1.
Lysergic acid diethylamide
–
Lysergic acid diethylamide, also known as acid, is a psychedelic drug known for its psychological effects. This may include altered awareness of the surroundings, perceptions, and feelings as well as sensations and it is used mainly as a recreational drug and for spiritual reasons. LSD is typically either swallowed or held under the tongue and it is often sold on blotter paper, a sugar cube, or gelatin. However, adverse reactions such as anxiety, paranoia. LSD is in the ergoline family, LSD is sensitive to oxygen, ultraviolet light, and chlorine, though it may last for years if it is stored away from light and moisture at low temperature. In pure form it is odorless and clear or white in color, as little as 20–30 micrograms can produce an effect. LSD was first made by Albert Hofmann in Switzerland in 1938 from ergotamine, the laboratory name for the compound was the acronym for the German Lyserg-säure-diäthylamid, followed by a sequential number, LSD-25. Hofmann discovered its properties in 1943. LSD was introduced as a medication under the trade-name Delysid for various psychiatric uses in 1947. In the 1950s, officials at the U. S, Central Intelligence Agency thought the drug might be useful for mind control and chemical warfare and tested the drug on young servicemen and students, and others without their knowledge. The subsequent recreational use by youth culture in the Western world as part of 1960s counterculture resulted in its prohibition, LSD currently has no approved uses in medicine. LSD is commonly used as a recreational drug, LSD is considered an entheogen because it can catalyze intense spiritual experiences, during which users may feel they have come into contact with a greater spiritual or cosmic order. Users sometimes report out of body experiences, in 1966, Timothy Leary established the League for Spiritual Discovery with LSD as its sacrament. LSD can cause pupil dilation, reduced appetite, and wakefulness, other physical reactions to LSD are highly variable and nonspecific, some of which may be secondary to the psychological effects of LSD. The most common immediate psychological effects of LSD are visual hallucinations and illusions, trips usually start within 20–30 minutes of taking LSD by mouth, peak three to four hours after ingestion, and last up to 12 hours. It is impossible to predict when a bad trip will occur, some users, including Albert Hofmann, report a strong metallic taste for the duration of the effects. LSD causes an animated experience of senses, emotions, memories, time. Generally beginning within 30 to 90 minutes after ingestion, the user may experience anything from subtle changes in perception to overwhelming cognitive shifts, changes in auditory and visual perception are typical
Lysergic acid diethylamide
–
Pink elephant blotters containing LSD
Lysergic acid diethylamide
–
Lysergic acid diethylamide
Lysergic acid diethylamide
–
White on White blotters (WoW)
Lysergic acid diethylamide
–
Albert Hofmann in 2006
2.
Zinc
–
Zinc is a chemical element with the symbol Zn and atomic number 30. It is the first element in group 12 of the periodic table, in some respects zinc is chemically similar to magnesium, both elements exhibit only one normal oxidation state, and the Zn2+ and Mg2+ ions are of similar size. Zinc is the 24th most abundant element in Earths crust and has five stable isotopes, the most common zinc ore is sphalerite, a zinc sulfide mineral. The largest workable lodes are in Australia, Asia, and the United States, Zinc is refined by froth flotation of the ore, roasting, and final extraction using electricity. Zinc metal was not produced on a large scale until the 12th century in India and was unknown to Europe until the end of the 16th century, the mines of Rajasthan have given definite evidence of zinc production going back to the 6th century BC. To date, the oldest evidence of pure zinc comes from Zawar, in Rajasthan, alchemists burned zinc in air to form what they called philosophers wool or white snow. The element was named by the alchemist Paracelsus after the German word Zinke. German chemist Andreas Sigismund Marggraf is credited with discovering pure metallic zinc in 1746, work by Luigi Galvani and Alessandro Volta uncovered the electrochemical properties of zinc by 1800. Corrosion-resistant zinc plating of iron is the application for zinc. Other applications are in batteries, small non-structural castings. A variety of compounds are commonly used, such as zinc carbonate and zinc gluconate, zinc chloride, zinc pyrithione, zinc sulfide. Zinc is an essential mineral perceived by the public today as being of exceptional biologic and public health importance, Zinc deficiency affects about two billion people in the developing world and is associated with many diseases. In children, deficiency causes growth retardation, delayed sexual maturation, infection susceptibility, enzymes with a zinc atom in the reactive center are widespread in biochemistry, such as alcohol dehydrogenase in humans. Consumption of excess zinc can cause ataxia, lethargy and copper deficiency, Zinc is a bluish-white, lustrous, diamagnetic metal, though most common commercial grades of the metal have a dull finish.6 pm. The metal is hard and brittle at most temperatures but becomes malleable between 100 and 150 °C, above 210 °C, the metal becomes brittle again and can be pulverized by beating. Zinc is a conductor of electricity. For a metal, zinc has relatively low melting and boiling points, the melting point is the lowest of all the transition metals aside from mercury and cadmium. Many alloys contain zinc, including brass, Other metals long known to form binary alloys with zinc are aluminium, antimony, bismuth, gold, iron, lead, mercury, silver, tin, magnesium, cobalt, nickel, tellurium, and sodium
Zinc
–
Zinc, 30 Zn
Zinc
–
Sphalerite (ZnS)
Zinc
–
Zinc acetate
Zinc
–
Zinc chloride
3.
