by A.J. deLange
The question of pH often arises in brewing circles — what is it, why is it important, and do we as brewers really need to worry about it? And if we do need to worry about pH, do we need a pH meter to measure it or will simple paper test strips do? And if we do need a pH meter, what kind should we get, how expensive are they, and how are they used, anyway? This two-part article will answer all these questions, and perhaps a few others.
This first article lays the basic foundation for understanding pH, beginning with a definition of pH as the concentration of hydrogen ions in a solution, discussing the question of where these ions come from in brewing, and introducing the important topic of buffering. The article then reviews the various steps in the brewing process in which pH is a factor, with discussion of when and why we make pH measurements in each of them.
The second and final installment, to appear in the next issue of BrewingTechniques, will look at the practical methods available for measuring pH, with emphasis on electronic pH meters. That article will describe their construction, give some guidelines on their use and maintenance, and offer some pointers on what to look for in buying one.
First, it is important to say that pH measurement is not something that all brewers necessarily need to worry about. pH control is often a fine-tuning operation that results in small changes in the resulting beer. Careful attention to pH may help you get a little more extraction or a bit more protein coagulation in the boil. These benefits are significant to commercial brewers, for whom pH is as important as temperature (for reasons we will explore), and to advanced home brewers who try to emulate commercial practices as far as they can with their limited equipment and resources. pH measurement is also often used as a guideline to reassure the brewer that the brew session is on track; once you know the normal pH values for your recipes and procedures, detecting odd-ball values during the course of a brew session can alert you to problems early on.
pH is a fairly sophisticated subject, and this article, particularly Part I, dives quite deep into technical waters. It assumes, for example, that readers will remember enough high school or college chemistry to be familiar with the basic properties of ions and to do arithmetic on numbers using scientific notation (that is, numbers expressed in the form 6.02 X 1023). Nothing beyond the introductory level of chemistry is required to understand this article, however, so give it a try. The application sections that follow the discussion of chemistry fundamentals are no doubt more familiar territory for brewers.
If you bog down, get that old yellowing textbook out of the attic for a little review. Key chemical terms and other significant terms are printed in italic type. Looking for them in your chemistry text would be a good way to start the review.
Beer and its antecedents water, mash, and wort are all aqueous solutions that contain a fairly large variety of ion types. Even the purest water, water to which nothing has been added, contains 6.02 X 1016 hydrogen ions and 6.02 X 1016 hydroxyl ions per liter. The pH of an aqueous solution is the negative logarithm (to the base 10) of the concentration of hydrogen ions in the solution expressed in moles per liter.
A mole is 6.02 X 1023 “things.” The “things” in the case of pH are hydrogen ions, but in other applications they could be atoms of some element or molecules of some compound. Thus, if the concentration of hydrogen ions is 0.01 mol/L, the pH is 2. If it is 0.002 the pH is 2.7. For pure water it is 7 (divide 6.02 X 1016 ions/L by 6.02 X 1023 ions/mol to get 1 X 10–7 mol/L). To reach this solution, follow the directions given in the following paragraph.
You can find the pH from the concentration by entering the concentration into a multifunction calculator (or a spreadsheet) and pressing the log key (or using the log function of the spreadsheet). Make sure to use the key that gives the log to the base 10, usually labeled “log,” and not the key that gives the log to the base e, which is usually labeled “In.” Reverse the sign of the result. (Because pH is really only meaningful when the concentration of hydrogen is less than 1, the pH values will be positive because logs of numbers less than 1 are negative. The smaller the number, the more negative the logarithm and, after sign reversal, the more acidic.)
Some readers may have already gathered that the hydrogen ion concentration of a solution is simply 10–pH. You can calculate concentration from pH by changing the sign of the pH and using the calculator key (sometimes called the antilog key) that raises 10 to that power. It is this aspect of pH that caused the German biochemist S.P.L. Sörensen to choose the abbreviation pH for this measure, which he introduced in 1909 (1). The letter p stands for the German word potenz, which means “power” in the mathematical sense.
