Weak acids indicators

For a weak acid indicator, indigo carmine, the acid form, HIn, is blue and the base form, In, is yellow. 

Fig. 5 Schematic representation of locations of a weakly acidic indicator, , adjacent to the cationic head groups, , and of a functional hydroxyl group in the water-rich region...

Notice that H2SO4 behaves as a strong acid (Step 1) and HS04 behaves as a weak acid, indicated by a double arrow (Step 2). 

HS denotes the solvent, H2S the lyonium and 1 the conjugate base of the weak acid indicator HI HI and 1 have different colors. The equilibrium constant of the acid-base interaction can be expressed in the following logarithmic form where yj- represents... 

In Section 9-6, we learned that an indicator is an acid or base whose various proton-ated species have different colors. For the weak-acid indicator HIn, the solution takes on the color of HIn when pH pA Hin 1 and has the color of In when pH pA Hin + 1. In the interval pATnin - 1 pH pA Hin + 1, a mixture of both colors is observed. 

Indicators change colors because they are either weak acids or weak bases. In solution, a weak-acid indicator (H7n) can be represented by the equation below, which is modeled in Figure 2.1. 

Epoxy ring opening is catalyzed by acids and bases. Evaluations of the role of ammonium ions as weak acids indicate that they do catalyze the epoxy cure. This can lead to a lowering of the Tg of the polymer. This compromises mechanical performance. TEM and mechanical properties indicate that the exfoliation and particle alignment of montmorillonite in the epoxy polymer in the above references are also significant variables. 

The addition of even a weak acid (such as ethanoic acid) to a nitrite produces nitrous acid which readily decomposes as already indicated. Hence a nitrite is distinguished from a nitrate by the evolution of nitrous fumes when ethanoic acid is added. 

The use of methyl-orange as an indicator is based on the fact that the sodium salt in aqueous solution furnishes a yellow anion, which when treated with acids (except weak acids such as H2CO3 and H SOj) apparently gives rise to a red... 

It has been shown that for most acid-base titrations the inflection point, which corresponds to the greatest slope in the titration curve, very nearly coincides with the equivalence point. The inflection point actually precedes the equivalence point, with the error approaching 0.1% for weak acids or weak bases with dissociation constants smaller than 10 , or for very dilute solutions. Equivalence points determined in this fashion are indicated on the titration curves in figure 9.8. 

The plT at which an acid-base indicator changes color is determined by its acid dissociation constant. For an indicator that is a monoprotic weak acid, ITIn, the following dissociation reaction occurs... 

The pH range of an indicator does not have to be equally distributed on either side of the indicator s piQ. For some indicators only the weak acid or weak base is colored. For other indicators both the weak acid and weak base are colored, but one form may be easier to see. In either case, the pH range is skewed toward those pH levels for which the less colored form of the indicator is present in higher concentration. 

Tartaric acid, H2C4H4O6, is a diprotic weak acid with a pK i of 3.0 and a pK 2 of 4.4. Suppose you have a sample of impure tartaric acid (%purity > 80) and that you plan to determine its purity by titrating with a solution of 0.1 M NaOH using a visual indicator to signal the end point. Describe how you would carry out the analysis, paying particular attention to how much sample you would use, the desired pH range over which you would like the visual indicator to operate, and how you would calculate the %w/w tartaric acid. 

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. 

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 weakly acidic character of acycHc polyhydric alcohols increases with the number of hydroxyl groups, as indicated by the piC values in aqueous solution at 18°C (13). 

Kinetic mles of oxidation of MDASA and TPASA by periodate ions in the weak-acidic medium at the presence of mthenium (VI), iridium (IV), rhodium (III) and their mixtures are investigated by spectrophotometric method. The influence of high temperature treatment with mineral acids of catalysts, concentration of reactants, interfering ions, temperature and ionic strength of solutions on the rate of reactions was investigated. Optimal conditions of indicator reactions, rate constants and energy of activation for arylamine oxidation reactions at the presence of individual catalysts are determined. 

The concentration of acid impurities is an important indication of the quality of petroleum products and the purity of organic solvents, plasticizers, mineral oils, food fats, and polymers. Methods are used to detect organic acids in such compounds have many disadvantages the alkalimetry - low sensitivity, especially in the determination of weak acids, the extraction-photometric method is laborious, instmmental methods are expensive. In addition, most of methods are commonly unsuitable for direct analysis. 

Examples include hydrochloric acid, nitric acid, and sulphuric acid. These are strong acids which are almost completely dissociated in water. Weak acids, such as hydrogen sulphide, are poorly dissociated producing low concentrations of hydrogen ions. Acids tend to be coiTosive with a sharp, sour taste and turn litmus paper red they give distinctive colour changes with other indicators. Acids dissolve metals such as copper and liberate hydrogen gas. They also react with carbonates to liberate carbon dioxide ... 

PK. — the negative logarithm of the equilibrium constant for acids or bases. This parameter is an indicator of the strength of an acid or base. Strong acids, such as H2SO4, and HCl, have low pK s (i.e., -1.0) while strong bases such as KOH and NaOH, have pK s close to 14.0. Weak acids and weak bases fall in the intermediate range. 

The reaction can be first or second order with respect to the H+ concentration. In weakly acidic solution it is first order in [H+], but second order in strongly acidic solution. This indicates that the monoprotonated as well as the diprotonated hydrazobenzene can undergo rearrangement. 

Direct titration [119,120] In order to know the number of exchangeable hydrogen ions at different dissociation stayes, various salts of weak acids were used. Data observed in Table 8 shows an increase in the exchange capacity with an increase in the pH of the solution, indicating the presence of weak acid capacity [118]. 

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