In this installment of “Advanced Brewing,” we’ll take a look at buffers and what advanced homebrewers should know about them. I will assume that you have knowledge of the basics of acids, bases and pH. (If not, see Chris Bible’s article, “The Principles of pH,” in the September 2007 issue of BYO.) For serious brewers, understanding pH can be an avenue to improving their beers. However, without an understanding of buffers, and buffering capacity, it can also cause them to venture down some blind alleys. This article will get fairly chem-nerdy, but I’ll summarize the key points as we go along and give a summary of the practical information at the end.
Armed with our knowledge of acids and pH, let’s conduct a quick thought experiment. Let’s imagine we have a large beaker of pure water — water with no minerals or anything else dissolved in it. Let’s also say we have a dropper bottle full of a strong acid (hydrochloric acid, HCl, for example) and a pH meter. Now, let’s imagine we take the pH of the pure water. It should be 7 if the meter is calibrated and we take the reading at room temperature. Now, let’s say we keep the pH electrode in the water and add drops of acid, checking on the pH after each drop. What you would see if you did this experiment is that the pH would lower each time a drop of acid was added.
Every college chemistry major has done (or at least should have done) this experiment. I remember sitting in lab with my beaker of deionized water, magnetic stir bar spinning, adding drop after drop of acid and making a graph of the nice smooth curve it generated.
Now let’s imagine a second experiment that adds a slight twist to the previous experiment. Let’s say we repeat the experiment, except we add a buffer to the water. What’s a buffer? The results of this experiment — that I did for real in introductory chemistry “a few” years ago — will show you. Initially, as the drops of acid are added, the pH drops at the same rate it did in the first experiment. However, at some point, the rate of pH change rapidly slows to a halt. (What point that is depends on what the buffer was.) For awhile, adding acid does not change the pH. Then, suddenly, the pH will begin to drop again, at the same rate it did in the pure water solution.
As the second experiment shows, a buffer is a substance that resists pH change within a limited pH range. It is usually a combination of a weak acid and its conjugate base (or sometimes weak base and its conjugate acid). The conjugate base is usually supplied as the salt of the acid. The salt of any acid is the acid with one or more of the acidic hydrogens replaced with something else. For example, phosphoric acid (H3PO4) has three sodium salts, monosodium phosphate (NaH2PO4), disodium phosphate (Na2HPO4) and trisodium phosphate (Na3PO4).
You can write the general equation of a buffer system as:
HX + H2O <------> H2O+ + X-
On the left side of this equation, HX is the acid and H2O is water. In the formula for the acid, H represents a hydrogen ion and X is any ion that can combine with a hydrogen ion to form a weak acid. On the right hand side of the equation, X- is the conjugate base of the acid and H3O+ is a hydronium ion. (When the acid breaks up into X- and H+, the H+ instantly interacts with a water molecule (H2O) and becomes H3O+.) For example, if you added acetic acid (CH3COOH) and sodium acetate (CH3COONa), you would have a buffer system as described above. The “X” would be a CH3COO– group. (The sodium (Na) does not play a role in the buffering, so it is ignored in the equation.)
The arrows in the equation represent an equilibrium. HX molecules would constantly be splitting apart at some rate, in the presence of water, to form H3O+ and X-ions. Likewise, these two ions would constantly be combining to form water and the weak acid (H2O and HX) at some other (usually much lesser) rate.
Recall that pH is the negative log of the hydronium ion content. So whatever the concentration of the hydronium ion is, pH = -log[H3O+]. So, the buffer solution would exist at a certain pH. This can be calculated as:
pH = pKa + log [X-]/[HX]
where pKa is acid dissociation constant of the weak acid and [X-] and [HX] are concentrations of X- and HX, respectively. dependent on the concentration of acid (HX) and concentration of the conjugate base (X-).
If you don’t know, don’t worry about what an acid dissociation constant is. The only thing you need to see in the above equation is this — by adding different amounts of weak acid and the base/salt to a buffer solution, you can change the pH and make a buffer solution for any pH value you want. (More on this later.)