Hydrochloric acid
–
Hydrochloric acid is a corrosive, strong mineral acid with many industrial uses. A colorless, highly pungent solution of chloride in water. Free hydrochloric acid was first formally described in the 16th century by Libavius, later, it was used by chemists such as Glauber, Priestley, and Davy in their scientific research. It has numerous applications, including household cleaning, production of gelatin and other food additives, descaling. About 20 million tonnes of acid are produced worldwide annually. It is also found naturally in gastric acid, Hydrochloric acid was known to European alchemists as spirits of salt or acidum salis. Both names are used, especially in other languages, such as German, Salzsäure, Dutch, Zoutzuur, Swedish, Saltsyra, Turkish, Tuz Ruhu, Polish, kwas solny and Chinese. Gaseous HCl was called marine acid air, the old name muriatic acid has the same origin, and this name is still sometimes used. The name hydrochloric acid was coined by the French chemist Joseph Louis Gay-Lussac in 1814, 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. Joseph Priestley of Leeds, England prepared pure hydrogen chloride in 1772, during the Industrial Revolution in Europe, demand for alkaline substances increased. A new industrial process developed by Nicolas Leblanc of Issoundun, France enabled cheap large-scale production of sodium carbonate, in this Leblanc process, common salt is converted to soda ash, using sulfuric acid, limestone, and coal, releasing hydrogen chloride as a by-product. Until the British Alkali Act 1863 and similar legislation in other countries, 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 already settled as an important chemical in numerous applications. After the year 2000, hydrochloric acid is made by absorbing by-product hydrogen chloride from industrial organic compounds production. Hydrochloric acid is the salt of hydronium ion, H3O+ and chloride and it is usually prepared by treating HCl with water. H C l + H2 O ⟶ H3 O + + C l − Hydrochloric acid can therefore be used to prepare salts called chlorides, Hydrochloric acid is a strong acid, since it is completely dissociated in water
Hydrochloric acid
–
Hydrochloric acid
4.
PH
–
In chemistry, pH is a numeric scale used to specify the acidity or basicity of an aqueous solution. It is approximately the negative of the base 10 logarithm of the concentration, measured in units of moles per liter. More precisely it is the negative of the logarithm to base 10 of the activity of the hydrogen ion, solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are basic. Pure water is neutral, at pH7, being neither an acid nor a base, contrary to popular belief, the pH value can be less than 0 or greater than 14 for very strong acids and bases respectively. The pH scale is traceable to a set of standard solutions whose pH is established by international agreement, the pH of aqueous solutions can be measured with a glass electrode and a pH meter, or an indicator. In the first papers, the notation had the H as a subscript to the p, as so. The exact meaning of the p in pH is disputed, but according to the Carlsberg Foundation and it has also been suggested that the p stands for the German Potenz, others refer to French puissance. Another suggestion is that the p stands for the Latin terms pondus hydrogenii, potentia hydrogenii and it is also suggested that Sørensen used the letters p and q simply to label the test solution and the reference solution. Currently in chemistry, the p stands for decimal cologarithm of, PH is defined as the decimal logarithm of the reciprocal of the hydrogen ion activity, aH+, in a solution. P H = − log 10 = log 10 For example and this definition was adopted because ion-selective electrodes, which are used to measure pH, respond to activity. For H+ number of electrons transferred is one and it follows that electrode potential is proportional to pH when pH is defined in terms of activity. The reference electrode may be a silver chloride electrode or a calomel electrode, the hydrogen-ion selective electrode is a standard hydrogen electrode. Reference electrode | concentrated solution of KCl || test solution | H2 | Pt Firstly, the cell is filled with a solution of hydrogen ion activity. Then the emf, EX, of the cell containing the solution of unknown pH is measured. PH = pH + E S − E X z The difference between the two measured emf values is proportional to pH and this method of calibration avoids the need to know the standard electrode potential. The proportionality constant, 1/z is ideally equal to 12.303 R T / F the Nernstian slope, to apply this process in practice, a glass electrode is used rather than the cumbersome hydrogen electrode. A combined glass electrode has a reference electrode. It is calibrated against buffer solutions of hydrogen ion activity
PH
–
Lemon juice tastes sour because it contains 5% to 6%
citric acid, which has a pH of 2.2.
PH
–
Chart showing the variation of color of universal indicator paper with pH
PH
–
pH values of some common substances
5.
Acid-base extraction
–
Acid-base extraction is a procedure using sequential liquid–liquid extractions to purify acids and bases from mixtures based on their chemical properties. Acid-base extraction is performed during the work-up after chemical syntheses and for the isolation of compounds. The product is free of neutral and acidic or basic impurities. It is not possible to separate chemically similar acids or bases using this simple method, the fundamental theory behind this technique is that salts, which are ionic, tend to be water-soluble while neutral molecules tend not to be. The addition of an acid to a mixture of a base and acid will result in the acid remaining uncharged. If the organic acid, such as an acid, is sufficiently strong. Conversely, the addition of a base to a mixture of an acid and base will result in the base remaining uncharged. Once again, the self-ionization of a base is suppressed by the added base. Usually, the mixture is dissolved in a solvent such as dichloromethane or diethyl ether. An aqueous solution of the acid or base is added, after shaking and allowing for phase separation, the phase containing the compound of interest is collected. The procedure is repeated with this phase at the opposite pH range. The order of the step is not important and the process can be repeated to increase the separation, however, it is often convenient to have the compound dissolved in the organic phase after the last step, so that the evaporation of the solvent yields the product. The procedure works only for acids and bases with a difference in solubility between their charged and their uncharged form. The procedure does not work for, Zwitterions with acidic and basic functional groups in the same molecule, very lipophilic amines that do not easily dissolve in the aqueous phase in their charged form, e. g. triphenylamine and trihexylamine. Very lipophilic acids that do not easily dissolve in the phase in their charged form. Lower amines like ammonia, methylamine, or triethanolamine which are miscible or significantly soluble in water at most pH, hydrophilic acids like acetic acid, citric acid, and most inorganic acids like sulfuric acid or phosphoric acid. Chromatography, a powerful but more complex procedure to separate compounds Extraction Multiphasic liquid Separating funnel Acid base extraction
Acid-base extraction
–
Overview of an acid-base extraction
6.
