6. I MPLEMENTACIÓN Y EVALUACIÓN
6.5. R ESULTADOS
The salt of the acid is also referred to as the proton acceptor (A) and the acid (HA) as the proton donor (D). 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 and the of HA are known. Also, the molar ratio of HA to that is necessary to prepare a buffer solution at a specific pH can be calculated if the is known.
A solution containing both HA and has the capacity to resist changes in pH; i.e., it acts as a buffer. If acid were added to the buffer solution, it would be neutralized by in solution:
>> (Eq. 3.7)
Base added to the buffer solution would be neutralized by reaction with HA:
>> (Eq. 3.8)
The most effective buffering system contains equal concentrations of the acid, HA, and its conjugate base, According to the Henderson-Hasselbalch equation (3.6), when is equal to [HA], pH equals Therefore, the of a weak acid-base system represents the center of the buffering region. The effective range of a buffer system is generally two pH units, centered at the value (Equation 3.9).
>> (Eq. 3.9)
Selection of a Biochemical Buffer
Virtually all biochemical investigations must be carried out in buffered aqueous solutions. The natural environment of biomolecules and cellular organelles is under strict pH control. When these components are extracted from cells, they are most stable if maintained in their normal pH range, usually 6 to 8. An artificial buffer system is found to be the best substitute for the natural cell milieu. It should also be recognized that many biochemical processes (especially some enzyme processes) produce or consume hydrogen ions. The buffer system neutralizes these solutions and maintains a constant chemical environment.
Effective pH range for a buffer = pKa ; 1
Although most biochemical solutions require buffer systems effective in the pH range 6 to 8, there is occasionally a need for buffering over the pH range 2 to 12.
Obviously, no single acid–conjugate base pair will be effective over this entire range, but several buffer systems are available that may be used in discrete pH ranges. Figure 3.2 compares the effective buffering ranges of common biological buffers. It should be noted that some buffers (phosphate, succinate, and citrate) have more than one value, so they may be used in different pH regions. Many buffer systems are effective in the usual biological pH range (6 to 8); however, there may be major problems in their use. Several characteristics of a buffer must be considered before a final selection is made. The molecular weights and pK values of several common buffer compounds are listed in Appendix III. Following is a discussion of the advantages and disadvantages of the commonly used buffers.
Phosphate Buffers
The phosphates are among the widely used buffers. These solutions have a high buffering capacity and are very useful in the pH range 6.5 to 7.5. Because phos-phate is a natural constituent of cells and biological fluids, its presence affords a more “natural” environment than many buffers. Sodium or potassium phos-phate solutions of all concentrations are commonly prepared with the use of the Henderson-Hasselbalch equation. The major disadvantages of phosphate solu-tions are: (1) precipitation or binding of common biological casolu-tions ( and ), (2) inhibition of some biological processes, including some enzymes, and (3) limited useful pH range.
Mg2+
Ca2+
pKa
Phosphate Phosphate Phosphate
0 1 2 3 4 5 6 7 8 9 10 11 12
pH Citrate Citrate
Succinate Acetate Formate
Borate Tris Imidazole
MES CAPS
HEPES
Histidine Glycine
Bis-Tris
Bicine
FIGURE 3.2 Effective buffering ranges of several common buffers.
See Table 3.1 for abbreviations.
STUDY EXERCISE 3.1
Henderson-Hasselbalch EquationHere we will describe the calculation method for preparation of a sodium phosphate buffer of 0.05 M total phosphate concentration, pH of 7.0, and a total, final volume of 1 liter. The Henderson-Hasselbalch equation (Equation 3.6) will be used to calculate the conjugate acid and conjugate base concentrations. At pH 7.0, the two forms of phos-phate present are:
We now use the Henderson-Hasselbalch equation with and (from Appendix III) to calculate the molar ratio:
This gives only one equation, but two unknowns, the concentrations of A and D; there-fore, we need another quantitative relationship between [A] and [D]. We know that the total phosphate concentration is 0.05 M, so:
Using these two equations for A and D, we determine that we need 0.031 mole of D and 0.019 mole of A. The appropriate phosphate reagent forms available commer-cially are:
monobasic sodium phosphate monohydrate, mole.
dibasic sodium phosphate heptahydrate, mole.