Now, let’s get to the interesting part. Imagine you add some strong acid to this buffered solution. In solution, the acid would dissociate and contribute H3O+ ions — H+, combined with water — and its conjugate base to the solution. (The conjugate base for the strong acid would differ from the conjugate base for the weak acid, the “X” above.) The “extra” H3O+ ions would combine with X- ions and form HX and H2O. The newly-formed HX would join the pool of HX in solution and break into H3O+ and X- at the same rate as before. The upshot of all these chemical interactions is that, due to the equilibrium between the reactions, the buffered solution would essentially “absorb” the acid and remain at nearly the same pH.
Of course, you could keep adding acid until the buffer is overwhelmed. Once this occurs, the pH would again change more quickly with the addition of more acid, almost as if you were adding the acid to a pure water solution.
Before we go on, let’s summarize everything up to this point. Buffers resist pH change, but this buffering capacity only exists within a certain pH range. I’ll also add one more comment that should be obvious, but is worth saying explicitly — the more buffering agent in solution, the more acid it takes to overcome its buffer capacity.
Two important buffers
In brewing liquor, wort and beer, there are a variety of different acids and ions that can form buffer pairs. If you have water that is high in carbonates, you can have a buffering solution composed of carbonic acid (H2CO3) and bicarbonate (HCO3-). Likewise, phosphoric acid (H3PO4) and various phosphates (including H2PO4-, NaHPO4-, Na2PO4-, etc.) can form buffer pairs. Both carbonates and phosphates exist in different forms at different pH values, and can thus form different buffering systems (at different pH values).
In water that is high in carbonates, you might need to add quite a bit of acid if you were adjusting the pH of your brewing liquor. But, the phosphates are a bit more interesting. Remember the equation for calculating the pH of a buffer solution? Remember how I said that, by varying the amount of the ingredients, you could make a buffer solution for any pH. Well, by mixing various phosphate compounds — which are safe for human consumption — you can make a buffer to hold your mash at any value you wish.
Since phosphoric acid has three hydrogens it can give up, the calculations are a little more involved than the equation on the previous page (which are based on a simple acid with a single donatable hydrogen). However, there are web-based calculators that will do the work for you. Search for “phosphate buffer” and “calculator” or “calculation” or “preparation” and you will find all sorts of help. For example, if you wished to make 10.0 mM solution with a pH of 5.2, you would mix 0.135% monosodium phosphate (monohydrate) and 0.0057% disodium phosphate (heptahydrate) into your buffer solution.
(The unit mM stands for millimolar, a measure of concentration. The hydrates mentioned are just the form of the chemical. Some “dry” powdered chemicals have water molecules associated with them. For example, disodium phosphate (heptahydrate) would have seven water molecules associated with each molecule of disodium phosphate.)
So let’s give a practical example. Let’s say you were planning to brew 5 gallons (19 L) of pale ale and planned to mash in with 4.0 gallons of water (15 L). To attempt to buffer your mash at a pH of 5.2, you wish to make a 10.0 mM solution for your brewing liquor. Fifteen liters of water weighs 15 kilograms. So you would need [0.00135 X 15 kg = ] 20 g of monosodium phosphate (monohydrate) and [0.000057 X 15 kg = ] 0.86 g of disodium phosphate (heptahydrate) to make this buffered brewing liquor. (I used metric units for these calculations, as that made things much easier.)
Given the mineral content of your water, composition of your grist and thickness of your mash, you may need to make the solution stronger or weaker (but not so strong it affects the flavor of your beer.) For solutions stronger or weaker than 10.0 mM, you would need to add proportionally more or less of the phosphate ingredients. The commercial product 5.2 pH stabilizer, by Five Star Chemicals — which claims to be “a food-grade blend of phosphate buffers” — recommends a usage rate of 1 tbsp. per 5 gallon (19 L) batch. Obviously, the units have shifted here, but that’s in the ballpark of the amounts calculated above. If you try this, test the pH of your mash with a pH meter and, of course, see if you taste any difference in your finished beer. Also, when mixing buffer solutions, it is often necessary to fine tune the pH with a little acid or base. A little phosphoric acid should be used if your pH is just a little bit above the target.
