Thiol is an organosulfur compound of the form R-SH, where R represents an alkyl or other organic substituent. Thiols are the sulfur analogue of alcohols, the word is a portmanteau of "thion" + "alcohol," with the first word deriving from Greek θεῖον = "sulfur"; the -- SH functional group itself is referred to as either a sulfhydryl group. Many thiols have strong odors resembling that of rotten eggs. Thiols are used as odorants to assist in the detection of natural gas, the "smell of natural gas" is due to the smell of the thiol used as the odorant. Thiols are sometimes referred to as mercaptans; the term "mercaptan" was introduced in 1832 by William Christopher Zeise and is derived from the Latin mercurium captāns because the thiolate group bonds strongly with mercury compounds. Thiols and alcohols have similar connectivity; because sulfur is a larger element than oxygen, the C–S bond lengths – around 180 picometres in length – is about 40 picometers longer than a typical C–O bond. The C -- S -- H angles approach 90 °.
In the solid or liquids, the hydrogen-bonding between individual thiol groups is weak, the main cohesive force being van der Waals interactions between the polarizable divalent sulfur centers. The S-H bond is much weaker than the O-H bond as reflected in their respective bond dissociation energy. For CH3S-H, the BDE is 366 kJ/mol. Due to the small difference in the electronegativity of sulfur and hydrogen, an S–H bond is polar. In contrast, O-H bonds in hydroxyl groups are more polar. Thiols have a lower dipole moment relative to the corresponding alcohol. There are several ways to name the alkylthiols: The suffix -thiol is added to the name of the alkane; this method is nearly identical to naming an alcohol and is used by the IUPAC, e.g. CH3SH would be methanethiol; the word mercaptan replaces alcohol in the name of the equivalent alcohol compound. Example: CH3SH would be methyl mercaptan, just as CH3OH is called methyl alcohol; the term sulfanyl or mercapto is used as e.g. mercaptopurine. Many thiols have strong odors resembling that of garlic.
The odors of thiols those of low molecular weight, are strong and repulsive. The spray of skunks consists of low-molecular-weight thiols and derivatives; these compounds are detectable by the human nose at concentrations of only 10 parts per billion. Human sweat contains /-3-methyl-3-sulfanylhexan-1-ol, detectable at 2 parts per billion and having a fruity, onion-like odor. Methanethiol is a strong-smelling volatile thiol detectable at parts per billion levels, found in male mouse urine. Lawrence C. Katz and co-workers showed that MTMT functioned as a semiochemical, activating certain mouse olfactory sensory neurons, attracting female mice. Copper has been shown to be required by a specific mouse olfactory receptor, MOR244-3, responsive to MTMT as well as to various other thiols and related compounds. A human olfactory receptor, OR2T11, has been identified which, in the presence of copper, is responsive to the gas odorants ethanethiol and t-butyl mercaptan as well as other low molecular weight thiols, including allyl mercaptan found in human garlic breath, the strong-smelling cyclic sulfide thietane.
Thiols are responsible for a class of wine faults caused by an unintended reaction between sulfur and yeast and the "skunky" odor of beer, exposed to ultraviolet light. Not all thiols have unpleasant odors. For example, furan-2-ylmethanethiol contributes to the aroma of roasted coffee, whereas grapefruit mercaptan, a monoterpenoid thiol, is responsible for the characteristic scent of grapefruit; the effect of the latter compound is present only at low concentrations. The pure mercaptan has an unpleasant odor. Natural gas distributors were required to add thiols ethanethiol, to natural gas after the deadly New London School explosion in New London, Texas, in 1937. Many gas distributors were odorizing gas prior to this event. Most gas odorants utilized contain mixtures of mercaptans and sulfides, with t-butyl mercaptan as the main odor constituent in natural gas and ethanethiol in liquefied petroleum gas. In situations where thiols are used in commercial industry, such as liquid petroleum gas tankers and bulk handling systems, an oxidizing catalyst is used to destroy the odor.
