A carboxylic acid is an organic compound that contains a carboxyl group. The general formula of a carboxylic acid is R–COOH, with R referring to the rest of the molecule. Carboxylic acids occur widely. Important examples include acetic acid. Deprotonation of a carboxyl group gives a carboxylate anion. Important carboxylate salts are soaps. Carboxylic acids are identified by their trivial names, they have the suffix -ic acid. IUPAC-recommended names exist. For example, butyric acid is butanoic acid by IUPAC guidelines. For nomenclature of complex molecules containing a carboxylic acid, the carboxyl can be considered position one of the parent chain if there are other substituents, for example, 3-chloropropanoic acid. Alternately, it can be named as a "carboxy" or "carboxylic acid" substituent on another parent structure, for example, 2-carboxyfuran; the carboxylate anion of a carboxylic acid is named with the suffix -ate, in keeping with the general pattern of -ic acid and -ate for a conjugate acid and its conjugate base, respectively.
For example, the conjugate base of acetic acid is acetate. Carboxylic acids are polar; because they are both hydrogen-bond acceptors and hydrogen-bond donors, they participate in hydrogen bonding. Together the hydroxyl and carbonyl group forms the functional group carboxyl. Carboxylic acids exist as dimers in nonpolar media due to their tendency to "self-associate". Smaller carboxylic acids are soluble in water, whereas higher carboxylic acids have limited solubility due to the increasing hydrophobic nature of the alkyl chain; these longer chain acids tend to be rather soluble in less-polar solvents such as ethers and alcohols. Hydrophobic carboxylic acids react aqueous sodium hydroxide to give water soluble sodium salts. For example, enathic acid has a small solubility in water, but its sodium salt is soluble in water: Carboxylic acids tend to have higher boiling points than water, not only because of their increased surface area, but because of their tendency to form stabilised dimers through hydrogen bonds.
For boiling to occur, either the dimer bonds must be broken or the entire dimer arrangement must be vaporised, both of which increase the enthalpy of vaporization requirements significantly. Carboxylic acids are Brønsted -- Lowry acids, they are the most common type of organic acid. Carboxylic acids are weak acids, meaning that they only dissociate into H3O+ cations and RCOO− anions in neutral aqueous solution. For example, at room temperature, in a 1-molar solution of acetic acid, only 0.4% of the acid are dissociated. Electron-withdrawing substituents, such as -CF3 group, give stronger acids. Electron-donating substituents give weaker acids Deprotonation of carboxylic acids gives carboxylate anions; each of the carbon–oxygen bonds in the carboxylate anion has a partial double-bond character. The carbonyl carbon's partial positive charge is weakened by the -1/2 negative charges on the 2 oxygen atoms. Carboxylic acids have strong sour odors. Esters of carboxylic acids tend to have pleasant odors, many are used in perfume.
Carboxylic acids are identified as such by infrared spectroscopy. They exhibit a sharp band associated with vibration of the C–O vibration bond between 1680 and 1725 cm−1. A characteristic νO–H band appears as a broad peak in the 2500 to 3000 cm−1 region. By 1H NMR spectrometry, the hydroxyl hydrogen appears in the 10–13 ppm region, although it is either broadened or not observed owing to exchange with traces of water. Many carboxylic acids are produced industrially on a large scale, they are pervasive in nature. Esters of fatty acids are the main components of lipids and polyamides of aminocarboxylic acids are the main components of proteins. Carboxylic acids are used in the production of polymers, pharmaceuticals and food additives. Industrially important carboxylic acids include acetic acid and methacrylic acids, adipic acid, citric acid, ethylenediaminetetraacetic acid, fatty acids, maleic acid, propionic acid, terephthalic acid. In general, industrial routes to carboxylic acids differ from those used on smaller scale because they require specialized equipment.
