Aqueous pH scale

The activity of the hydrogen ion is affected by the properties of the solvent in which it is measured. Scales of pH only apply to the medium, ie, the solvent or mixed solvents, eg, water—alcohol, for which the scales are developed. The comparison of the pH values of a buffer in aqueous solution to one in a nonaqueous solvent has neither direct quantitative nor thermodynamic significance. Consequently, operational pH scales must be developed for the individual solvent systems. In certain cases, correlation to the aqueous pH scale can be made, but in others, pH values are used only as relative indicators of the hydrogen-ion activity. 

Labib, M.E. and Williams, R., An experimental comparison between the aqueous pH scale and the electron donocity scale. Colloid Polym. Sci., 264, 533, 1986. Morris, G.E. et al.. Surface chemistry and rheological behaviour of titania pigment suspensions. Colloids Sutf. A, 155, 27, 1999. 

Measurement of pH in a nonaqueous solvent when the electrode is standardized with an aqueous solution has little significance in terms of possible hydrogen ion activity because of the unknown liquid-junction potential, which can be rather large, depending on the solvent. Measurements made in this way are usually referred to as apparent pH. pH scales and standards for nonaqueous solvents have been suggested using an approach similar to the one for aqueous solutions. These scales have no rigorous relation to the aqueous pH scale, however. You are referred to the book by Bates (Ref. 3) for a discussion of this topic. See also M. S. Frant, How to Measure pH in Mixed Nonaqueous Solutions, Today s Chemist at Work, American Chemical Society, June, 1995, p. 39. 

Many organic reactions involve acid concentrations considerably higher than can be accurately measured on the pH scale, which applies to relatively dilute aqueous solutions. It is not difficult to prepare solutions in which the formal proton concentration is 10 M or more, but these formal concentrations are not a suitable measure of the activity of protons in such solutions. For this reason, it has been necessaiy to develop acidity functions to measure the proton-donating strength of concentrated acidic solutions. The activity of the hydrogen ion (solvated proton) can be related to the extent of protonation of a series of bases by the equilibrium expression for the protonation reaction. 

Using Environmental Examples to Teach About Acids. Acid-base reactions are usually presented to secondary students as examples of aqueous equilibrium (2). In their study of acids and bases, students are expected to master the characteristic properties and reactions. They are taught to test the acidity of solutions, identify familiar acids and label them as strong or weak. The ionic dissociation of water, the pH scale and some common reactions of acids are also included in high school chemistry. All of these topics may be illustrated with examples related to acid deposition (5). A lesson plan is presented in Table I. 

To establish the operational pH scale, the pH electrode can be cahbrated with a single aqueous pH 7.00 phosphate buffer, with the ideal Nernst slope assumed. As Eqs. (2a)-(2d) require the free hydrogen ion concentration, an addihonal electrode standardization step is necessary. That is where the operational scale is converted to the concentration scale pcH (= -log [H ]) as described by Avdeef and Bucher [24] ... 

Oumada et al. [148] described a new chromatographic method for determining the aqueous pKa of dmg compounds that are sparingly soluble in water. The method uses a rigorous intersolvent pH scale in a mobile phase consisting of a mixture of aqueous buffer and methanol. A glass electrode, previously standardized with common aqueous buffers, was used to measure pH online. The apparent ionization constants were corrected to a zero-cosolvent pH scale. Six sparingly soluble nonsteroidal antiinflammatory weak acids (diclofenac, flurbiprofen, naproxen, ibu-profen, butibufen, fenbufen) were used successfully to illustrate the new technique. 

The acidity function is determined by successive use of a range of indicators. Hammett and Deyrup started with p-nitroaniline, the pKAl of which in a dilute aqueous solution is 1.11 (the solvolysis constant has been identified with the acidity constant). Since in a dilute aqueous solution, Yb = Ybh+ 1, the acidity function for the aqueous media goes over to the PH scale. By means of p-nitroaniline, the acidity constant of another somewhat more acidic indicator is obtained under conditions such that both forms of each indicator are present at measurable concentrations. Then H0 as well as the p A2 of the other indicator is determined by using Eq. (1.4.40). By means of this indicator, values of H0 not accessible with p-nitroaniline may be reached. The H0 scale is further extended by using a third indicator and its p A, is determined in the same way (see Fig. 1.11). The concentration ratios are determined photometrically in the visual or ultraviolet region. Figure 1.12 shows the dependence of H0 on the composition of the H2S04-H20 mixture and was obtained as indicated above. 

Thus, these relationships can be used to define a pH scale for non-aqueous protic media, consistent with the pH scale for aqueous solutions. For standard hydrogen pressure, the potential of the hydrogen electrode depends on the pH(s) according to the relationship... 

