Buffer tables, calculation 72-74

Phosphate buffer, pH6.0-7.7, 0.01-1.0M Table 7.4 gives the molar fraction x of the two basic salts of the buffer system H2 POj/HPOj for different pH and molarities M. The weighted portion g in grams of the salts Me2HP04 and MeH2P04 (Me sodium or potassium) per 1000 ml buffer is calculated by... 

If the buffered solution is dilute, this is its hydrogen-ion concentration. Because the activities of ions are affected by other ions, however, there is appreciable deviation from the calculated values in salt solutions as concentrated as 0,1 M, This fact accounts for the small discrepancies beuveen the pH values calculated from equilibrium con status and those gi en in the buffer tables. 

A citrate buffer has recently been added to the list of standard buffers. Table 2-4. Using published pK° values of 3.13,4.76, and 6.40 and activity coefficients, calculate the pH (activity) expected at 25° and compare with the NBS value. 

Table 2.2 Dilution values (ApHy,) for equimolar acid buffer solutions (calculated from Davies equation)...

Although available buffer tables are extensive, they may fail to meet particular requirements, such as a specified ionic strength or a nominated buffer species. Where pAa values of the buffer substances are known the compositions of the solutions can be calculated as indicated below. Considerations governing the design of new types of buffer systems are also discussed. 

Tables 10.1 and 10.2 enable buffer tables to be constructed for use at a constant ionic strength of 0.1 for monoacidic bases or monobasic acids, given the appropriate practical pA a. Where necessary, this can be calculated from the thermodynamic value in Appendix 111 by using Table 2.3.

However, there are very many substances that could be, or have been, of use as buffers, and Appendix III lists some of these. Chapter 5 shows how new pH-buffer tables can be constructed from the thermodynamic values, and some simple computer programmes are included to facilitate the necessary calculations. Tables and worked examples are given for use if a computer is not available. 

Aqueous solutions buffered to a pH of 5.2 and containing known total concentrations of Zn + are prepared. A solution containing ammonium pyrrolidinecarbodithioate (APCD) is added along with methyl isobutyl ketone (MIBK). The mixture is shaken briefly and then placed on a rotary shaker table for 30 min. At the end of the extraction period the aqueous and organic phases are separated and the concentration of zinc in the aqueous layer determined by atomic absorption. The concentration of zinc in the organic phase is determined by difference and the equilibrium constant for the extraction calculated. 

Calibrating the electrode presents a third complication since a standard with an accurately known activity for H+ needs to be used. Unfortunately, it is not possible to calculate rigorously the activity of a single ion. For this reason pH electrodes are calibrated using a standard buffer whose composition is chosen such that the defined pH is as close as possible to that given by equation 11.18. Table 11.6 gives pH values for several primary standard buffer solutions accepted by the National Institute of Standards and Technology. 

From Eq. (4) and the data of Tables 1, 2, and 3 it can easily be calculated that the rate of hydrolysis of these enamines should rapidly reach a maximum value in weakly acidic solutions at decreasing pH, assuming that no buffer is used. The raTe should then be constant and pH-independent. For enamine (1) Eq. (4) reduces to k = Kkn,ou and this rate must be the same as that at the intersection of the straight lines in Fig. 1. This appeared to be true for the observed rate at pH 2 15). 

Table 5-2. Selected rate constants and half-livese) for some reactions of substituted benzenediazonium ions with buffer solutions (pH 9.00) at 25 °C (rate constants from Virtanen and Kuok-kanen, 1977 half-lives calculated by the present author).

In Figure 2 the solubility and speciation of plutonium have been calculated, using stability data for the hydroxy and carbonate complexes in Table III and standard potentials from Table IV, for the waters indicted in Figure 2. Here, the various carbonate concentrations would correspond to an open system in equilibrium with air (b) and closed systems with a total carbonate concentration of 30 mg/liter (c,e) and 485 mg/liter (d,f), respectively. The two redox potentials would roughly correspond to water in equilibrium wit air (a-d cf 50) and systems buffered by an Fe(III)(s)/Fe(II)(s)-equilibrium (e,f), respectively. Thus, the natural span of carbonate concentrations and redox conditions is illustrated. 

