Dr. Jon Atwater
Adapted from Tyner 3/96
Lab Skill: Week 6 Preparation of Buffers
Biological systems are very sensitive to changes in pH because the concentration
of hydrogen ion can affect cellular chemical reactions. Most chemical reactions in cells
occur with the help of biologic catalysts called enzymes. The three-dimensional shape,
and the charge of certain amino acid side chains within the enzyme are important for their
functional activity. Changes in pH can affect the shape and/or charged groups on an
enzyme, affecting its function. Typically, enzymes exhibit peak activity within a very
narrow pH range, thus the importance of maintaining the proper pH. The interior of cells,
the blood, and other fluids are buffered in living systems (in vivo). Buffered solutions
resist a change in pH even when H+ ions are added to or removed from the system. In the
laboratory, we use buffered solutions to maintain the proper pH for enzymatic reactions
in vitro (in the test tube).
A buffer system is composed of a weak acid and its anion (conjugate base). We
observed a buffer system in action last lab when we titrated acetic acid with NaOH. Look
at the graph that you generated last week. Note that between 0.1 and 0.9 OH- equivalents
added (remember adding OH- is the same as removing H+), the pH didn’t change
drastically, the solution is acting as a buffer. Recall the equilibrium equation for acetic
HC2H3O2 === H+ + C2H3O2-
As H+ is removed by adding OH-, more acetic acid will dissociate to produce H+ and
acetate ion in accordance with its equilibrium constant. Likewise, if H+ were added to
the system, the equilibrium would shift to the left, forming acetic acid. If we shift the
reaction too far in either direction, we loose the buffering capacity by depleting the acid
or the conjugate base, as we observe in the titration curve. A very important equation can
be derived from rearranging the equilibrium equation:
[H+] = Ka [Acid] , Take the –log of each side; –log[H+] = -log Ka [Acid]
Note that: -log Ka = pKa, then, pH = pKa + (-log [Acid])
or pH = pKa + log [Anion] This is known as the Henderson Hasselbach equation.
This equation is extremely useful in the lab because it can be used to figure out how to
make a buffer. Notice that when the [Acid] = [Anion], the pH = pKa. When a buffer has
a pH equal to the pKa, it has its optimal buffering capacity. Therefore, when selecting a
buffer you would want to select one that has a pKa near he pH that you want to maintain.
As a general rule, the pH range of a buffer should be no more than +/- 1 of the pKa. Last
week, you experimentally determined the pKa for acetic acid (pKa = 4.75), although you
didn’t know it! The useful range for an acetate buffer is then pH 3.75-5.75.
There are two ways to prepare a buffer solution of a specific concentration and
pH. The first way is to use the Henderson Hasselbach equation to determine the
concentrations of acid and anion that we need to obtain a particular pH. Lets try an
example. Suppose we wanted to make a 100ml of 0.5M Acetate buffer pH 4.
Since the Ka for acetic acid is 1.76 x 10-5, pKa = -log 1.76 x 10-5 = 4.75. Plugging this in
to the Henderson Hasselbach equation, 4 = 4.75 + log [Anion]
or 10 = [Anion] = 0.1778, 0.1778[Acid] = [Anion]
From this we can calculate how much acid and anion we need to prepare the buffer!
Lets say we had stock solutions of 1M acetic acid and 1M sodium acetate, using
C1V1 = C2V2 : Since the final concentration is to be0.5M, the contribution from acetic
acid and sodium acetate must total 0.5M. Therefore, if the final concentration of acetic
acid is x, then the final concentration of sodium acetate must be (0.5-x). Plugging this in
to our equation gives:
0.1778(x) = 0.5-x, or x = 0.424M = [acid], and the final [anion]
would be 0.5 M– 0.424M = 0.076M. Now we can solve for V1 for acetic acid and sodium
1M(Vacetic acid) = 0.424M(0.1L), V = 0.0424L = 42.4 ml acetic acid
1M(Vacetate) = 0.076M(0.1L) = 0.0076L = 7.6ml sodium acetate
Now, bring to volume with 50ml H2O.
The second method is to simply dissolve the amount of buffer needed to make a
given volume and concentration in water, and then titrate the solution with a strong acid
or base until the desired pH is achieved. Then, bring the buffer solution up to the proper
volume. This method works for many buffers, however there are some buffers whose pH
changes with changes in concentration, so for these buffers this method is of no use.
There are many other considerations when choosing the proper buffer for a
particular use. Some buffers are sensitive to temperature. For example Tris buffer
becomes more acidic at higher temperatures. Some buffers may be inappropriate because
they interact with other components in the reaction or assay system and adversely affect
the experiment. An important example of this in the biotechnology lab is the use of Tris-
borate-EDTA (TBE) buffer versus Tris-acetate-EDTA (TAE) buffer. TBE buffer is
commonly used for separating DNA fragments by gel electrophoresis because it gives
superior resolution of the DNA fragments. However, if the goal is to extract the
separated DNA fragments from the gel, TAE is the buffer of preference. This is because
the borate interacts with the gel matrix in such a way that it is difficult to recover the
DNA from the gel (I know, because I have made his mistake!)
1. Using the Henderson Hasselbach equation to prepare a buffer solution.
2. Making a buffer solution by the titration method.
3. More pH meter experience.
1. Calculate how to make 100mls of your assigned buffer using 1M Acetic acid and
1M sodium acetate.
2. Prepare your buffer, and measure the pH. Make sure you write out your procedure
step by step in your notebook.
3. Prepare 100ml of 1M Tris-HCl pH 8 by the titration method. Write out your
procedure, and record how much 1M HCl it takes to pH your buffer. Measure the
pH again after you bring the solution up to volume.
4. How many drops of 0.1M HCl does it take to lower the pH of 50ml water 1 pH
unit? How many drops of 0.1M HCl does it take to lower the pH of 50ml 0.1M
Tris pH 8 1 pH unit?
For next week, please read Chapter 19 pp. 371-384, and Chapter 20