Acid strength weak acids

A major development in the study of EGB s is the recently reported measurements of rates of protonation by acids of known pK. The correlation of such rates with pK, the Bronsted relationship, also enables bases of determined pK to be used in the measurement of kinetic acidities of weak acids. This quantitative approach will eventually lead to the optimisation of reaction conditions for preparative reactions by providing data which can be used to match the acid/base pairs more exactly. In many organic reactions involving bases the base chosen is stronger than is strictly neccessary and consequently such reactions are often complicated by side reactions such as condensation reactions and isomerisations. The advantage of an EGB of moderate strength has been seen in the vitamin A preparation described in Scheme 18, where the facile cisftrans isomerisation is avoided. 

Acidification of these nitronic acid salts may lead to a variety of products depending upon the nature and strength of the acid employed. Weak acids, such as acetic or carbonic acids, simply regenerate the nitro-paraffin. Warm, concentrated mineral acids hydrolyze the salts of primary aci-nitroparaffins to produce fatty acids and salts of hydroxyl-amine. This is in reality a reaction of the nitronic acids and occurs when the primary nitroparaffins themselves are warmed with concentrated mineral acid.17... 

THE PROTOLYSIS OF ACIDS. STRENGTHS OF ACIDS AND BASES It is of interest to examine the processes which take place when an acid is dissolved in a solvent, first of all in water. According to the Br0nsted-Lowry theory this dissolution is accompanied by a protolytic reaction, in which the solvent (water) acts as a base. To elucidate these processes, let us examine what happens if a strong acid (hydrochloric acid) and a weak acid (acetic acid) undergo protolysis. 

Notably, as water has both acid and base properties it is termed amphiprotic and is the most common solvent for acid-base reactions. There are strong acids and weak acids and strong bases and weakbases . The primary feature determining the so-called strength of an acid or base relates to its degree of dissociation in solution, that is, the fraction that produces ions in solution. Importantly, acid or... 

Reactions of acids or bases with water are only one aspect of solvent effects. Any acid will react with a basic solvent and any base will react with an acidic solvent, with the extent of the reaction varying with their relative strengths. Eor example, acetic acid (a weak acid) will react with water to a very slight extent, but hydrochloric acid (a strong acid) reacts completely, both forming H3O, together with the acetate ion and chloride ion, respectively. 

The A5 double bond in steroids is relatively inert towards hydrogenation under neutral conditions, but can be activated by addition of acids. The catalytic effect is determined by the strength of the acid. Thus weak acids (pKa > 3) are ineffective, while strong acids (pA < 3) are efficient catalysts16. This allows regioselective hydrogenation of other double bonds which can be reduced under neutral conditions. Over platinum, the A3 double bond is hydrogenated nearly exclusively from the ot-face. The 5/1-product is only formed in the presence of 3a-substituents. 

The term strong add has nothing to do with the corrosive properties of an acid. It is also important to distinguish between acid strength (strong acid, weak acid) and... 

It is clear from the name that carboxylic acids are acidic. The hydrogen of the carboxyl group (-COOH) is acidic. The strength of these acids can be explained in terms of the relative stabilities of their conjugate bases. Weaker the conjugate base is, the stronger the acid. Keep in mind the fact that carboxylic acids are weak acids. 

So now we can expand our chart of acid and base strengths to include the important classes of alcohols, phenols, and carboxylic acids. They conveniently, and memorably, have piCa values of about 0 for the protonation of alcohols, about 5 for the deprotonation of carboxylic acids, about 10 for the deprotonation of phenols, and about 15 for the deprotonation of alcohols. The equilibria above each piCa shows that at approximately that pH, the two species each form 50% of the mixture. You can see that carboxylic acids are weak acids, alkoxide ions (RO ) are strong bases, and that it will need a strong acid to protonate an alcohol. 

Equation (5.14) shows that [H3O+] depends on C and also on its strength. It is through the Ka value that the acid s identity can be discerned. Contrary to strong acids, two weak acids at the same analytical concentration do not exhibit the same pH value apart from the accidental case in which both acids have the same pKa. The pH values of solutions of hydrocyanic acid (pKa = 932), acetic acid (pKa = 4.15), and hydrofluoric acid (pKa = 3.17) in relation with -log C are given in Fig. 5.2. 

