Acid-base indicator defined

What is an acid-base indicator Define the equivalence (stoichiometric) point and the end point of a titration. Why should you choose an indicator so that the two points coincide Do the pH values of the two points have to be within 0.01 pH unit of each other Explain. 

Use an acid—base indicator, which marks the endpoint of a titration by changing color. Although the equivalence point of a titration, defined by the stoichiometry, is not necessarily the same as the endpoint (where the indicator changes color), careful selection of the indicator will ensure only negligible error. 

Q2.32 What makes a good acid-base indicator Q2.33 Define pKa. 

The importance of the equilibrium law is widespread throughout chemistry. We have seen in Chapter 7 that defining the optimum conditions for certain key industrial gas phase reactions is dependent on a thorough understanding of the factors that determine the proportions of reactants and products in an equilibrium mixture, and we will return to a consideration of the Haber process later in this chapter. However, chemists use the equilibrium law to represent the extent to which a weak acid or base ionizes or dissociates, defining terms such as the dissociation constants and in relation to these effects. The behaviour of acid-base indicators is also explained in terms of the equilibria involved and the application of Le Chatelier s principle. 

From a practical standpoint, acids can be identified by their sour taste, their ability to react with a variety of metals and carbonate minerals, and the effect they have on the colors of substances called acid-base indicators. Methyl red is an acid-base indicator that appears red in acidic environments and yellow otherwise (see Figure 5-10). From a chemist s point of view, however, an acid can be defined as a substance that provides hydrogen ions (H ) in aqueous solution. This definition was first proposed by Svante Arrhenius in 1884. 

Conductometric titrations. Van Meurs and Dahmen25-30,31 showed that these titrations are theoretically of great value in understanding the ionics in non-aqueous solutions (see pp. 250-251) in practice they are of limited application compared with the more selective potentiometric titrations, as a consequence of the low mobilities and the mutually less different equivalent conductivities of the ions in the media concerned. The latter statement is illustrated by Table 4.7108, giving the equivalent conductivities at infinite dilution at 25° C of the H ion and of the other ions (see also Table 2.2 for aqueous solutions). However, in practice conductometric titrations can still be useful, e.g., (i) when a Lewis acid-base titration does not foresee a well defined potential jump at an indicator electrode, or (ii) when precipitations on the indicator electrode hamper its potentiometric functioning. 

H0 is defined so as to be similar to pH, and to reduce to it in dilute solution, i.e. to pH = pKa — log I. The idea is that versions of equation (8) can be written for weak base indicators that protonate to different extents in the same add solutions (overlapping indicators indicators because they indicate the solution acidity) subtracting two of these (say for indicators A and B) leads to equation (9), and if the activity coefficients for A and B, and for AH+ and BH+, approximately cancel, the value of pA bh+ can be calculated from the measured ionization ratios for A and B if pAah+ is known ... 

In conclusion, the C/E ratios for donors (acids) indicate whether hardness or softness is most important in interactions of a particular donor (acid), but softness or hardness so defined does not enable one to predict even the relative strength of interaction towards a soft or hard acid (base) because the magnitude of the C and E numbers are lost when the ratio is taken. 

The study of the reactivity of the nucleic acid bases utilizes indices based on the knowledge of the molecular electronic structure. There are two possible approaches to the prediction of the chemical properties of a molecule, the isolated and reacting-molecule models (or static and dynamic ones, respectively). Frequently, at least in the older publications, the chemical reactivity indices for heteroaromatic compounds were calculated in the -electron approximation, but in principle there is no difficulty to define similar quantities in the all-valence or allelectron methods. The subject is a very broad one, and we shall here mention only a new approach to chemical reactivity based on non-empirical calculations, namely the so-called molecular isopotential maps. 

So what are acids and bases Vinegar is actually a dilute solution of acetic acid in water, about a 5 percent solution, but it rather nicely displays the characteristic properties of acids they are sour, they turn purple-cabbage indicator red or pink, and they react with bases to form water. A solution of sodium bicarbonate nicely displays several of the characteristics of basic solutions it tastes bitter, it turns purple-cabbage indicator blue, and it reacts with acids to form water. The last property, listed for both acid and base, the ability to react with each other, is really the defining property because acid-base reactions, like redox reactions, occur in tandem one substance acts as an acid and one substance acts as a base. Acid neutralizes base and base neutralizes acid. 

FIGURE 1.1 loiiization of acids and bases. An acid is defined as a chemical that dissociates and donates a proton to water. A base is defined as a chemical that can accept a proton. The double arrows indicate that the ionization process occurs in the forward and backward directions. The term equilibrium means that the rate of the forward reaction is equal to die rate of the backward reaction, and that no net accumulation of products or reactants occurs over time. 

The addition of organic solvents to water should modify acid-base phenomena, but assessment of such effects poses many problems, as only the measured pH of aqueous solutions can be interpreted in terms of hydrogen ion concentrations. The quantitative comparison of the acidities of partially aqueous solutions is therefore a problem of far greater complexity than the measurements of pH values in aqueous media. As mentioned earlier, a proton activity (paH) is defined in such a way that — log paH is equal to pH when the medium is water, and its value can be measured both by the electromotive force of a cell with liquid junction and by the spectrophotometry of colored indicators. 

The following clinical terms are used to describe the acid-base status. Addemia is defined as an arterial blood pH <7.35 d.nd alkalemia indicates an arterial blood pH >7.45. Acidosis and alkalosis refer to patliological states that lead to acidemia or alkalemia. For example, in common acid-base disorders... 

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