melting point may be determined by heating the material in a black body furnace and measuring the black-body temperature with an optical pyrometer. For the highest melting materials, this may require extrapolation by several hundred degrees.
The spectral radiance from an incandescent body is known to be a function of its temperature. An optical pyrometer matches the radiance of a body under study to the radiance of a source, calibrated as a function of temperature. In this way, the measurement of the absolute magnitude of the intensity of radiation is unnecessary. However, known temperatures must be used to determine the calibration of the pyrometer. For temperatures above the calibration range of the source, an extrapolation technique must be employed; this extrapolation is accomplished by using Planck's law of radiation. The constants in this equation are not known with sufficient accuracy, causing errors in the extrapolation to become larger at higher temperatures. However, standard techniques have been developed to perform this extrapolation. Consider the case of using gold as the source. In this technique, the current through the filament of the pyrometer is adjusted until the light intensity of the filament matches that of a black-body at the melting point of gold.
This establishes the primary calibration temperature and can be expressed in terms of current through the pyrometer lamp. With the same current setting, the pyrometer is sighted on another black-body at a higher temperature. An absorbing medium of known transmission is inserted between this black-body; the temperature of the black-body is adjusted until a match exists between its intensity and that of the pyrometer filament. The true higher temperature of the black-body is determined from Planck's Law; the absorbing medium is removed and the current through the filament is adjusted to match the filament intensity to that of the black-body. This establishes a second calibration point for the pyrometer; this step is repeated to carry the calibration to hi
The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor. The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water at 93.4 °C at 1,905 metres altitude. For a given pressure, different liquids will boil at different temperatures; the normal boiling point of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere. At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid; the standard boiling point has been defined by IUPAC since 1982 as the temperature at which boiling occurs under a pressure of 1 bar.
The heat of vaporization is the energy required to transform a given quantity of a substance from a liquid into a gas at a given pressure. Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid. A saturated liquid contains as much thermal energy. Saturation temperature means boiling point; the saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition. If the pressure in a system remains constant, a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy is removed.
A liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied. The boiling point corresponds to the temperature at which the vapor pressure of the liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Boiling points may be published with respect to the NIST, USA standard pressure of 101.325 kPa, or the IUPAC standard pressure of 100.000 kPa. At higher elevations, where the atmospheric pressure is much lower, the boiling point is lower; the boiling point increases with increased pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point; the boiling point decreases with decreasing pressure until the triple point is reached. The boiling point cannot be reduced below the triple point. If the heat of vaporization and the vapor pressure of a liquid at a certain temperature are known, the boiling point can be calculated by using the Clausius–Clapeyron equation, thus: T B = − 1, where: T B is the boiling point at the pressure of interest, R is the ideal gas constant, P is the vapour pressure of the liquid at the pressure of interest, P 0 is some pressure where the corresponding T 0 is known, Δ H vap is the heat of vaporization of the liquid, T 0 is the boiling temperature, ln is the natural logarithm.
Saturation pressure is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased, so is saturation temperature. If the temperature in a system remains constant, vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. A liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased. There are two conventions regarding the standard boiling point of water: The normal boiling point is 99.97 °C at a pressure of 1 atm. The IUPAC recommended standard boiling point of water at a standard pressure of 100 kPa is 99.61 °C. For comparison, on top of Mount Everest, at 8,848 m elevation, the pressure is about 34 kPa and the boiling point of water is 71 °C; the Celsius temperature scale was defined until 1954 by two points: 0 °C being defined by the wate
Rhodamine is a family of related chemical compounds, fluorone dyes. Examples are Rhodamine 6G and Rhodamine B, they are used as dye laser gain media. They are used as a tracer dye within water to determine the rate and direction of flow and transport. Rhodamine dyes fluoresce and can thus be detected and inexpensively with instruments called fluorometers. Rhodamine dyes are used extensively in biotechnology applications such as fluorescence microscopy, flow cytometry, fluorescence correlation spectroscopy and ELISA. Rhodamine dyes are toxic, are soluble in water and ethanol; the laser dye rhodamine 123 is used in biochemistry to inhibit mitochondrion function. Rhodamine 123 seems to bind to the mitochondrion membranes and inhibit transport processes the electron transport chain, thus slowing down inner respiration, it is a substrate of P-glycoprotein, overexpressed in cancer cells. Recent reports indicate that rhodamine 123 may be a substrate of multidrug resistance-associated protein, or more MRP1.