In the early days, pH was measured by a hydrogen cell, which was an elaborate device containing a platinum wire in a tube immersed in the solution whose pH was to be determined. Hydrogen gas was bubbled into the solution around the wire, where some of it was oxidized (oxidation and reduction here refer to the transfer of electrons) to hydrogen ion. Some hydrogen ion in the solution was reduced to hydrogen gas. The relative amount of reduction/oxidation resulted in a net production or consumption of electrons, depending on the amount of hydrogen ions present in the solution. This caused an electrical potential (voltage) proportional to the log of the concentration of hydrogen ions in the solution.
DeClerck provided a fascinating description of the equipment and procedures for pH measurement using the hydrogen electrode (2). Today, pH measurement is still best done electrically, but the electrodes have improved considerably. The principals behind pH measurement, however, are unchanged (more on that in Part II).
Sources of Hydrogen Ions
Because this article focuses on the measurement of hydrogen ion concentration, it is reasonable to ask where the hydrogen ions come from in brewing. They come from three sources: the water itself, weak acids, and the salts of weak acids.
The chemistry of weak acids is reasonably involved and won’t be discussed much here, but it is important to define weak acids. Before we can define a weak acid, however, we must define an acid, which can be done in several ways.
For our purposes, the simplest definition (there are several) of an acid will do: a weak acid is a substance that gives up hydrogen ions in an aqueous solution.
What happens is that the chemical bond between the hydrogen atom and chlorine atom is broken, and the molecule separates into two parts. The chlorine, which has a tremendous affinity for electrons, keeps the hydrogen atom’s electron and so acquires a negative charge. The hydrogen atom, devoid of its electron, thus has a net positive charge. It is attracted to a nearby water molecule to form the hydronium ion, H3O+, which is simply a hydrogen ion riding on a water molecule. If we were being formal here we would be using the term hydronium ion instead of hydrogen ion throughout. HCl in water is thus an acid: hydrochloric acid.
Hydrogen chloride totally dissociates into ions when put into water — it is a strong acid. This means that if we dissolve 0.1 mmol (millimole) of HCl into 1 L of water, 100% will dissociate, giving 0.1 mmol of hydrogen ions and 0.1 mmol of chloride ions. Converting the millimoles to moles by dividing by 1,000 gives 0.0001 mol/L, or 1 X 10–4. The water itself had 1 X 10–7 to start, so the total is now 1.001 X 10–4, and the pH of such a solution is easily calculated (to sufficient accuracy) to be 4.
A weak acid is one that does not completely separate into its component ions. Carbonic acid, which is found in beer because of the carbonate salts in the water and dissolved carbon dioxide from air and yeast, is an example of a weak acid. The extent to which a weak acid dissociates depends on the dissociation constant. The dissociation constant is often reported as the pK, which is the negative of its logarithm, thus putting it on the same basis as pH, to which it is often compared.
If we dissolve 0.1 mmol carbon dioxide in 1 L of pure water, 0.1 mmol of carbonic acid is formed. Of this, only 6.2% or 0.0062 mmol dissociates into hydrogen ions and bicarbonate ions. Converting to moles (divide by 1,000) gives 6.2 X 10–6 mol/L hydrogen ion from the carbonic acid, which, added to the 1 X 10–7 already present from the water, gives 6.3 X 10–6 mol/L, for a pH of 5.2.
Each weak acid has one or more pKs (carbonic acid has 2, phosphoric acid has 3). Each pK is associated with the loss of one hydrogen ion. One very important fact that must be understood is that the pKs change with temperature. This means different proportions of the acids dissociate and different amounts of hydrogen ion are released to the solution at different temperatures. Thus, a solution’s pH will be a function of its temperature. It is also true that an electronic pH meter’s response to a given pH changes with temperature, and the two effects are often confused.
The most important inorganic acids for brewers are carbonic acid (and salts) in brewing water and phosphoric acid (and salts) from malt. Probably the most important organic acid is phytic acid from malt, but many others exist as well.
Buffers are mixrures of ions that can absorb (buffer) chemical stresses that would otherwise cause a change in the pH. They are of interest to brewers for two reasons. First, pH meters are calibrated using buffers prepared to hold pH 4 and pH 7. Second, mash acts as a buffer, as does liquor. Calibration buffers will be discussed in Part II, in the section on the use of pH meters. Mash buffering is discussed in this article, in the section on mash pH. For the present, it is sufficient to illustrate the important properties of buffers.