Carbonylation of alcohols as illustrated by the Cativa process for production of acetic acid. Formic acid is prepared by a different carbonylation pathway starting from methanol. Oxidation of aldehydes with air using cobalt and manganese catalysts; the required aldehydes are obtained from alkenes by hydroformylation. Oxidation of hydrocarbons using air. For simple alkanes, this method is inexpensive but not selective enough to be useful. Allylic and benzylic compounds undergo more selective oxidations. Alkyl groups on a benzene ring are oxidized to the carboxylic acid, regardless of its chain length. Benzoic acid from toluene, terephthalic acid from para-xylene, phthalic acid from ortho-xylene are illustrative large-scale conversions. Acrylic acid is generated from propene. Base-cata
Citric acid is a weak organic acid that has the chemical formula C6H8O7. It occurs in citrus fruits. In biochemistry, it is an intermediate in the citric acid cycle, which occurs in the metabolism of all aerobic organisms. More than a million tons of citric acid are manufactured every year, it is used as an acidifier, as a flavoring and chelating agent. A citrate is a derivative of citric acid. An example of the former, a salt is trisodium citrate; when part of a salt, the formula of the citrate ion is written as C6H5O3−7 or C3H5O3−3. Citric acid exists in greater than trace amounts in a variety of fruits and vegetables, most notably citrus fruits. Lemons and limes have high concentrations of the acid; the concentrations of citric acid in citrus fruits range from 0.005 mol/L for oranges and grapefruits to 0.30 mol/L in lemons and limes. Within species, these values vary depending on the cultivar and the circumstances in which the fruit was grown. Industrial-scale citric acid production first began in 1890 based on the Italian citrus fruit industry, where the juice was treated with hydrated lime to precipitate calcium citrate, isolated and converted back to the acid using diluted sulfuric acid.
In 1893, C. Wehmer discovered. However, microbial production of citric acid did not become industrially important until World War I disrupted Italian citrus exports. In 1917, American food chemist James Currie discovered certain strains of the mold Aspergillus niger could be efficient citric acid producers, the pharmaceutical company Pfizer began industrial-level production using this technique two years followed by Citrique Belge in 1929. In this production technique, still the major industrial route to citric acid used today, cultures of A. niger are fed on a sucrose or glucose-containing medium to produce citric acid. The source of sugar is corn steep liquor, hydrolyzed corn starch or other inexpensive sugary solutions. After the mold is filtered out of the resulting solution, citric acid is isolated by precipitating it with calcium hydroxide to yield calcium citrate salt, from which citric acid is regenerated by treatment with sulfuric acid, as in the direct extraction from citrus fruit juice.
In 1977, a patent was granted to Lever Brothers for the chemical synthesis of citric acid starting either from aconitic or isocitrate/alloisocitrate calcium salts under high pressure conditions. This produced citric acid in near quantitative conversion under what appeared to be a reverse non-enzymatic Krebs cycle reaction. In 2007, worldwide annual production stood at 1,600,000 tons. More than 50% of this volume was produced in China. More than 50% was used as an acidity regulator in beverages, some 20% in other food applications, 20% for detergent applications and 10% for related applications other than food, such as cosmetics, pharmaceutics and in the chemical industry. Citric acid was first isolated in 1784 by the chemist Carl Wilhelm Scheele, who crystallized it from lemon juice, it can exist either as a monohydrate. The anhydrous form crystallizes from hot water, while the monohydrate forms when citric acid is crystallized from cold water; the monohydrate can be converted to the anhydrous form at about 78 °C.
Citric acid dissolves in absolute ethanol at 15 °C. It decomposes with loss of carbon dioxide above about 175 °C. Citric acid is considered to be a tribasic acid, with pKa values, extrapolated to zero ionic strength, of 5.21, 4.28 and 2.92 at 25 °C. The pKa of the hydroxyl group has been found, by means of 13C NMR spectroscopy, to be 14.4. The speciation diagram shows that solutions of citric acid are buffer solutions between about pH 2 and pH 8. In biological systems around pH 7, the two species present are the citrate ion and mono-hydrogen citrate ion; the SSC 20X hybridization buffer is an example in common use. Tables compiled for biochemical studies are available. On the other hand, the pH of a 1 mM solution of citric acid will be about 3.2. The pH of fruit juices from citrus fruits like oranges and lemons depends on the citric acid concentration, being lower for higher acid concentration and conversely. Acid salts of citric acid can be prepared by careful adjustment of the pH before crystallizing the compound.
See, for example, sodium citrate. The citrate ion forms complexes with metallic cations; the stability constants for the formation of these complexes are quite large because of the chelate effect. It forms complexes with alkali metal cations. However, when a chelate complex is formed using all three carboxylate groups, the chelate rings have 7 and 8 members, which are less stable thermodynamically than smaller chelate rings. In consequence, the hydroxyl group can be deprotonated, forming part of a more stable 5-membered ring, as in ammonium ferric citrate, 5Fe2·2H2O. Citric acid can be esterified at one or more of the carboxylic acid functional groups on the molecule, to form any of a variety of mono-, di-, tri-, mixed esters. Citrate is an intermediate in the TCA cycle, a central metabolic pathway for animals and bacteria. Citrate synthase catalyzes the condensation of oxaloacetate with acetyl CoA to form citrate. Citrate acts as the substrate for aconitase and is converted into aconitic acid.