There are many rhodamine derivatives used for imaging purposes, for example Carboxytetramethylrhodamine, tetramethylrhodamine and its isothiocyanate derivative and, sulforhodamine 101 and Rhodamine Red. TRITC is the base rhodamine molecule functionalized with an isothiocyanate group, replacing a hydrogen atom on the bottom ring of the structure; this derivative is reactive towards amine groups on proteins inside cells. A succinimidyl-ester functional group attached to the rhodamine core, creating NHS-rhodamine, forms another common amine-reactive derivative. Other derivatives of rhodamine include newer fluorophores such as Alexa 546, Alexa 633, DyLight 550 and DyLight 633, HiLyte fluor 555 HiLyte 594 have been tailored for various chemical and biological applications where higher photostability, increased brightness, different spectral characteristics, or different attachment groups are needed. Absorption and Emission Spectra of Rhodamine B Absorption and Emission Spectra of Rhodamine 6G Absorption and Emission Spectra of Rhodamine 123 Berlier et al. 2003 J. Histochem Cytochem refers to Alexa 633 as a rhodamine derivative
Redox is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes. Redox reactions include all chemical reactions; the chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms: Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion. Reduction is a decrease in oxidation state by a molecule, atom, or ion; as an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon, oxidized. Although oxidation reactions are associated with the formation of oxides from oxygen molecules, oxygen is not included in such reactions, as other chemical species can serve the same function; the reaction can occur slowly, as with the formation of rust, or more in the case of fire.
There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide or the reduction of carbon by hydrogen to yield methane, more complex processes such as the oxidation of glucose in the human body. "Redox" is a portmanteau of the words "reduction" and "oxidation". The word oxidation implied reaction with oxygen to form an oxide, since dioxygen was the first recognized oxidizing agent; the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. The meaning was generalized to include all processes involving loss of electrons; the word reduction referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was "reduced" to metal. Antoine Lavoisier showed. Scientists realized that the metal atom gains electrons in this process; the meaning of reduction became generalized to include all processes involving a gain of electrons. Though "reduction" seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, its historical meaning.
Since electrons are negatively charged, it is helpful to think of this as reduction in electrical charge. The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes when they occur at electrodes; these words are analogous to protonation and deprotonation, but they have not been adopted by chemists worldwide. The term "hydrogenation" could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions in organic chemistry and biochemistry. But, unlike oxidation, generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance; the word "redox" was first used in 1928. The processes of oxidation and reduction occur and cannot happen independently of one another, similar to the acid–base reaction; the oxidation alone and the reduction alone are each called a half-reaction, because two half-reactions always occur together to form a whole reaction.
When writing half-reactions, the gained or lost electrons are included explicitly in order that the half-reaction be balanced with respect to electric charge. Though sufficient for many purposes, these general descriptions are not correct. Although oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur; the oxidation state of an atom is the fictitious charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is best defined as an increase in oxidation state, reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" though no electron transfer occurs. In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, the oxidant or oxidizing agent gains electrons and is reduced.
The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidizing form, e.g. Fe2+/Fe3+ Substances that have the ability to oxidize other substances are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers; that is, the oxidant removes electrons from another substance, is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is called an electron acceptor. Oxygen is the quintessential oxidizer. Oxidants are chemical substances with elements in high oxidation states, or else electronegative elements that can gain extra electrons by oxidizing another substance. Substances that have the ability to reduce other substances are said to be reductive or reducing and are known as
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.
Barium is a chemical element with symbol Ba and atomic number 56. It is a soft, silvery alkaline earth metal; because of its high chemical reactivity, barium is never found in nature as a free element. Its hydroxide, known in pre-modern times as baryta, does not occur as a mineral, but can be prepared by heating barium carbonate; the most common occurring minerals of barium are barite and witherite, both insoluble in water. The name barium originates from the alchemical derivative "baryta", from Greek βαρύς, meaning "heavy." Baric is the adjectival form of barium. Barium was identified as a new element in 1774, but not reduced to a metal until 1808 with the advent of electrolysis. Barium has few industrial applications, it was used as a getter for vacuum tubes and in oxide form as the emissive coating on indirectly heated cathodes. It is a component of YBCO and electroceramics, is added to steel and cast iron to reduce the size of carbon grains within the microstructure. Barium compounds are added to fireworks to impart a green color.