One common way to make a simple buffer is to find an acid with a pK close to the desired pH,* and then dissolve an equal number of moles of the acid and a salt of the acid in water. For example, if we put 15 mmol of sodium bicarbonate (salt) into 1 L of water and add 15 mmol of carbonic acid, the pH of the resulting solution would be 6.38, which is the first pK of carbonic acid.
Now let’s add a strong acid, such as hydrochloric, to our buffer until the pH reaches 4. Next add very small measured amounts of alkali (a strong base, such as lye) to the solution until the pH reaches, say, 4.2. Record the amount of alkali added. Then add more alkali until the solution reaches pH 4.4, and record the total amount. We can repeat this process, called titration, until a suitably high pH is reached, and then plot the data. They should fall along a curve close to the solid curve in Figure 1.
The dotted line in Figure 1 is the slope of the solid line at each value of pH, and its value can be read against the righthand vertical axis in units of the number of mmol required for a change of 1 unit of pH. At pH 4, small amounts of alkali will cause a substantial change in the pH. At pHs between 5 and 6, appreciably more is required, and, as the dotted curve makes clear, the amount of alkali per unit change of pH is a maximum at pH 6.38, which is the first pK for carbonic acid.
As we continue to add alkali, the pH will continue to increase but at a faster rate; that is, we see a bigger change per millimole of added alkali, or less alkali required to achieve a unit of change. The pH changes fastest (least alkali required per unit change) at pH 8.35, which is halfway between pK1 (6.38) and pK2 (10.32), the second pK for carbonic acid. The behavior near pK2 is similar to that near pK1; that is, we see buffering action again, and a lot of alkali required to change the pH. As brewers, we are not concerned with the action of this buffer near pK2 because it is out of the range of our brewing applications.
*pKs are published in tables in, for example, the Handbook of Chemistry and Physics. These tables have data on “standard buffers,” which are the basis for the strict definition of pH.
Thus far we have discussed adding alkali and moving along the titration curve to the right (towards higher pH). We can equally move back along the curve to the left by adding strong acid (such as hydrochloric acid), which neutralizes the alkali, in effect taking it out of the picture.
Let’s look at this pH-lowering process to illustrate the practical effect of buffering. Suppose we are at pH 7. We see from the solid line on Figure 1 that we have added about 25 mmol of alkali to get to this point. Now let’s add 16 mmol of strong acid. This would neutralize 16 mmol of the added alkali, and we would be back at the point where we had added only 9 mmol alkali (pH 6). Thus a 16 mmol acid or base addition is required to move this buffer back and forth between pH 6 and 7, the values of which span pK1.
Now suppose we had pure water at pH 7 and added 16 mmol of strong acid to it. The hydrogen ion concentration would jump from 1 X 10–7to 1.6 X 10-2, and the pH would be 1.8. Rather than the 1 pH unit change in the case of the buffer, the water has changed 5.2 pH units.
The amount of acid or base required to effect a particular pH change is called the buffering capacity of the buffer. The dotted line in Figure 1 is the buffering capacity of our example buffer. The shape of the buffering capacity curve is set by the pK or pKs of the acid involved. Its amplitudes are proportional to the concentration of the buffer acid and salt. If we doubled the amounts of carbonic acid and sodium bicarbonate used to make the buffer (we cannot actually do this because carbon dioxide gas would escape from the solution at low pH), the buffering capacity would be doubled at all values of pH. As we will see, the importance of buffer systems in brewing is based on the buffer system formed by the salts and acids in the malt and by the salts in the water. Adjusting mash pH is a matter of modifying the buffer system to the desired mash pH.
Why Brewers Are Concerned with pH
Many chemical and biochemical reactions take place during the making of beer. Many if not all of these reactions are sensitive to pH. For example, the coagulation of proteins in the boil and duting decoctions is influenced by pH. Also, many brewing reactions involve enzymes (proteins) that serve as catalysts; that is, they are not themselves consumed in the reaction, but have a profound effect on the rate at which the reaction takes place. These enzymes typically are able to catalyze reactions at acceptable rates only within certain regions of temperature and pH.