The cycle ends with regeneration of oxaloacetate. This series
Oxalic acid is an organic compound with the formula C2H2O4. It is a colorless crystalline solid, its condensed formula is HOOCCOOH. Its acid strength is much greater than that of acetic acid. Oxalic acid is a reducing agent and its conjugate base, known as oxalate, is a chelating agent for metal cations. Oxalic acid occurs as the dihydrate with the formula C2H2O4·2H2O, it occurs in many foods, but excessive ingestion of oxalic acid or prolonged skin contact can be dangerous. Its name comes from the fact that early investigators isolated oxalic acid from wood-sorrel flowering plants; the preparation of salts of oxalic acid from plants had been known, at the latest, since 1745, when the Dutch botanist and physician Herman Boerhaave isolated a salt from sorrel. By 1773, François Pierre Savary of Fribourg, Switzerland had isolated oxalic acid from its salt in sorrel. In 1776, Swedish chemists Carl Wilhelm Scheele and Torbern Olof Bergman produced oxalic acid by reacting sugar with concentrated nitric acid.
By 1784, Scheele had shown that oxalic acid from natural sources were identical. In 1824, the German chemist Friedrich Wöhler obtained oxalic acid by reacting cyanogen with ammonia in aqueous solution; this experiment may represent the first synthesis of a natural product. Oxalic acid is manufactured by the oxidation of carbohydrates or glucose using nitric acid or air in the presence of vanadium pentoxide. A variety of precursors can be used including glycolic ethylene glycol. A newer method entails oxidative carbonylation of alcohols to give the diesters of oxalic acid: 4 ROH + 4 CO + O2 → 2 2 + 2 H2OThese diesters are subsequently hydrolyzed to oxalic acid. 120,000 tonnes are produced annually. Oxalic acid was obtained by using caustics, such as sodium or potassium hydroxide, on sawdust. Although it can be purchased, oxalic acid can be prepared in the laboratory by oxidizing sucrose using nitric acid in the presence of a small amount of vanadium pentoxide as a catalyst; the hydrated solid can be dehydrated by azeotropic distillation.
Developed in the Netherlands, an electrocatalysis by a copper complex helps reduce carbon dioxide to oxalic acid. Anhydrous oxalic acid exists as two polymorphs; because the anhydrous material is both acidic and hydrophilic, it is used in esterifications. Oxalic acid is a strong acid, despite being a carboxylic acid: Oxalic acid undergoes many of the reactions characteristic of other carboxylic acids, it forms esters such as dimethyl oxalate. It forms. Oxalate, the conjugate base of oxalic acid, is an excellent ligand for metal ions, e.g. the drug oxaliplatin. Oxalic acid and oxalates can be oxidized by permanganate in an autocatalytic reaction. At least two pathways exist for the enzyme-mediated formation of oxalate. In one pathway, oxaloacetate, a component of the Krebs citric acid cycle, is hydrolyzed to oxalate and acetic acid by the enzyme oxaloacetase: 2− + H2O → C2O2−4 + CH3CO−2 + H+It arises from the dehydrogenation of glycolic acid, produced by the metabolism of ethylene glycol. Calcium oxalate is the most common component of kidney stones.
Early investigators isolated oxalic acid from wood-sorrel. Members of the spinach family and the brassicas are high in oxalates, as are sorrel and umbellifers like parsley. Rhubarb leaves contain about 0.5% oxalic acid, jack-in-the-pulpit contains calcium oxalate crystals. The Virginia creeper, a common decorative vine, produces oxalic acid in its berries as well as oxalate crystals in the sap, in the form of raphides. Bacteria produce oxalates from oxidation of carbohydrates. Plants of the genus Fenestraria produce optical fibers made from crystalline oxalic acid to transmit light to subterranean photosynthetic sites. Carambola known as starfruit contains oxalic acid along with caramboxin; the formation of occurring calcium oxalate patinas on certain limestone and marble statues and monuments has been proposed to be caused by the chemical reaction of the carbonate stone with oxalic acid secreted by lichen or other microorganisms. Oxidized bitumen or bitumen exposed to gamma rays contains oxalic acid among its degradation products.
