Buffer Preparation:

Choices you have when you go to prepare a buffer:

  1. How much volume of buffer do you wish to prepare?
  2. What pH do you want?
  3. Usually dictated by the system you are studying.
  4. What is the "buffer concentration" (by definition, buffer concentration is [weak acid]+[weak base]).?
  5. Increasing the buffer concentration increases the buffer capacity (generally a good thing), but also increases the cost (a bad thing) and the ionic strength of the buffer.
  6. "Typical" buffer concentrations are 0.02 to 0.10 M.
  7. What weak acid/conjugate base do you wish to use?
  8. For a reasonable buffer capacity, choose a weak acid with pKa within 1 of the pH of the buffer you are making.
  9. Other factors to consider are economy (cheap is better), toxicity (bad) and any specific interations of the acid or its conjugate base with the chemical system being studied (usually undesirable because they complicate things).
  10. What ionic strength do you want? (Note: this is somewhat of an advanced topic, and is sometimes ignored even by practicing scientists. It is included here only for completeness.)
  11. Ionic strength can be critical for the stability of biomolecules. Proteins, for instance, often unfold when the ionic strength gets above about 0.5 M and are usually most stable at the ionic stength found in cells--about 0.15 M.
  12. If all ions in the buffer are monocations or monoanions, then the ionic strength is just the concentration of cations (which is also equal to the concentration of anions).
  13. Buffers employing a dianion (such as a H2PO4- /HPO42- buffer) have higher ionic strength for a given buffer concentration. The Debye-Huckel equation can be used to calculated ionic strength for these buffers.
  14. The ionic strength of a buffer can be increased by adding sodium chloride or other neutral salt.

Preparing Buffers

Approach # 1 (by weighing out both the acid and its conjugate base):

  1. Using pen and paper (ideally your lab notebook) and a calculator, solve the following two equations for the two unknowns, [acid] and [base], which are the desired concentrations of weak acid and its conjugate weak base, respectively:

PH = pKa + log ([base]/ [acid]) (Henderson-Hasselbach equation)

Buffer concentration = [acid] + [base]

  1. Convert the [acid] to moles of acid using the desired final volume of the buffer, and then convert this to grams using the formula weight of the solid form of the acid. The formula weight includes any ions (sodium, for instance) or water of hydration which co-crystallize with the weak acid. Weigh out the weak acid.
  2. Repeat the above calculation to get grams of the weak base. Weigh out the weak base.
  3. Dissolve the weak acid and weak base together in the desired volume of deionized water.

Approach # 2a (practical approach, starting with the weak acid):

  1. Weigh out the number of moles of weak acid equal to the (buffer concentration) x (desired final volume).
  2. In a beaker or large bottle, add deionized water to the weak acid until you reach about 90% of the desired final volume.
  3. While stirring with a stir bar and measuring the pH with a calibrated meter, add sodium hydroxide dropwise until the pH reaches the desired value. The precise molarity of the sodium hydroxide is not important, but 6 M is a convenient concentration I use in my lab.
  4. Add deionized water to the desired volume.

Approach # 2b (practical approach, starting with the weak base):

This method is preferred over #2a when the buffer involves a neutral weak base and a cationic conjugate acid (which applies to amines and imines such as ammonia, TRIS and imidazole).

  1. Weigh out the number of moles of weak base equal to the (buffer concentration) x (desired final volume).
  2. In a beaker or large bottle, add deionized water to the weak base until you reach about 90% of the desired final volume.
  3. While stirring with a stir bar and measuring the pH with a calibrated meter, add hydrochloric acid dropwise until the pH reaches the desired value. The precise molarity of the hydrochloric acid is not important, but 6 M is a convenient concentration I use in my lab.
  4. Add deionized water to the desired volume.

Describing Buffers

In the biochemical literature, buffers are described by designations best illustrated by example. For instance, "0.10 M, pH 5.0 sodium acetate buffer" means that the "buffer concentration" ([acid]+[base]) is 0.10 M and that the buffer is either prepared from acetic acid and sodium acetate by approach #1, or by adding sodium hydroxide to acetic acid using approach #2a. Approach #2b (adding HCl to sodium acetate) probably should be avoided, since it introduces chloride ions, which are not indicated by the buffer designation.

A web site for creating buffer recipes

As part of JaMBW, the Java based Molecular Biologist's Workbench, Rob Beynon has produced a very useful web site for creating buffer recipes using any of the commonly used acid/base pairs. In addition to allowing the user to make the five choices mentioned above, it also allows for the possibility that the buffer will be used at a different temperature than the temperature at which it is prepared. The address is

This page maintained by , Last updated 3/9/06 .