Every all-grain brewer is aware that one of the brewer’s most important duties is to control temperature to a value where a particular enzyme or enzymes can do a particular job. It is also important to see to it that the pH is within specified limits. Most small-scale brewers, however, are less likely to consider pH, if only because they have no means of accurately measuring it.
Luckily, in most cases the pH naturally hovers around the right levels because of the presence of natural buffers in the beer. It is only when things go wrong — for example when extraction efficiency is disappointing — that we may begin to wonder whether pH was involved. When this is indeed the case, the pH excursion is usually large enough that it is detectable by approximate measurement tools such as pH papers. Having detected the pH discrepancy, the brewer can correct it by adding acid to lower pH or calcium carbonate to raise it to the proper range.
A brewer who is suitably equipped and who wishes a finer degree of control over the brewing process — or if not control, reassurance that all is well throughout the brew day — will take many pH measurements. The sections that follow describe some of these pH measurement steps (though the list is probably not exhaustive).
pH in the Brewhouse
Water treatment: As brewers, we are concerned that our brewing water contains ions in quantities that are appropriate for the style of beer we’re making and the methods we’re using. We want to be sure, for example, that we have the correct level of sulfate to give the desired hop character to the beer. We also want to be sure that when we dough in with this water, the pH will be where we want it to be. The parameter that tells us this is the alkalinity, which is defined as the amount of hydrochloric acid per liter required to change the pH of water to pH 4.3 (3).
First, consider pure water, which has a pH of 7. We have already shown that if we add 0.1 mmol/L of hydrochloric acid to pure water, its pH drops to 4. The log of 0.5 is –0.3, so that 0.05 mmol/L of hydrochloric acid would get us to pH 4.3; thus, by this definition the alkalinity of distilled water is 0.05 mmol/L.
We can approximate the mineral content of Dortmund water by adding 365 mg/L (ppm) gypsum, 212 mg/L chalk, 247 mg/L epsom salts, 270 mg/L baking soda, and 3.1 mmol/L hydrochloric acid to pure water. Let’s use this mix as a second example. Its alkalinity is 4.24 mmol/L (84 times that of pure water), but its pH is also 7.
Note that the two example waters are very different and will result in significantly different performance in the mash tun, but they have the same pH. Obviously, then, the brewer cannot plan water treatment on the basis of pH alone. The key is the alkalinity.
Fortunately, alkalinity is one water parameter that is easy to measure by anyone equipped with a pH meter. The procedure is to titrate a 100-mL sample of the water with 0.1 N hydrochloric acid (0.1 N means that each milliliter of the liquid contains 0.1 mmol hydrogen ion). In this application of titration, you would add small measured quantities of the acid to the sample until the pH reaches 4.3 (3). The alkalinity is the number of milliliters of acid used, but the milliliters are usually multiplied by 50 to express the alkalinity in units of “parts per million” (CaCO3), which is the same as “millligrams per liter” (CaCO3).
If the water contains carbonate, the titration data would be similar to those in Figure 1; carbonate waters constitute a buffer, with increased buffering capacity at pH 6.38. The alkalinity measurement is in reality a measurement of the buffering capacity of the water.
In the earlier example, we determined how much acid would be required to move the pH of the buffer from pH 7 to pH 6. Here we calculate the amount of acid required to move the water from its native pH to pH 4.3.
Once you know the alkalinity of your liquor, you can decide whether water treatment is required or not. This decision depends on whether the acidity in the grist can overcome the buffering capacity of the water. Water treatment schemes are beyond the scope of this article, but they are discussed in many home brewing books (for example, reference 4) and in most brewing textbooks (for example, reference 5). As might be expected, many water treatment schemes are designed to reduce alkalinity, usually by chemically precipitating carbonate. Some brewers use other methods, such as ion exchange and reverse osmosis.
At this point, let’s assume you have a reservoir of liquor at hand and are ready to dough in. What do you do first? Check the pH. Why? Because you know what the pH of your dough-in water should be. If it is something other than what you expect, you know you have a potential problem. Your treatment (if used) was ineffective, or, in the case of untreated water, something changed in the water supply. In either event, you have been alerted to a situation that may lead to a problem with the beer.