Oxalic acid may increase the leaching of radionuclides conditioned in bitumen for radioactive waste disposal. The conjugate base of oxalic acid is the hydrogenoxalate anion, its conjugate base is a competitive inhibitor of the lactate dehydrogenase enzyme. LDH catalyses the conversion of pyruvate to lactic acid oxidising the coenzyme NADH to NAD+ and H+ concurrently. Restoring NAD+ levels is essential to the continuation of anaerobic energy metabolism through glycolysis; as cancer cells preferentially use anaerobic metabolism inhibition of LDH has been shown to inhibit tumor formation and growth, thus is an interesting potential course of cancer treatment. About 25% of produced oxalic acid will be used as a mordant in dyeing processes, it is used in bleaches for pulpwood. It is used in baking powder and as a third reagent in silica analysis instruments. Oxalic acid's main applications include cleani
Lewis acids and bases
A Lewis acid is a chemical species that contains an empty orbital, capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base is any species that has a filled orbital containing an electron pair, not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base. Trimethylborane is a Lewis acid. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, the lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3; the terminology refers to the contributions of Gilbert N. Lewis; the terms nucleophile and electrophile are more or less interchangeable with Lewis base and Lewis acid, respectively. However, these terms their abstract noun forms nucleophilicity and electrophilicity, emphasize the kinetic aspect of reactivity, while the Lewis basicity and Lewis acidity emphasize the thermodynamic aspect of Lewis adduct formation.
In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me3B←NH3. Some sources indicate the Lewis base with a pair of dots, which allows consistent representation of the transition from the base itself to the complex with the acid: Me3B +:NH3 → Me3B:NH3A center dot may be used to represent a Lewis adduct, such as Me3B•NH3. Another example is boron trifluoride diethyl etherate, BF3•Et2O. Although there have been attempts to use computational and experimental energetic criteria to distinguish dative bonding from non-dative covalent bonds, for the most part, the distinction makes note of the source of the electron pair, dative bonds, once formed, behave as other covalent bonds do, though they have considerable polar character. Moreover, in some cases, the use of the dative bond arrow is just a notational convenience for avoiding the drawing of formal charges.
In general, the donor–acceptor bond is viewed as somewhere along a continuum between idealized covalent bonding and ionic bonding. Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR3 where R can be an organic substituent or a halide. For the purposes of discussion complex compounds such as Et3Al2Cl3 and AlCl3 are treated as trigonal planar Lewis acids. Metal ions such as Na+, Mg2+, Ce3+, which are invariably complexed with additional ligands, are sources of coordinatively unsaturated derivatives that form Lewis adducts upon reaction with a Lewis base. Other reactions might be referred to as "acid-catalyzed" reactions; some compounds, such as H2O, are both Lewis acids and Lewis bases, because they can either accept a pair of electrons or donate a pair of electrons, depending upon the reaction. Lewis acids are diverse. Simplest are those, but more common are those. Examples of Lewis acids based on the general definition of electron pair acceptor include: the proton and acidic compounds onium ions, such as NH4+ and H3O+ high oxidation state transition metal cations, e.g. Fe3+.
Again, the description of a Lewis acid is used loosely. For example, in solution, bare protons do not exist; some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior: BF3 + F− → BF4−In this adduct, all four fluoride centres are equivalent. BF3 + OMe2 → BF3OMe2Both BF4− and BF3OMe2 are Lewis base adducts of boron trifluoride. In many cases, the adducts violate the octet rule, such as the triiodide anion: I2 + I− → I3−The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I2. In some cases, the Lewis acid is capable of binding two Lewis base, a famous example being the formation of hexafluorosilicate: SiF4 + 2 F− → SiF62− Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX3, but as aggregates and polymers that must be degraded by the Lewis base.
A simpler case is the formation of adducts of borane. Monomeric BH3 does not exist appreciably, so the adducts of borane are generated by degradation of diborane: B2H6 + 2 H− → 2 BH4−In this case, an intermediate B2H7− can be isolated. Many metal complexes serve as Lewis acids, but only after dissociating a more weakly bound Lewis base water. 2+ + 6 NH3 → 2+ + 6 H2O The proton is one of the strongest but is one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is solvated (bound to solvent
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
Phenol is an aromatic organic compound with the molecular formula C6H5OH. It is a white crystalline solid, volatile; the molecule consists of a phenyl group bonded to a hydroxy group. It requires careful handling due to its propensity for causing chemical burns. Phenol was first extracted from coal tar, it is an important industrial commodity as a precursor to useful compounds. It is used to synthesize plastics and related materials. Phenol and its chemical derivatives are essential for production of polycarbonates, Bakelite, detergents, herbicides such as phenoxy herbicides, numerous pharmaceutical drugs. Phenol is an organic compound appreciably soluble in water, with about 84.2 g dissolving in 1000 mL. Homogeneous mixtures of phenol and water at phenol to water mass ratios of ~2.6 and higher are possible. The sodium salt of phenol, sodium phenoxide, is far more water-soluble. Phenol is weakly acidic and at high pHs gives the phenolate anion C6H5O−: PhOH ⇌ PhO− + H+ Compared to aliphatic alcohols, phenol is about 1 million times more acidic, although it is still considered a weak acid.