Most methods of pKa measurement were developed using water-soluble samples. However, many drugs are poorly soluble in water alone, and require the presence of a water-miscible co-solvent to keep them in aqueous solution. The solvent affects the pKa in two ways (i) it causes the pH scale to shift and (ii) it causes the pKa to shift. The consequence is that apparent pKa values measured in the presence of solvent are different from aqueous values. 

In 1909, Sorenson described the development of the pH scale based on the work of Arrhenius and the characteristics of water. Experiments and the resulting calculations show that water dissociates into hydrogen ions (H+) and hydroxide ions (HO-) and that the product of their concentrations equals close to 10"14 ions in aqueous solution. From this, a pH scale from 0 to 14 was developed and the scale describing this relationship using the abbreviation pH was developed (in the older literature [13], one may encounter both the p and the h capitalized, i.e., as PH). Today, it is universally designated as pH [13,22]. 

The potentiometric technique has proved to be of great significance and utility for determining endpoints of titrations in a non-aqueous media. The mV scale rather than the pH scale of the potentiometer must be used for obvious reasons, namely ... 

You can describe the acidity of an aqueous solution quantitatively by stating the concentration of the hydronium ions that are present. [HsO" ] is often, however, a very small number. The pH scale was devised by a Danish biochemist named Spren Sorensen as a convenient way to represent acidity (and, by extension, basicity). The scale is logarithmic, based on 10. Think of the letter p as a mathematical operation representing -log. The pH of a solution is the exponential power of hydrogen (or hydroni-um) ions, in moles per litre. It can therefore be expressed as follows ... 

The lower the pH, the more acidic the solution the higher the pH, the more basic the solution. The pH scale only applies to aqueous solutions, and is only a measure of the acidity of the solution. It does not indicate how strong the acid is (that is a function of pA a) and the pH of an acid will change as we alter its concentration. For instance, dilution will decrease the H3O+ concentration, and thus the pH will increase. 

The acidity of an aqueous solution is determined by the concentration of HsO ions. Thus, the pH of a solution indicates the concentration of hydrogen ions in the solution. The concentration of hydrogen ions can be indicated as [H ] or its solvated form in water as [H3O ]. Because the [HsO ] in an aqueous solution is typically quite small, chemists have found an equivalent way to express [H30 ] as a positive number whose value normally lies between 0 and 14. The lower the pH, the more acidic is the solution. The pH of a solution can be changed simply by adding acid or base to the solution. Do not confuse pH with pK. The pH scale is used to describe the acidity of a solution. The pK is characteristic of a particular compound, and it tells how readily the compound gives up a proton. 

The pH scale in water is widely used as a measure of acid-base properties in aqueous solutions. It is defined by pH=-log a(H+). In Section 3.1, we dealt with the poH value, defined by poH=-log a(H+), for solutions in amphiprotic and aprotic solvents of high permittivity. Recently, however, the symbol pH has also been used for the value of -log o(H+) in such non-aqueous solutions. Therefore, hereafter, the symbol pH is used instead of paH-... 

In aprotic solvents, in which pKSn values are not definitive, it is difficult to estimate the activity of the lyate ion from the pH of the solution. It is different from the case in water, in which a(OI I ) can be estimated from the pH value. Although the pH scale in aprotic solvents has such a disadvantage, it is still useful to quantitatively understand the acid-base aspects of chemical phenomena. Wider use of the pH concept in non-aqueous solutions is desirable. 

The pH scale in non-aqueous solutions was briefly discussed in Chapter 3. In this section, practical methods of pH measurements in non-aqueous systems are considered, with emphasis on the differences from those in aqueous solutions. 

Note that the pH scale is logarithmic, not arithmetic. To say that two solutions differ in pH by 1 pH unit means that one solution has ten times the H+ concentration of the other, but it does not tell us the absolute magnitude of the difference. Figure 2-15 gives the pH of some common aqueous fluids. A cola drink (pH 3.0) or red wine (pH 3.7) has an H+ concentration approximately 10,000 times that of blood (pH 7.4). 

Strong bases in dry solvents are usually used in organic synthesis to generate reactive enol anions from ketones. Nevertheless, the kinetic studies discussed here were mostly performed on aqueous solutions. Apart from the relevance of this medium for biochemical reactions and green chemistry, it has the advantage of a well-defined pH-scale permitting quantitative studies of acid and base catalysis. 

It was S0rensen s idea7 to use this relationship, which can be considered as a basis to the modern definition of the pH scale of acidity for aqueous solutions. The pH of a dilute solution of acid is related to the concentration of the solvated proton from Eq. (1.8). Depending on the dilution, the proton can be further solvated by two or more solvent molecules. 

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