Acid dissociation constant
–
An acid dissociation constant, Ka, is a quantitative measure of the strength of an acid in solution. It is the constant for a chemical reaction known as dissociation in the context of acid–base reactions. In the example shown in the figure, HA represents acetic acid, and A− represents the acetate ion, the chemical species HA, A− and H3O+ are said to be in equilibrium when their concentrations do not change with the passing of time. The definition can then be more simply H A ⇌ A − + H +, K a = This is the definition in common usage. A weak acid has a pKa value in the approximate range −2 to 12 in water, pKa values for strong acids can, however, be estimated by theoretical means. The definition can be extended to non-aqueous solvents, such as acetonitrile and dimethylsulfoxide. Denoting a solvent molecule by S H A + S ⇌ A − + S H +, K a = When the concentration of solvent molecules can be taken to be constant, K a =, as before. The value of pKa also depends on structure of the acid in many ways. For example, Pauling proposed two rules, one for successive pKa of polyprotic acids, and one to estimate the pKa of oxyacids based on the number of =O and −OH groups. Other structural factors that influence the magnitude of the dissociation constant include inductive effects, mesomeric effects. Hammett type equations have frequently applied to the estimation of pKa. The quantitative behaviour of acids and bases in solution can be only if their pKa values are known. These calculations find application in different areas of chemistry, biology, medicine. Acid dissociation constants are essential in aquatic chemistry and chemical oceanography. In living organisms, acid–base homeostasis and enzyme kinetics are dependent on the pKa values of the acids and bases present in the cell. According to Arrheniuss original definition, an acid is a substance that dissociates in solution, releasing the hydrogen ion H+. The equilibrium constant for this reaction is known as a dissociation constant. Brønsted and Lowry generalised this further to an exchange reaction
Acid dissociation constant
7.
Acid strength
–
The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely ionizes in a solution, in water, one mole of a strong acid HA dissolves yielding one mole of H+ and one mole of the conjugate base, A−. Essentially, none of the non-ionized acid HA remains. Examples of strong acids are hydrochloric acid, hydroiodic acid, hydrobromic acid, perchloric acid, nitric acid, in aqueous solution, each of these essentially ionizes 100%. In contrast, an acid only partially dissociates. Examples in water include carbonic acid and acetic acid, at equilibrium, both the acid and the conjugate base are present in solution. Stronger acids have an acid dissociation constant and a smaller logarithmic constant than weaker acids. The stronger an acid is, the more easily it loses a proton, two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base, while Ka measures the strength of an acidic molecule, the strength of an aqueous acid solution is measured by pH, which is a function of the concentration of hydronium ions in solution. The pH of a solution of an acid in water is determined by both Ka and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, for concentrated solutions of strong acids with a pH less than about zero, the Hammett acidity function is a better measure of acidity than the pH. Sulfonic acids, which are organic oxyacids, are a class of strong acids, a common example is p-toluenesulfonic acid. Unlike sulfuric acid itself, sulfonic acids can be solids, in fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable. Superacids are acid solutions that are more acidic than 100% sulfuric acid, examples of superacids are fluoroantimonic acid, magic acid and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium salts and they can also quantitatively stabilize carbocations. In aqueous solution, an acid is an acid that ionizes completely by losing one proton.74. An example is HCl for which pKa = −6.3 and this generally means that, in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution. No acid species stronger than H3O+ can exist in aqueous solution, the acidity of the stronger acid is said to be leveled to the acidity of hydronium ion. Strong acids are distinguished from weak acids, in which dissociation is incomplete and is represented as an equilibrium, the typical definition of a weak acid is any acid that does not dissociate completely
Acid strength
–
butanoic acid or
butyric acid
8.
Buffer solution
–
For an individual weak acid or weak base component, see Buffering agent. For uses not related to chemistry, see Buffer. A buffer solution is an aqueous solution consisting of a mixture of an acid and its conjugate base. Its pH changes very little when a small amount of acid or base is added to it. Buffer solutions are used as a means of keeping pH at a constant value in a wide variety of chemical applications. In nature, there are systems that use buffering for pH regulation. For example, the buffering system is used to regulate the pH of blood. Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A−. HA ⇌ H+ + A− When some strong acid is added to a mixture of the weak acid and its conjugate base. Because of this, the ion concentration increases by less than the amount expected for the quantity of strong acid added. Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added, the effect is illustrated by the simulated titration of a weak acid with pKa =4.7. The relative concentration of undissociated acid is shown in blue and of its base in red. The pH changes relatively slowly in the region, pH = pKa ±1. OH− + H+ → H2O Once the acid is more than 95% deprotonated the pH rises rapidly because most of the alkali is consumed in the neutralization reaction. Buffer capacity, β, is a measure of the resistance of a buffer solution to pH change on addition of hydroxide ions. It can be defined as follows, β = d n d where dn is an infinitesimal amount of added base and d is the resulting infinitesimal change in the cologarithm of the hydrogen ion concentration. With this definition the buffer capacity of an acid, with a dissociation constant Ka. This equation shows that there are three regions of raised buffer capacity, at very low p the term in the denominator predominates and buffer capacity rises exponentially with decreasing pH
Buffer solution
–
Simulated
titration of an acidified solution of a weak acid (pK a = 4.7) with alkali.
9.
