Indicators, acid-base constant

Another problem is that different workers make their calculations of second-order rate constants in the micelle in different ways. For example, the surface potential of a micelle may be specifically included in the calculation in order to estimate ion binding, but there is a circularity in this argument because surface potentials are often estimated from micellar effects upon indicator acid-base equilibria which themselves depend upon ion-binding to micelles (Fernandez and Fromherz, 1977). 

Acid-base constants/proton-affinity constants These parameters cannot, at present, be established experimentally. Because there are probably several distinct sites and electrostatic effects, the overall surface-charge curves do not give any hint as to the nature and reactivity of these sites. Experimental back-titration data by Schulthess and Sparks [103], which indicate steps in the proton adsorption isotherm (i.e., the surface-charge density versus pH curve), were interpreted by these authors to be indicative of individual sites with individual reactivities. However, such data are not reproduced by the bulk of research groups. In the previous subsection, the back-titration technique has been discussed in some detail and, in particular, the results on goethite obtained with base titrations at low pH were found to show significant scatter compared to data obtained with the coulometric approach. 

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 acidity constant for an acid-base indicator was determined by preparing three solutions, each of which has a total indicator concentration of 5.00 X 10- M. The first solution was made strongly acidic with HCl and has an absorbance of 0.250. The second solution was made strongly basic and has an absorbance of 1.40. The pH of the third solution was measured at 2.91, with an absorbance of 0.662. What is the value of K, for the indicator ... 

A double end point, acid—base titration can be used to determine both sodium hydrosulfide and sodium sulfide content. Standardized hydrochloric acid is the titrant thymolphthalein and bromophenol blue are the indicators. Other bases having ionization constants in the ranges of the indicators used interfere with the analysis. Sodium thiosulfate and sodium thiocarbonate interfere quantitatively with the accuracy of the results. Detailed procedures to analyze sodium sulfide, sodium hydro sulfide, and sodium tetrasulfide are available (1). 

It turns out that in low-viscosity blending the acdual result does depend upon the measuring technique used to measure blend time. Two common techniques, wliich do not exhaust the possibilities in reported studies, are to use an acid-base indicator and inject an acid or base into the system that will result in a color change. One can also put a dye into the tank and measure the time for color to arrive at uniformity. Another system is to put in a conductivity probe and injecl a salt or other electrolyte into the system. With any given impeller type at constant power, the circulation time will increase with the D/T ratio of the impeller. Figure 18-18 shows that both circulation time and blend time decrease as D/T increases. The same is true for impeller speed. As impeller speed is increased with any impeller, blend time and circulation time are decreased (Fig. 18-19). 

The initial goal of the kinetic analysis is to express k as a function of [H ], pH-independent rate constants, and appropriate acid-base dissociation constants. Then numerical estimates of these constants are obtained. The theoretical pH-rate profile can now be calculated and compared with the experimental curve. A quantitative agreement indicates that the proposed rate equation is consistent with experiment. It is advisable to use other information (such as independently measured dissociation constants) to support the kinetic analysis. 

We see in Table 11-IV that the equilibrium view of acid strengths suggests that we regard water itself as a weak acid. It can release hydrogen ions and the extent to which it does so is indicated in its equilibrium constant, just as for the other acids. We shall see that this type of comparison, stimulated by our equilibrium considerations, leads us to a valuable generalization of the acid-base concept. 

The indicator colour change is affected by the hydrogen ion concentration of the solution, and no account of this has been taken in the above expression for the formation constant. Thus solochrome black, which may be written as H2In , exhibits the following acid-base behaviour ... 

Discussion. The dissociation of an acid-base indicator is well suited to spectrophotometric study the procedure involved will be illustrated by the determination of the acid dissociation constant of methyl red (MR). The acidic (HMR) and basic (MR-) forms of methyl red are shown below. 

The authors studied, as they call it, "acid-base equilibria in glacial acetic acid however, as they worked at various ratios of indicator-base concentration to HX or B concentration, we are in fact concerned with titration data. In this connection one should realize also that in solvents with low e the apparent strength of a Bronsted acid varies with the reference base used, and vice versa. Nevertheless, in HOAc the ionization constant predominates to such an extent that overall the picture of ionization vs. dissociation remains similar irrespective of the choice of reference see the data for I and B (Py) already given, and also those for HX, which the authors obtained at 25° C with I = p-naphthol-benzein (PNB) and /f B < 0.0042, i.e., for hydrochloric acid K C1 = 1.3 102, jjrfflci 3 9. IQ-6 an jjHC1 2.8 10 9 and for p-toluenesulphonic acid Kfm° = 3 7.102( K ms 4 0.10-6) Kmt = 7 3.10-9... 

An indication of the nature of the transition state in aromatic substitution is provided by the existence of some extrathermodynamic relationships among rate and acid-base equilibrium constants. Thus a simple linear relationship exists between the logarithms of the relative rates of halogenation of the methylbenzenes and the logarithms of the relative basicities of the hydrocarbons toward HF-BFS (or-complex equilibrium).288 270 A similar relationship with the basicities toward HC1 ( -complex equilibrium) is much less precise. The jr-complex is therefore a poorer model for the substitution transition state than is the [Pg.150]

Words that can be used as topics in essays acid-base indicator ion-product constant... 

With reference to a solvent, this term is usually restricted to Brpnsted acids. If the solvent is water, the pH value of the solution is a good measure of the proton-donating ability of the solvent, provided that the concentration of the solute is not too high. For concentrated solutions or for mixtures of solvents, the acidity of the solvent is best indicated by use of an acidity function. See Degree of Dissociation Henderson-Hasselbalch Equation Acid-Base Equilibrium Constants Bronsted Theory Lewis Acid Acidity Function Leveling Effect... 

Reflectance measurements provided an excellent means for building an ammonium ion sensor involving immobilization of a colorimetric acid-base indicator in the flow-cell depicted schematically in Fig. 3.38.C. The cell was furnished with a microporous PTFE membrane supported on the inner surface of the light window. The detection limit achieved was found to depend on the constant of the immobilized acid-base indicator used it was lO M for /7-Xylenol Blue (pAT, = 2.0). The response time was related to the ammonium ion concentration and ranged from 1 to 60 min. The sensor remained stable for over 6 months and was used to determine the analyte in real samples consisting of purified waste water, which was taken from a tank where the water was collected for release into the mimicipal waste water treatment plant. Since no significant interference fi-om acid compounds such as carbon dioxide or acetic acid was encountered, the sensor proved to be applicable to real samples after pH adjustment. The ammonium concentrations provided by the sensor were consistent with those obtained by ion chromatography, a spectrophotometric assay and an ammonia-selective electrode [269]. 

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