Preparing a Buffer

Chemical supply houses sell buffers having a variety of pH values and concentrations, but chemists or lab technicians often have to prepare a buffer solution for a specific environmental or biomedical application. Several steps are required to prepare a buffer  

Decide on the conjugate acid-base pair. The choice is determined mostly by the desired pH. Remember that a buffer is most effective when the buffer-component concentration ratio is close to 1 in that case, the pH is close to the pAf of the acid. Convert pK to K, choose the acid from a list, such as that in Appendix C, and use the sodium salt as the conjugate base. 

Find the ratio of [A ]/[HA] that gives the desired pH, using the Henderson-Hasselbalch equation. Note that, because HA is a weak acid, and thus dissociates very little, the equilibrium concentrations are approximately equal to the initial concentrations that is, 

Choose the buffer concentration and calculate the amounts to mix. Remember that the higher the concentration, the greater the buffer capacity. For most laboratory applications, concentrations from 0.05 M to 0.5 M are suitable. From a given amount (usually in the form of concentration and volume) of one component, find the amount of the other component using the buffer-component concentration ratio. 

Mix the amounts together and adjust the buffer pH to the desired value. Add small amounts of strong acid or strong base, while monitoring the solution with a pH meter. 

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Suppose you need to prepare a buffer with a pH of 9.36. Using the Henderson-Hasselbalch equation, you calculate the amounts of acetic acid and sodium acetate needed and prepare the buffer. When you measure the pH, however, you find that it is 9.25. If you have been careful in your calculations and measurements, what can account for the difference between the obtained and expected pHs In this section, we will examine an important limitation to our use of equilibrium constants and learn how this limitation can be corrected. 

To prepare a buffer, we can mix solutions of a weak acid HB and the sodium salt of that acid NaB, which consists of Na+ and B ions. This mixture can react with either strong base... 

Borate-gluconate eluant. Prepare a buffer concentrate by dissolving the following substances in water and making up to 1 L with distilled, de-ionised... 

Prepare a buffer solution (pH 10) by dissolving 8.0 g ammonium nitrate in 65 mL of de-ionised water and adding 35 mL of concentrated ammonia solution (sp. gr. 0.88). 

Procedure. Prepare a buffer solution (pH 7) by dissolving 3.7 g pure citric acid and 23.6 g of potassium dihydrogenphosphate in 1 L of distilled water. 

Prepare a buffer solution (pH 4.5) by dissolving 6 g acetic (ethanoic) acid and 13,6 g sodium acetate in 1 L of distilled water. Pipette 10 mL of a commercial sample of citrus fruit juice into a 1 L graduated flask and make up to the mark with oxygen-free water,... 

Suppose we are culturing bacteria that require an acid environment and want to prepare a buffer close to pH = 4. We prepare a buffer solution that is 0.040 M NaCH. C02(aq) and 0.080 M CH,COOH(aq) at 25°C. What is the pH of the buffer solution ... 

This very simple result makes it easy to make an initial choice of a buffer we just select an acid that has a pfC, close to the pH that we require and prepare an equimolar solution with its conjugate base. When we prepare a buffer for pH > 7 we have to remember that the acid is supplied by the salt, that the conjugate base is the base itself, and that the pKa is that of the conjugate acid of the base (and hence related to the pKh of the base by pR l + pKh = p/conjugate acid and base have unequal concentrations—such as those considered in Examples 11.1 and 11.2—are buffers, but they may be less effective than those in which the concentrations are nearly equal (see Section 11.3). Table 11.1 lists some typical buffer systems. 

Self-Test 11.3A Which of the buffer systems listed in Table 11.1 would be a good choice to prepare a buffer with pH close to 5 ... 

Now consider the overall shape of the pH curve. The slow change in pH about halfway to the stoichiometric point indicates that the solution acts as a buffer in that region (see Fig. 11.3). At the halfwayr point of the titration, [HA] = [A ] and pH = pfCa. In fact, one way to prepare a buffer is to neutralize half the amount of weak acid present with strong base. The flatness of the curve near pH = pKa illustrates very clearly the ability of a buffer solution to stabilize the pH of the solution. Moreover, we can now see how to determine pKa plot the pH curve during a titration, identify the pH halfway to the stoichiometric point, and set pKa equal to that pH (Fig. 11.8). To obtain the pfCh of a weak base, we find pK3 in the same way but go on to use pKa -1- pfq, = pKw. The values recorded in Tables 10.1 and 10.2 were obtained in this way. 

A second way to prepare a buffer is by adding strong base to a solution of a weak acid. This produces a buffer solution if the number of moles of strong base is about half the number of moles of weak acid. As a simple example, if 1 L of 0.5 M NaOH is mixed with 1 L of 1.0 M CH3 CO2 H, hydroxide anions react quantitatively... 