A copper-based oxidation catalyst neutralizes the volatile thiols and transforms them into inert products. Thiols show little association both with water molecules and among themselves. Hence, they have lower boiling points and are less soluble in water and other polar solvents than alcohols of similar molecular weight. For this reason thiols and corresponding thioether functional group isomers have similar solubility characteristics and boiling points, whereas the same is not true of alcohols and their corresponding isomeric ethers; the S-H bond in thiols is weak compared to the O-H bond in alcohols. For CH3X-H, the bond enthalpies are 365.07 for X = S and 440.2 kcal/mol for X = O. H-atom abstraction from a thiol gives a thiyl radical with the formula RS. where R = alkyl or aryl. Volatile thiols are and unerringly detected by their distinctive odor. S-specific analyzers for gas chromatographs are useful. Spectroscopic indicators are the D2O-exchangeable SH signal in the 1H NMR spectrum; the νSH band appears near 2400 cm−
Nucleophile is a chemical species that donates an electron pair to form a chemical bond in relation to a reaction. All molecules or ions with a free pair of electrons or at least one pi bond can act as nucleophiles; because nucleophiles donate electrons, they are by definition Lewis bases. Nucleophilic describes the affinity of a nucleophile to the nuclei. Nucleophilicity, sometimes referred to as nucleophile strength, refers to a substance's nucleophilic character and is used to compare the affinity of atoms. Neutral nucleophilic reactions with solvents such as alcohols and water are named solvolysis. Nucleophiles may take part in nucleophilic substitution, whereby a nucleophile becomes attracted to a full or partial positive charge; the terms nucleophile and electrophile were introduced by Christopher Kelk Ingold in 1933, replacing the terms anionoid and cationoid proposed earlier by A. J. Lapworth in 1925; the word nucleophile is derived from philos for love. In general, in a row across the periodic table, the more basic the ion the more reactive it is as a nucleophile.
Within a series of nucleophiles with the same attacking element, the order of nucleophilicity will follow basicity. Sulfur is in general a better nucleophile than oxygen. Many schemes attempting to quantify relative nucleophilic strength have been devised; the following empirical data have been obtained by measuring reaction rates for a large number of reactions involving a large number of nucleophiles and electrophiles. Nucleophiles displaying the so-called alpha effect are omitted in this type of treatment; the first such attempt is found in the Swain–Scott equation derived in 1953: log 10 = s n This free-energy relationship relates the pseudo first order reaction rate constant, k, of a reaction, normalized to the reaction rate, k0, of a standard reaction with water as the nucleophile, to a nucleophilic constant n for a given nucleophile and a substrate constant s that depends on the sensitivity of a substrate to nucleophilic attack. This treatment results in the following values for typical nucleophilic anions: acetate 2.7, chloride 3.0, azide 4.0, hydroxide 4.2, aniline 4.5, iodide 5.0, thiosulfate 6.4.
Typical substrate constants are 0.66 for ethyl tosylate, 0.77 for β-propiolactone, 1.00 for 2,3-epoxypropanol, 0.87 for benzyl chloride, 1.43 for benzoyl chloride. The equation predicts that, in a nucleophilic displacement on benzyl chloride, the azide anion reacts 3000 times faster than water; the Ritchie equation, derived in 1972, is another free-energy relationship: log 10 = N + where N+ is the nucleophile dependent parameter and k0 the reaction rate constant for water. In this equation, a substrate-dependent parameter like s in the Swain–Scott equation is absent; the equation states that two nucleophiles react with the same relative reactivity regardless of the nature of the electrophile, in violation of the reactivity–selectivity principle. For this reason, this equation is called the constant selectivity relationship. In the original publication the data were obtained by reactions of selected nucleophiles with selected electrophilic carbocations such as tropylium or diazonium cations: or ions based on Malachite green.