The ionization of an acid, HA, is reflected in the acid dissociation constant, K. It depends on both the strength of the H—A bond and the stability of the conjugate base, A , in the solvent. Although acetic acid and other carboxylic acids are weak acids, they are far more acidic than alcohols or phenols. For example, the of acetic acid is about 10" times larger than the AT of ethanol. 

The strength of an acid is measured by the value of its dissociation constant, strong acids, e.g. HCl, HNO3. being substantially fully ionized in solution and weak acids predominately unionized. 

Since the hydrogen-element bond energy decreases from sulphur to tellurium they are stronger acids than hydrogen sulphide in aqueous solution but are still classified as weak acids—similar change in acid strength is observed for Group Vll hydrides. 

In dilute aqueous solution hydrogen fluoride is a weak acid but the acid strength increases with the concentration of hydrogen fluoride. 

The strength of a weak acid is measured by its acid dissociation constant, which IS the equilibrium constant for its ionization m aqueous solution... 

Note that the concentration of H2O is omitted from the expression because its value is so large that it is unaffected by the dissociation reaction.The magnitude of provides information about the relative strength of a weak acid, with a smaller corresponding to a weaker acid. The ammonium ion, for example, with a Ka of 5.70 X 10 °, is a weaker acid than acetic acid. 

All other things being equal, the strength of a weak acid increases if it is placed in a solvent that is more basic than water, whereas the strength of a weak base increases if it is placed in a solvent that is more acidic than water. In some cases, however, the opposite effect is observed. For example, the pKb for ammonia is 4.76 in water and 6.40 in the more acidic glacial acetic acid. In contradiction to our expectations, ammonia is a weaker base in the more acidic solvent. A full description of the solvent s effect on a weak acid s piQ or on the pKb of a weak base is beyond the scope of this text. You should be aware, however, that titrations that are not feasible in water may be feasible in a different solvent. 

Directions are provided in this experiment for determining the dissociation constant for a weak acid. Potentiometric titration data are analyzed by a modified Gran plot. The experiment is carried out at a variety of ionic strengths and the thermodynamic dissociation constant determined by extrapolating to zero ionic strength. 

Because they are weak acids or bases, the iadicators may affect the pH of the sample, especially ia the case of a poorly buffered solution. Variations in the ionic strength or solvent composition, or both, also can produce large uncertainties in pH measurements, presumably caused by changes in the equihbria of the indicator species. Specific chemical reactions also may occur between solutes in the sample and the indicator species to produce appreciable pH errors. Examples of such interferences include binding of the indicator forms by proteins and colloidal substances and direct reaction with sample components, eg, oxidising agents and heavy-metal ions. 

The bottoms from the stripper (40—60 wt % acid) are sent to an acid reconcentration unit for upgrading to the proper acid strength and recycling to the reactor. Because of the associated high energy requirements, reconcentration of the diluted sulfuric acid is a cosdy operation. However, a propylene gas stripping process, which utilizes only a small amount of added water for hydrolysis, has been described (63). In this modification, the equiUbrium quantity of isopropyl alcohol is stripped so that acid is recycled without reconcentration. Kquilibrium is attained rapidly at 50°C and isopropyl alcohol is removed from the hydrolysis mixture. Similarly, the weak sulfuric acid process minimizes the reconcentration of the acid and its associated corrosion and pollution problems. 

Effect on Oxide—Water Interfaces. The adsorption (qv) of ions at clay mineral and rock surfaces is an important step in natural and industrial processes. SiUcates are adsorbed on oxides to a far greater extent than would be predicted from their concentrations (66). This adsorption maximum at a given pH value is independent of ionic strength, and maximum adsorption occurs at a pH value near the piC of orthosiUcate. The pH values of maximum adsorption of weak acid anions and the piC values of their conjugate acids are correlated. This indicates that the presence of both the acid and its conjugate base is required for adsorption. The adsorption of sihcate species is far greater at lower pH than simple acid—base equihbria would predict. 

The characteristics of soluble sihcates relevant to various uses include the pH behavior of solutions, the rate of water loss from films, and dried film strength. The pH values of sihcate solutions are a function of composition and concentration. These solutions are alkaline, being composed of a salt of a strong base and a weak acid. The solutions exhibit up to twice the buffering action of other alkaline chemicals, eg, phosphate. An approximately linear empirical relationship exists between the modulus of sodium sihcate and the maximum solution pH for ratios of 2.0 to 4.0. 

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