Although liquor pH that differs appreciably from its expected value indicates a change in the water, a water pH check close to the expected value is not proof positive that something has not changed, though it does give some reassurance. In many cases, the pH measurements taken during the brewing process are like familiar milestones along the road to a good beer, and little more than that.
The mash: During the mash, proteins and starches are broken down into simpler (lower molecular weight) molecules. As previously mentioned, the rates of these reactions are controlled by a variety of enzymes, each of which has its optimum temperature and pH range. The proteases, for example, are generally most active between 104 and 140 °F (40–60 °C) and in the pH range of 4.2–5.3 (6); β-glucanase is most active between 95 and 104 °F (35–40 °C) in the broad pH range of 4–6 (6,3); α-amylase between 149 and 153 °F (65–67 °C) (7) and in the pH range of 5.3–5.7 (5); and β-amylase between 126 and 144 °F (52–62 °C) (8) and pH 5.1–5.3 (5).
The optimum pH and optimum temperature for each of these enzymes are interrelated; that is, the optimum temperature is a function of the pH, and vice versa. As most brewers already know, the character of the finished beer can be varied by adjusting these parameters. Small-scale brewers attempt fine temperature adjustment, but usually are content if the pH of the mash is maintained between 5.2 and 5.6 (some say 5.7), a range that represents a fair compromise for all the enzymes involved in mashing. If it gets out of this range, most brewers take steps to bring it to a value between these limits.
The most critical pH measurement the brewer makes is probably the one made at the completion of dough-in. When very pale (Pilsener) malt is doughed in with water of low alkalinity, the mash pH is usually around 5.8; the malts contain several acids that together form a buffer system with peak capacity at around pH 5.8. If the liquor has high alkalinity, its buffering capacity exceeds that of the malt and the mash pH will be higher than 5.8. If the grist contains highly kilned malts, the buffering capacity will be stronger and the pH lower. Brewers want to keep mash pH to between 5.2 and 5.7 during the saccharification step, which can be done as long as the mash-in pH is between about 5.3 and 5.8. Noonan indicates that mashing may proceed if the pH is between 4.7 and 6.2 (4).
If the grist contains highly kilned malts, it will contain additional acids that will lower the pH further. These acids are formed from sugars in the grain. It is surprising how much a small addition of patent malt can affect mash pH.
Usually, when the pH is out of the desirable range it is too high, the result of alkaline liquor and insufficient acid in the grist. The usual response in such cases is to add gypsum, which contains calcium that reacts as indicated above.
It is important not to add too much gypsum for some styles of beer (such as Pilseners) because gypsum also contains sulfate, which gives a dry bite to the hop bitterness, ruining the style. These beers, brewed with low-kilned pale malts, are of course the ones most likely to require pH adjustment because of the low buffering capacity of the malts. In low-kilned pale malts, some phytase survives malting and will free more phytic acid over time, provided that some calcium is present. The purpose of the acid rest is to permit acidification of the mash. If the rest is long enough, lactic bacteria spores will germinate and a lactic fermentation will start. The lactic acid produced lowers the pH nicely, but the procedure is time consuming and a bit risky.
Some brewers use a mineral acid (hydrochloric, sulfuric, or phosphoric) or an organic acid (frequently lactic) to lower pH. If you use these acids, you must be aware of the flavor effects of the anion (chloride, sulfate, phosphate, lactate) that you are adding to the brew, and purists will be bothered by this violation of the Reinheitsgebot, the traditional German beer purity law that prohibits adding anything to the mash (although the Biersteuergesetz does allow mineral acids in some cases).
Calcium chloride avoids the hop harshness effect by pairing calcium with chloride rather than sulfate, but it is difficult for home brewers to obtain and is difficult to weigh because it picks up water from the air. If you are brewing a Burton ale, you want exactly the dry bitterness that sulfate produces, and gypsum is fine for pH correction in such beers.
When the pH is too low, as would likely be the case when dark beers are brewed with liquor that is not sufficiently alkaline (the malt buffer system is stronger), the usual action is to add chalk (calcium carbonate) to neutralize the malt acids.