Self-ionization of water
–
The self-ionization of water is an ionization reaction in pure water or in an aqueous solution, in which a water molecule, H2O, deprotonates to become a hydroxide ion, OH−. The hydrogen nucleus, H+, immediately protonates another water molecule to form hydronium and it is an example of autoprotolysis, and exemplifies the amphoteric nature of water. Chemically pure water has a conductivity of 0.055 µS/cm. According to the theories of Svante Arrhenius, this must be due to the presence of ions, the ions are produced by the self-ionization reaction H2O + H2O ⇌ H3O+ + OH− This equilibrium applies to pure water and any aqueous solution. When the equilibrium constant is written as a product of concentrations it is necessary to make corrections to the value of K w depending on ionic strength, at 25 °C and zero ionic strength, Kw is equal to 1. 0×10−14. For most practical purposes, the molal and molar concentrations are equal near the ambient temperature and pressure. g. for seawater, or at elevated temperatures and we can also define pKw = −log10 Kw =14 at 25 °C. This is analogous to the notations pH and pKa for an acid dissociation constant, the logarithmic form of the equilibrium constant equation is pKw = pH + pOH. The dependence of the water ionization on temperature and pressure has been investigated thoroughly. The value of pKw decreases as temperature increases from the point of ice to a minimum at c.250 °C. With electrolyte solutions, the value of pKw is dependent on ionic strength of the electrolyte, values for sodium chloride are typical for a 1,1 electrolyte. With 1,2 electrolytes, MX2, pKw decreases with increasing ionic strength, the value of Kw is usually of interest in the liquid phase. Example values for superheated steam and supercritical water fluid are given in the table, the values are for supercritical fluid except those marked, a at saturation pressure corresponding to 350 °C. In water–heavy water mixtures equilibria several species are involved, H2O, HDO, D2O, H3O+, D3O+, H2DO+, HD2O+, the rate of reaction for the ionization reaction 2 H2O → H3O+ + OH− depends on the activation energy, ΔE‡. According to the Boltzmann distribution the proportion of water molecules that have sufficient energy, thus some dissociation can occur because sufficient thermal energy is available. The following sequence of events has been proposed on the basis of electric field fluctuations in liquid water, within 1 picosecond, however, a second reorganization of the hydrogen bond network allows rapid proton transfer down the electric potential difference and subsequent recombination of the ions. This timescale is consistent with the time it takes for hydrogen bonds to reorientate themselves in water, the inverse recombination reaction H3O+ + OH− →2 H2O is among the fastest chemical reactions known, with a reaction rate constant of 1. 3×1011 M−1 s−1 at room temperature. Such a rapid rate is characteristic of a reaction, in which the rate is limited by the speed of molecular diffusion. Water molecules dissociate into equal amounts of H3O+ and OH−, so their concentrations are equal to 1. 00×10−7 mol dm−3 at 25 °C, a solution in which the H3O+ and OH− concentrations equal each other is considered a neutral solution
Self-ionization of water
–
Temperature dependence of the water ionization constant at 25 MPa
10.
Lewis acids and bases
–
A Lewis acid is a chemical species that reacts with a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that donates a pair of electrons to a Lewis acid to form a Lewis adduct, for example, OH− and NH3 are Lewis bases, because they can donate a lone pair of electrons. In the adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, usually the terms Lewis acid and Lewis base are defined within the context of a specific chemical reaction. For example, in the reaction of Me3B and NH3 to give Me3BNH3, Me3B acts as a Lewis acid, the terminology refers to the contributions of Gilbert N. Lewis. Another example is boron trifluoride etherate, BF3•Et2O, the center dot is also used to represent hydrate coordination in various crystals, as in MgSO4•7H2O for hydrated magnesium sulfate. In general, however, the bond is viewed as simply 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, for the purposes of discussion, even complex compounds such as Et3Al2Cl3 and AlCl3 are treated as trigonal planar Lewis acids. Other reactions might simply 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. Simplest are those that react directly with the Lewis base, but more common are those that undergo a reaction prior to forming the adduct. Again, the description of a Lewis acid is used loosely. For example, in solution, bare protons do not exist, BF3 + OMe2 → BF3OMe2 Both BF4− and BF3OMe2 are Lewis base adducts of boron trifluoride. 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, 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 usually only after dissociating a more weakly bound Lewis base, 2+ +6 NH3 → 2+ +6 H2O The proton is one of the strongest but is also one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is heavily solvated, the key step is the acceptance by AlCl3 of a chloride ion lone-pair, forming AlCl4− and creating the strongly acidic, that is, electrophilic, carbonium ion. RCl +AlCl3 → R+ + AlCl4− A Lewis base is an atomic or molecular species where the highest occupied molecular orbital is highly localized, typical Lewis bases are conventional amines such as ammonia and alkyl amines. Other common Lewis bases include pyridine and its derivatives, some of the main classes of Lewis bases are amines of the formula NH3−xRx where R = alkyl or aryl
Lewis acids and bases
–
Diagram of some Lewis bases and acids
Lewis acids and bases
–
Lewis diagrams showing the formation of the
ammonium ion (diagram should show
adduct having a positive charge)
11.
Organic acid
–
An organic acid is an organic compound with acidic properties. The most common organic acids are the carboxylic acids, whose acidity is associated with their carboxyl group –COOH, sulfonic acids, containing the group –SO2OH, are relatively stronger acids. Alcohols, with –OH, can act as acids but they are very weak. The relative stability of the base of the acid determines its acidity. Other groups can also confer acidity, usually weakly, the thiol group –SH, the group. In biological systems, organic compounds containing these groups are referred to as organic acids. In general, organic acids are acids and do not dissociate completely in water. Lower molecular mass organic acids such as formic and lactic acids are miscible in water, on the other hand, most organic acids are very soluble in organic solvents. P-Toluenesulfonic acid is a strong acid used in organic chemistry often because it is able to dissolve in the organic reaction solvent. Exceptions to these solubility characteristics exist in the presence of other substituents that affect the polarity of the compound. The pKa a logarithmic measure of the dissociation constant, categorizes the strength of an acid, the lower or more negative the number. It should not be confused with pH, the measure of actual hydrogen ion concentration. Phenol C6H5OH The 19th century name carbolic acid came from the original German name karbolsäure and it is a weak, non-carboxylic acid, but is corrosive and causes chemical burns on the skin due to a protein-degenerating effect. Uric acid C5H4N4O3 Uric acid is a purine derivative which is a diprotic acid but not a carboxylic one. Taurine C2H7NO3S One of the few natural sulfonic acids, discovered in bile, p-Toluenesulfonic acid CH3C6H4SO3H A strong acid that, unlike some strong mineral acids, is non-oxidizing. Trifluoromethanesulfonic acid CF3SO3H One of the strongest acids in existence, organic or inorganic, about a thousand times stronger than sulfuric acid, first synthesized in 1954, it also resists oxidation/reduction reactions and does not sulfonate substrates like sulfuric or chlorosulfonic acid. Aminomethylphosphonic acid A strong phosphonic acid Simple organic acids like formic or acetic acids are used for oil and these organic acids are much less reactive with metals than are strong mineral acids like hydrochloric acid or mixtures of HCl and hydrofluoric acid. For this reason, organic acids are used at temperatures or when long contact times between acid and pipe are needed
Organic acid
–
The general structure of a few organic acids. From left to right:
carboxylic acid,
sulfonic acid. The acidic hydrogen in each molecule is colored red.
