Buffer solutions calculation

Explain the behavior of a buffer solution. Calculate its pH from the concentrations of its conjugate acid-base pair (Section 15.5, Problems 43-46). 

This chapter describes polyfunctional acid and base systems, including buffer solutions. Calculations of pH and titration curves are also described. 

A glass electrode was determined to have a potential of 0.395 V when measured against the SCE in a standard pH 7.00 buffer solution. Calculate the pH of the unknown solution for which the following potential readings were obtained (the potential decreases with increasing pH) ... 

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

EXAMPLE 16.1 Calculating the pH of a Buffer Solution Calculate the pH of a buffer solution that is O.KX) M in HC2H3O2 and 0.100 M in NaC2H302 SOLUTION ... 

Henderson-Hasselbach equation A simplified version of the relationships used in calculations on buffer solutions. 

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. 

Calculate the solubility (g/100 mL) of iron(n) hydroxide in buffered solutions with the following pH s. 

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).

We can adjust the pH of a buffer solution by adding some acid to lower it or some base to raise it. Another way to adjust the pH of a buffer by adding more salt (which supplies the conjugate acid or base). Example 11.2 shows how to calculate the effect of added acid or base on the pH of a buffer. 

Suppose we dissolve 1.2 g of sodium hydroxide (0.030 mol NaOH) in 500. mL of the buffer solution described in Example 11.1. Calculate the pH of the resulting solution and the change in pH. Assume that the volume of the solution remains unchanged. 

Commercially available buffer solutions can be purchased for virtually any desired pH. A buffer solution commonly used to calibrate pH meters contains 0.025 m Na2HP04(aq) and 0.025 M KH2P04(aq) and has pH = 6.87 at 25°C. However, the method demonstrated in Example 11.1 would give pH = 7.2 for this solution. Because these calculations interpret activities as molarities, not effective molarities, they ignore ion—ion interactions so the values calculated are onl) approximate. 

Calculate the pH change when acid or base is added to a buffer solution (Example 1 1.2). 

The question asks if this is a buffer solution. A buffer solution contains both a weak acid and its conjugate base as major species. Thus, to answer the question, we must calculate the concentrations of acetate anions and acetic acid in the solution. We use a compressed version of the seven-step method to obtain these concentrations. 

The sum of the moles of the products is equal to the number of initial moles of acetate, and the final amounts of acid and base are comparable. Thus, the calculations support the conclusion that the final mixture is a buffer solution. 

CI8-OOO2. Calculate the concentrations of the major species present in a buffer solution prepared by adding 1.40 g NaOH to 140. mL of 0.750 M NH4 Cl. 

The buffer equation, which is often called the Henderson-Hasselbalch equation, is used to calculate the equilibrium pH of a buffer solution directly from initial concentrations. The approximation is valid as long as the difference between initial concentrations and equilibrium concentrations is negligibly small. As a rule of thumb, the buffer equation can be applied when initial concentrations of H j4 and A differ by less than a factor of 10. Example provides an illustration of the use of the buffer equation. 

Buffer capacity is determined by the amounts of weak acid and conjugate base present in the solution. If enough H3 O is added to react completely with the conjugate base, the buffer is destroyed. Likewise, the buffer is destroyed if enough OH is added to consume all of the weak acid. Consequently, buffer capacity depends on the overall concentration as well as the volume of the buffer solution. A buffer solution whose overall concentration is 0.50 M has five times the capacity as an equal volume of a buffer solution whose overall concentration is 0.10 M. Two liters of 0.10 M buffer solution has twice the capacity as one liter of the same buffer solution. Example includes a calculation involving buffer capacity. 

Use the seven-step strategy to calculate the pH of the buffer solution using the buffer equation. Then compare the amount of acid in the solution with the amount of added base. Buffer action is destroyed if the amount of added base is sufficient to react with all the acid.The buffering action of this solution is created by the weak acid H2 PO4 and its conjugate base HP04. The equilibrium constant for this... 

C18-0009. Addition of 5.25 g of NaOH to the buffer solution described in Example would exceed the capacity of the buffer. Calculate the concentration of excess hydroxide ion and the pH of the solution. 

For the preparation of a buffer solution of pH = 5.00, sodium acetate and acetic acid should be added to pure water in a molar ratio of 1.8 1.0. The exact amounts must be calculated using the volume and concentration ofthe solution, as Example Illustrates. 

Because we know we are dealing with a buffer solution made from a specific conjugate acid-base pair, we can work directly with the buffer equation. We need to calculate the ratio of concentrations of conjugate base and acid that will produce a buffer solution of the desired pH. Then we use mole-mass-volume relationships to translate the ratio into actual quantities. 

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. 

C18-0095. The pH of a formic acid/formate buffer solution is 4.04. Calculate the acid/conjugate base ratio for this solution. Draw a molecular picture that shows a small region of the buffer solution. (You may omit spectator ions and water molecules.) Use the following symbols ... 

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