Dichloromethane is a geminal organic compound with the formula CH2Cl2. This colorless, volatile liquid with a moderately sweet aroma is used as a solvent. Although it is not miscible with water, it is polar, miscible with many organic solvents. Natural sources of dichloromethane include oceanic sources, macroalgae and volcanoes. However, the majority of dichloromethane in the environment is the result of industrial emissions. DCM is produced by treating either chloromethane or methane with chlorine gas at 400–500 °C. At these temperatures, both methane and chloromethane undergo a series of reactions producing progressively more chlorinated products. In this way, an estimated 400,000 tons were produced in the US, Japan in 1993. CH4 + Cl2 → CH3Cl + HCl CH3Cl + Cl2 → CH2Cl2 + HCl CH2Cl2 + Cl2 → CHCl3 + HCl CHCl3 + Cl2 → CCl4 + HClThe output of these processes is a mixture of chloromethane, dichloromethane and carbon tetrachloride; these compounds are separated by distillation. DCM was first prepared in 1839 by the French chemist Henri Victor Regnault, who isolated it from a mixture of chloromethane and chlorine, exposed to sunlight.
DCM's volatility and ability to dissolve a wide range of organic compounds makes it a useful solvent for many chemical processes. Carbon diselenide is produced by reacting selenium powder with dichloromethane vapor near 550°C. 2 Se + CH2Cl2 → CSe2 + 2 HClIt is used as a paint stripper and a degreaser. In the food industry, it has been used to decaffeinate coffee and tea as well as to prepare extracts of hops and other flavorings, its volatility has led to its use as an aerosol spray propellant and as a blowing agent for polyurethane foams. The chemical compound's low boiling point allows the chemical to function in a heat engine that can extract mechanical energy from small temperature differences. An example of a DCM heat engine is the drinking bird; the toy works at room temperature. DCM chemically welds certain plastics. For example, it is used to seal the casing of electric meters. Sold as a main component of plastic welding adhesives, it is used extensively by model building hobbyists for joining plastic components together.
It is referred to as "Di-clo." It is used in the garment printing industry for removal of heat-sealed garment transfers, its volatility is exploited in novelty items: bubble lights and jukebox displays. DCM is used in the material testing field of civil engineering. DCM is the least toxic of the simple chlorohydrocarbons, but it is not without health risks, as its high volatility makes it an acute inhalation hazard, it can be absorbed through the skin. Symptoms of acute overexposure to dichloromethane via inhalation include difficulty concentrating, fatigue, headaches, numbness and irritation of the upper respiratory tract and eyes. More severe consequences can include suffocation, loss of consciousness and death. DCM is metabolized by the body to carbon monoxide leading to carbon monoxide poisoning. Acute exposure by inhalation has resulted in optic hepatitis. Prolonged skin contact can result in DCM dissolving some of the fatty tissues in skin, resulting in skin irritation or chemical burns, it may be carcinogenic, as it has been linked to cancer of the lungs and pancreas in laboratory animals.
Other animal studies showed salivary gland cancer. Research is not yet clear as to. DCM crosses the placenta. Fetal toxicity in women who are exposed to it during pregnancy, has not been proven. In animal experiments, it was fetotoxic at doses that were maternally toxic but no teratogenic effects were seen. In people with pre-existing heart problems, exposure to DCM can cause abnormal heart rhythms and/or heart attacks, sometimes without any other symptoms of overexposure. People with existing liver, nervous system, or skin problems may worsen after exposure to methylene chloride. In many countries, products containing DCM must carry labels warning of its health risks. In February 2013, the U. S. Occupational Safety and Health Administration and the National Institute for Occupational Safety and Health warned that at least 14 bathtub refinishers have died since 2000 from DCM exposure; these workers had been working alone, in poorly ventilated bathrooms, with inadequate or no respiratory protection, no training about the hazards of DCM.
