Indicators, acid-base range

The lower the value of this constant, the larger the deferences in acidity indices (pH) between the standard solutions of strong acids and bases, that results in a wider acid-base range for the solvent. This refers not only to the acid-base equilibria in aqueous solutions but also applies to any donor-acceptor interaction in molecular solvents which are prone to heterolytic dissociation with the formation of acidic and basic particles, as provided by an appropriate definition of acids and bases. It follows from equations (1.1.3) and (1.1.4) that the Arrhenius definition can only be used for the description of acid-base interactions in aqueous solutions, since the reaction between the acid of solvent and the base of solvent can result in the formation only of the solvent molecules. In the case considered, this solvent is water. 

As follows from Part 1, the ionic melts based on molten alkali metal halides are referred to the solvents of the Second Kind (Kind II), and, therefore, the acid-base ranges for these media are half-open (see Fig. 1.1.1, scheme N3). Therefore, to form an idea of the relative oxoacidic properties of the studied chloride melts it is enough to know their oxobasicity indices. The necessary experimental parameters obtained at 600 °C are presented in Table 1.3.1. The data in this Table show that the KCl-LiCl eutectic melt possesses appreciable acidic properties, the corresponding oxobasicity index being equal to 3.2. 

Redox indicator acid-base indicator (pH range 5.7-11.4 colour change red - blue), used as a 0.02% aq. soln. of Na salt. Dark green cryst. powder. Sol. H2O, alkalis, EtOH, Me2CO. pK 5.5. E° + 0.639V (30°). 

Acid—Base Indicators— Substances whose colors depend on the pH of a solution are known as acid-base indicators. Acid-base indicators exist in solution as a weak acid (HIn) and its conjugate base (InH- Each form has a different color and the proportions of the two forms determine the color of the solution, which in turn depends on the pH of the solution. The pH range over which an acid-base indicator changes color (Fig. 17-7) is determined by the of the specific indicator. 

Standardization—External standards, standard additions, and internal standards are a common feature of many quantitative analyses. Suggested experiments using these standardization methods are found in later chapters. A good project experiment for introducing external standardization, standard additions, and the importance of the sample s matrix is to explore the effect of pH on the quantitative analysis of an acid-base indicator. Using bromothymol blue as an example, external standards can be prepared in a pH 9 buffer and used to analyze samples buffered to different pHs in the range of 6-10. Results can be compared with those obtained using a standard addition. 

A list of several common acid-base indicators, along with their piQs, color changes, and pH ranges, is provided in the top portion of Table 9.4. In some cases. 

Ladder diagram showing the range of pH levels over which a typical acid-base indicator changes color. 

Indicator Acid Color Base Color pH Range P/Ca... 

The indicator method is especially convenient when the pH of a weU-buffered colorless solution must be measured at room temperature with an accuracy no greater than 0.5 pH unit. Under optimum conditions an accuracy of 0.2 pH unit is obtainable. A Hst of representative acid—base indicators is given in Table 2 with the corresponding transformation ranges. A more complete listing, including the theory of the indicator color change and of the salt effect, is also available (1). 

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). 

Another definition of acids and bases is due to G. N. Lewis (1938). From the experimental point of view Lewis regarded all substances which exhibit typical acid-base properties (neutralisation, replacement, effect on indicators, catalysis), irrespective of their chemical nature and mode of action, as acids or bases. He related the properties of acids to the acceptance of electron pairs, and bases as donors of electron pairs, to form covalent bonds regardless of whether protons are involved. On the experimental side Lewis definition brings together a wide range of qualitative phenomena, e.g. solutions of BF3, BC13,... 

For some purposes it is desirable to have a sharp colour change over a narrow and selected range of pH this is not easily seen with an ordinary acid-base indicator, since the colour change extends over two units of pH. The required result may, however, be achieved by the use of a suitable mixture of indicators these are generally selected so that their pK ln values are close together and the overlapping colours are complementary at an intermediate pH value. A few examples will be given in some detail. 

Strong acid and strong base. For 0.1 M or more concentrated solutions, any indicator may be used which has a range between the limits pH 4.5 and pH 9.5. With 0.01 M solutions, the pH range is somewhat smaller (5.5—8.5). If carbon dioxide is present, either the solution should be boiled while still acid and the solution titrated when cold, or an indicator with a range below pH 5 should be employed. 

Aminopyridines, aminopyridine oxides, and 3-aminoquinoline are obviously diazotized by analogous mechanisms. Kalatzis (1967 b) studied the diazotization of 4-aminopyridine over a very large range of acid concentrations (0.0025-5.0 m HC104). This compound is comparable to 2-aminothiazole in its acid-base properties the heterocyclic nitrogen is easily protonated at pH 10, whereas the amino group is a very weak base (pKa = -6.5). Therefore, the kinetics indicate that the (mono-protonated) 4-aminopyridinium ion reacts with the nitrosyl ion. The... 

It follows that at half-neutralization pH - pKa, while at A = 1/11 the pH = pjFira - land at A = 10/11 the pH = pKa + 1 in fact this means that the whole titration takes place within 2 pH units, which agrees with the maximum pH range of acid-base colour indicators. 

Although litmus paper, cabbage juice, and phenolphthalein can indicate whether a substance is acidic or basic, they have limitations in that they cannot determine an exact pH. To do this, an acid-base indicator called universal indicator can be used. Universal indicator is actually a mixture of several different acid-base indicators (usually phenolphthalein, methyl red, bromthymol blue, and thymol blue). This mixture produces a wide range of colors to indicate different pHs. Under very acidic conditions, universal indicator is red. It turns orange and then yellow between the pHs of 3 to 6. It is green at neutral pH and turns greenish-blue as a solution becomes more alkaline. In very basic conditions, universal indicator turns a dark purple color. 

Fig. 2.1. Lipophilicity profiles for diclofenac (acid), propranolol (base) and morphine (ampholyte). Dashed lines indicate the pH range where molecule may partition in its ionized form.

Indicator Absorbance maximum color of acid/base forms or pKa value or pH range... 

Table 5.1 summarizes the details of some useful acid-base indicators. Exact agreement with the pH range expressed by equation (5.5) is by no means always observed. This is because some colour changes are easier to see than others and so the general approximation made in deriving equation (5.5) is not uniformly close. Structurally, the indicators form three groups phthaleins (e g. phenolphthalein) sulphonephthaleins (e.g. phenol red) and azo compounds (e.g. methyl orange). 

Table 5.1 A range of visual indicators for acid-base titrations...

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