Many other reaction types have since been described. Typical Ritchie N+ values are: 0.5 for methanol, 5.9 for the cyanide anion, 7.5 for the methoxide anion, 8.5 for the azide anion, 10.7 for the thiophenol anion. The values for the relative cation reactivities are −0.4 for the malachite green cation, +2.6 for the benzenediazonium cation, +4.5 for the tropylium cation. In the Mayr-Patz equation: log = s The second order reaction rate constant k at 20 °C for a reaction is related to a nucleophilicity parameter N, an electrophilicity parameter E, a nucleophile-dependent slope parameter s; the constant s is defined as 1 with 2-methyl-1-pentene as the nucleophile. Many of the constants have been derived from reaction of so-called benzhydrylium ions as the electrophiles: and a diverse collection of π-nucleophiles:. Typical E values are +6.2 for R = chlorine, +5.90 for R = hydrogen, 0 for R = methoxy and -7.02 for R = dimethylamine. Typical N values with s in parenthesis are -4.47 for electrophilic aromatic substitution to toluene, -0.41 for electrophilic addition to 1-phenyl-2-propene, 0.96 for addition to 2-methyl-1-pentene, -0.13 for reaction with triphenylallylsilane, 3.61 for reaction with 2-methylfuran, +7.48 for reaction with isobutenyltributylstannane and +13.36 for reaction with the enamine 7.
The range of organic reactions include SN2 reactions: With E = -9.15 for the S-methyldibenzothiophenium ion, typical nucleophile values N are 15.63 for piperidine, 10.49 for methoxide, 5.20 for water. In short, nucleophilicities towards sp2 or sp3 centers follow the same pattern. In an effort to unify the above described equations the Mayr equation is rewritten as: log = s E s N ( N + E
In organic chemistry and biochemistry, a side chain is a chemical group, attached to a core part of the molecule called the "main chain" or backbone. The side chain is a hydrocarbon branching element of a molecule, attached to a larger hydrocarbon backbone, it is one factor in determining reactivity. A side chain is known as a pendant chain, but a pendant group has a different definition; the placeholder R is used as a generic placeholder for alkyl group side chains in chemical structure diagrams. To indicate other non-carbon groups in structure diagrams, X, Y, or Z are used; the R symbol was introduced by 19th-century French chemist Charles Frédéric Gerhardt, who advocated its adoption on the grounds that it would be recognizable and intelligible given its correspondence in multiple European languages to the initial letter of one or more words used to denote the concept and sharing the meaning "root" or "residue": French racine and résidu, these terms' respective English translations along with radical, Latin radix and residuum, German Rest.
In polymer science, the side chain of an oligomeric or polymeric offshoot extends from the backbone chain of a polymer. Side chains have noteworthy influence on a polymer's properties its crystallinity and density. An oligomeric branch may be termed a short-chain branch, a polymeric branch may be termed a long-chain branch. Side groups are different from side chains. In proteins, which are composed of amino acid residues, the side chains are attached to the alpha-carbon atoms of the amide backbone; the side chain connected to the alpha-carbon is specific for each amino acid and is responsible for determining charge and polarity of the amino acid. The amino acid side chains are responsible for many of the interactions that lead to proper protein folding and function. Amino acids with similar polarity are attracted to each other, while nonpolar and polar side chains repel each other. Nonpolar/polar interactions can still play an important part in stabilizing the secondary structure due the large amount of them occurring throughout the protein.
Alkyl Backbone chain Branching Functional group Pendant group Substituent
Urea known as carbamide, is an organic compound with chemical formula CO2. This amide has two –NH2 groups joined by a carbonyl functional group. Urea serves an important role in the metabolism of nitrogen-containing compounds by animals and is the main nitrogen-containing substance in the urine of mammals, it is a colorless, odorless solid soluble in water, non-toxic. Dissolved in water, it is neither alkaline; the body uses it in most notably nitrogen excretion. The liver forms it by combining two ammonia molecules with a carbon dioxide molecule in the urea cycle. Urea is used in fertilizers as a source of nitrogen and is an important raw material for the chemical industry. Friedrich Wöhler's discovery in 1828 that urea can be produced from inorganic starting materials was an important conceptual milestone in chemistry, it showed for the first time that a substance known only as a byproduct of life could be synthesized in the laboratory without biological starting materials thereby contradicting the held doctrine of vitalism.