Barium sulfate is used as an insoluble additive to oil well drilling fluid, as well as in a purer form, as X-ray radiocontrast agents for imaging the human gastrointestinal tract. The soluble barium ion and soluble compounds are poisonous, have been used as rodenticides. Barium is a silvery-white metal, with a slight golden shade when ultrapure; the silvery-white color of barium metal vanishes upon oxidation in air yielding a dark gray oxide layer. Barium has good electrical conductivity. Ultrapure barium is difficult to prepare, therefore many properties of barium have not been measured yet. At room temperature and pressure, barium has a body-centered cubic structure, with a barium–barium distance of 503 picometers, expanding with heating at a rate of 1.8×10−5/°C. It is a soft metal with a Mohs hardness of 1.25. Its melting temperature of 1,000 K is intermediate between those of the lighter strontium and heavier radium; the density is again intermediate between those of radium. Barium is chemically similar to magnesium and strontium, but more reactive.
It always exhibits the oxidation state of +2, except in a few rare and unstable molecular species that are only characterised in the gas phase such as BaF. Reactions with chalcogens are exothermic. Reactions with other nonmetals, such as carbon, phosphorus and hydrogen, are exothermic and proceed upon heating. Reactions with water and alcohols are exothermic and release hydrogen gas: Ba + 2 ROH → Ba2 + H2↑ Barium reacts with ammonia to form complexes such as Ba6; the metal is attacked by most acids. Sulfuric acid is a notable exception because passivation stops the reaction by forming the insoluble barium sulfate on the surface. Barium combines with several metals, including aluminium, zinc and tin, forming intermetallic phases and alloys. Barium salts are white when solid and colorless when dissolved, barium ions provide no specific coloring, they are denser than the calcium analogs, except for the halides. Barium hydroxide was known to alchemists. Unlike calcium hydroxide, it absorbs little CO2 in aqueous solutions and is therefore insensitive to atmospheric fluctuations.
This property is used in calibrating pH equipment. Volatile barium compounds burn with a green to pale green flame, an efficient test to detect a barium compound; the color results from spectral lines at 455.4, 493.4, 553.6, 611.1 nm. Organobarium compounds are a growing field of knowledge: discovered are dialkylbariums and alkylhalobariums. Barium found in the Earth's crust is a mixture of seven primordial nuclides, barium-130, 132, 134 through 138. Barium-130 undergoes slow radioactive decay to xenon-130 by double beta plus decay, barium-132 theoretically decays to xenon-132, with half-lives a thousand times greater than the age of the Universe; the abundance is ≈ 0.1 %. The radioactivity of these isotopes is so weak. Of the stable isotopes, barium-138 composes 71.7% of all barium. In total, barium has about 40 known isotopes, ranging in mass between 114 and 153; the most stable artificial radioisotope is barium-133 with a half-life of 10.51 years. Five other isotopes have half-lives longer than a day.
Barium has 10 meta states, of which barium-133m1 is the most stable with a half-life of about 39 hours. Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral baryte were found in volcanic rock near Bologna, so were called "Bologna stones." Alchemists were attracted to them. The phosphorescent properties of baryte heated with organics were described by V. Casciorolus in 1602. Carl Scheele determined that baryte contained a new element in 1774, but could not isolate barium, only barium oxide. Johan Gottlieb Gahn isolated barium oxide two year
Hydrochloric acid or muriatic acid is a colorless inorganic chemical system with the formula H2O:HCl. Hydrochloric acid has a distinctive pungent smell, it is classified as acidic and can attack the skin over a wide composition range, since the hydrogen chloride dissociates in aqueous solution. Hydrochloric acid is the simplest chlorine-based acid system containing water, it is a solution of hydrogen chloride and water, a variety of other chemical species, including hydronium and chloride ions. It is an important chemical reagent and industrial chemical, used in the production of polyvinyl chloride for plastic. In households, diluted hydrochloric acid is used as a descaling agent. In the food industry, hydrochloric acid is used in the production of gelatin. Hydrochloric acid is used in leather processing. Hydrochloric acid was discovered by the alchemist Jabir ibn Hayyan around the year 800 AD, it was called acidum salis and spirits of salt because it was produced from rock salt and "green vitriol" and from the chemically similar common salt and sulfuric acid.