OSHA has since issued a DCM standard. In the European Union, the European Parliament voted in 2009 to ban the use of DCM in paint-strippers for consumers and many professionals; the ban took effect in December 2010. In Europe, the Scientific Committee on Occupational Exposure Limit Values recommends for DCM an occupational exposure limit of 100 ppm and a short-term exposure limit of 200 ppm. Concerns about its health effects have led to a search for alternatives in many of these applications. On March 15, 2019, the U. S. Environmental Protection Agency issued a final rule to prohibit the manufacture and distribution of methylene chloride in all paint removers for consumer use, effective in 180 days. Dichloromethane is not classified as an ozone-depleting substance by the Montreal Protocol; the U. S. Clean Air Act does not regulate dichloromethane as an ozone depleter. According to the EPA, the atmospheric lifetime of dichloromethane is short, such that the substance decomposes before reaching the ozone layer.
Ozone concentrations measured at the midlatitudes from the g
Nitromethane is an organic compound with the chemical formula CH3NO2. It is the simplest organic nitro compound, it is a polar liquid used as a solvent in a variety of industrial applications such as in extractions, as a reaction medium, as a cleaning solvent. As an intermediate in organic synthesis, it is used in the manufacture of pharmaceuticals, explosives and coatings. Nitromethane is used as a fuel additive in e.g.. Top Fuel drag racing and miniature internal combustion engines in radio control, control line and free flight model aircraft. Nitromethane is produced industrially by combining propane and nitric acid in the gas phase at 350–450 °C; this exothermic reaction produces the four industrially significant nitroalkanes: nitromethane, nitroethane, 1-nitropropane, 2-nitropropane. The reaction involves free radicals, including the alkoxyl radicals of the type CH3CH2CH2O, which arise via homolysis of the corresponding nitrite ester; these alkoxy radicals are susceptible to C—C fragmentation reactions, which explains the formation of a mixture of products.
It can be prepared in other methods. The reaction of sodium chloroacetate with sodium nitrite in aqueous solution produces this compound: ClCH2COONa + NaNO2 + H2O → CH3NO2 + NaCl + NaHCO3 The principal use of nitromethane is as a stabilizer for chlorinated solvents, which are used in dry cleaning, semiconductor processing, degreasing, it is used most as a solvent or dissolving agent for acrylate monomers, such as cyanoacrylates. It is used as a fuel in some forms of racing, it can be used as an explosive. This type of mixture is called PLX. Another used mixtures are ANNM and ANNMAl – explosive mixtures of ammonium nitrate and aluminium powder. Nitromethane is a acidic carbon acid, it has a pKa of 17.2 in DMSO solution. This value indicates an aqueous pKa of about 11; the reason of that being so acidic is due to the resonance structure below: It is slow to deprotonate. Protonation of the conjugate base O2NCH2-, nearly isosteric with nitrate, occurs at oxygen. In organic synthesis nitromethane is employed as a one carbon building block.
Its acidity allows it to undergo deprotonation, enabling condensation reactions analogous to those of carbonyl compounds. Thus, under base catalysis, nitromethane adds to aldehydes in 1,2-addition in the nitroaldol reaction; some important derivatives include the pesticides chloropicrin, beta-nitrostyrene, trisnitromethane. Reduction of the latter gives trisaminomethane, 3CNH2, better known as tris, a used buffer. In more specialized organic synthesis, nitromethane serves as a Michael donor, adding to α,β-unsaturated carbonyl compounds via 1,4-addition in the Michael reaction. Nitromethane is used as a fuel in motor racing drag racing, as well as for radio-controlled models. In this context, nitromethane is referred to as "nitro", is the principal ingredient for fuel used in the "Top Fuel" category of drag racing; the oxygen content of nitromethane enables it to burn with much less atmospheric oxygen. 4 CH3NO2 + 3 O2 → 4 CO2 + 6 H2O + 2 N2The amount of air required to burn 1 kg of gasoline is 14.7 kg, but only 1.7 kg of air is required for 1 kg of nitromethane.
Since an engine's cylinder can only contain a limited amount of air on each stroke, 8.6 times more nitromethane than gasoline can be burned in one stroke. Nitromethane, has a lower specific energy: gasoline provides about 42–44 MJ/kg, whereas nitromethane provides only 11.3 MJ/kg. This analysis indicates that nitromethane generates about 2.3 times the power of gasoline when combined with a given amount of oxygen. Nitromethane can be used as a monopropellant, i.e. a fuel that burns without added oxygen. The following equation describes this process: 2 CH3NO2 → 2 CO + 2 H2O + H2 + N2Nitromethane has a laminar combustion velocity of 0.5 m/s, somewhat higher than gasoline, thus making it suitable for high-speed engines. It has a somewhat higher flame temperature of about 2,400 °C; the high heat of vaporization of 0.56 MJ/kg together with the high fuel flow provides significant cooling of the incoming charge, resulting in reasonably low temperatures Nitromethane is used with rich air–fuel mixtures because it provides power in the absence of atmospheric oxygen.