Converting the number of moles of each reagent to grams:
To prepare the solution, dissolve the appropriate amount of each reagent in about 975 mL of purified water in a 1-liter beaker. Check the pH of the solution and adjust to 7.0 by careful, dropwise addition (with stirring) of dilute NaOH or dilute phosphoric acid. Transfer the solution to a 1-liter volumetric flask, add water to the mark, mix well, and do a final check of the pH.
In the mid-1960s, N. E. Good and his colleagues recognized the need for a set of buffers specifically designed for biochemical studies (Good and Izawa, 1972). He and others noted major disadvantages of the established
STUDY EXERCISE 3.2
Buffer Made by TitrationWe will illustrate the preparation of a Tris buffer using the titration method. (In Study Exercise 3.1, we used the calculation method to prepare the phosphate buffer.) Because Tris is commercially available in highly purified crystals, Tris Base the ap-propriate amount may be weighed, dissolved in water, and titrated to the desired pH with an acid, usually HCl. For 1 liter of 0.1 M Tris-HCl buffer, pH 7.0, dissolve 12.11 g (0.1 mole) of Tris Base in about 975 mL of purified water in a 1-liter beaker. Adjust the pH (originally about 10) to 7.0 by careful, dropwise addition of concentrated hydrochloric acid. This should be done in a 1-liter beaker with gentle stirring. Transfer the solution to a 1-liter vol-umetric flask, add water to the mark, mix well, and make a final check of the pH.
As you can see by comparing the two procedures, the titration method for buffer preparation is faster and much more convenient than the calculation method. Is it possible to prepare the phosphate buffer in Study Exercise 3.1 by using the titration method? If so, describe the reagent(s) needed and the procedure you would use.
1MW = 121.12,
buffer systems. Good outlined several characteristics essential in a biological buffer system:
1. between 6 and 8.
2. Highly soluble in aqueous systems.
3. Exclusion or minimal transport by biological membranes.
4. Minimal salt effects.
5. Minimal effects on dissociation due to ionic composition, concentration, and temperature.
6. Buffer–metal ion complexes nonexistent or soluble and well defined.
7. Chemically stable.
8. Insignificant light absorption in the ultraviolet and visible regions.
9. Readily available in purified form.
Good investigated a large number of synthetic zwitterionic buffers and found many of them to meet these criteria. Table 3.1 lists several of these buffers and their properties. Good’s buffers are widely used, but their main disadvantage is high cost.
Some zwitterionic buffers, such as Tris, HEPES, and PIPES, have been shown to produce radicals under a variety of experimental conditions, so they should be avoided if biological redox processes or radical-based reactions are being studied.
Radicals are not produced from MES or MOPS buffers.
The use of the synthetic zwitterionic buffer Tris [Tris(hydroxymethyl)-aminomethane] is now probably greater than that of phosphate. It is useful in the pH range 7.5 to 8.5. Tris is available in a basic form as highly purified crystals, which makes buffer preparation especially convenient. Although Tris is a primary amine, it causes minimal interference with biochemical processes and does not precipitate calcium ions. However, Tris has several disadvantages, including:
(1) pH dependence on concentration, since the pH decreases 0.1 pH unit for each 10-fold dilution; (2) interference with some pH electrodes; and (3) a large compared to most other buffers. Most of these drawbacks can be mini-mized by: (1) adjusting the pH after dilution to the appropriate concentration, (2) purchasing electrodes that are compatible with Tris, and (3) preparing the buffer at the temperature at which it will be used.