Free hydrochloric acid was first formally described in the 16th century by Libavius. It was used by chemists such as Glauber and Davy in their scientific research. Unless pressurized or cooled, hydrochloric acid will turn into a gas if there is around 60% or less of water. Hydrochloric acid is known as hydronium chloride, in contrast to its anhydrous parent known as hydrogen chloride, or dry HCl. Hydrochloric acid was known to European alchemists as spirits of acidum salis. Both names are still used in other languages, such as German: Salzsäure, Dutch: Zoutzuur, Swedish: Saltsyra, Turkish: Tuz Ruhu, Polish: kwas solny, Bulgarian: солна киселина, Russian: соляная кислота, Chinese: 鹽酸, Korean: 염산, Taiwanese: iâm-sng. Gaseous HCl was called marine acid air; the old name muriatic acid has the same origin, this name is still sometimes used. The name hydrochloric acid was coined by the French chemist Joseph Louis Gay-Lussac in 1814. Hydrochloric acid has been an important and used chemical from early history and was discovered by the alchemist Jabir ibn Hayyan around the year 800 AD.
Aqua regia, a mixture consisting of hydrochloric and nitric acids, prepared by dissolving sal ammoniac in nitric acid, was described in the works of Pseudo-Geber, a 13th-century European alchemist. Other references suggest that the first mention of aqua regia is in Byzantine manuscripts dating to the end of the 13th century. Free hydrochloric acid was first formally described in the 16th century by Libavius, who prepared it by heating salt in clay crucibles. Other authors claim that pure hydrochloric acid was first discovered by the German Benedictine monk Basil Valentine in the 15th century, when he heated common salt and green vitriol, whereas others argue that there is no clear reference to the preparation of pure hydrochloric acid until the end of the 16th century. In the 17th century, Johann Rudolf Glauber from Karlstadt am Main, Germany used sodium chloride salt and sulfuric acid for the preparation of sodium sulfate in the Mannheim process, releasing hydrogen chloride gas. Joseph Priestley of Leeds, England prepared pure hydrogen chloride in 1772, by 1808 Humphry Davy of Penzance, England had proved that the chemical composition included hydrogen and chlorine.
During the Industrial Revolution in Europe, demand for alkaline substances increased. A new industrial process developed by Nicolas Leblanc of Issoudun, France enabled cheap large-scale production of sodium carbonate. In this Leblanc process, common salt is converted to soda ash, using sulfuric acid and coal, releasing hydrogen chloride as a by-product; until the British Alkali Act 1863 and similar legislation in other countries, the excess HCl was vented into the air. After the passage of the act, soda ash producers were obliged to absorb the waste gas in water, producing hydrochloric acid on an industrial scale. In the 20th century, the Leblanc process was replaced by the Solvay process without a hydrochloric acid by-product. Since hydrochloric acid was fully settled as an important chemical in numerous applications, the commercial interest initiated other production methods, some of which are still used today. After the year 2000, hydrochloric acid is made by absorbing by-product hydrogen chloride from industrial organic compounds production.
Since 1988, hydrochloric acid has been listed as a Table II precursor under the 1988 United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances because of its use in the production of heroin and methamphetamine. Hydrochloric acid is the salt of H3O + and chloride, it is prepared by treating HCl with water. HCl + H 2 O ⟶ H 3 O + + Cl − However, the speciation of hydrochloric acid is more complicated than this simple equation implies; the structure of bulk water is infamously complex, the formula H3O+ is a gross oversimplification of the true nature of the solvated proton, H+, present in hydrochloric acid. A combined IR, Raman, X-ray and neutron diffraction study of concentrated solutions of hydrochloric acid revealed that the primary form of H+ in these solutions is H5O2+, along with the chloride anion, is hydrogen-bonded to neighboring wa