When rich air–fuel mixtures are used and carbon monoxide are two of the combustion products. These gases ignite, sometimes spectacularly, as the very rich mixtures of the still burning fuel exits the exhaust ports. Rich mixtures are necessary to reduce the temperature of combustion chamber hot parts in order to control pre-ignition and subsequent detonation. Operational details depend on engine characteristics. A small amount of hydrazine blended in nitromethane can increase the power output further. With nitromethane, hydrazine forms an explosive salt, again a monopropellant; this unstable mixture poses a severe safety hazard and the Academy of Model Aeronautics does not permit its use in competitions. In model aircraft and car glow fuel, the primary ingredient is methanol with some nitromethane (0% to 65%, but over 30%, 10–20% lubricants. Moderate amounts of nitromethane tend to increase the power created by the engine, making the engine easier t
In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end. Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms. A polar molecule with two or more polar bonds must have a geometry, asymmetric in at least one direction, so that the bond dipoles do not cancel each other. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension and melting and boiling points. Not all atoms attract electrons with the same force; the amount of "pull" an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities – such as fluorine and nitrogen – exert a greater pull on electrons than atoms with lower electronegativities such as alkali metals and alkaline earth metals. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity.
Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is smaller than a fundamental charge, they are called partial charges, denoted as δ+ and δ−; these symbols were introduced by Sir Christopher Ingold and Dr. Edith Hilda Ingold in 1926; the bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges. These dipoles within molecules can interact with dipoles in other molecules, creating dipole-dipole intermolecular forces. Bonds can fall between one of two extremes – being nonpolar or polar. A nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A polar bond is more called an ionic bond, occurs when the difference between electronegativities is large enough that one atom takes an electron from the other.
The terms "polar" and "nonpolar" are applied to covalent bonds, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used. Bond polarity is divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the Pauling scale: Nonpolar bonds occur when the difference in electronegativity between the two atoms is less than 0.5 Polar bonds occur when the difference in electronegativity between the two atoms is between 0.5 and 2.0 Ionic bonds occur when the difference in electronegativity between the two atoms is greater than 2.0Pauling based this classification scheme on the partial ionic character of a bond, an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond, predominantly ionic.
As a quantum-mechanical description, Pauling proposed that the wave function for a polar molecule AB is a linear combination of wave functions for covalent and ionic molecules: ψ = aψ + bψ. The amount of covalent and ionic character depends on the values of the squared coefficients a2 and b2. While the molecules can be described as "polar covalent", "nonpolar covalent", or "ionic", this is a relative term, with one molecule being more polar or more nonpolar than another. However, the following properties are typical of such molecules. A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms. A molecule may be polar either as a result of polar bonds due to differences in electronegativity as described above, or as a result of an asymmetric arrangement of nonpolar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital. A polar molecule has a net dipole as a result of the opposing charges from polar bonds arranged asymmetrically.
Water is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other. The dipoles do not cancel out resulting in a net dipole. Due to the polar nature of the water molecule itself, polar molecules are able to dissolve in water. Other examples include sugars, which have many polar oxygen–hydrogen groups and are overall polar. If the bond dipole moments of the molecule do not cancel, the molecule is polar. For example, the water molecule contains two polar O−H bonds in a bent geometry; the bond dipole moments do not cancel, so that the molecule forms a molecular dipole with its negative pole at the oxygen and its positive pole midway between the two hydrogen atoms. In the figure each bond joins the central O atom with a negative charge to an H atom with a positive charge; the hydrogen fluoride, HF, molecule is polar by virtue of polar covalent bonds – in the covalent bond electrons are displaced toward the more electronegative fluorine atom.
Ammonia, NH3, molecule. The molecule has two lone electrons in an orbital, that points towards the fourth apex of the approximate tetrahedron; this orbital is not participating in covalent bonding.
