Laboratory 3: Biological Buffers

The exercises in this laboratory are designed to illustrate the ionization properties of amino acids as a basis for understanding the functional properties of proteins that will be illustrated later in the course. In addition, you will be responsible for calculating the components and preparing a biological buffer of a known pH. This exercise is intended to provide you with practical experience in basic laboratory procedures and illustrate the buffer properties of common biochemical solutions.

pH and the preparation of buffered solutions

pH

Most biological processes within the cell take place in aqueous environments. Water is an amphoteric substance that serves as a proton donor (acid) or a proton acceptor (base). It exists as an equilibrium represented by the following equation:

H2O « H+ + OH-

In pure water, [H+] = [OH-] = 10-7 M. However, the concentration of H+ and OH- in aqueous solutions can vary greatly because many solutes contribute (acids) or absorb (bases) protons. Although the ratio of [H+]/[OH-] can vary greatly in aqueous solutions, the product of the proton and hydroxide ion concentration in aqueous solutions is always 10-14 M (i.e. [H+] x [OH-] = 10-14 M). Therefore, the [OH-] of  an aqueous solution can be calculated if the [H+] is determined experimentally, and vice versa. As a convenient method of expressing the concentration of [H+] in a solution, chemists use the term, pH. The alkalinity or acidity of a solution is expressed in pH units. pH is defined as:

pH = -log[H+] or pH = log1/[H]

Therefore, the pH of pure water is 7.

However, many acidic and basic solutes in cells react with the protons or hydroxyl groups. This could potentially shift the equilibrium of the above reaction and therefore dramatically change the pH within the cell. As one might imagine, such wild fluctuations in cellular pH would seriously effect normal physiological processes. In order to maintain a constant pH, cells utilize natural buffering agents that absorb or contribute protons to maintain a constant pH. To study biological reactions in vitro, it is necessary for investigators to use artificial buffers to maintain a constant pH during a reaction or when culturing cells.

A buffer is a solution that resists pH change upon the addition of acid or base. All laboratory pH buffers can be thought of as weak acids and their dissociation in water can be described by the following equation:

HA  «  A- +  H+ (Equation 1)

HA = any weak acid

The degree of the dissociation of HA into A- and H+ can be described by an equilibrium constant, Ka.

Ka = [A-][H+]/[HA] (Equation 2)

By taking the reciprocal of both sides of the equation

1/Ka = [HA]/[A-][H+]

and by taking the log of both sides

log1/Ka = log1/[H+] + log[HA]/[A-]

The log1/[H+] = pH, and by definition, we can define log1/Ka = pKa.

Therefore,

pKa = pH + log[HA]/[A-]

By reorganizing the terms in this equation, we can derive an equation that describes the pH of a solution containing a solute with a known pKa.

pH = pKa + log[A-]/[HA] (Equation 3)

This is known as the Henderson-Hasselbach equation. This equation will be useful throughout the course and during your scientific careers for making buffer solutions.

When the concentrations of the conjugate acid and base of a buffering compound are equal, then pH = pKa. This is the point of maximum buffering capacity of the solution. This is most obvious when a titration curve of the substance is generated by titering a known concentration of the buffer with varying amounts of a strong acid or base. For a monoprotic buffer (a molecule that contributes only a single proton) like acetic acid, the titration curve will look something like figure 1 but the volume of base required will vary.

Figure 1. Titration of 0.1 M acetic acid with NaOH.

EXAMPLE:

Calculate the volume of glacial acetic acid (17.6 N) and weight of sodium acetate (formula weight = 82) that is required to make 100 ml of 0.2 M buffer at pH 3.9. The pKa of acetic acid is 4.8.

Solution:

For acetic acid pH = pKa + log[CH3COO-]/[CH3COOH]

Let [CH3COOH] = (x) mol/l

Then [CH3COO-] = (0.2 - x) mol/l

Therefore, pH = pKa + log(0.2 - x/x)

3.9 = 4.8 + log(0.2 - x/x)

-0.9 = log(0.2 - x/x)

Solving for x:

x = 0.178 mol/l = [CH3COOH]

0.2 - x = 0.022 mol/l = [CH3COO-]

For 100 ml of solution:

0.1L x 0.178 mol/l x l/17.6 mol = 0.00101 L CH3COOH

0.1 x 0.022 mol/l x 82 g/mol = 0.18 g CH3COONa

Mix the above and bring to volume with water.

Many laboratory buffers are multiprotic (e.g. phosphoric acid and glycine). Therefore, these buffers have multiple pKa's corresponding to the dissociation of protons from each functional group. For phosphoric acid the proton dissociation equations can be written as:

H3PO4  «  H2PO4- + H+ pKa1 = 2.12

H2PO4-  « HPO4-2 + H+ pKa2 = 7.21

HPO4-2   « PO4-3 + H+ pKa3 = 12.32

To use the Henderson-Hasselbach equation with multiprotic buffers, you must select the pKa and corresponding dissociation equation that is closest to the pH of the solution you desire.

EXAMPLE:

Calculate the pH of a solution prepared by mixing 1.83 g of KH2PO4 (F.W. = 136) and 1.16 g of K2HPO4 (F.W. = 174) in 0.2 l of water.

Solution:

The applicable dissociation equation for this solution corresponds to pKa2 (see above).