When adding these corrective substances for the first time, experiment. Add a small amount of the substance, mix it into the mash very thoroughly, and wait several minutes before checking the pH. If x mL of a sample gives y units of pH change, then another x mL should give another y units of pH change. The ratio of x and y is the buffering capacity of the mash. As we have seen, buffering capacity varies with pH. Figure 2 illustrates the fact that this variance is less dramatic with malt than with a simple buffer.
Figure 2 shows what happens when a small amount of malt is added to water and titrated with sodium hydroxide. The figure shows data for two titrations. In the first titration, 5 g of ground patent malt was added to 250 mL of reverse-osmosis water. For the second titration, 10 g of Pilsener malt were used with a like amount of water. The vertical axis is labeled in milliequivalents (a milliequivalent and a millimole are the same for sodium hydroxide) per kilogram of malt, so that the reader can get a rough idea of the total amount of alkali necessary to move a kilogram of either malt (which would typically be mashed in 2 L of water) one pH unit. Note that the small addition of black patent malt lowered the pH of the water to below 5.
In the same way that we defined alkalinity for liquor in terms of the amount of acid required to titrate to a given pH, we can also define the acidity of the malt in terms of an amount of alkali. In this context, acidiy is usually referred to as the titratable acidity. If we define titratable acidity as the amount of alkali required to neutralize the acidity in the mash — that is, to bring the pH to 7 — Figure 2 shows that the patent malt has a titratable acidity of 60 mEq/kg, as opposed to 20 mEq/kg for the Pilsener malt, and so the patent malt is, by this definition, three times as acidic as the Pilsener malt. Also note the slopes of the curves that reflect buffering capacity. At low pH, the buffering capacity of the patent malt is about the same as that of the Pilsener malt. At higher pH, the buffering capacity of the patent malt is about double that of the Pilsener malt.
Some readers have probably already deduced that you can plan your water treatment by comparing malt titratable acidity to liquor alkalinity, and they are correct, though that approach will not be developed further in this article. Only large breweries would want to bother with the calculations, and even there they would have to experiment to validate the results. As an alternative to calculation, you may experiment with acids or calcium carbonate on a small test mash until you achieve suitable mash-in pH, and then scale the acid or carbonate addition to the full brew length.
With decoction mashing, portions of the mash are removed from the mash vessel and boiled. Decoction produces a variety of effects, but one of them is that the boiling precipitates calcium as calcium phosphate. This precipitation drives the equilibrium in equation 3 to the right, resulting in the release of more hydrogen ions; the boiled decoction therefore becomes more acidic than the rest mash. Upon remixing, the pH of the rest mash lowers as well. In the classical triple decoction schedule, two decoctions precede the saccharification rest, providing two steps of pH reduction.
The brewer’s goal is to get the mash pH between 5.2 and 5.6 for the saccharification rest. By taking measurements during each temperature step, you can see whether you are likely to reach that goal. Most brewers know from experience what levels of pH are expected and desired. It is best to correct any problems at dough-in because proper pH is needed as much for protein and gum digestion as for the lysis of starch.
Lautering and sparging: Brewers monitor pH during sparging because tannin and silicate extraction from malt husks increases substantially if the pH of the runoff is allowed to rise above 6. Some brewers monitor runoff pH and gravity. Others acidify the sparge water to a pH of 6 or a little less (phosphoric acid is popular for this), and others know from experience that when the gravity of the runoff reaches a certain level, the pH is near 6 and it is time to stop collection.
It has also been reported that grain bed permeability and hence lauter and sparge flow rates increase as pH is reduced to within the 5.2 to 5.8 range (8).
The boil: Some of the same reactions mentioned in connection with decoctions take place in the boil, resulting in reduction of the pH. Proteins coagulate, and humulone and related compounds isomerize. Although these reactions are not catalyzed by enzymes, they nonetheless depend upon the pH of the solution in which they occur. The idea that lower pH facilitates protein coagulation is familiar to anyone who adds vinegar to the water used for poaching eggs. At the same time, however, it is possible to have too much of a good thing — protein coagulation is reduced if pH drops below 5 (5).
