Buffers
Most biological systems will function only within a quite narrow range of
conditions, and their activity can vary widely within that range. The acidity, or free
proton concentration, of the environment is an important parameter. To prevent the
proton concentration of a solution from changing, a buffer system resists changes in pH
upon addition of an acid or base. That is, adding a large amount of hydrogen ions or
hydroxide ions to a buffer produces only a minor change in its hydrogen ion
concentration because the components of the buffer react with the added H+ or OH- and
remove them from the solution.
The most commonly used buffers are solutions of a weak acid and its salt or a weak
base and its salt. Consider first a weak acid and its salt, e.g., a mixture of acetic acid and
sodium acetate. Since acetic acid is weak, it is largely unionized, while sodium acetate
ionizes completely to produce acetate ions; thus, the active components of the buffer are
actually acetic acid and acetate ions. If a strong base such as NaOH is added, acetic acid
reacts with the hydroxide ions and replaces them with acetate ions. If a strong acid such
as HCl is added, acetate ions react with the hydronium ions and replaces them with acetic
acid (acetic acid is a much weaker acid than H3O+).
The role of the weak acid (acetic acid in this example) and its conjugate base
(acetate ion in this example) is clear: the weak acid is a proton donor, which neutralizes
added OH- ions, and the acetate ion is a proton acceptor, which neutralizes added H3O+.
Both OH- and H3O+ are converted to water. The pH therefore tends to remain unaltered.
Furthermore, in order for the buffer to resist increases and decreases in pH equally well,
the molar concentration of the weak acid HA and its conjugate base A- should be equal.
The optimal pH of the buffer (i.e., the pH at which the buffer best resists changes in pH
in either direction) is thus the pH of this equimolar mixture, which is in fact the pKa of
the weak acid.
Designing a buffer
The desired buffer pH determines what compounds are used to make buffer. To
make a buffer with acidic pH, (pH less than 7), the solution is made using a weak acid
and a soluble salt of its anion. Typically, the range for an acid buffer is centered around
the value of the dissociation constant for the weak acid. The pH is adjusted by controlling
the ratio of the moles of weak acid and the moles of soluble salt. To make a buffer
solution with basic pH, (pH is more than 7), the solution is typically made using a weak
base and a soluble salt of its cation.
To make a buffer you must first pick a compound whose pKa is close to the
pH you want for the solution, and then decide what the buffer concentration should be
depending on the desired ionic strength and the buffering capacity required. If the pH is
expected to decrease during the experiment, choose a buffer with a pKa slightly below the
working pH. Conversely, if the pH is expected to increase during the experiment, select a
buffer with a pKa slightly above the working pH
After choosing a buffer, the next step is to decide its concentration. The buffer
concentration must be sufficient to maintain the pH within acceptable limits with the
changes in proton concentration expected to occur. For biological systems, the total
buffer concentration ([Buffer] total) is within a range from 1 mM to 200 mM.
Practical ways to make a buffer
Using the buffer pKa - calculate the amounts (in moles) of acid/salt or base/salt
present in the buffer at the desired pH. If both forms (i.e., the acid and the salt) are
available, convert the amount required from moles to grams, using the molecular weight
of that component, and the weigh out the correct amounts of both forms. If only one
form is available, you can prepare the buffer by adding all of the buffer as one form, and
then adding acid or base to convert some of the added buffer to the other form. Decide
what the total concentration of buffer will be in the solution, and convert the
concentration to amount (in moles) using the volume of solution, and then to grams,
using the molecular weight of the buffer form available. Then calculate the amounts (in
moles) of each form that will be present in the final solution, using the buffer pKa and the
desired pH. Then calculate how much strong acid or base must be added to convert
enough of the buffer form added to the other form, to give the correct amounts of each
form at the pH of the final solution. Dissolve the buffer and strong acid or base in
slightly less water than is required for the final solution volume. Check the pH and
correct if necessary. Add water to the final volume.
The Two Solution Method - make separate solutions of the acid form and base form
of the buffer, both solutions having the same buffer concentration (and ionic strength, if
required) as the concentration of total buffer in the final solution. To obtain the desired
pH, add one solution to the other while monitoring the pH with a pH meter.
Acid or base titration method - prepare a buffer composed of an acid and its salt by
dissolving the acid form of the buffer in about ~70% of the water required for the final
solution volume. Adjust the pH using a strong base, such as NaOH. When preparing a
buffer composed of a base and its salt, start with the base form and adjust the pH with
strong acid, such as HCl. After the pH is correct, dilute to just under the final solution
volume.
Examples:
1. Preparation of a 0.1M Phosphate buffer, pH 7.4, 1 Liter:
The Henderson-Hasselbach equation will be used to determine the amount in
grams
to be used in sodium phosphate monobasic and sodium phosphate dibasic to make the
phosphate buffer (The pKa of H2PO4- = 7.2).
pH = pKa + log
[A-]
--------[HA]
[Proton acceptor]
--------------------[Proton donor]
To prepare the 0.1M phosphate buffer pH 7.4, first weigh out 8.62 grams of sodium
phosphate dibasic and 5.42 grams of sodium phosphate monobasic on the balance and
add 1L of distilled water. Measure the pH. If the pH is below 7.4, adjust it by adding
several drops of NaOH. If the pH is above 8.3, adjust it with several drops of HCl.
2. Preparation of a 0.1M Tris-HCl buffer, pH 8.0, 1 Liter, and 0.5M NaCl
To prepare the 0.1M Tris-HCl buffer pH 8.0, first weigh out 12.1g of Tris. Place it in
a 1L beaker and add distilled water to 60% and adjust the pH to 8.0 by HCl. Then
complete the volume to 1L in a volumetric flask.

How many grams of acetic acid (CH3CO2H) and sodium acetate are needed to
make up one liter of a 50 mM acetate buffer with a pH of 5.0?

For this buffer, which is greater, the buffer capacity with respect to addition of
acid, or the buffer capacity with respect of addition of base?

After making the calculations, you realize that there is no available sodium
acetate. However, you do have a 1 M NaOH solution. How is the buffer made?