12.
Strong acid
–
The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely ionizes in a solution, in water, one mole of a strong acid HA dissolves yielding one mole of H+ and one mole of the conjugate base, A−. Essentially, none of the non-ionized acid HA remains. Examples of strong acids are hydrochloric acid, hydroiodic acid, hydrobromic acid, perchloric acid, nitric acid, in aqueous solution, each of these essentially ionizes 100%. In contrast, an acid only partially dissociates. Examples in water include carbonic acid and acetic acid, at equilibrium, both the acid and the conjugate base are present in solution. Stronger acids have an acid dissociation constant and a smaller logarithmic constant than weaker acids. The stronger an acid is, the more easily it loses a proton, two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base, while Ka measures the strength of an acidic molecule, the strength of an aqueous acid solution is measured by pH, which is a function of the concentration of hydronium ions in solution. The pH of a solution of an acid in water is determined by both Ka and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, for concentrated solutions of strong acids with a pH less than about zero, the Hammett acidity function is a better measure of acidity than the pH. Sulfonic acids, which are organic oxyacids, are a class of strong acids, a common example is p-toluenesulfonic acid. Unlike sulfuric acid itself, sulfonic acids can be solids, in fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable. Superacids are acid solutions that are more acidic than 100% sulfuric acid, examples of superacids are fluoroantimonic acid, magic acid and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium salts and they can also quantitatively stabilize carbocations. In aqueous solution, an acid is an acid that ionizes completely by losing one proton.74. An example is HCl for which pKa = −6.3 and this generally means that, in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution. No acid species stronger than H3O+ can exist in aqueous solution, the acidity of the stronger acid is said to be leveled to the acidity of hydronium ion. Strong acids are distinguished from weak acids, in which dissociation is incomplete and is represented as an equilibrium, the typical definition of a weak acid is any acid that does not dissociate completely
Strong acid
–
butanoic acid or
butyric acid
13.
Weak acid
–
The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely ionizes in a solution, in water, one mole of a strong acid HA dissolves yielding one mole of H+ and one mole of the conjugate base, A−. Essentially, none of the non-ionized acid HA remains. Examples of strong acids are hydrochloric acid, hydroiodic acid, hydrobromic acid, perchloric acid, nitric acid, in aqueous solution, each of these essentially ionizes 100%. In contrast, an acid only partially dissociates. Examples in water include carbonic acid and acetic acid, at equilibrium, both the acid and the conjugate base are present in solution. Stronger acids have an acid dissociation constant and a smaller logarithmic constant than weaker acids. The stronger an acid is, the more easily it loses a proton, two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base, while Ka measures the strength of an acidic molecule, the strength of an aqueous acid solution is measured by pH, which is a function of the concentration of hydronium ions in solution. The pH of a solution of an acid in water is determined by both Ka and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, for concentrated solutions of strong acids with a pH less than about zero, the Hammett acidity function is a better measure of acidity than the pH. Sulfonic acids, which are organic oxyacids, are a class of strong acids, a common example is p-toluenesulfonic acid. Unlike sulfuric acid itself, sulfonic acids can be solids, in fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable. Superacids are acid solutions that are more acidic than 100% sulfuric acid, examples of superacids are fluoroantimonic acid, magic acid and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium salts and they can also quantitatively stabilize carbocations. In aqueous solution, an acid is an acid that ionizes completely by losing one proton.74. An example is HCl for which pKa = −6.3 and this generally means that, in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution. No acid species stronger than H3O+ can exist in aqueous solution, the acidity of the stronger acid is said to be leveled to the acidity of hydronium ion. Strong acids are distinguished from weak acids, in which dissociation is incomplete and is represented as an equilibrium, the typical definition of a weak acid is any acid that does not dissociate completely
Weak acid
–
butanoic acid or
butyric acid
14.
Solid acid
–
Solid acids are acids that do not dissolve in the reaction medium. They are often used in heterogeneous catalysts, examples include oxides, which function as Lewis acids including silico-aluminates, and sulfated zirconia. Many transition metal oxides are acidic, including titania, zirconia, such acids are used in cracking. Many solid Bronsted acids are also employed industrially, including sulfonated polystyrene, solid phosphoric acid, niobic acid, solid acids are used in catalysis in many industrial chemical processes, from large-scale catalytic cracking in petroleum refining to the synthesis of various fine chemicals. One large scale application is alkylation, e. g. the combination of benzene and ethylene to give ethylbenzene, another application is the rearrangement of cyclohexanone oxime to caprolactam. Many alkylamines are prepared by amination of alcohols, catalyzed by solid acids, solid acids can be used as electrolytes in fuel cells
Solid acid
–
In acid-catalyzed
Fischer esterification, the
proton binds to oxygens and functions as a
Lewis acid to activate the ester
carbonyl (top row) as an
electrophile, and converts the
hydroxyl into the good
leaving group water (bottom left). Both lower the kinetic barrier and speed up the attainment of
chemical equilibrium.
15.
Molecule
–
A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge, however, in quantum physics, organic chemistry, and biochemistry, the term molecule is often used less strictly, also being applied to polyatomic ions. In the kinetic theory of gases, the molecule is often used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are in fact monoatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one element, as with oxygen, or it may be heteronuclear. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances and they also make up most of the oceans and atmosphere. Also, no typical molecule can be defined for ionic crystals and covalent crystals, the theme of repeated unit-cellular-structure also holds for most condensed phases with metallic bonding, which means that solid metals are also not made of molecules. In glasses, atoms may also be together by chemical bonds with no presence of any definable molecule. The science of molecules is called molecular chemistry or molecular physics, in practice, however, this distinction is vague. In molecular sciences, a molecule consists of a system composed of two or more atoms. Polyatomic ions may sometimes be thought of as electrically charged molecules. The term unstable molecule is used for very reactive species, i. e, according to Merriam-Webster and the Online Etymology Dictionary, the word molecule derives from the Latin moles or small unit of mass. Molecule – extremely minute particle, from French molécule, from New Latin molecula, diminutive of Latin moles mass, a vague meaning at first, the vogue for the word can be traced to the philosophy of Descartes. The definition of the molecule has evolved as knowledge of the structure of molecules has increased, earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties. Molecules are held together by covalent bonding or ionic bonding. Several types of non-metal elements exist only as molecules in the environment, for example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements, a covalent bond is a chemical bond that involves the sharing of electron pairs between atoms
Molecule
–
Atomic force microscopy image of a
PTCDA molecule, which contains five carbon rings in a non-linear arrangement.