Proteins are made up of chains of amino acids, weak acids that dissociate according to the pH of the wort. Thus, at lower pH they have a net positive electric charge, at a pH called the iso-electric point (symbolized by pI) they have zero net electric charge, and at high pH (pH > pI) a net negative electric charge. Similarly charged particles repel one another, and oppositely charged particles attract, so it is evident that protein agglutination properties will vary with pH because of the effect of pH on the electric charges of the protein molecules.
Hough et al. provide data on the expected drop in pH during the boil, depending on the original pH and the duration of the boil (5). For an original pH of 6.06 and a boil time of 3 hours, the pH can be expected to drop 0.23 units; for an initial pH of 5.63, pH can be expected to drop 0.24 units; and for an initial pH of 5.09, pH can be expected to drop 0.1 units. Because we would hope the wort to be between pH 5.2 and 5.6 going into the kettle, we would expect it to finish the boil at somewhere between pH 5.1 and 5.4.
Fermentation: Fermentation is, biochemically speaking, an extremely complex process. If we consider the pathways that produce fusels, VDKs, esters, sterols, and so on, fermentation involves the action of probably close to a hundred enzymes. Again, each of these enzymes has its optimum temperature and pH range.
The yeast establish multiple buffer systems of their own and probably use these buffer systems to regulate their internal metabolism. They like an acid pH, and the pH drops rapidly by the end of the exponential growth phase. In lager beers, the pH is usually 4.2–4.4, whereas the pH with some ales can go as low as 3.8 (2). This drop is caused in part by saturation of the wort with carbon dioxide (but this cannot bring the pH below 4.0) and in part by the production of organic acids such as succinic acid; succinic acid’s pK1 of 4.16 is close to the pH expected at this point, which means maximum buffering action from the acid.
After the initial abrupt drop, the pH of the fermenting beer should change little. It may increase or decrease slightly as the fermentation continues. This is another example of those situations in which the brewer should know what to expect and be alert to deviations from the norm. A slower than normal drop in pH may indicate underpitching or underoxygenation, or both. An excessive drop in pH or pH continuing to drop would probably indicate infection with undesired organisms, some of which, in particular the lactic acid bacteria Lactobacillus and Pediococcus, are capable of growth at pHs as low as 3.5, though they prefer pH around 5.5.
In some special cases, such as wit beers, Iambics, or Berliner Weisse, sourness is an important part of the style and is traditionally brought about by controlled use of lactic acid bacteria. One well-known commercial brewer of wit beers, for example, introduces Lactobacillus at some point during the fermentation and monitors the pH. When the pH reaches the desired value, the beer is pasteurized to stop fermentation and lactic acid production.
Finished beer: A particular style of beer made in a particular way from a given set of ingredients and fermented with a known strain of yeast under controlled conditions should have a particular pH just as it should have a particular specific gravity. If it does not, then something went wrong somewhere along the line. If it does, it gives you a good indication that the process went as desired and that the product will be what you wanted.
In the case of the small-scale brewer trying to make one of the special beers like a wit but without the means or the desire to use the live organisms, an alternative is to add lactic (or tartaric or citric) acid when the beer is finished. This should be done to taste, but it is important to record the amounts of additions and the associated pHs that give the desired result. In future batches, the acid can be added until the previously determined pH is reached.
Yeast washing: Brewers often wash yeast collected from a previous fermentation in a solution of phosphoric, citric, sulfuric, or tartaric acid to kill bacteria and thus render the yeast suitable for repitching. It has been determined that the yeast can survive at pH 2, whereas the bacteria of concern cannot. The process of yeast washing involves monitoring the pH of the slurry and adding acid until the pH measures 2.0 (3).
Measuring Your Steps
As you can see, pH plays a major role in most brewing processes. Although the brewing process is blessed with natural buffering systems that tend to keep pH in acceptable ranges, measuring, monitoring, and adjusting pH can be of benefit. In the next issue of BrewingTechniques, we’ll review the various methods available for measuring pH and take a close look at electronic pH meters.
All contents copyright 2016 by MoreFlavor Inc. All rights reserved. No part of this document or the related files may be reproduced or transmitted in any form, by any means (electronic, photocopying, recording, or otherwise) without the prior written permission of the publisher.