An acid is a molecule or ion capable of donating a hydron, or, capable of forming a covalent bond with an electron pair. The first category of acids is the proton donors or Brønsted acids. In the special case of aqueous solutions, proton donors form the hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid contains a hydrogen atom bonded to a chemical structure, still energetically favorable after loss of H+. Aqueous Arrhenius acids have characteristic properties which provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, react with bases and certain metals to form salts; the word acid is derived from the Latin acidus/acēre meaning sour. An aqueous solution of an acid has a pH less than 7 and is colloquially referred to as'acid', while the strict definition refers only to the solute. A lower pH means a higher acidity, thus a higher concentration of positive hydrogen ions in the solution.
Chemicals or substances having the property of an acid are said to be acidic. Common aqueous acids include hydrochloric acid, acetic acid, sulfuric acid, citric acid; as these examples show, acids can be solutions or pure substances, can be derived from acids that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid; the second category of acids are Lewis acids. An example is boron trifluoride, whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia. Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons into the solution, which accept electron pairs. However, hydrogen chloride, acetic acid, most other Brønsted-Lowry acids cannot form a covalent bond with an electron pair and are therefore not Lewis acids.
Conversely, many Lewis acids are not Brønsted-Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists always refer to a Lewis acid explicitly as a Lewis acid. Modern definitions are concerned with the fundamental chemical reactions common to all acids. Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, so the Arrhenius and Brønsted-Lowry definitions are the most relevant; the Brønsted-Lowry definition is the most used definition. Hydronium ions are acids according to all three definitions. Although alcohols and amines can be Brønsted-Lowry acids, they can function as Lewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms; the Swedish chemist Svante Arrhenius attributed the properties of acidity to hydrogen ions or protons in 1884. An Arrhenius acid is a substance that, when added to water, increases the concentration of H+ ions in the water. Note that chemists write H+ and refer to the hydrogen ion when describing acid-base reactions but the free hydrogen nucleus, a proton, does not exist alone in water, it exists as the hydronium ion, H3O+.
Thus, an Arrhenius acid can be described as a substance that increases the concentration of hydronium ions when added to water. Examples include molecular substances such as acetic acid. An Arrhenius base, on the other hand, is a substance which increases the concentration of hydroxide ions when dissolved in water; this decreases the concentration of hydronium because the ions react to form H2O molecules: H3O+ + OH− ⇌ H2O + H2ODue to this equilibrium, any increase in the concentration of hydronium is accompanied by a decrease in the concentration of hydroxide. Thus, an Arrhenius acid could be said to be one that decreases hydroxide concentration, while an Arrhenius base increases it. In an acidic solution, the concentration of hydronium ions is greater than 10−7 moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a pH of less than 7. While the Arrhenius concept is useful for describing many reactions, it is quite limited in its scope.
In 1923 chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently recognized that acid-base reactions involve the transfer of a proton. A Brønsted-Lowry acid is a species. Brønsted-Lowry acid-base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid, the organic acid that gives vinegar its characteristic taste: CH3COOH + H2O ⇌ CH3COO− + H3O+ CH3COOH + NH3 ⇌ CH3COO− + NH+4Both theories describe the first reaction: CH3COOH acts as an Arrhenius acid because it acts as a source of H3O+ when dissolved in water, it acts as a Brønsted acid by donating a proton to water. In the second example CH3COOH undergoes the same transformation, in this case donating a proton to ammonia, but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. CH3COOH is
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
Diethyl ether, or ether, is an organic compound in the ether class with the formula 2O, sometimes abbreviated as Et2O. It is a colorless volatile flammable liquid, it is used as a solvent in laboratories and as a starting fluid for some engines. It was used as a general anesthetic, until non-flammable drugs were developed, such as halothane, it has been used as a recreational drug to cause intoxication. Most diethyl ether is produced as a byproduct of the vapor-phase hydration of ethylene to make ethanol; this process uses solid-supported phosphoric acid catalysts and can be adjusted to make more ether if the need arises. Vapor-phase dehydration of ethanol over some alumina catalysts can give diethyl ether yields of up to 95%. Diethyl ether can be prepared both in laboratories and on an industrial scale by the acid ether synthesis. Ethanol is mixed with a strong acid sulfuric acid, H2SO4; the acid dissociates in the aqueous environment producing hydronium ions, H3O+. A hydrogen ion protonates the electronegative oxygen atom of the ethanol, giving the ethanol molecule a positive charge: CH3CH2OH + H3O+ → CH3CH2OH2+ + H2OA nucleophilic oxygen atom of unprotonated ethanol displaces a water molecule from the protonated ethanol molecule, producing water, a hydrogen ion and diethyl ether.
