PH scale and

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. 

Figure 1.3. The pH scale showing the region unavailable to chemists before the development of the pH scale and the pH meter.

Because the concentration of the hydronium ion, H30, can vary tremendously in solutions of acids and bases, a scale to easily represent the acidity of a solution was developed. It is called the pH scale and is related to the [H30 ] ... 

The pH of pure water is 7.00. On the pH scale this is called neutral. A solution whose [H30+] is greater than in pure water will have a pH less than 7 00 and is called acidic. A solution whose [H30+] is less than in pure water will have a pH greater than 7.00 and is called basic. Figure 15.2 shows the pH scale and the pH values of some common substances. 

In this section, you learned about the relationship between the pH scale and the concentrations of the ions that form when water and weak acids dissociate. In the next section, you will learn that the equilibrium of weak bases is similar to the equilibrium of weak acids. As you will see, solutions that contain a mixture of a weak acid and a salt of its conjugate base have properties with important biochemical and industrial applications. 

The principles involved in these equilibria are presented in detail in numerous works and will not be repeated here. For a thorough treatment, including the definition of pH scales and methods of measurement, the reader is referred to such works as those of Edsall and Wyman (1958) and Bates (1964). Applications to milk are discussed by Walstra and Jenness (1984). It should suffice here to present some of the basic relationships in equation form. 

Accuracy and Interpretation of Measured pH Values. To define the pH scale and pertnil the calibration of pH measurement systems, a scries of reference buffer solutions have been certified hy the U.S. National Institute of Standards and Technology iNIST). The acidity function which is the experimental basis for the assignment of pH. is reproducible within about O.IKl.I pH unit from It) to 40T. However, errors in the standard potential of the cell, in the composition of the buffer materials, and in the preparation of the solutions may raise the uncertainty to 0 005 pH unit. The accuracy of ihe practical scale may he furthei reduced to (I.Ot)X-(l.(ll pH unit as a result of variations in the liquid-junction potential. 

The pH scale and pH values for some common substances are shown in Figure 15.3. Because the pH is the negative log of [H30 + ], the pH decreases as [H30 + ] increases. Thus, when [H30 + ] increases from 10-7M to 1CT6M, the pH decreases from 7 to 6. As a result, acidic solutions have pH less than 7, and basic solutions have pH greater than 7. 

Dilution experiments using universal indicator and 0.1 mol dm 3 solutions of hydrochloric acid, sodium hydroxide, ethanoic acid and ammonia. These will establish the pH scale and promote an understanding of the distinction between strong and weak electrolytes. 

Espinosa, S., Bosch, E., and Roses, M. Retention of ionizable compounds on HPLC 5 pH scales and the retention of acids and bases with acetonittile-water mobile phases. Anal Chem. 2000, 72, 5193-5200. 

Dickson, A. G. (1984) pH Scales and Proton Transfer Reactions in Saline Media Such as Seawater, Geochim. Cosmochim. Acta 48, 2299-2308. 

In order to solve the equations and determine pH and the concentrations of the species that make up the alkalinity, the apparent equilibrium constants, F, must be accurately known. These constants have been evaluated and re-evaluated in seawater over the past 50 y. The pH scales and methods of measuring pH during these experiments have been different, and this has complicated comparisons of the data until recently, when many have been converted to a common scale. Equations for the best fit to carbonate system equilibrium constants as a function of temperature and salinity are presented by Luecker et al. (2000), DoE (1994) and Millero (1995) (see Appendix 4.2). 

Probably many readers of the present book do not expect that something as self-evident as pH measurements might cause a problem. Yet, apparent easiness and obviousness is a pitfall. Different aspects of the pH scale and of pH measurements have been discussed in numerous handbooks of chemistry. Reference [204] is a special monograph devoted solely to pH measurements. A few aspects of the pH of solutions, not directly related to surface charging or adsorption, are discussed in this section. 

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. 

Because the concentration of ions in solution generally is quite small one has for the sake of convenience chose to express a solutions acidness based on the decimal logarithm to the concentration of BT ions completely analogous to the principles that the acid constant Ka was expressed as an acid exponent pKa. this is known as the pH scale and the pH values are defined as ... 

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. 

Values of ApH should be in the range of 0.08 to 0.18. It empirically corrects for differences between the two pH scales and for measurement errors associated with the electrode pair. The pH(X) of samples measured using lUPAC aqueous buffers, can be converted to pH. or pHp using the appropriate measured ApH ... 

For example, a solution of pH 3 has an H concentration of 10 M, which is 10 times that of a solution of pH 4 ([H ] = 10 M) and 100 times that of a solution of pH 5. This is illustrated in Table 16.2. Also note from Table 16.2 that the pH decreases as the [H ] increases. That is, a lower pH means a more acidic solution. The pH scale and the pH values for several common substances are shown in Figure 16.3. 

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