PH-scale

As a result of a variable liquid-junction potential, the measured pH may be expected to differ seriously from the determined from cells without a liquid junction in solutions of high acidity or high alkalinity. Merely to affirm the proper functioning of the glass electrode at the extreme ends of the pH scale, two secondary standards are included in Table 8.14. In addition, values for a 0.1 m solution of HCl are given to extend the pH scale up to 275°C [see R. S. Greeley, Anal. Chem. 32 1717 (I960)] ... 

Values based on the conventional activity pH scale as defined by the National Bureau of Standards (U.S.) and pertain to a temperature of 25°C [Ref Bower and Bates, J. Research Natl. Bur. Standards U.S., 55 197 (1955) and Bates and Bower, Anal. Chem., 28 1322 (1956)]. Buffer value is denoted by column headed /3. 

Equation 6.10 also allows us to develop a pH scale that indicates the acidity of a solution. When the concentrations of H3O+ and OH- are equal, a solution is neither acidic nor basic that is, the solution is neutral. Letting... 

The pH of an acidic solution, therefore, must be less than 7.00. A basic solution, on the other hand, will have a pH greater than 7.00. Figure 6.3 shows the pH scale along with pH values for some representative solutions. 

Thermodynamically, the activity of a single ionic species is an inexact quantity, and a conventional pH scale has been adopted that is defined by reference to specific solutions with assigned pH(5) values. These reference solutions, in conjunction with equation 3, define the pH( of the sample solution. 

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. 

Acidity is defined in terms of the pH scale, where pH is the negative logarithm of the hydrogen ion [H ] concentration. 

Fig. 10-11. The pH scale is a measure of hydrogen ion concentration. The pH of common substances is shown with various values along the scale. The Adirondack Lakes are located in the state of New York and are considered to be receptors of acidic deposition. Source U.S. Environmental Protection Agency, Acid Rain—Research Summary," EPA-600/8-79-028, Cincinnati, 1979.

At the pH = Jt there is a balance of charge and there is no migration in an electric field. This is referred to as the isoelectric point and is determined by the relative dissociation constants of the acidic and basic side groups and does not necessarily correspond to neutrality on the pH scale. The isoelectric point for casein is about pH = 4.6 and at this point colloidal stability is at a minimum. This fact is utilised in the acid coagulation techniques for separating casein from skimmed milk. 

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. 

Acid catalysis is an important kinetic phenomenon, and its study often requires the use of concentrated acid solutions, in which the conventional pH scale is not applicable. In sueh solutions (e.g., sulfuric acid-water mixtures covering the full range of compositions) the acid component simultaneously functions both as an acid and as a solvent thus, a medium effect is superimposed on the acidity effect. In this section we briefly describe the acidity function approach to coping with this problem. (A comparable approach can be taken to the study of highly... 

Table 2.2 gives the pH scale. Note again the reciprocal relationship between [H ] and [OH ]. Also, because the pH scale is based on negative logarithms, low pH values represent the highest H concentrations (and the lowest OH concentrations, as K, specifies). Note also that... 

The pH scale is widely used in biological applications because hydrogen ion concentrations in biological fluids are very low, about 10 M or 0.0000001 M, a value more easily represented as pH 7. The pH of blood plasma, for example, is 7.4 or 0.00000004 M H. Certain disease conditions may lower the plasma pH level to 6.8 or less, a situation that may result in death. At pH 6.8, the H concentration is 0.00000016 M, four times greater than at pH 7.4. 

At pH 7, [H ] = [OH ] that is, there is no excess acidity or basicity. The point of neutrality is at pH 7, and solutions having a pH of 7 are said to be at neutral pH. The pH values of various fluids of biological origin or relevance are given in Table 2.3. Because the pH scale is a logarithmic scale, two solutions whose pH values differ by one pH unit have a 10-fold difference in [H ]. For example, grapefruit juice at pH 3.2 contains more than 12 times as much H as orange juice at pH 4.3. 

FIGURE 2.13 The titration curves of several weak electrolytes acetic acid, Imidazole, and ammonlnm. Note that the shape of these different curves Is Identical. Only their position along the pH scale Is displaced. In accordance with their respective affinities for ions, as reflected In their differing values. 

Hydrogen was recognized as the essential element in acids by H. Davy after his work on the hydrohalic acids, and theories of acids and bases have played an important role ever since. The electrolytic dissociation theory of S. A. Arrhenius and W. Ostwald in the 1880s, the introduction of the pH scale for hydrogen-ion concentrations by S. P. L. Sprensen in 1909, the theory of acid-base titrations and indicators, and J. N. Brdnsted s fruitful concept of acids and conjugate bases as proton donors and acceptors (1923) are other land marks (see p. 48). The di.scovery of ortho- and para-hydrogen in 1924, closely followed by the discovery of heavy hydrogen (deuterium) and... 

C 0.5012 mol 2 at 15°C). It is clearly unwise to associate a pH meter reading too closely with pH unless under very controlled conditions, and still less sensible to relate the reading to the actual hydrogen-ion concentration in solution. For further discussion of pH mea.surements, see Pure Appl. Chem. 57, 531-42 (1985) Definition of pH Scales, Standard Reference Values, Measurement of pH and Related Terminology. Also C E News, Oct. 20. 1997. p. 6. 

Fig. 2.8 Potential-pH diagram calculated for Fe-H20 system at 250 C. The pH scale refers to the solution measured at 25°C and then raised to 250°C (after Ashworth )...

Several complications are involved in the calculation of potential-pH equilibrium diagrams for temperatures other than 25°C including the fact that the pH scale itself varies with temperature thus, diagrams in which the pH scale refers to the temperature for which the equilibria are calculated are probably preferable for most purposes . The most notable consequence of increasing temperature on the equilibria appears to be a widening of the pH range within which the hydroxide Ni(OH)2 is thermodynamically stable. 

Click Coached Problems for a self-study module on the pH scale. 

Solvent properties and dipoles, 313 Sorbitol, 423 Sprensen pH scale, 190 Space, interstellar, 448 Spectrograph mass, 242 simple, 247 Spectroscopy, 187 infrared, 249 microwave, 249 X-ray, 248 Spectrum... 

The two hydrogen electrodes may be replaced by a single glass electrode which is transferred from one cell to the other. The pH difference thus determined is a pure number. The pH scale is defined by specifying the nature of the standard solution and assigning a pH value to it. 

Consider the two diagrams in Fig. 6-3. Both are said to feature upward bends. In terms used earlier, Fig. 6-3a represents a transition with increasing [H+] from zeroth to first order in [H+], and Fig. 6-3b from negative first to zeroth. (Of course, the use of the pH scale reverses their appearance.) As noted previously (Rule 8, Section 6.2), this observation signifies a mechanism with two independent pathways. 

The pH scale was introduced by the Danish chemist Soren Sorensen in 1909 in the course of his work on quality control in the brewing of beer and is now used throughout science, medicine, agriculture, and engineering. 

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. 

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