PH calculation for

Suppose we are titrating the triprotic acid H P04 with a solution of NaOH. The experimentally determined pH curve is shown in Fig. 11.13. Notice that there are three stoichiometric points (B, D, and F) and three buffer regions (A, C, and E). In pH calculations for these systems, we assume that, as we add the hydroxide solution, initially NaOH reacts completely with the acid to form the diprotic conjugate base... 

A flow chart summarizes the major species in solution and the pH calculations for the four key regions of a weak acid titration curve. 

Figure 9-1 shows the pH calculated for different concentrations of strong base or strong acid in water. There are three regions ... 

The results of pH calculations for the titration of 0.100 M CH3C02H with 0.100 M NaOH are plotted in Figure 16.7. Comparison of the titration curves for the weak acid-strong base titration and the strong acid-strong base case shows several significant differences ... 

Both types of logarithms follow the same mathematical rules, but we have to be careful as to when to use each. (Common logs are used with pH calculations, for example, in Section 19.3, and natural logs are generally used with half-life problems in Section 21.2.) Common logs may be converted to natural logs if necessary, and vice versa, using the ratio of the two ... 

Typically, pH calculations for solutions of weak bases are very similar to those for weak acids. 

Although we might expect the pH calculations for solutions of polyprotic acids to be complicated, the most common cases are surprisingly straightforward. To illustrate, we will consider a typical case, phosphoric acid, and a unique case, sulfuric acid. 

This means that in a solution prepared by dissolving H3P04 in water, only the first dissociation step makes an important contribution to [H+], This greatly simplifies the pH calculations for phosphoric acid solutions, as is illustrated in Example 7.8. 

The pH Calculations for an Aqueous Solution of a Weak Acid HA (major species HA and H20)... 

Buffered solutions are simply solutions of weak acids or bases containing a common ion. The pH calculations for buffered solutions require exactly the same procedures previously introduced in Chapter 7. This is not a new type of problem. 

Detailed mechanistic analysis of any reaction in HF is complicated by the presence of multiple fluoride species. In relatively dilute solutions, HF, H, F , and HF2 need to be considered [35, 56-58]. The equilibria and rate constants for these species are given in Table 5. Figure 6a and b shows the relative concentrations of HF, HF, and F as a function of pH calculated for 1 M and 10 M total fluoride concentration, respectively. As can be seen from these figures, at low pH HF molecules are the dominant species whereas at high pH F dominates. At intermediate values of pH, fluoride electrolytes are dominated by HF2 with a maximum mole fraction at about pH 3. 

Henley, R,W. 1984. pH calculations for hydrothermal fluids. In Fluid-mineral equilibria in hydrothermal systems. Rev. in Econ. Geology I, pp. 93-98. 

Precipitation pHs calculated for different concentrations Ct of dodecyl amine are given in Table 2.6. 

The approach needed to perform pH calculations for these systems is virtually identical to that used above. This makes sense because, as is true of all buffered solutions, a weak acid (BH ) and a weak base (B) are present. A typical case is illustrated in Example 15.5. 

DNA thermodegradation is analyzed by incubating 0.5 xg of DNA in 20 p.1 incubation mixture covered with 150 p,l of HaO-saturated paraffin oil to prevent evaporation. The pH and temperature of the incubation mixtures are controlled through temperature and pH probes. The pH determined by probes can be slightly different from the pH calculated for pH-temperature compensation, depending on salt present or buffer strength. 

Figure 13. Lower Species distribution in 2-10 mol dm citric acid solution as calculated by using K = 1.2-10, Ki = 4.9-10, and Ki, = 2.5-10 in the presence of I-IO mol dm NaN03. Upper Surface concentration (or coverage, 0) of citric acid on hematite as a function of the pH, calculated for the adsorption of its HX + X anions (solid line). The circles represent the experimental data taken from ref [25]. Dashed lines give the calculated contributions of individual ions to the total adsorption density [26].

Figure 4-8. Combination of Figures 4-6 and 4-7 as a pH calculator for 0.01 M solutions of acetic acid (intersection s ) and of sodium acetate (intersection sj at ionic strength 0.1 M.

The calculations are also quite sensitive to the amount of base present in the system, which is reflected by the absolute values of the pH. This is illustrated by the following values of pH — pH" calculated for the case where pC02 = 40 mm. and C = 9.4 X lO" (as in the solutions studied by Ferguson)... 

Thus, all the values of pH calculated for all the concentrations of inhibitor added to the solution, between 100 ppm and 1000 ppm CeClj-VH O, are higher than the maximum value of pH possible in the cathodic areas. Because of the reduction of O, there is little probability that, in the experimental conditions used (Arenas and de Damborenea, 2003), precipitation of CeO occurs, as proposed by this mechaiusm from the aqueous ions of Ce(OH)/. ... 

In another type of calculation, we are given the pH of a solution and asked to determine the [H3O+]. This is a reverse of the pH calculation. For whole number pH values, the negative pH value is the power of 10 in the H30 concentration. 

---