Acids and their salts aren’t the only things with the ability to buffer. Next we’ll look at amino acids, biological molecules that are found in relatively high concentrations in wort.
One class of molecules found in wort are amino acids. Amino acids are the building blocks of proteins and come from the malt and the breakdown of proteins in the malt. An amino acid can be pictured as a central carbon atom with 4 “arms.” One arm is a hydrogen atom. Another arm varies, depending on the amino acid and is often symbolized with a “–R.” The R group can be as simple as a hydrogen atom, in the case of the amino acid glycine, or a much larger collection of atoms (as in the heterocyclic R group in tryptophan). The third arm is a carboxyl group, which consists of a carbon atom, two oxygen atoms and a hydrogen atom (in chemical short hand, –COOH). The final arm would be an amino group, consisting of a nitrogen atom and two hydrogens (–NH2).
Like many molecules, amino acids are either positively charged, neutral or negatively charged, depending on the pH of the solution they are dissolved in. In a high pH solution, the carboxyl group is prone to losing a hydrogen ion and leaving behind a –COO- group, leaving the amino acid with a net negative charge. In a low pH solution, the amino group is prone to picking up a hydrogen ion to form a –NH3+ group. This gives the molecule a net positive charge. At a certain pH, called the isoelectric point, the amino has both a –COO- and a –NH3+ group. The molecule has a net neutral charge, but positive and negative portions of the molecule. The isoelectric point varies for each different amino acid.
Amino acids and other molecules that have positively and negatively charged ends are called zwitterions. (The term comes from the German word “zwitter” which means hermaphrodite and is pronounced “tsvitter-ion.”) Just as a weak acid and its conjugate base will buffer a solution at its pKa, a zwitterion will buffer a solution at its isoelectric point (pI). If acid is added to a solution, the “minus end” will absorb some of the hydrogen ions. If a base is added, the “plus end” will absorb some of the negation ion.
The take-home message
So let’s combine what we’ve learned here with some other stuff we should know and see what the take-home message of this article should be.
Unless it is completely devoid of dissolved minerals, your brewing water is likely to have some things dissolved in it that will act as a buffer. These buffers will resist a change in pH within their given range. For this reason, it’s impossible to say “add X fl. oz. (mL) of acid to Y quarts (L) of water to change the pH by Z units.” Unless you know exactly what’s dissolved in your water, you don’t know what pH values it’s buffered at and how strong the buffer is. When adjusting the pH of water, it’s easier to just add acid (or base) and check the result with a pH meter. The concentration of acids, bases and buffers in tap water is tiny compared with the concentration of acids and buffers in wort. Amino acids are strong buffers and are present in wort in concentrations greater than any mineral dissolved in your brewing liquor. Thus, the pH of your brewing liquor is relatively unimportant. The pH of your brewing liquor does not determine your mash pH. In fact, almost all brewing waters will yield a mash pH in the low to mid 5 range, regardless of water pH. For this reason, adjusting your water pH is of little use (unless you know from experience that adjusting your water to a certain pH yields a given pH in the mash). Likewise, the pH of your sparge water does not determine the pH of the wort you collect. Even near the end of wort collection, the weak wort in your grain bed is more heavily buffered than your water.
You can mix up buffer solutions to a given pH. In brewing, phosphate buffers are used by some brewers to control mash pH. At the concentrations they are used, phosphate buffers are not harmful and won’t alter your beer’s flavor. Finding the minimum amount of buffer solution required may take a few trials.
The concepts of buffers and their effects on pH can be a little intimidating when you first encounter them. But, they do have some practical value (and you do get to say “zwitterion” a lot).
Chris Colby is Editor of Brew Your Own.