Buffer helps to maintain constant pH. Buffer solutions are widely used in microfluidic experiments. The main functions of a buffer are (i) to set the solution pH, and (ii) to stabilize the solution chemistry by maintaining pH and conductivity during and between experiments. Persat and et al. 2009, present a list of common buffers with their molecular formulas, pKa’s, temperature dependences, valences, electrophoretic mobilities, and other relevant properties.
In microchip electrokinetic experiments, pH and ionic strength affect most major aspects of system performance including separation efficiency and migration times, binding constants, Joule heating, biomolecular adsorption characteristics, power requirements, and overall reproducibility of the assay. For example, a pH change of 2 units and 100 fold factor in ionic strength can each change zeta potential (and electroosmotic mobility) by approximately two folds (e.g. for the case of silica between pH 6 and 8, and 1 mMto 100 mM ionic strength).
Which buffer?
The first step is the selection of pH, which is typically entirely driven by the application. For example, Tris hydrochloride buffer at pH = 8.0 is excellent for yield and specificity in polymerase chain reaction; however, an electroosmotic pump device may benefit from the 9.2 pH of a borate buffer, despite the typical problems with borate in biochemical studies.
Other concerns include temperature dependence, solubility, availability, cost, cross-reactions, etc. For example, the organic propionic acid has low solubility but its pKa is less sensitive to temperature changes than, say, Tris.
The second step is determining ionic strength of the electrolyte which at least in part determines Debye length, zeta potential, surface charge density, and conductivity among other key parameters. Designing for both specified pH and ionic strength typically requires computational tools. The most straightforward design approach is to fix weak electrolyte identity (given pH needs), then simultaneously and proportionally vary both weak electrolyte total concentration and titrant concentration (although this is just approximate).
For example, the online tool by
Beynon and Easter allows for specification of both ionic strength and pH given a weak electrolyte. The results can then be checked for buffer capacity, temperature dependence, etc. A second way of setting ionic strength is to add a socalled ‘‘neutral salt’’ such as KCl or NaCl. This allows for quick hand calculations, but such are again only approximate since ionic strength affects pKa . Also, the resulting buffer has less buffering capacity than, say, a weak electrolyte buffer and titrated with a strong electrolyte with equal ionic strength. Specifying both conductivity and pH is even more complex and typically always requires iteration since ionic strength affects both pKa and electrophoretic mobility.
Important advice in buffer preparation:
Prepare your own buffers. Many chemical suppliers offer premade buffer stock solutions such as Tris acetate EDTA (TAE) for agarose gel electrophoresis, Tris borate EDTA for polyacrylamide and agarose gel electrophoresis, phosphate buffered saline for cell culture, Tris hydrochloride for polymerase chain reaction, Tris glycine for SDS page, etc.
We do not advise using premade buffers for quantitative microchip electrokinetics work for reasons including: (i) premade buffers often contain additives and preservatives (e.g., EDTA, SDS, salts); (ii) the experimentalist does not have full control of pH and ionic strength; (iii) the manufacturer does not always provide exact content; (iv) the name of the buffer can be misleading (e.g., typical Tris EDTA buffer (TE) contains chloride ions); and (v) these buffers are not always cheap.
We advise preparation of buffers by hand starting from solid crystals or pure liquid electrolytes of (at least) reagent grade.
We advise preparation of a large amount of buffer stock solution for better reproducibility. Overly frequent preparation of fresh buffer may yield variations in pH. For example, causes include the temperature dependence of pH (despite temperature correction of some pH meters), and inaccuracy of weighing scales. However, this should be weighed against other practical considerations such as the possibility of culturing bacteria in the buffer, time kept in refrigeration, etc. For example, a Tris HCl buffer at pH 8.2 can likely be kept out of a refrigerator for months without bacterial growth; but a pH 7.3 phosphate buffer requires refrigeration and ultimately replacement.
To ensure that quoted buffer concentrations are correct, aliquots of stock solutions can be titrated with a strong acid or base to confirm that the buffer behaves as predicted. Accuracy in buffer pH also requires proper use of the pH meter. A pH meter is not trivial to use, or to keep calibrated over many measurements. Before each set of titrations (or final verification of predicted pH), the pH meter must be calibrated with fresh standard solutions.
We also advise titration of a large volume of weak electrolyte with a high concentration titrant. The initial volume of weak electrolyte should be on the same order of magnitude as the final volume (e.g., start with 50 mL of Tris for a final volume of 100 mL); this for convenience of measurement and accuracy. Once the pH reaches the correct value, addition of solvent to reach the final volume may slightly modify the pH as ionic strength decreases. The pH of the stock solution is, of course, the pH measured at the end of this process, and not the targeted pH value. Avoid the practice of ‘‘backpedaling’’ the titration pH after an overshoot, to keep tight control of ionic strength.
If buffer chemistry is solely determined based on calculations (and the buffer is not actually titrated), then the final measured pH should be reported (and ideally the predicted pH). Roughly speaking, if you care about pH, you should titrate your buffer. If you want to fix pH and ionic strength independently (e.g., and care about individual mobility values) you can validate your buffer calculations (with experiments), use calculations to specify buffer contents; and then verify and report the resulting measured pH. Detailed description of buffer preparation and titration procedures are available in several textbooks.29,42 Also, see the next section for suggestions on how to report buffers.
How to report buffers?
Common incorrect reporting includes not reporting pH, not reporting buffer concentration, not reporting titrant, and reporting a ‘‘standard’’ chemistry which leaves significant ambiguity. Common incorrect assumptions include assuming that ‘‘pure’’ deionized water has a pH 7 and/or a Debye length of 1 mm, buffers used well beyond their buffering range (e.g., pH 1 + units from pKa), ideal, ‘‘pure’’ water as an electrolyte or buffer and (low) buffer concentrations which are clearly on the order of expected bicarbonate and/or carbonate ion concentrations.
For a titrated buffer: Report weak electrolyte concentration, titrant, and pH. An example would be ‘‘100 mM Tris titrated with hydrochloric acid to pH 8.5.’’
For a buffer prepared by quantifying components: Report weak electrolyte concentration, titrant concentration, and the final measured pH. An example: ‘‘Buffer was 100 mM Tris and 100 mM HEPES with a measured pH 7.8.’’
We also suggest researchers consider reporting several useful items including temperature at which pH was measured, conductivity measurement, specific manufacturer and model of pH meter, and predicted ionic strength along with associated assumptions. If the stock solution is to be diluted significantly (i.e., 10-fold), it is useful to report pH after dilution.
Lastly, we note that some typical buffers have a standard composition. Examples include Tris acetate EDTA (TAE) or Tris borate EDTA (TBE), phosphate buffered saline (PBS). The concentrations of these is by convention described by a dilution factor from a standard concentration (e.g., 1 TAE or 0.5 TBE). We believe this is an acceptable report of a buffer, provided the buffer is well known. However, if in doubt, they should be reported fully. For example, so-called ‘‘TES’’ (Tris EDTA sodium) buffer is not ‘‘obvious,’’ as it is easily confused with the weak base TES (i.e., 3-{[tris(hydroxymethyl)methyl] amino}-ethanesulfonic acid).
