Acid-basicity equilibrium

Fig. 12.2 Factors that influence the pKj. (A) The acid-basic equilibrium governed by the /fa- (B) The acid-base equilibrium in the case...

The materialization of an electric charge in a side chain of a given ELP due to acid-basic equilibrium has been considered in the literature as a highly ef-... 

By comparison, all nutrients are taken up in the liquid state by biomass, as are the products released (Figure 6). The gas-liquid transfer as well as the acid-basic equilibrium of C02,iiq/HC03, iiq does not take place inside the cell but at the interface of the liquid and gaseous phases in the reactor. The heat related to these transfer phenomena is produced outside the cells but inside the reactor boundary. The uptake or release of species in the liquid phase (in dynamic equilibrium with the gaseous phase) is due to the cellular metabolism. 

The small differences m basicity between ammonia and alkylammes and among the various classes of alkylammes (primary secondary tertiary) come from a mix of effects Replacing hydrogens of ammonia by alkyl groups affects both sides of the acid-base equilibrium m ways that largely cancel... 

In their acidity, basicity, and the directive influence exerted on electrophilic substitution reactions in benzenoid nuclei, acylamino groups show properties which are intermediate between those of free amino and hydroxyl groups, and, therefore, it is at first surprising to find that the tautomeric behavior of acylaminopyridines closely resembles that of the aminopyridines instead of being intermediate between that of the amino- and hydroxy-pyridines. The basicities of the acylaminopyridines are, indeed, closer to those of the methoxy-pyridines than to those of the aminopyridines, the position of the tautomeric equilibrium being determined by the fact that the acyl-iminopyridones are strong bases like the iminopyridones and unlike the pyridones themselves. Thus, relative to the conversion of an... 

Basic hydrolysis occurs by nucleophilic addition of OH- to the amide carbonyl group, followed by elimination of amide ion (-NH2) and subsequent deprotonation of the initially formed carboxylic acid by amide ion. The steps are reversible, with the equilibrium shifted toward product by the final deprotonation of the carboxylic acid. Basic hydrolysis is substantially more difficult than the analogous acid-catalyzed reaction because amide ion is a very poor leaving group, making the elimination step difficult. 

All sodium salts are soluble, and so are all nitrate salts, so It makes sense that neither of these ions participates in a solubility equilibrium. Furthermore, nitrate and sodium cations are neither acidic nor basic, so it makes sense that neither participates in an acid-base equilibrium. 

C17-0025. Write the acid-base equilibrium that determines the pH of aqueous solutions of each of the following salts, and state whether the resulting solution is acidic, basic, or neither (a) NH4 I (b) NaClOq and (c) NaCHg CO2. ... 

The acidic and basic equilibrium constants Ka and Kb are given in Table 3. 

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]

Our goal in this chapter is to help you understand the equilibrium systems involving acids and bases. If you don t recall the Arrhenius acid-base theory, refer to Chapter 4 on Aqueous Solutions. You will learn a couple of other acid-base theories, the concept of pH, and will apply those basic equilibrium techniques we covered in Chapter 14 to acid-base systems. In addition, you will need to be familiar with the log and 10 functions of your calculator. And, as usual, in order to do well you must Practice, Practice, Practice. 

Our goal in this chapter is to help you continue learning about acid-base equilibrium systems and, in particular, buffers and titrations. If you are a little unsure about equilibria and especially weak acid-base equilibria, review Chapters 14 and 15. You will also learn to apply the basic concepts of equilibria to solubility and complex ions. Two things to remember (1) The basic concepts of equilibria apply to all the various types of equilibria, and (2) Practice, Practice, Practice. 

If one wants to understand why such changes occur, one can look at a few of the basic equilibrium properties of such complexes. Figure 1 illustrates the trends which occur when a sample is titrated with copper, monitoring three different parameters. The black dots indicate the relative amount of bound copper as indicated by free copper ions sensed with an ion-selective electrode (Xc of left ordinate). The triangles represent the change of the absorbance of the solution at 465 nm (right ordinate). The curve with the open circles is the relative quenching of the fulvic acid fluorescence (Q of left ordinate). We see that we are able to probe several different types of sites with different types of probes for this multidentate system. 

In aqueous solutions, the peak potentials of the oxidation of thiols vary with pH (Aiip/ApH = 60 mV), reflecting the position of the acid-base equilibrium affecting the SH group. In basic solutions. 

If the profile of the observed or the intrinsic rate constant plotted against pH resembles the profile for an acid-base titration curve, this strongly suggests that one of the reactants is involved in an acid-base equilibrium in that pH range. Such behavior is ftiirly common and is illustrated by the second-order reaction between the Co(II)-trien complex and O2 (Fig. 1.12). The limiting rate constants at the higher and low acidities correspond to the acidic and basic forms of the Co(II) reactant, probably. 