Molecule
Molecule
–
A
scanning tunneling microscopy image of
pentacene molecules, which consist of linear chains of five carbon rings.
Molecule
–
Arrangement of
polyvinylidene fluoride molecules in a
nanofiber –
transmission electron microscopy image.
16.
Hydron (chemistry)
–
Unlike other ions, the hydron consists only of a bare atomic nucleus. The hydron is too reactive to occur in many liquids, even though it is sometimes visualized to do so by students of chemistry, a free hydron would react with a molecule of the liquid to form a more complicated cation. Examples are the ion in water-based acids, and H 2F+, the unstable cation of fluoroantimonic acid. For this reason, in such liquids including liquid acids, hydrons diffuse by contact from one complex cation to another, the hydrated form of the hydrogen cation, the hydronium ion H 3O+, is a key object of Arrhenius definition of acid. Other hydrated forms, the Zundel cation H 5O+2 which is formed from a proton, the hydron itself is crucial in more general Brønsted–Lowry acid–base theory, which extends the concept of acid–base chemistry beyond aqueous solutions. The negatively charged counterpart of the hydron is the anion, H−. Proton, having the symbol p or 1H+, is the +1 ion of protium. Deuteron, having the symbol 2H+ or D+, is the +1 ion of deuterium, 2H or D. Triton, having the symbol 3H+ or T+, is the +1 ion of tritium, other isotopes of hydrogen are too unstable to be relevant in chemistry. The name proton refers to isotopically pure 1H+, on the other hand, referring to the hydron as simply hydrogen ion is not recommended because hydrogen anions also exist. The term hydron was defined by IUPAC in 1988, traditionally, the term proton was and is used in place of hydron. The latter term is only used in the context where comparisons between the various isotopes of hydrogen is important. Otherwise, referring to hydrons as protons is still considered acceptable, the transfer of H+ in an acid-base reaction is usually referred to as proton transfer. Acid and bases are referred to as proton donors and acceptors correspondingly, however, although 99. 9844% of natural hydrogen nuclei are protons, the remainder are deuterons. Deprotonation Superacid Dihydrogen cation Trihydrogen cation Hydrogen ion cluster
Hydron (chemistry)
–
Hydron
17.
Covalent bond
–
A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, for many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration. Covalent bonding includes many kinds of interactions, including σ-bonding, π-bonding, metal-to-metal bonding, agostic interactions, bent bonds, the term covalent bond dates from 1939. In the molecule H2, the atoms share the two electrons via covalent bonding. Covalency is greatest between atoms of similar electronegativities, thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they be of comparable electronegativity. Covalent bonding that entails sharing of electrons more than two atoms is said to be delocalized. The term covalence in regard to bonding was first used in 1919 by Irving Langmuir in a Journal of the American Chemical Society article entitled The Arrangement of Electrons in Atoms and Molecules. Langmuir wrote that we shall denote by the term covalence the number of pairs of electrons that an atom shares with its neighbors. The idea of covalent bonding can be traced several years before 1919 to Gilbert N. Lewis and he introduced the Lewis notation or electron dot notation or Lewis dot structure, in which valence electrons are represented as dots around the atomic symbols. Pairs of electrons located between atoms represent covalent bonds, multiple pairs represent multiple bonds, such as double bonds and triple bonds. An alternative form of representation, not shown here, has bond-forming electron pairs represented as solid lines, Lewis proposed that an atom forms enough covalent bonds to form a full outer electron shell. In the diagram of methane shown here, the atom has a valence of four and is, therefore, surrounded by eight electrons, four from the carbon itself. Each hydrogen has a valence of one and is surrounded by two electrons – its own one electron plus one from the carbon, walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond in 1927. Their work was based on the valence bond model, which assumes that a bond is formed when there is good overlap between the atomic orbitals of participating atoms. Atomic orbitals have specific directional properties leading to different types of covalent bonds, sigma bonds are the strongest covalent bonds and are due to head-on overlapping of orbitals on two different atoms. A single bond is usually a σ bond, pi bonds are weaker and are due to lateral overlap between p orbitals. A double bond between two given atoms consists of one σ and one π bond, and a bond is one σ. Covalent bonds are also affected by the electronegativity of the atoms which determines the chemical polarity of the bond
Covalent bond
–
Early concepts in covalent bonding arose from this kind of image of the molecule of
methane. Covalent bonding is implied in the
Lewis structure by indicating electrons shared between atoms.
18.
Electron pair
–
In chemistry, an electron pair or a Lewis pair consists of two electrons that occupy the same orbital but have opposite spins. The electron pair concept was introduced in a 1916 paper of Gilbert N. Lewis, because electrons are fermions, the Pauli exclusion principle forbids these particles from having exactly the same quantum numbers. Therefore, the way to occupy the same orbital, i. e. have the same orbital quantum numbers, is to differ in the spin quantum number. This limits the number of electrons in the orbital to exactly two. The pairing of spins is often energetically favorable and electron pairs therefore play a large role in chemistry. They can form a bond between two atoms, or they can occur as a lone pair of valence electrons. They also fill the core levels of an atom, because the spins are paired, the magnetic moment of the electrons cancels and the contribution of the pair to the magnetic properties will in general be a diamagnetic one. Although a strong tendency to pair off electrons can be observed in chemistry, a very special case of electron pair formation occurs in superconductivity, the formation of Cooper pairs. Polyhedral skeletal electron pair theory Jemmis mno rules Lewis acids and bases Nucleophile Electron pair production
Electron pair
–
MO diagrams depicting covalent (left) and polar covalent (right) bonding in a
diatomic molecule. In both cases a bond is created by the formation of an electron pair.