More than 90% of world industrial production of urea is destined for use as a nitrogen-release fertilizer. Urea has the highest nitrogen content of all solid nitrogenous fertilizers in common use. Therefore, it has the lowest transportation costs per unit of nitrogen nutrient. Many soil bacteria possess the enzyme urease, which catalyzes conversion of urea to ammonia or ammonium ion and bicarbonate ion, thus urea fertilizers transform to the ammonium form in soils. Among the soil bacteria known to carry urease, some ammonia-oxidizing bacteria, such as species of Nitrosomonas, can assimilate the carbon dioxide the reaction releases to make biomass via the Calvin cycle, harvest energy by oxidizing ammonia to nitrite, a process termed nitrification. Nitrite-oxidizing bacteria Nitrobacter, oxidize nitrite to nitrate, mobile in soils because of its negative charge and is a major cause of water pollution from agriculture. Ammonium and nitrate are absorbed by plants, are the dominant sources of nitrogen for plant growth.
Urea is used in many multi-component solid fertilizer formulations. Urea is soluble in water and is therefore very suitable for use in fertilizer solutions, e.g. in'foliar feed' fertilizers. For fertilizer use, granules are preferred over prills because of their narrower particle size distribution, an advantage for mechanical application; the most common impurity of synthetic urea is biuret. Urea is spread at rates of between 40 and 300 kg/ha but rates vary. Smaller applications incur lower losses due to leaching. During summer, urea is spread just before or during rain to minimize losses from volatilization; because of the high nitrogen concentration in urea, it is important to achieve an spread. The application equipment must be calibrated and properly used. Drilling must not occur on contact with or close to seed, due to the risk of germination damage. Urea dissolves in water for application through irrigation systems. In grain and cotton crops, urea is applied at the time of the last cultivation before planting.
In high rainfall areas and on sandy soils and where good in-season rainfall is expected, urea can be side- or top-dressed during the growing season. Top-dressing is popular on pasture and forage crops. In cultivating sugarcane, urea is side-dressed after planting, applied to each ratoon crop. In irrigated crops, urea can be applied dry to the soil, or dissolved and applied through the irrigation water. Urea dissolves in its own weight in water, but becomes difficult to dissolve as the concentration increases. Dissolving urea in water is endothermic—the solution temperature falls when urea dissolves; as a practical guide, when preparing urea solutions for fertigation, dissolve no more than 3 g urea per 1 L water. In foliar sprays, urea concentrations of between 0.5% and 2.0% are used in horticultural crops. Low-biuret grades of urea are indicated. Urea absorbs moisture from the atmosphere and therefore is stored either in closed or sealed bags on pallets or, if stored in bulk, under cover with a tarpaulin.
As with most solid fertilizers, storage in a cool, well-ventilated area is recommended. Overdose or placing urea near seed is harmful. Urea is a raw material for the manufacture of two main classes of materials: urea-formaldehyde resins and urea-melamine-formaldehyde used in marine plywood. Urea can be used to make urea nitrate, a high explosive, used industrially and as part of some improvised explosive devices, it is a stabilizer in nitrocellulose explosives. Urea is used in SNCR and SCR reactions to reduce the NOx pollutants in exhaust gases from combustion from Diesel, dual fuel, lean-burn natural gas engines; the BlueTec system, for example, injects a water-based urea solution into the exhaust system. The ammonia produced by the hydrolysis of the urea reacts with the nitrogen oxide emissions and is converted into nitrogen and water within the catalytic converter. Trucks and cars using these catalytic converters need to carry a supply of diesel exhaust fluid, a solution of urea in water. Urea in concentrations up to 10 M is a powerful protein denaturant as it disrupts the noncovalent bonds in the proteins.