This system has been almost completely dropped in favour of using pATa throughout the acidity-basicity scale. To measure the strength of a base, we use the pATa of its conjugate acid, i.e. we consider the equilibrium... 

Enols and enolization feature prominently in some of the basic biochemical pathways (see Chapter 15). Biochemists will be familiar with the terminology enol as part of the name phosphoenolpyruvate, a metabolite of the glycolytic pathway. We shall here consider it in non-ionized form, i.e. phosphoenolpyruvic acid. As we have already noted (see Section 10.1), in the enolization between pyruvic acid and enolpyruvic acid, the equilibrium is likely to favour the keto form pyruvic acid very much. However, in phosphoenolpyruvic acid the enol hydroxyl is esterified with phosphoric acid (see Section 7.13.2), effectively freezing the enol form and preventing tautomerism back to the keto form. 

Nitrosamines (257) are amphoteric their acidic character is shown by their dissolution in alkalis, producing intensely dark-colored liquids from which the starting material may be regenerated by acids.170 Due to their weakly basic properties, they also dissolve in strong sulfuric and phosphoric acid and may be extracted therefrom, at least partially, by ether. In these acidic media, equilibrium mixtures are formed with the corresponding diazonium salt which may be precipitated under suitable conditions (see Section III.H, 1). 3-... 

Indirect Exchange Rates. In this case, the line shape is indirectly related to the acid-base equilibrium. Besides measuring intermolecular processes like the proton exchange rates, DNMR often has been used to measure intramolecular processes like conformational changes that occur on the same time scale. When the activation energy of such a process is very different in the acidic and basic forms for an indicator, DNMR can be used to measure the ionization ratio. 

Since solid acid catalysts are used extensively in chemical industry, particularly in the petroleum field, a reliable method for measuring the acidity of solids would be extremely useful. The main difficulty to start with is that the activity coefficients for solid species are unknown and thus no thermodynamic acidity function can be properly defined. On the other hand, because the solid by definition is heterogeneous, acidic and basic sites can coexist with variable strength. The surface area available for colorimetric determinations may have widely different acidic properties from the bulk material this is especially true for well-structured solids like zeolites. It is also not possible to establish a true acid-base equilibrium. 

SOLUTION The acid-base equilibrium for this reaction is shown in the following figure. Determining the relative amounts of the acid and conjugate base requires direct substitutions in the Henderson-Hasselbalch equation (Equation 9.1). The pH is 7.0, and the pk-., is 4.2. The equation evaluates to reveal that the ratio of the conjugate base to the acid is 630 1. As a check on this result, since pH 7.0 is more basic than the pKa of benzoic acid, the conjugate base should predominate. 

Proton transfer (PT), i.e., the kinetic aspects of heteroaromatic prototropic tautom-erism, is an important and somewhat neglected topic or, at least, much less studied than the thermodynamic aspects (equilibrium constants, acidity, basicity, pK, etc.). Intermolecular proton transfer between two heterocycles, one protonated and one neutral, occurs along a hydrogen bond (see Sect. 3.5). When the proton transfer occurs in a crystal or in an amorphous solid, we speak of SSPT (vide supra). 

FIGURE 8.3 pH rate profile for reactions exhibiting a single acid-base equilibrium. The solid curves represent cases where both acidic and basic forms are reactive. The broken lines depict limiting cases with only one reactive form. 

For the best effect, the stabilizers must be added to the reactants before the point of acid/base equilibrium is reached. Catalysts are normally not added to the prepolymer, as they may be rather basic and may cause the reactions to proceed too swiftly. 

In aqueous solution, organic acids and bases exist in equilibrium mixtures in their neutral and ionic forms. Because the neutral and ionic forms will not have the same partition coefficient, the amount extracted depends on the acid-base equilibrium. For an efficient extraction, the analyte should be at least 95% in the extracable form. This would usually mean either as its free acid or free base. Figure 2.1 is a nomogram relating pK values to percentage of ionization at various pH values [21]. In most cases, pH adjustment of the sample to pH = pK — 2 for acidic compounds or pH = pK + 2 for basic compounds is sufficient. 

Despite dramatic differences in coordination environment, neutral iridium(i) allyl complex 158 reacts with acetone to form iridacyclobutane complex 159 (Equation 79) <19940M1592>. This transformation may proceed by nucleophilic addition of the enol tautomer, as the authors suggest, or, plausibly, an acid/base equilibrium induced by the basic Ir(l) center the cationic hydridoiridium intermediate thus formed is alkylated at the central allyl carbon by the enolate. 

---