Therefore, pH = pKa2 + log[K2HPO4]/[KH2PO4]

[K2HPO4] = 1.16 g x mol/174 g x 1/0.2 l = 0.033 M

[KH2PO4] = 1.83 g x mol/136 g x 1/0.2 l = 0.067 M

pH = 7.21 + log(0.033/0.067)

pH = 6.92

Experimental procedures 

Week 1: Preparation of an aqueous buffer solution

1. Prepare 100 ml of 0.1 M NaPO4, pH 7.5. Start with solid NaH2PO4 and Na2HPO4 and H2O. Use the Henderson-Hasselbach equation to calculate the molar amount of each form of NaPO4 that you need to add to the solution. Convert the molar amount to a weight amount using the formula weight given on the containers of each chemical.

2. Weigh out the appropriate amount of each component on the balance.

3. Set up a 100 ml beaker on a stir plate. Add 80 ml of H2O and a stir bar and start stirring the liquid. Dissolve the solid NaPO4 components by slowly adding each powder to the beaker. Allow the solid to dissolve completely. When they are dissolved, adjust the volume to 100 ml using a 100 ml graduated cylinder.

4. Transfer the solution back to the beaker and measure the pH of the solution using the pH meter. Does it agree with the desired pH of 7.5? If not, can you explain why?

Week 2: Experimental procedures

Amino acids as buffers

Amino acids are the basic building blocks of protein molecules. All amino acids have two or more different functional groups. Two groups found in all amino acids are the a-COOH and a-NH3+ groups. The a-COOH and a-NH3+ groups are weak acid groups that ionize in aqueous solutions to by donating a proton. These groups have pKa values of 2-3 and 9-10, respectively. In addition some amino acids (i.e. lysine and aspartic acid) have side chains that also act as weak acids and bases. The ability of these groups to donate and accept protons is essential to the chemical properties of proteins, and the ionization of these groups often take part in the formation of protein structure and the catalytic activities of enzymes.

pH Measurement

The pH of a solution can be measured by a number of methods. The two most common methods are the use of pH paper and the use of a pH meter. pH paper is a paper that is impregnated with mixtures of dyes. The dyes in the paper vary in color depending on the pH of the surrounding solution. The pH of a unknown solution can be determined by placing a drop of a solution on the paper and comparing the paper's color to a set of color standards generated from known pH values. This method is accurate only within 0.5 pH units because of the judgement involved in comparing the colors of the unknown and standard papers. pH paper will give an estimate of the pH, but cannot be used to quantitatively determine the pH of a solution. However, pH paper is useful in situations where a biochemist needs to quickly determine the approximate pH of a solution or in situations where a pH meter cannot be used.

The pH meter is used to determine the exact pH of a solution. It is an electronic instrument that consists of a voltimeter attached to an electrode that is permeable to H+ ions. Submerging the electrode in a solution will result in a voltage reading that is proportional to the concentration of H+ in the solution. The display of the voltimeter is usually altered to give a direct reading of the pH rather than a voltage. The pH meter will give an exact measurement of the H+ concentration of a solution. In this exercise, you will use the pH meter to prepare solutions of a given pH. The use of the instruments is described on an operations guide sheet that is placed next to each instrument.

Titration of the amino acid, glycine

To illustrate the ionization properties of the functions groups of amino acids, you will perform a titration of the a-COOH and a-NH3+ groups of the simplest amino acid, glycine. From these data, you will determine the pKa values of these two functional groups.

1. Standardize the pH meter according to the instructions supplied with the meter.

2. Transfer 50 ml of a 0.1 M solutions of glycine to a clean 250 ml beaker and add 50 ml of H2O. Place a stir bar in the beaker and place the beaker on a stir plate next to the pH meter.

3. Rinse the electrode of the pH meter and place it in the glycine solution. Stir the solution very gently, and measure the pH.

4. Fill a 25 ml buret with a solution of 1 M NaOH and adjust the buret over the glycine solution.

5. Begin titrating the glycine solution by adding 0.2 ml increments of 1 M NaOH while gently stirring the solution. Record the precise volume of NaOH added at each step. Be sure that the electrode remains submerged during the procedure. Stop between the addition of each increment of 1 M NaOH and record the pH. Be sure that you allow the pH to stabilize before you record the value.

6. Continue the titration until the pH reaches 12-13 units.

7. Tabulate the data as given below in the example. Calculate the equivalents of base added at each step of the titration, and calculate the accumulated equivalents at each step.

An equivalent refers to the moles of OH- or H+ added with each addition of NaOH. It is determined by multiplying the volume of NaOH added by the concentration of the NaOH.

8. Plot the accumulated equivalents of NaOH added vs. the pH on a linear graph similar to that shown in Figure 1.

9. From your plot, estimate the pKa values for the a-COOH and a-NH3+ groups of glycine.

Preparation of an aqueous buffer solution

1. Prepare 50 ml of 0.3 M NaPO4, pH 7.5. Start with solid NaH2PO4 and Na2HPO4 and H2O. Use the Henderson-Hasselbach equation to calculate the molar amount of each form of NaPO4 that you need to add to the solution. Convert the molar amount to a weight amount using the formula weight given on the containers of each chemical.

2. Weigh out the appropriate amount of each component on the balance.

3. Set up a 100 ml beaker on a stir plate. Add 30 ml of H2O and a stir bar and start stirring the liquid. Dissolve the solid NaPO4 components by slowly adding each powder to the beaker. Allow the solid to dissolve completely. When they are dissolved, adjust the volume to 50 ml using a 100 ml graduated cylinder.

4. Transfer the solution back to the beaker and measure the pH of the solution using the pH meter. Does it agree with the desired pH of 7.5? If not, can you explain why?

5. One of the most common buffers used in cell biology and biochemistry is called phosphate buffered saline (PBS). It is composed of 0.01 M NaPO4, pH 7.5, and 0.15 M NaCl. Use the 0.3 M NaPO4 solution that you just prepared and solid NaCl to prepare 100 ml of PBS. Measure the pH using the pH meter.