CH3CH2OH2+ + CH3CH2OH → H2O + H+ + CH3CH2OCH2CH3This reaction must be carried out at temperatures lower than 150 °C in order to ensure that an elimination product is not a product of the reaction. At higher temperatures, ethanol will dehydrate to form ethylene; the reaction to make diethyl ether is reversible, so an equilibrium between reactants and products is achieved. Getting a good yield of ether requires that ether be distilled out of the reaction mixture before it reverts to ethanol, taking advantage of Le Chatelier's principle. Another reaction that can be used for the preparation of ethers is the Williamson ether synthesis, in which an alkoxide performs a nucleophilic substitution upon an alkyl halide, it is important as a solvent in the production of cellulose plastics such as cellulose acetate. Diethyl ether has a high cetane number of 85–96 and is used as a starting fluid, in combination with petroleum distillates for gasoline and Diesel engines because of its high volatility and low flash point.
Ether starting fluid is sold and used in countries with cold climates, as it can help with cold starting an engine at sub-zero temperatures. For the same reason it is used as a component of the fuel mixture for carbureted compression ignition model engines. In this way diethyl ether is similar to one of its precursors, ethanol. Diethyl ether is a common laboratory aprotic solvent, it has limited solubility in water and dissolves 1.5 g/100 g water at 25 °C. This, coupled with its high volatility, makes it ideal for use as the non-polar solvent in liquid-liquid extraction; when used with an aqueous solution, the diethyl ether layer is on top as it has a lower density than the water. It is a common solvent for the Grignard reaction in addition to other reactions involving organometallic reagents. Due to its application in the manufacturing of illicit substances, it is listed in the Table II precursor under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances as well as substances such as acetone and sulfuric acid.
William T. G. Morton participated in a public demonstration of ether anesthesia on October 16, 1846 at the Ether Dome in Boston, Massachusetts. However, Crawford Williamson Long, is now known to have demonstrated its use as a general anesthetic in surgery to officials in Georgia, as early as March 30, 1842, Long publicly demonstrated ether's use as a surgical anesthetic on six occasions before the Boston demonstration. British doctors were aware of the anesthetic properties of ether as early as 1840 where it was prescribed in conjunction with opium. Diethyl ether supplanted the use of chloroform as a general anesthetic due to ether's more favorable therapeutic index, that is, a greater difference between an effective dose and a toxic dose. Diethyl ether increases tracheobronchial secretions. Diethyl ether could be mixed with other anesthetic agents such as chloroform to make C. E. mixture, or chloroform and alcohol to make A. C. E. Mixture. In the 21st century, ether is used; the use of flammable ether was displaced by nonflammable fluorinated hydrocarbon anesthetics.
Halothane was the first such anesthetic developed and other used inhaled anesthetics, such as isoflurane and sevoflurane, are halogenated ethers. Diethyl ether was found to have undesirable side effects, such as post-anesthetic nausea and vomiting. Modern anesthetic agents reduce these side effects. Prior to 2005 it was on the World Health Organization's List of Essential Medicines for use as an anesthetic. Ether was once used in pharmaceutical formulations. A mixture of alcohol and ether, one part of diethyl ether and three parts of ethanol, was known as "Spirit of ether", Hoffman's Anodyne or Hoffman's Drops. In the United States this concoction was removed from the Pharmacopeia at some point prior to June 1917, as a study published by William Procter, Jr. in the American Journal of Pharmacy as early as 1852 showed that there were differences in formulation to be found between commercial manufacturers, between international pharmacopoeia, from Hoffman's original recipe. The anesthetic and intoxicating effects of ether have made it a recreational drug.
Diethyl ether in anesthetic dosage is an inhalant which has a long history
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