It reacts with aqueous NaOH to lose H+, giving the salt sodium phenoxide, whereas most alcohols react only partially. One explanation for the increased acidity over alcohols is resonance stabilization of the phenoxide anion by the aromatic ring. In this way, the negative charge on oxygen is delocalized on to the ortho and para carbon atoms through the pi system. An alternative explanation involves the sigma framework, postulating that the dominant effect is the induction from the more electronegative sp2 hybridised carbons. In support of the second explanation, the pKa of the enol of acetone in water is 10.9, making it only less acidic than phenol. Thus, the greater number of resonance structures available to phenoxide compared to acetone enolate seems to contribute little to its stabilization. However, the situation changes. A recent in silico comparison of the gas phase acidities of the vinylogues of phenol and cyclohexanol in conformations that allow for or exclude resonance stabilization leads to the inference that about 1⁄3 of the increased acidity of phenol is attributable to inductive effects, with resonance accounting for the remaining difference.
The phenoxide anion has a similar nucleophilicity to free amines, with the further advantage that its conjugate acid does not become deactivated as a nucleophile in moderately acidic conditions. Phenolate esters are more stable toward hydrolysis than acid anhydrides and acyl halides but are sufficiently reactive under mild conditions to facilitate the formation of amide bonds. Phenol exhibits keto-enol tautomerism with its unstable keto tautomer cyclohexadienone, but only a tiny fraction of phenol exists as the keto form; the equilibrium constant for enolisation is 10−13, which means only one in every ten trillion molecules is in the keto form at any moment. The small amount of stabilisation gained by exchanging a C=C bond for a C=O bond is more than offset by the large destabilisation resulting from the loss of aromaticity. Phenol therefore exists entirely in the enol form. Phenoxides are enolates stabilised by aromaticity. Under normal circumstances, phenoxide is more reactive at the oxygen position, but the oxygen position is a "hard" nucleophile whereas the alpha-carbon positions tend to be "soft".
Phenol is reactive toward electrophilic aromatic substitution as the oxygen atom's pi electrons donate electron density into the ring. By this general approach, many groups can be appended to the ring, via halogenation, acylation and other processes. However, phenol's ring is so activated—second only to aniline—that bromination or chlorination of phenol leads to substitution on all carbon atoms ortho and para to the hydroxy group, not only on one carbon. Phenol reacts with dilute nitric acid at room temperature to give a mixture of 2-nitrophenol and 4-nitrophenol while with concentrated nitric acid, more nitro groups get substituted on the ring to give 2,4,6-trinitrophenol, known as picric acid. Aqueous solutions of phenol are weakly acidic and turn blue litmus to red. Phenol is neutralized by sodium hydroxide forming sodium phenate or phenolate, but being weaker than carbonic acid, it cannot be neutralized by sodium bicarbonate or sodium carbonate to liberate carbon dioxide. C6H5OH + NaOH → C6H5ONa + H2OWhen a mixture of phenol and benzoyl chloride are shaken in presence of dilute sodium hydroxide solution, phenyl benzoate is formed.
This is an example of the Schotten-Baumann reaction: C6H5OH + C6H5COCl → C6H5OCOC6H5 + HClPhenol is reduced to benzene when it is distilled with zinc dust, or when phenol vapour is passed over granules of zinc at 400 °C: C6H5OH + Zn → C6H6 + ZnOWhen phenol is reacted with diazomethane in the presence of boron trifluoride, anisole is obtained as the main product and nitrogen gas as a byproduct. C6H5OH + CH2N2 → C6H5OCH3 + N2When phenol reacts with iron chloride solution, an intense violet-purple solution is formed; because of phenol's commercial importance, many methods have been developed for its production. The dominant current route, accounting for 95% of production, is the cumene process, which uses benzene and propene as feedstock and involves the partial oxidation of cumene vi