This property can be exploited to increase the solubility of some proteins. A mixture of urea and choline chloride is used as
H2 antagonists, sometimes referred to as H2RA and called H2 blockers, are a class of medications that block the action of histamine at the histamine H2 receptors of the parietal cells in the stomach. This decreases the production of stomach acid. H2 antagonists can be used in the treatment of dyspepsia, peptic ulcers and gastroesophageal reflux disease, they have been surpassed by proton pump inhibitors. H2 antagonists are a type of antihistamine, although in common use the term "antihistamine" is reserved for H1 antagonists, which relieve allergic reactions. Like the H1 antagonists, some H2 antagonists function as inverse agonists rather than receptor antagonists, due to the constitutive activity of these receptors; the prototypical H2 antagonist, called cimetidine, was developed by Sir James Black at Smith, Kline & French – now GlaxoSmithKline – in the mid-to-late 1960s. It was first marketed in 1976 and sold under the trade name Tagamet, which became the first blockbuster drug; the use of quantitative structure-activity relationships led to the development of other agents – starting with ranitidine, first sold as Zantac, which has fewer adverse effects and drug interactions and is more potent.
Cimetidine ranitidine famotidine nizatidine roxatidine lafutidine Cimetidine was the prototypical histamine H2-receptor antagonist from which drugs were developed. Cimetidine was the culmination of a project at Smith, Kline & French by James W. Black, C. Robin Ganellin, others to develop a histamine receptor antagonist that would suppress stomach acid secretion. In 1964, it was known that histamine stimulated the secretion of stomach acid, that traditional antihistamines had no effect on acid production. From these facts the SK&F scientists postulated the existence of two different types of histamine receptors, they designated the one acted upon by the traditional antihistamines as H1, the one acted upon by histamine to stimulate the secretion of stomach acid as H2. The SK&F team used a classical design process starting from the structure of histamine. Hundreds of modified compounds were synthesised in an effort to develop a model of the then-unknown H2 receptor; the first breakthrough was Nα-guanylhistamine, a partial H2-receptor antagonist.
From this lead, the receptor model was further refined, which led to the development of burimamide, a specific competitive antagonist at the H2 receptor. Burimamide is 100 times more potent than Nα-guanylhistamine, proving its efficacy on the H2 receptor; the potency of burimamide was still too low for oral administration. And efforts on further improvement of the structure, based on the structure modification in the stomach due to the acid dissociation constant of the compound, led to the development of metiamide. Metiamide was an effective agent, it was proposed that the toxicity arose from the thiourea group, similar guanidine analogues were investigated until the discovery of cimetidine, which would become the first clinically successful H2 antagonist. Ranitidine was developed by Glaxo, in an effort to match the success of Smith, Kline & French with cimetidine. Ranitidine was the result of a rational drug design process utilising the by-then-fairly-refined model of the histamine H2 receptor and quantitative structure-activity relationships.
Glaxo refined the model further by replacing the imidazole-ring of cimetidine with a furan-ring with a nitrogen-containing substituent, in doing so developed ranitidine. Ranitidine was found to have a far-improved tolerability profile, longer-lasting action, ten times the activity of cimetidine. Ranitidine was introduced in 1981 and was the world's biggest-selling prescription drug by 1988; the H2-receptor antagonists have since been superseded by the more effective proton pump inhibitors, with omeprazole becoming the biggest-selling drug for many years. The H2 antagonists are competitive antagonists of histamine at the parietal cell's H2 receptor, they suppress the normal secretion of acid by parietal cells and the meal-stimulated secretion of acid. They accomplish this by two mechanisms: Histamine released by ECL cells in the stomach is blocked from binding on parietal cell H2 receptors, which stimulate acid secretion. H2-antagonists are used by clinicians in the treatment of acid-related gastrointestinal conditions, including: Peptic ulcer disease Gastroesophageal reflux disease Dyspepsia Prevention of stress ulcer Prevention of aspiration pneumonitis during surgery.