19.
Hydronium ion
–
In chemistry, hydronium is the common name for the aqueous cation H 3O+, the type of oxonium ion produced by protonation of water. It is the ion present when an Arrhenius acid is dissolved in water. It is the amount of hydronium ions relative to hydroxide ions that determines a solutions pH, at 25 °C, water has a pH of 7. A pH value less than 7 indicates an acidic solution, according to IUPAC nomenclature of organic chemistry, the hydronium ion should be referred to as oxonium. Hydroxonium may also be used unambiguously to identify it, a draft IUPAC proposal also recommends the use of oxonium and oxidanium in organic and inorganic chemistry contexts, respectively. An oxonium ion is any ion with a trivalent oxygen cation, for example, a protonated hydroxyl group is an oxonium ion, but not a hydronium ion. Since O+ and N have the number of electrons, H 3O+ is isoelectronic with ammonia. As shown in the images above, H 3O+ has a pyramidal molecular geometry with the oxygen atom at its apex. The H–O–H bond angle is approximately 113°, and the center of mass is close to the oxygen atom. Because the base of the pyramid is made up of three hydrogen atoms, the H 3O+ molecules symmetric top configuration is such that it belongs to the C3v point group. Because of this symmetry and the fact that it has a dipole moment, the transition dipole lies along the c-axis and, because the negative charge is localized near the oxygen atom, the dipole moment points to the apex, perpendicular to the base plane. Hydronium is the cation that forms from water in the presence of hydrogen ions and these hydrons do not exist in a free state, they are extremely reactive and are solvated by water. An acidic solute is generally the source of these hydrons, however and this special case of water reacting with water to produce hydronium ions is commonly known as the self-ionization of water. The resulting hydronium ions are few and short-lived, PH is a measure of the relative activity of hydronium and hydroxide ions in aqueous solutions. In acidic solutions, hydronium is the active, its excess proton being readily available for reaction with basic species. The hydronium ion is very acidic, at 25 °C, its pKa is 0 and it is the most acidic species that can exist in water, any stronger acid will ionize and protonate a water molecule to form hydronium. PH was originally conceived to be a measure of the ion concentration of aqueous solution. We now know that all such free protons quickly react with water to form hydronium
Hydronium ion
–
Zundel cation
Hydronium ion
20.
Acid-base reaction
–
An acid–base reaction is a chemical reaction that occurs between an acid and a base. Their importance becomes apparent in analyzing acid–base reactions for gaseous or liquid species, the first of these concepts was provided by the French chemist Antoine Lavoisier, around 1776. The first scientific concept of acids and bases was provided by Lavoisier in around 1776, however, Davy failed to develop a new theory, concluding that acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances. One notable modification of oxygen theory was provided by Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals, in 1838, Justus von Liebig proposed that an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal. This redefinition was based on his work on the chemical composition of organic acids. Liebigs definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition, the first modern definition of acids and bases in molecular terms was devised by Svante Arrhenius. As defined by Arrhenius, an Arrhenius acid is a substance that dissociates in water to hydrogen ions, that is. This causes the protonation of water, or the creation of the hydronium ion, thus, in modern times, the symbol H+ is interpreted as a shorthand for H3O+, because it is now known that a bare proton does not exist as a free species in aqueous solution. An Arrhenius base is a substance that dissociates in water to form ions, that is. The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, under this definition, pure H2SO4 and HCl dissolved in toluene are not acidic, and molten NaOH and solutions of calcium amide in liquid ammonia are not alkaline. The reaction of an acid with a base is called a neutralization reaction, the products of this reaction are a salt and water. Acid + base → salt + water In this traditional representation an acid–base neutralization reaction is formulated as a double-replacement reaction. For example, the reaction of acid, HCl, with sodium hydroxide, NaOH, solutions produces a solution of sodium chloride, NaCl. HCl + NaOH → NaCl + H2O The modifier in this equation is important and it was implied by Arrhenius, not included explicitly. It indicates that the substances are dissolved in water, in fact though all three substances, HCl, NaOH and NaCl are capable of existing as pure compounds, in aqueous solutions they are fully dissociated into the ions H+, Cl−, Na+ and OH−. An acid–base reaction is, thus, the removal of an ion from the acid. The removal of an ion from an acid produces its conjugate base. The reception of a proton by a base produces its conjugate acid, in this approach, acids and bases are fundamentally different in behavior from salts, which are seen as electrolytes, subject to the theories of Debye, Onsager, and others
Acid-base reaction
–
Svante Arrhenius
Acid-base reaction
–
Johannes Nicolaus Brønsted and
Thomas Martin Lowry
Acid-base reaction
21.
Thomas Martin Lowry
–
Lowry was born in Low Moor, Bradford, West Yorkshire, England, in a Cornish family. He was the son of the Reverend E. P. Lowry. He was educated at Kingswood School, Bath, Somerset, during those years he realized that he wanted to be a chemist. He studied chemistry under Henry Edward Armstrong, an English chemist whose interests were primarily in organic chemistry, from 1896 to 1913 Lowry was assistant to Armstrong, and between 1904 and 1913 worked as Lecturer in Chemistry at the Westminster Training College. From 1920 till his death, Lowry served as the Chair of Physical Chemistry at the University of Cambridge and he married a daughter of the Rev. C. Wood in 1904 and was survived by his widow, two sons and a daughter. Since the establishment of the Faraday Society in 1903, Lowry had been its active member, in 1914 he was elected a Fellow of the Royal Society. During and after the World War I, Lowry acted as Director of Shell-filling and worked for the Trench Warfare Committee, Chemical Warfare Committee, for this service, he was awarded the Order of the British Empire and the Order of Saints Maurice and Lazarus. In 1898, Lowry noted the change in rotation on nitro-d-camphor with time. He studied changes in optical rotation caused by acid- and base-catalyzed reactions of camphor derivatives and this led in 1923 to his formulation of the protonic definition of acids and bases, now known as Brønsted–Lowry acid-base theory, independently of the work by Johannes Nicolaus Brønsted. Lowry published a few hundred papers and several books and his 1935 monograph on Optical Rotatory Power has long been regarded as a standard work on the subject
Thomas Martin Lowry
–
Thomas Martin Lowry
22.