A buffer soiution must contain both the acid and its conjugate base, so at least two reagents must be added to water to prepare a buffer soiution. An acetate buffer can be prepared, for example, from pure water, concentrated acetic acid, and an acetate salt. The cation contained in the salt should not have acid-base properties of its own, so sodium acetate would be an appropriate choice, but ammonium acetate would not. 

A technician wants to prepare a buffer solution at pH - 9.00 with an overall concentration of 0.125 moi/L. The technician has soiutions of 1.00 M HCl and NaOH and bottles of all common salts. What reagents should be used, and in what quantities, to prepare 1.00 L of a suitable buffer ... 

C18-0119. In a biochemistry laboratory, you are asked to prepare a buffer solution to be used as a solvent for isolation of an enzyme. On the shelf are the following solutions, all 1.00 M formic acid (Za = 1.8 X 10 ), acetic acid = 1.8 x 10 ), sodium formate (NaHC02), and sodium acetate (NaCH3 CO2). Describe how you would prepare 1.0 L of a pH = 4.80 buffer solution... 

Because a buffer contains approximately equal amounts of a weak base (NH3) and its conjugate acid (NH4+), to prepare a buffer we simply add an amount of HC1 equal to approximately half the amount of NItyaq) initially present. 

The pATa s of the three acids help us choose the one to be used in the buffer. It is the acid with a pATa within 1.00 pH unit of 3.50. pATa = 3.74 for HCH02, pATa = 4.74 for HC2H302, and p/sTj =2.15 for H3P04. Thus, we choose HCH02 and NaCH02 to prepare a buffer with pH = 3.50. The Henderson-Hasselbalch equation is used to determine the relative amounts of each component present in the buffer solution. 

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). 

Worked Example 6.12 We need to prepare a buffer of pH 9.8 by mixing solutions of ammonia and ammonium chloride solution. What volumes of each are required Take the Ka of the ammonium ion as 6 x 10 l0. Assume the two solutions have the same concentration before mixing. 

Using appropriate buffer equations, one can calculate (a) the pH value of a buffer system when pA" and molar ratio (or concentration) of buffer components are known, and (b) the molar ratio of buffer components required to prepare a buffer with a desired pH value. 

Using the buffer (Henderson Hasselbalch) equation, one can calculate the molar ratio of buffer components needed to prepare a buffer of desired pH. This is explained in the following examples. 

What is the molar ratio of base to salt required to prepare a buffer solution having a pH of 10 Assume that the pKb of the base is 4.62 at 25°C. 

What ratio of the concentrations of acetic acid to sodium acetate is needed to prepare a buffer solution of pH = 4.00 The Ka of acetic acid is 1.76 X 10... 

Prepare a buffer solution that is 20 mM tris-(hydroxymethyl)amino methane (TRIS or THAM), 6 mM sodium acetate, and 1 mM disodium EDTA. Adjust to pH = 7.9. 

This seemingly simple series of events does not address all the requirements. If the device is preparing media does that mean it prepares a buffer to be diluted or only degasses the premixed media When media is dispensed, is there a need to perform a preliminary dispense to assure removal of the previous media If samples are to be read on-line is dilution required prior to reading Systems intended for method development (MD) will have many different requirements than one intended for QA. The value of the automation to the user may be very different for each of these two areas. In fact the MD user may not appreciate the need to automate more than one run at a time and will prefer a semiautomated system, since the MD user may have many different experiments to perform that may be labor intensive. Just a few... 

What substance can be added to an aqueous solution of propanoic acid to prepare a buffer solution ... 

A student wishes to prepare a buffer solution that has a pH of 10.0. The following chart lists five weak bases with their K, values. Which base would be most appropriate to use in preparing the buffer ... 

Which of the following acids would be most suitable for preparing a buffer of pH 3.10 (i) hydrogen peroxide (ii) propanoic acid (iii) cyanoacetic acid (iv) 4-aminobenzenesulfonic acid ... 

Go back and read the Example Preparing a Buffer in a Diprotic System" on page 187. See if it makes more sense now. 

How many grams of sodium acetate must be added to 250 ml of 0 200 M HC2H3C)2 in order tQ prepare a buffer with pH = 5 0(P... 

Equation 2.6 is the familiar Henderson-Hasselbalch equation, which defines the relationship between pH and the ratio of acid and conjugate base concentrations. The Henderson-Hasselbalch equation is of great value in buffer chemistry because it can be used to calculate the pH of a solution if the molar ratio of buffer ions ([A-]/[HA]) and the pKa of HA are known. Also, the molar ratio of HA to A- that is necessary to prepare a buffer solution at a specific pH can be calculated if the pKa is known. 

El 1. You need to prepare a buffer for biochemistry lab. The required solution is 0.5 M sodium phosphate, pH 7.0. Use the Henderson-Hasselbalch equation to calculate the number of moles and grams of monobasic sodium phosphate (NaH2P04) and dibasic sodium phosphate (Na2HP04) necessary to make 1 liter of the solution. 

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