Oral H2-antagonists reduce gastric acidity and volume and have shown to reduce the frequency of aspiration pneumonitis. People who suffer from infrequent heartburn may take either antacids or H2-receptor antagonists for treatment; the H2-antagonists offer several advantages over antacids, including longer duration of action, greater efficacy, ability to be used prophylactically before meals to reduce the chance of heartburn occurring. Proton pump inhibitors, are the preferred treatment for erosive esophagitis since they have been shown to promote healing better than H2-antagonists. H2 antagonists are, in general, well-tolerated, except for cimetidine
Imidazole is an organic compound with the formula C3N2H4. It is a white or colourless solid, soluble in water, producing a mildly alkaline solution. In chemistry, it is an aromatic heterocycle, classified as a diazole, has non-adjacent nitrogen atoms. Many natural products alkaloids, contain the imidazole ring; these imidazoles feature varied substituents. This ring system is present in important biological building blocks, such as histidine and the related hormone histamine. Many drugs contain an imidazole ring, such as certain antifungal drugs, the nitroimidazole series of antibiotics, the sedative midazolam; when fused to a pyrimidine ring, it forms a purine, the most occurring nitrogen-containing heterocycle in nature. The name "imidazole" was coined in 1887 by the German chemist Arthur Rudolf Hantzsch. Imidazole is a planar 5-membered ring, it exists in two equivalent tautomeric forms, because hydrogen can be bound to one or the other nitrogen atom. Imidazole is a polar compound, as evidenced by its electric dipole moment of 3.67 D.
It is soluble in water. The compound is classified as aromatic due to the presence of a planar ring containing 6 π-electrons; some resonance structures of imidazole are shown below: Imidazole is amphoteric. That is, it can function as both an acid and as a base; as an acid, the pKa of imidazole is 14.5, making it less acidic than carboxylic acids and imides, but more acidic than alcohols. The acidic proton is the one bound to nitrogen. Deprotonation gives the imidazole anion, symmetrical; as a base, the pKa of the conjugate acid is 7, making imidazole sixty times more basic than pyridine. The basic site is the nitrogen with the lone pair. Protonation gives the imidazolium cation, symmetrical. Imidazole was first reported in 1858 by the German-British chemist Heinrich Debus, although various imidazole derivatives had been discovered as early as the 1840s, it was shown that glyoxal and ammonia condense to form imidazole. This synthesis, while producing low yields, is still used for generating C-substituted imidazoles.
In one microwave modification, the reactants are benzil and ammonia in glacial acetic acid, forming 2,4,5-triphenylimidazole. Imidazole can be synthesized by numerous methods besides the Debus method. Many of these syntheses can be applied to substituted imidazoles by varying the functional groups on the reactants; these methods are categorized by the number of reacting components. One componentThe or bond can be formed by the reaction of an imidate and an α-aminoaldehyde or α-aminoacetal; the example below applies to imidazole. Two componentThe and bonds can be formed by treating a 1,2-diaminoalkane, at high temperatures, with an alcohol, aldehyde, or carboxylic acid. A dehydrogenating catalyst, such as platinum on alumina, is required; the and bonds can be formed from N-substituted α-aminoketones and formamide with heat. The product will be a 1,4-disubstituted imidazole, but here since R1 = R2 = hydrogen, imidazole itself is the product; the yield of this reaction is moderate, but it seems to be the most effective method of making the 1,4 substitution.
Three componentThis method proceeds in good yields for substituted imidazoles. An adaptation of the Debus method, it is called the Debus-Radziszewski imidazole synthesis; the starting materials are substituted glyoxal, aldehyde and ammonia or an ammonium salt. Formation from other heterocyclesImidazole can be synthesized by the photolysis of 1-vinyltetrazole; this reaction will give substantial yields only if the 1-vinyltetrazole is made efficiently from an organotin compound, such as 2-tributylstannyltetrazole. The reaction, shown below, produces imidazole. Imidazole can be formed in a vapor-phase reaction; the reaction occurs with formamide and hydrogen over platinum on alumina, it must take place between 340 and 480 °C. This forms a pure imidazole product. Van Leusen reactionThe Van Leusen reaction can be employed to form imidazoles starting from TosMIC and an aldimine; the Van Leusen Imidazole Synthesis allows the preparation of imidazoles from aldimines by reaction with tosylmethyl isocyanide.