Aqueous solution
–
An aqueous solution is a solution in which the solvent is water. It is usually shown in chemical equations by appending to the relevant chemical formula, for example, a solution of table salt, or sodium chloride, in water would be represented as Na+ + Cl−. The word aqueous means pertaining to, related to, similar to, as water is an excellent solvent and is also naturally abundant, it is a ubiquitous solvent in chemistry. Substances that are hydrophobic often do not dissolve well in water, an example of a hydrophilic substance is sodium chloride. Acids and bases are aqueous solutions, as part of their Arrhenius definitions, the ability of a substance to dissolve in water is determined by whether the substance can match or exceed the strong attractive forces that water molecules generate between themselves. If the substance lacks the ability to dissolve in water the molecules form a precipitate, reactions in aqueous solutions are usually metathesis reactions. Metathesis reactions are another term for double-displacement, that is, when a cation displaces to form a bond with the other anion. The cation bonded with the latter anion will dissociate and bond with the other anion, aqueous solutions that conduct electric current efficiently contain strong electrolytes, while ones that conduct poorly are considered to have weak electrolytes. Those strong electrolytes are substances that are ionized in water. Nonelectrolytes are substances that dissolve in water yet maintain their molecular integrity, examples include sugar, urea, glycerol, and methylsulfonylmethane. When writing the equations of reactions, it is essential to determine the precipitate. To determine the precipitate, one must consult a chart of solubility, soluble compounds are aqueous, while insoluble compounds are the precipitate. Remember that there may not always be a precipitate, when performing calculations regarding the reacting of one or more aqueous solutions, in general one must know the concentration, or molarity, of the aqueous solutions. Solution concentration is given in terms of the form of the prior to it dissolving. Metal ions in aqueous solution Solubility Dissociation Acid-base reaction theories Properties of water Zumdahl S.1997, 4th ed. Boston, Houghton Mifflin Company
Aqueous solution
–
The first
solvation shell of a sodium ion dissolved in water.
23.
Litmus
–
Litmus is a water-soluble mixture of different dyes extracted from lichens. It is often absorbed onto filter paper to one of the oldest forms of pH indicator. Litmus was used for the first time about 1300 AD by Spanish alchemist Arnaldus de Villa Nova, from the 16th century on, the blue dye was extracted from some lichens, especially in the Netherlands. Litmus can be found in different species of lichens, currently, the main sources are Roccella montagnei and Dendrographa leucophoea. The main use of litmus is to test whether a solution is acidic or basic, wet litmus paper can also be used to test for water-soluble gases that affect acidity or alkalinity, the gas dissolves in the water and the resulting solution colors the litmus paper. For instance, ammonia gas, which is alkaline, colors the red litmus paper blue. Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic or alkaline conditions, litmus can also be prepared as an aqueous solution that functions similarly. Under acidic conditions, the solution is red, and under basic conditions, chemical reactions other than acid-base can also cause a color change to litmus paper. For instance, chlorine gas turns blue litmus paper white – the litmus dye is bleached and this reaction is irreversible, so the litmus is not acting as an indicator in this situation. The litmus mixture has the CAS number 1393-92-6 and contains 10 to 15 different dyes, most of the chemical components of litmus are likely to be the same as those of the related mixture known as orcein, but in different proportions. In contrast with orcein, the constituent of litmus has an average molecular mass of 3300. Acid-base indicators on litmus owe their properties to a 7-hydroxyphenm, noxazone chromophore, some fractions of litmus were given specific names including erythrolitmin, azolitmin, spaniolitmin, leucoorcein, and leucazolitmin. Azolitmin shows nearly the same effect as litmus, red litmus contains a weak diprotic acid. When it is exposed to a compound, the hydrogen ions react with the added base. The conjugated base, formed from the acid, has a blue color. Nitrazine strips, measures pH about 4.5 to 7.5 with more precision Phenolphthalein Bromothymol blue Methyl orange Universal indicator, the color chart is usually included with the strips
Litmus
–
Litmus powder
Litmus
–
Parmelia sulcata
Litmus
–
Litmus paper being used
![{\displaystyle K_{a}={\frac {\ce {[H+][A^{-}]}}{\ce {[HA]}}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/089cad3017a8be4c4b9490538bd001032c84bc22)
![{\displaystyle {\begin{aligned}\alpha _{\ce {H2A}}&={\frac {\ce {[H+]^{2}}}{{\ce {[H+]^{2}}}+[{\ce {H+}}]K_{1}+K_{1}K_{2}}}={\frac {\ce {[H2A]}}{\ce {{[H2A]}+{[HA^{-}]}+{[A^{2-}]}}}}\\\alpha _{\ce {HA^{-}}}&={\frac {[{\ce {H+}}]K_{1}}{{\ce {[H+]^{2}}}+[{\ce {H+}}]K_{1}+K_{1}K_{2}}}={\frac {\ce {[HA^{-}]}}{\ce {{[H2A]}+{[HA^{-}]}+{[A^{2-}]}}}}\\\alpha _{\ce {A^{2-}}}&={\frac {K_{1}K_{2}}{{\ce {[H+]^{2}}}+[{\ce {H+}}]K_{1}+K_{1}K_{2}}}={\frac {\ce {[A^{2-}]}}{\ce {{[H2A]}+{[HA^{-}]}+{[A^{2-}]}}}}\end{aligned}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/6fd9c193f75cc3a359a8c95f0d66dd21e7d76a8f)
![{\displaystyle \alpha _{{\ce {H}}_{n-i}A^{i-}}={{[{\ce {H+}}]^{n-i}\displaystyle \prod _{j=0}^{i}K_{j}} \over {\displaystyle \sum _{i=0}^{n}{\Big [}[{\ce {H+}}]^{n-i}\displaystyle \prod _{j=0}^{i}K_{j}}{\Big ]}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/5a82c39009cacb7d891f1f06a46a6558c3c91e76)
