A former colleague and friend of mine worked in technical support, taking calls from scientists. Most of the calls came from life science researchers frustrated by failed experiments. My friend would listen for a bit, then almost always exclaim in very charming but booming Russian-accented English, “No! You must add buffer to the experiment! It is very important that you add the buffer!”
Why was my colleague so frustrated? Why were these scientists having such trouble with their experiments? Clearly, they needed to know more about buffers (or just when to use one). So, here’s a quick lesson in how buffers work.
It’s good to share—a proton
A buffer is rather simple by definition: it’s a solution that contains an acid matched with an equal-strength, or conjugate, base. Remember that acids donate protons, and bases accept protons. The stronger the acid, the more readily it will dissociate into protons and anions (in water, where most biological reactions take place). Stronger bases are more likely to dissociate into cations, or proton acceptors. The strength of an acid or base is measured by its dissociation constant, or pKa. The lower the pKa, the more acidic the molecule.
A good way to look at this dissociation is in this figure:
H2O –> OH– + H+
This of course is water; OH– is a strong base, and H+ is a strong acid (the strongest, being a proton). A buffering action occurs here because water can dissociate into a strong acid and a strong base, but balance each other out. If you were to add more H+ to this solution, more H2O would dissociate to generate a matching amount of OH–. The more H+ you add, the more dissociation of water occurs, moving the equation to the right. Thus, the water to the left is adding more OH– to the right to match the amount of strong acid you’ve just added. A buffer, then, does not keep the reaction solidly at a given pH, but does prevent wild swings in the acid-base balance.
What about us?
Biological systems are notable for having peak activity in a very pH narrow range (at a pH of about 7); thus, it is no coincidence that biological acids and bases are weak. But even a pH that is only slightly lower or higher than 7.0 can be very important when considering a buffer (remember, pH and buffer strength are calculated on a logarithmic scale). Egg whites and seawater usually exist at a pH approaching 8.0; blood, sweat and tears are closer to 7.3; milk is about 7.0. Gastric juices, however, heavy in HCl, have a pH just below 2.0.
Too much of a good thing?
It’s not only important to add buffers – it is just as important to know how much to add. To achieve that, you need to know your conjugate acid or base’s dissociation constant (pKa). You can calculate this using the following equation:
Ka = [H+][A–]/[HA].
Then, take the negative logarithm of your Ka to get pKa, the most commonly used dissociation constant figure.
This number allows you to match the pH of your suspension/solution to a buffer with the right pKa (you want them to be very close). If your pH is close to the pKa, then adding extra acid or base will do little to change your pH. For example, if you are working with a molecule that normally exists in blood (which has a pH of about 7.2), you’ll want to add a buffer solution like carbonic acid H2CO3 (and its conjugate base, bicarbonate), because it has a pKa that’s very close: 6.37.
Finally, you can calculate the relationship between pH and pKa in your solution using the famous Henderson-Hasselbalch Equation:
pH = pKa + log ([A–]/[HA])
Using the right buffer could lead to more successful experiments…and at least it will keep my friend from screaming!
Mathews, Van Holde, and Ahern. (2000). Biochemistry. Addison Wesley: San Francisco.
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