Solution Chemistry—pH Scale

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. 

The approach to proton adsorption from aqueous solution must be different from the approach to adsorption of other solutes, because water molecules can provide or absorb a practically unlimited number of protons (higher by several orders of magnitude than the concentration of any other species in solution and the concentration of surface sites) to balance the changes induced by adsorption. Thus, adsorption isotherms based on the concept of a distribution of a limited amount of adsorbate molecules between solution and surface are not applicable. Most authors accept this obvious fact, but a few others have used the same formalism for proton adsorption as is used for other solutes. For example, in [205], the surface charging of alumina is discussed in terms of adsorption isotherms (amount adsorbed vs. equilibrium concentration). Positive adsorption of protons is equivalent to negative adsorption of OH , and vice versa. In adsorption experiments, uptake of protons and release of OH cannot be distinguished. Only the net result of uptake/releasc of H and OH can be obtained, and independent curves of 11 and OH adsorption reported in the literature [206,207] must be based on measurements of other quantities. 

The procedure of pH measurement is not limited to insertion of a combination electrode into a solution (dispersion) and waiting until a constant value is displayed. Some pH value will be displayed, even when the rules given in every user manual are disobeyed. Typical errors are inadequate calibration of the pH electrode, use of outdated pH buffers, old electrodes (2 years is a typical lifetime), insufficient flow in the salt bridge between reference electrode and solution (or incorrect level and/or composition of the solution in the bridge), and insufficient electrical contact between solution and electrode. These examples of carelessness are commonplace in scientific laboratories, and typically induce errors in the range of a few tenths of a pH unit. 

It is not obvious to all scientists that the pH reported in (pH) or ao(pH) plots (from which the lEP or PZC is determined) is the equilibrium pH of the dispersion used for the measurements. The following description was found in [208], The authors equilibrated their particles in a solution 1 of pH 1, 1.9, 3, 5, 7, 8, 11, or 13. The particles were then separated from solution 1 and redispersed in pure water. The new dispersion (particles in solution 2) was used to measure the electrophoretic mobility. Obviously, the pH of the solution 2 formed by equilibration of pretreated particles with water was different from the pH of solution 1, and most scientists would have plotted the potential against the pH of solution 2 to 

FIGURE 1.18 Electrokinetic curve with one outstanding point. In order to explain the difference between the outstanding point and the smooth curve connecting all other points solely in terms of the error in the potential, broad margin of error has to be allowed. The same discrepancy may be explained by allowing relatively narrow margin of error in the pH. 

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