PH scale standardization

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. 

Covington, A. K., R. G. Bates and R. A. Durst, Definition of pH scales, standard reference values, measurement of pH and related terminology, Pure Appl. Chem., 57, 531 (1985). 

Covington, A.K., Bates, R.G. and Durst, R.A., Definition of pH Scales, Standard Reference Values, Measurement of pH and Related Terminology, Pure Appl. Chem. 57 (1985) 531-542. 

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. 

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. 

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. 

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... 

It has been emphasized repeatedly that the individual activity coefficients cannot be measured experimentally. However, these values are required for a number of purposes, e.g. for calibration of ion-selective electrodes. Thus, a conventional scale of ionic activities must be defined on the basis of suitably selected standards. In addition, this definition must be consistent with the definition of the conventional activity scale for the oxonium ion, i.e. the definition of the practical pH scale. Similarly, the individual scales for the various ions must be mutually consistent, i.e. they must satisfy the relationship between the experimentally measurable mean activity of the electrolyte and the defined activities of the cation and anion in view of Eq. (1.1.11). Thus, by using galvanic cells without transport, e.g. a sodium-ion-selective glass electrode and a Cl -selective electrode in a NaCl solution, a series of (NaCl) is obtained from which the individual ion activity aNa+ is determined on the basis of the Bates-Guggenheim convention for acr (page 37). Table 6.1 lists three such standard solutions, where pNa = -logflNa+, etc. 

Table 9-2 lists pA-., values for common buffers that are widely used in biochemistry. The measurement of pH with glass electrodes, and the buffers used by the U.S. National Institute of Standards and Technology to define the pH scale, are described in Chapter 15. 

The quantity F is most variable when the experimental design calls for titrations over a relatively wide range of pH and is best known at those points on the pH scale where the pH meter has been calibrated with primary standard buffers. 

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 International Union of Pure and Applied Chemistry (IUPAC) recommendation [3] for the definition of pH scales has formed the basis for the standardisation of pH measurements since 1985. IUPAC recommended two different approaches to derive the pH values of pH standard buffer solutions. They yield two different pH values for one solution [4],... 

Measurements of the first dissociation constant for HjS in seawater have been made by a number of workers. The measurements of Savenko (12) and Goldhaber and Kaplan (2Q) were made using the National Bureau of Standards (N.B.S.) pH scale (21)... 

Worldwide acceptance of analytical results requires reliable, traceable, and comparable measurements. A key property of a reliable result is its traceability to a stated reference. Traceability basically means that a laboratory knows what is being measured and how accurately it is measured. It is also an important parameter where comparability of results is concerned and is usually achieved by linking the individual result of chemical measurements to a commonly accepted reference or standard. The result can therefore be compared through its relation to that reference or standard. Every link in the traceability chain must be based on the comparison of an unknown value with a known value. The stated reference might be an International System of unit (SI) or a conventional reference scale such as the pH scale, the delta scale for isotopic measurements, or the octane number scale for petroleum fuel. In order to be able to state the uncertainty of the measurement result, the uncertainty of the value assigned to that standard must be known. Therefore a traceability chain should be designed and then demonstrated using the value of the respective standard with its uncertainty.11... 

It is appropriate at this point to discuss the "apparent" pH, which results from the sad fact that electrodes do not truly measure hydrogen ion activity. Influences such as the surface chemistry of the glass electrode and liquid junction potential between the reference electrode filling solution and seawater contribute to this complexity (see for example Bates, 1973). Also, commonly used NBS buffer standards have a much lower ionic strength than seawater, which further complicates the problem. One way in which this last problem has been attacked is to make up buffered artificial seawater solutions and very carefully determine the relation between measurements and actual hydrogen ion activities or concentrations. The most widely accepted approach is based on the work of Hansson (1973). pH values measured in seawater on his scale are generally close to 0.15 pH units lower than those based on NBS standards. These two different pH scales also demand their own sets of apparent constants. It is now clear that for very precise work in seawater the Hansson approach is best. 

If the soil suspension were instead an aqueous solution, a scale of activity values for Na+ could be defined in terms of emf data obtained for standard reference solutions of prescribed (Na+), in exactly the same way as the scale of (H) values (the operational pH scale) is defined (arbitrarily) in terms of emf data for standard buffer solutions.39,40 However, the success of this extrathermo-dynamic calibration technique depends entirely on the extent to which E, and B in the standard reference solutions are the same as E, and B in the solution of interest. For the case of a soil suspension, the presence of colloidal material may cause these two parameters to differ very much from what they would be in a reference aqueous solution. If the difference is indeed large, the value of (Na+), m, or any other ionic activity estimated with the help of standard solutions and an equation like Eq. s2.23 would be of no chemical significance. 

A very considerable portion of the text has been wholly rewritten, and the entire text has been subjected to a revision and rearrangement. Specific new exercises and discussions which have been introduced include such topics as the determinations of vapor density and molecular weight, the standardization of acids and the titration of acids and bases, Faraday s law, and the use of the pH scale of hydrogen-ion concentration. Several new preparar tions have been introduced, and a few of the old ones have been discontinued. A complete list of apparatus and chemicals required in the course has been added to the Appendix. 

The data obtained for acetic acid at 25° arc plotted in Fig. 90 the results are seen to fall approximately on a straight line, and from the intercept at zero ionic strength pKa is seen to be 4.72. The difference between this value and that given previously is to be attributed to an incorrect standardization of the pH scale (cf. footnote, p. 349). 

The British standard for the pH scale is an aqueous solution of potassium hydrogen phthalate (0.05 M), which has a pH of 4.001 at 20 °C and is often used as a calibration solution for pH meters. 

To establish a pH scale, Sorensen chose a dilute hydrochloric acid solution for a standard. He took the concentration of hydrogen ions in such a solution to be given by aC, where C is the concentration of hydrochloric acid and a is a degree of dissociation determined from conductance measurements. His procedure had drawbacks first, there is evidence that the extrapolation procedure does not actually reduce the liquid-junction potential of Cell (3-6) to zero second, the hydrochloric acid is completely dissociated (dissociation constant about 1.6 x 10 ), and therefore the concentration of H is C rather than a somewhat smaller quantity. 

A pH scale could be defined in terms of a single standard buffer. Every practical pH reading, however, involves a liquid-junction potential, and the variation in this potential between the readings of Cell (3-9) for the unknown and the standard is tacitly included in Equation (3-10). The liquid-junction potential is essentially constant for solutions of intermediate pH values, say 3 to 11, but beyond these limits it varies considerably because the concentrations of the imusually mobile H" " or OH ions become appreciable. Although this variation does not affect the definition of pH by (3-10), the measured value of pH deviates appreciably from the known best values of —log %+ for solutions of high acidity or alkalinity. For this reason, a series of standards covering a range of values of pH, has been adopted. For an experimental... 

A different set of pH values must be adopted at each temperature in effect, there is a different standard state for hydrogen ions, and therefore a different pH scale, for each temperature. Selected values of pHj for several standards recommended by the U.S. National Bureau of Standards are shown in Table 3-1. 

For most solvents, standard buffers have not been established for standardizing the pH scale. Many measurements of pH in nonaqueous solvents have therefore been made by extension of the pH scale in water by use of aqueous buffers. Even for hydrogen-bonding solvents similar to water, extrapolation procedures are uncertain since the effect of both liquid-junction potential and transfer activity coefficient are... 

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