In its simplest form, a buffered solution contains a mixture of a weak acid and its conjugate base.

The position of acid/base equilibrium is represented by the acid dissociation constant, Ka. This number is large if the acid is stronger and equilibrium tends toward dissociation. It is small for an equilibrium that tends toward proton capture. Buffers used in life science tend to range from 10-4 to 10-10 in their Ka values.

Ka is usually expressed as its negative logarithm, pKa:

A buffer with a Ka of 10-2 has a pKa of 2, favoring dissociation. A buffer with a Ka of 10-12 has a pKa of 12, favoring proton capture.

After a few math operations the expression for the dissociation quotient becomes a very useful form, the Henderson-Hasselbalch equation:

This equation is central to buffer design. The formula states that the pH of a buffered solution will differ from the buffer's pKa by an amount which is determined by the ratio of the base to the acid forms in solution. If these concentrations are equal then the pH = pKa. If the concentration of the base form is greater than the acid then the pH > pKa. If the concentration of acid form is greater than the base then pH < pKa.

A buffer maintains a nearly constant pH by absorbing protons released by other sources in solution or releasing protons if another species is depleting them. For example, a small amount of strong acid introduced to pure water causes the pH to plummet five or six points. However, if that same amount of strong acid were introduced to a concentrated buffer solution, it merely causes the change of some of the buffer's weak base form to the weak acid form. By the Henderson Hasselbalch equation, to drop the pH of the solution a full point below the pKa, enough base would have to be consumed for the concentration of the acid form to become ten times greater than the base. If the buffer is sufficiently concentrated, the small amount of strong acid in this example would barely alter the pH. In electrophoresis, the buffer is able to maintain a relatively constant pH as long as either the acid or the base do not become exhausted.

The Henderson-Hasselbalch equation gives information in addition to how the pH is a function of the relative predominance of acid base forms of a buffer. For certain instances, it is useful to think of this equation the other way around. In predicting the state of ionization of minor solution components, for example, it can be useful to think of the relative predominance of various forms as a function of an externally determined pH. An instructive example might be a solution whose pH is maintained by a concentrated buffer. The degree of ionization of a minor component in this solution, such as a protein, would depend on the pH of its environment, determined by the buffer. Suppose an ionizable group upon a protein has a pKa of 11, an amine group, for example, which might be either the nonionized amine form (base) or the ammonium ion form (acid). If the solution is buffered to a stable pH of 8, the Henderson-Hasselbalch equation tells you that the ammonium ion form will predominate at a ratio of 1000 to 1. In native protein electrophoresis, this concept is essential to understanding how buffers control the state of ionization of sample molecules.

The buffer system in electrophoresis controls the pH of the gel, preventing damage to sample molecules and, in certain cases, controlling the ionization state of the molecules. A second, though no less significant function derives from the fact that the vast majority of current flowing through the electrophoresis gel is carried by the buffer ions. For a homogeneous system like denaturing PAGE electrophoresis of DNA, in which the type and concentration of buffers in the tank and gel are the same, the buffer prevents wide swings in pH and controls the conductivity of the gel. For the native electrophoresis of proteins, the buffer pH has the added function of controlling the state of ionization of the samples. In this second case, even slight changes in pH can result in large effects on the relative mobility of sample components. In a multiphasic system, such as SDS-PAGE electrophoresis of proteins, where buffers in the tank and gel are different, the considerations of buffer design can take on an even greater level of complexity.

In all cases, the ionic strength of the buffer in the gel must be sufficient to keep the sample in solution and to provide sufficient buffering capacity. Higher concentrations of gel buffer will also slow the diffusion of samples, and result in sharper bands. The benefits of higher buffer concentrations must, however, be balanced against the fact that the more concentrated the buffer in the gel, the higher the electrical conductivity. With higher concentration, at a given voltage, the current will be greater and more heat will be generated. Because of the problems caused by excessive heating, high buffer concentrations must be accompanied by a low voltage gradient.

Several factors to consider when choosing a buffer include:

1) pKa value - A buffer should be chosen with a pKa that is very close to the desired pH, preferably within a half point. The buffer will have the greatest capacity both to absorb or release protons with the acid and the base form well represented in solution. It should be noted that pKa is not constant for all conditions but is a function of the total ionic strength and the temperature, so the stoichiometry should be modeled after actual running conditions. Amines are particularly susceptible to changes in pKa with temperature, because, with amines, there is no net increase in ions in solution with the dissociation of ammonium species, so there is little inherent entropy change due to changes in the ordering of water molecules in the solution. This leads to an increase in the significance of the entropy change due to heat flow (the ratio of enthalpy change to temperature) in determining the position of equilibrium.

Usually, in native protein electrophoresis, basic proteins are best separated at acid pH. However, the vast majority of proteins have isoelectric points below 7.5 and are best separated in slightly alkaline conditions, pH 8-9. This pH range has also proven efficacious for most forms of DNA electrophoresis. Thus, buffers with a pKa in the range of 7-9 are best suited for most electrophoretic applications.

2) Formal charges of buffer species - Generally, buffers which form ions of high charge magnitude (+2, +3, -3, etc.) are more difficult candidates with which to work. This type of buffer yields high ionic strength without providing high buffering capacity. At relatively low concentrations, the gel conducts too much current. Furthermore, with ions moving quickly through the gel, the buffer may become depleted. One of the reasons Tris-borate is a popular buffer for electrophoresis is that both Tris base and borate are uncharged part of the time at the desired pH, which reduces their electrophoretic mobility. Reduced mobility of the buffer ions allows for high concentrations of the buffer solution to be employed with the consequent benefits to buffering capacity and sample stability without producing unacceptably high conductivity.

3) Molecular size - In addition to its low charge, Tris base moves slowly in electrophoresis because of its relatively large molecular size. Having a low charge to mass ratio, Tris moves much more slowly than small ions such as chloride or phosphate. Buffer ions are not sieved by the matrix, so their migration rates are determined solely by their charge to mass ratio.

These three factors is by no means exhaustive. Other factors to consider when choosing a buffer would include toxicity, solubility, UV absorption and the possibility of interaction with other species present in the solution.

Among the most commonly used electrophoresis buffers are those shown in the table below:

NEXT TOPIC: Homogeneous Buffer Systems