¢pKa>C°
pKa
Buffer Name Abbreviation
Structure (All structures
shown in salt form) pKa120°C2
Useful
TABLE 3.1 Structures and Properties of Several Synthetic Zwitterionic Buffers
Buffer Name Abbreviation
Structure (All structures shown
in salt form) pKa120°C2
Useful pH
Range ≤pKa/C°
Concentration of a Saturated
Solution 1M, 0°C2
3-(N-Morpholino)-propanesulfonic acid
MOPS
O NHCH+ 2CH2CH2SO3– 7.20 6.5–7.9 — —
Piperazine- -bis-2-ethanesulfonic acid
N,N¿ PIPES
–O2SCH2CH2N NHCH+ 2CH2SO3– 6.8 6.4–7.2 -0.0085 —
N-Tris(hydroxymethyl)- methyl-2-aminoethane-sulfonic acid
TES 1HOCH223N+HCH2CH2SO-3 7.5 7.0–8.0 -0.020 2.6
N-Tris(hydroxymethyl)-methylglycine
Tricine 1HOCH223CN+H2CH2CH2COO- 8.15 7.5–8.5 -0.021 0.8
Tris(hydroxymethyl)-aminomethane
Tris 1HOCH223CN+H3 8.3 7.5–9.0 -0.031 2.4
(continued )
Carboxylic Acid Buffers
The most widely used buffers in this category are acetate, formate, citrate, and succinate. This group is useful in the pH range 3 to 6, a region that offers few other buffer choices. All of these acids are natural metabolites, so they may inter-fere with the biological processes under investigation. Also, citrate and succinate may interfere by binding metal ions ( etc.). Formate buffers are especially useful because they are volatile and can be removed by evaporation under reduced pressure.
Borate Buffers
Buffers of boric acid are useful in the pH range 8.5 to 10. Borate has the major dis-advantage of complex formation with many metabolites, especially carbohydrates.
Amino Acid Buffers
The most commonly used amino acid buffers are glycine (pH 2 to 3 and 9.5 to 10.5), histidine (pH 5.5 to 6.5), glycine amide (pH 7.8 to 8.8), and glycylglycine (pH 8 to 9). These provide a more “natural” environment similar to cellular components and extracts; however, they may interfere with some biological processes, as do the carboxylic acid and phosphate buffers.
Buffer Dilutions
In biochemistry lab activities, buffers of many different concentrations and volumes are required. It is usually not practical, economical, or convenient to prepare all needed buffers from scratch. However, it is possible to have available a few con-centrated stock buffer solutions that may be diluted to produce a new buffer solu-tion of desired concentrasolu-tion and amount.
Dilutions are usually defined in the following way: A dilution of 1 mL to 10 mL (also 1:10; or 1 to 10) means that you take 1 mL of stock solution and add water to a final total volume of 10 mL. It is important to note that the amount of the buffer components taken from the stock solution is equal to the amount of buffer components in the diluted solution. However, the concentration of the buffer components has been changed. If the stock solution above had a buffer concentration of 0.5 M, then the 1:10 diluted solution has a buffer concentration of 0.05 M.
Two kinds of dilutions, linear and serial, are used most often in biochemistry.
• Linear dilutions Table 3.2 shows the results of a linear dilution process in which a stock solution of 1 M is used to produce samples of different concentrations. Note that there is a linear (progressive) decrease in the concentration of the diluted samples, but the final volumes remain the same (10 mL). The volume of stock solution to use in each case is calculated from the following relationship:
volume1stock2 = volume1diluted2 * concentration 1diluted2 concentration1stock2
Fe3+, Ca2+, Zn2+, Mg2+,
Sample calculation for the first diluted solution in Table 3.2:
• Serial dilutions Here dilution starts with a stock buffer solution and then each diluted solution produced is used to prepare the next, etc. For example:
2.0 mL of Solution I is diluted to a final volume of 100 mL (Solution II) 2.0 mL of Solution II is diluted to a final volume of 100 mL (Solution III) 2.0 mL of Solution III is diluted to a final volume of 100 mL (Solution IV) If Solution I has a buffer concentration of 10 M and water is used for dilu-tion in each case, what is the buffer concentradilu-tion of Soludilu-tion II?
The equation above for linear dilution may also be used here:
For Solution II: concentration1diluted2 = 2.0 mL * 10 M
100 mL = 0.2 M volume1stock2 = 10 mL * 0.80 M
1.0M = 8.0 mL TABLE 3.2 Example of a Linear Dilution Process
Stock (M) Final dilution (M) Stock (mL) Water (mL)* Final volume (mL)
1.0 0.80 8.0 2.0 10
1.0 0.60 6.0 4.0 10
1.0 0.40 4.0 6.0 10
1.0 0.20 2.0 8.0 10
1.0 0.00 0.0 10.0 10
*Because all solutions are not additive, this number is the approximate amount of water needed to bring the final volume to 10 mL. It may be necessary to add a few drops of water to bring the final volume to 10 mL.