Buffer solutions stoichiometric concentrations

Buffers are usually prepared by dissolving the required amount of buffering substance and adjusting the pH to the desired value with strong acid (e.g., HQ) or alkali (e.g., NaOH). Normally, there is a free choice of add or alkali, and thus the counterion. However, for certain applications, for example with electrophoresis buffers, the counterion is specified. The molar concentration of a buffer solution is usually expressed as the stoichiometric concentration of the buffering compound. 

A formula for computing the buffer value may be obtained as follows. If to a liter of a solution of a weak acid of concentration C, B equivalents of a strong base are added, the stoichiometric concentrations of add and its salt will be respectively, (C B) and B. 

For this buffer solution, the standard approximations hold i.e. the actual concentration equals the stoichiometric concentration ... 

Determination of ozone in aqueous solution is perhaps the most problematic for a variety of reasons (1) ozone is unstable (2) ozone is volatile and easily lost from solution and (3) ozone reacts with many organic compounds to form products such as ozonides and hydrogen peroxide that are also good oxidants. Careful study of the use of iodometric methods for the determination of ozone in aqueous solution has revealed that the stoichiometric ratio of ozone reacted with iodine produced in the reaction varies from 0.65 to 1.5, depending on pFI, buffer composition and concentration, iodide ion concentration, and other reaction conditions. As a result, iodometric methods are not recommended. Ozone can be determined iodimetrically by addition of an excess of a standard solution of As(III), followed by titration of the excess As(III) with a standard solution of iodine to a starch endpoint. Methods using DPD, syringaldazine, and amperometric titrations have also been developed. 

An important point worth noting in Example 17-4 is that if a solution is to be an effective buffer, the assumptions c — x) cand (c -I- x) c will always be valid (c represents the numerical part of an expression of molarity). That is, the equilibrium concentrations of the buffer components will be very nearly the same as their stoichiometric concentrations. As a result, in Example 17-4 we could have gone directly from the stoichiometric concentrations of the buffer components to the expression... 

To calculate how the pH of a buffer solution changes when small amounts of a strong acid or base are added, we must first use stoichiometric principles to establish how much of one buffer component is consumed and how much of the other component is produced. Then the new concentrations of weak acid (or weak base) and its salt can be used to calculate the pH of the buffer solution. Essentially, this problem is solved in two steps. First, we assume that the neutralization reaction proceeds to completion and determine new stoichiometric concentrations. Then these new stoichiometric concentrations are substituted into the equilibrium constant expression and the expression is solved for [H30 ], which is converted to pH. This method is applied in Example 17-6 and illustrated in Figure 17-6. 

Direct Titrations. The most convenient and simplest manner is the measured addition of a standard chelon solution to the sample solution (brought to the proper conditions of pH, buffer, etc.) until the metal ion is stoichiometrically chelated. Auxiliary complexing agents such as citrate, tartrate, or triethanolamine are added, if necessary, to prevent the precipitation of metal hydroxides or basic salts at the optimum pH for titration. Eor example, tartrate is added in the direct titration of lead. If a pH range of 9 to 10 is suitable, a buffer of ammonia and ammonium chloride is often added in relatively concentrated form, both to adjust the pH and to supply ammonia as an auxiliary complexing agent for those metal ions which form ammine complexes. A few metals, notably iron(III), bismuth, and thorium, are titrated in acid solution. 

When plotted on a graph of pH vs. volume of NaOH solution, these six points reveal the gross features of the titration curve. Adding additional calculated points helps define the pH curve. On the curve shown here, the red points A-D were calculated using the buffer equation with base/acid ratios of 1/3 and 3/1. Point E was generated from excess hydroxide ion concentration, 2.00 mL beyond the second stoichiometric point. You should verify these additional five calculations. 

FIGURE 8.5 Multiple-sensor respirometry. Representative calibration traces of PNOS (thin line, left ordinate) and PHSS (thick line, right ordinate) operating simultaneously in PBS, pH 7.3 at 37°C, with 50pM DTPA in a closed chamber respirometer. After NO additions were made, the chamber solution was replaced with fresh buffer, to which Na2S stock solutions were then injected in a stepwise manner. The stable POS signal shown at 2 pM 02 demonstrates that the POS does not respond to NO or H2S. Injections of anoxic buffered NO and H2S stocks are shown with concentrations at arrows, as are additions of Lucina pectinata ferric hemoglobin I (metHb I), which stoichiometrically binds to H2S (after [41]). 

Molar ratio (MR) between sulfuric acid and sodium sulfate at constant sodium sulfate concentration. The second dissociation constant of sulfuric acid is rather low, in the range of 0.01. Consequently, in a solution containing both sulfuric acid and sodium sulfate at MR < 1, substantially all the sulfuric acid reacts with the stoichiometric amount of sodium sulfate to give sodium bisulfate (buffer action). Hence, the actual concentration of free protons (H ) is directly proportional to the actual concentration of sodium bisulfate and inversely proportional to that of the unreacted sodium. sulfate. This type of dependence indicates that the actual concentration of free protons should increase quickly when MR exceeds a certain critical value (ca. 0.5). At higher MR values the current transported by the protons becomes significant at the expense of that transported by the sodium ions, and the cathodic efficiency shows a sharp decrease. 

Fig. 5. Enthalpy of interaction of native calf-thymus DNA with (a) actinomycin D (full symbols), (b) daunomycin hydrochloride (open symbols). - Solvent 0.01 M phosphate buffer, pH 7.0. r is the final stoichiometric ratio between the molar concentration of antibiotic and the molar (phosphorous base) concentration of DNA. F= 25 °C. In two experiments with daunomycin, solutions of DNA already containing a given amount of the drug were added of more daunomycin to span the r indicated by the arrows.

However, these four coupled equilibria from Scheme 12.1 are not the full story. Boronic acids readily form stable complexes with buffer conjugate bases (phosphate, citrate and imidazole) [20], In fact, both binary boronate-X complexes are formed with Lewis bases (X), as well as ternary boronate-X-saccharide complexes. In some cases, these previously unrecognized species persist into acidic solution and tmder some stoichiometric conditions they can be the dominant components of the solution. These complexes suppress the boronate and boronic acid concentrations, leading to a decrease in the measured apparent formation constants (f pp). As a consequence, the scope of the simple diol-boronate recognition system is greatly expanded over the simple picture of Scheme 12.1. 

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