The reaction has been expanded to a two-step synthesis in which the aldimine is generated in situ: the Van Leusen Three-Component Reaction. Imidazole is incorporated into many important biological molecules; the most pervasive is the amino acid histidine. Histidine is present in many proteins and enzymes and plays a vital part in the structure and binding functions of hemoglobin. Imidazole-based histidine compounds play a important role in intracellular buffering. Histidine can be decarboxylated to histamine, a common biological compound. Histamine can cause urticaria; the relationship between histidine and histamine is shown below: One of the applications of imidazole is in the purification of His-tagged proteins in immobilised metal affinity chromatography. Imidazole is used to elute tagged proteins bound to nickel ions attached to the surface of beads in the chromatography column. An excess of imidazole is passed through the column, which displaces the His-tag from nickel coordination, freeing the His-tagged proteins.
Imidazole has become an important part of many pharmace
Acid dissociation constant
An acid dissociation constant, Ka, is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid–base reactions. K a =; the chemical species HA, A−, H+ are said to be in equilibrium when their concentrations do not change with the passing of time, because both forward and backward reactions are occurring at the same fast rate. The chemical equation for acid dissociation can be written symbolically as: HA ↽ − − ⇀ A − + H + where HA is a generic acid that dissociates into A−, the conjugate base of the acid and a hydrogen ion, H+, it is implicit in this definition that the quotient of activity coefficients, Γ, Γ = γ A − γ H + γ A H is a constant that can be ignored in a given set of experimental conditions. For many practical purposes it is more convenient to discuss the logarithmic constant, pKa p K a = − log 10 The more positive the value of pKa, the smaller the extent of dissociation at any given pH —that is, the weaker the acid.
A weak acid has a pKa value in the approximate range −2 to 12 in water. For a buffer solution consisting of a weak acid and its conjugate base, pKa can be expressed as: p K a = pH − log 10 The pKa for a weak monoprotic acid is conveniently determined by potentiometric titration with a strong base to the equivalence point and taking the pH value measured at one-half this volume as being equal to pKa; that is because at this half equivalence point, the number of moles of strong base added is one-half the number of moles of weak acid present, while the concentrations of the conjugate base and the remaining weak acid are the same. Acids with a pKa value of less than about −2 are said to be strong acids. In water, the dissociation of a strong acid in dilute solutions is complete such that the final concentration of the undissociated acid final is low. Consider a strong monoprotic acid, such as HCl; because of their 1:1 ratio, the final concentration of the conjugate base, final, is taken to be equal to the concentration of the hydronium ion, which can be directly measured by a pH meter.
For strong monoprotic acids like HCl, final and are both nearly equal to the initial concentration of initial placed into solution. With conventional acid-base titration methods it is difficult to measure the pH of a strong acid solution and, hence, to determine the or final, with a sufficient number of significant figures to and compute the low values encountered for final, which can be as low as 10-9 mol per liter for some strong acids. Furthermore, if 100% dissociation is assumed, final is zero and the fraction within parenthesis in the equation above becomes undefined; because the second expression on the right-hand side of the above equation is therefore indeterminable by conventional titration methods, the entire equation is not as useful a means of experimentally measuring pKa for strong acids as it is for weak acids. However, pKa and/or Ka values for strong acids can be estimated by theoretical means, such as computing gas phase dissociation constants and using Gibbs free energies of solvation for the molecular anions.
It is possible to use spectroscopy in some cases to determine the ratio of the concentrations of the conjugate base produced and the undissociated acid. For example, the Raman spectra of dilute nitric acid solutions contain signals of the nitrate ion and as the solutions become more concentrated signals of undissociated nitric acid molecules emerge; the acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction. The value of the pKa changes with temperature and can be understood qualitatively based on Le Châtelier's principle: when the reaction is endothermic, Ka increases and pKa decreases with