A practical problem in solution preparation usually requires a different strategy than our standard seven-step procedure. The technician must first identify a suitable conjugate acid-base pair and decide what reagents to use. Then the concentrations must be calculated, using pH and total concentration. Finally, the technician must determine the amounts of starting materials. The technician needs a buffer at pH = 9.00. Of the buffer systems listed in Table 18-1. the combination of NH3 and NH4 has the proper pH range for the required buffer solution. 

Figure 4, Decomposition of -n-butyU t -nitrosoacetamide by calf liver RNA in phosphate buffer (pH 7) at 25°C, The dashed line is the calculated rate based on coefficients given in Table II,...

Example. A solution of ethyl acetate in pH 9.5 buffer (25°C) was assayed in triplicate several times over a 20-hour period. The data obtained are presented in Table 2. The results were plotted on semilogarthmic graph paper as shown in Fig. 3. Calculate the psuedo-first-order rate constant for the hydrolysis of ethyl acetate at pH 9.5 (25°Q. 

In practice, a buffer solution (see Table 3. 4) is prepared by the partial neutralization of the selected weak acid or base with a suitable strong acid or base, or by the addition of the calculated amount of the corresponding salt. The assumption is made that the salt is completely dissociated in solution, e.g. an... 

We can prepare a buffer of almost any pH provided we know the pAa of the acid and such values are easily calculated from the Ka values in Table 6.5 and in most books of physical chemistry and Equation (6.50). We first choose a weak acid whose pKa is relatively close to the buffer pH we want. We then need to measure out accurately the volume of acid and base solutions, as dictated by Equation (6.50). 

A buffer solution of pH 3.00 is needed. From Table 5.1, select a weak acid-conjugate base combination that would give that pH and calculate the ratio of acid concentration to conjugate base concentration that would give that pH. 

The Gibbs free energy reaction profiles in Fig. 16 have been calculated from the results in Table 16 and the mechanism in (30) and refer to reaction in a 1 1 2-methylphenol buffer at buffer concentrations of 0.001 and 0.1 moldm" (Fig. 16(a) and (b), respectively). TS(1) is the transition state for opening of the intramolecular hydrogen bond and TS(2) is the transition... 

In the first type of study, pseudo first-order kinetics were observed in both the sediment and aqueous phases from t=0 through two half-lives in overall chlorpyrifos disappearance (total time -8 days). For these studies, computer calculations using the model illustrated in equations 7 were again used to calculate values for kj, k and kg, assuming a value of k equal to the pseudo first-order rate constant in distilled water buffered to the same pH. Values were also calculated for Obfi assuming kg 0 (equation 10) for comparison to the experimental kg values. The results of these calculations are shown in Table VII. 

Data in Table 4.2 corresponds to the application of the H-point standard addition method to a mixture of a commercial madder pigment diluted with silica, using morin as a reference compound. Calculations were performed by taking m/niR = 10.246, using square-wave voltammetric currents measured for sample-modified PIGEs in contact with an acetate buffer of pH 4.90. Linear plots of ii/ip(R) (squares) and i2/ip(R) (solid squares) vs. mA/mp for additions of purpurin are shown in Fig. 4.17. 

C.I. Direct Black 38, C.I. Direct Brown 95 and C.I. Direct Blue 6 were studied individually in the visible spectrum. A baseline was recorded using the phosphate reduction buffer in both cells. Subsequent additions of known amounts of dye and scanning allowed absorption maxima and molar absorptivity to be determined. Then 3 mg of sodium hydrosulfite was added to the dye-containing cell. The concentration of dye remaining after reduction was calculated using Beer s Law. The remaining dye varied from 0 to 6% of the original amount of dye added (Table I). Reduction was complete within 30 minutes. 

A more comprehensive analysis of the influences on the ozone solubility was made by Sotelo et al., (1989). The Henry s Law constant H was measured in the presence of several salts, i. e. buffer solutions frequently used in ozonation experiments. Based on an ozone mass balance in a stirred tank reactor and employing the two film theory of gas absorption followed by an irreversible chemical reaction (Charpentier, 1981), equations for the Henry s Law constant as a function of temperature, pH and ionic strength, which agreed with the experimental values within 15 % were developed (Table 3-2). In this study, much care was taken to correctly analyse the ozone decomposition due to changes in the pH as well as to achieve the steady state experimental concentration at every temperature in the range considered (0°C [Pg.86]

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