Article

The Principles of pH

pH is an important factor in brewing quality beer. The pH levels during various stages of the brewing process affect extract potential, beer color, hot-break formation, foam stability, hop oil extraction, hop bitterness and lauterability of the beer. It is also an important consideration for beer quality during storage as a low pH inhibits bacterial growth.

So What, Exactly, is pH?

The pH value of a solution is a way of expressing the acidity or alkalinity of that solution. Most homebrewers are familiar with the pH scale and know that values over 7 are basic (or alkaline) and values under seven are acidic (Assuming the pH reading is taken at 68 °F/20 °C). You might not know that Soren Sorensen, a Danish biochemist working for Carlsberg labs, is the man who established the concept of pH. But what, exactly, is pH? A good place to start explaining pH is with pure water. Pure water is a mixture of mostly H2O molecules, and a very small number of hydronium ions (H3O+) and hydroxyl ions (OH-). This is because, in pure water, a small number of water molecules spontaneously disassociate into H+ and OH- ions. The H+ ions almost immediately combine with a water molecule to form H3O+. The idea of pH can be approached nicely using the concept of the ionization constant of water, or Kw. Kw is defined as the product of the concentration of H3O+ and OH- ions present within the solution or:

Kw = [H3O+][OH-]

(In a chemical equation, brackets around an ion or molecule indicate the concentration of that substance.)
For pure water at 25 ºC (77 °F), Kw = 10-14, which can also be written as 0.00000000000001. Since, in pure water, each disassociated water molecule provides one ion of H3O+ and one ion of OH-, it follows that [H3O+] = [OH-]. Therefore:

Kw = [H3O+]2 = 10-14

By solving for [H3O+], we find that:

[H3O+] = 10-7 (or 0.0000001)

pH is defined as the negative log of the hydronium ion concentration, or equivalently:

pH = -log [H3O+]

So, in pure water at 25 °C (77 °F), the concentration of H3O+ is 10-7, and thus the pH is 7.

In pure water at 25 °C (77 °F), the pOH — the negative log of the hydroxl ion concentration — is also 7 (because the concentration of hydronium and hydroxl ions are equal). In addition, in any dilute aqueous solution, at any temperature, the following is true:

pH + pOH = pKw

This relationship becomes important when measuring the pH of solutions — such as mashing or boiling wort — at temperatures other than 25 °C (77 °F) as Kw varies with temperature.

The relative acidity or alkalinity of an aqueous solution depends on whether there are more H3O+ ions or OH- ions present within the solution. If there are more H3O+ ions present, the solution is acidic. If there are more OH- ions present, the solution is basic (or alkaline).

For example, if you add some acid to pure water at 25 °C (77 °F), the concentration of hydronium ions goes up. Consequently, the pH goes down because pH is the negative log of that concentration. In addition, since at this temperature pH + pOH = 14, the concentration of hydoxyl ions goes down — because some hydronium ions and hydoxyl ions react to form water molecules — and consequently the pOH goes up.

Because of the way that pH is defined, the pH scale is not linear. A solution that has a pH of 4 is ten times more acidic than a solution that has a pH of 5, and one hundred times more acidic than a solution that has a pH of 6.

Change During the Brewing Process

During the brewing process, the pH of the wort and beer changes. Water from most municipal water sources will have a pH over 7 (because it is treated to prevent corrosion of pipes). When combined with crushed malt, the pH of the grain and water mixture drops considerably compared to the initial pH of the water alone.

Natural pH Decrease

This observed pH decrease is the result of changing mineral composition within the solution. The principal change that happens during the mashing process is the precipitation of phosphates and amino acids derived from the malt. Phosphates, such as phosphoric acid, will disassociate. For example:

H3PO4 –> H+ + H2PO4-

H2PO4- –> H+ + HPO4-2

and

HPO4-2 –> H+ + PO4-3

If calcium ions are present, the phosphates will precipitate as calcium phosphate, leaving behind hydrogen ions:

3Ca+2 + 2H3PO4 –> 6H+ + Ca3(PO4)2

A similar reaction occurs if magnesium ions are present, but magnesium phosphate is more soluble than calcium phosphate, so the effect on pH is less dramatic.

A reaction will also occur if amino acids or polypeptides are present within the solution. The calcium ions will react with the amino acid group:

2(Amino Acid Group – COOH) + Ca+2 –> Ca (Amino Acid Group – COO)2 + 2H+

If calcium sulfate (CaSO4) is added to the brewing water, amino acids will form the insoluble precipitate as described above, leaving behind hydrogen ions (H+) — which, remember, instantly combine with water to form hydronium ions — and sulfate ions, (SO4-2).

These changes in mineral composition and the precipitation of calcium salts are responsible for the majority of the pH decrease that is observed prior to fermentation. However, the composition of the grain bill also influences pH. If the same water is used for two mashes, a mash with dark specialty malts would settle into a lower pH than a mash composed entirely of pale base malts. A mash of pale malts and a starchy adjunct — such as rice or corn — would have a pH higher than either of the previous two.

Interference With the Natural pH Decrease

The presence of other minerals within the brewing water can interfere with the pH decrease during the brewing process. Specifically, the carbonate (CO3-2) and bicarbonate (HCO3-) ions (the ions associated with temporary water hardness) can act as buffers to pH decrease. These ions interact with water molecules to form hydroxyl ions (OH-):

CO3-2 + H2O –> HCO3- + OH-

HCO3- + H2O –> H2CO3 + OH-

These extra OH- ions will then react with any H3O+ ions that they happen to encounter and produce water molecules. This effectively removes the extra H+ ions that are being generated by the brewing process and limits the natural pH decrease. This is why it is important to ensure that the ions responsible for temporary hardness are removed from the brewing water, especially when brewing light-colored beers.

Proper Mash pH

Optimally, the pH of an infusion mash should be in the 5.2–5.6 range, with the lower half of this range often cited as being preferable. This range is a compromise between the pH optima for a variety of processes. Mash pH affects many aspects of the brew including extract yield, fermentability, tannin extraction, lauterability and saccharification time.

In an infusion mash, the greatest extract yield is achieved when the pH of the mash is 5.2–5.4.  The most fermentable wort is obtained in the 5.3–5.4 range. Fastest conversion time is obtained in the 5.3–5.6 range.

If the pH during mashing is too high, starch and protein hydrolysis can be adversely affected. Also, high pH during the mash will increase the amount of dextrins present in the wort, resulting in a less fermentable wort.

The husks of malted barley contain compounds such as polyphenols (such as tannins) and silica compounds that are more soluble, and therefore more easily extracted, under high pH conditions. Polyphenols can contribute to colloidal instability and can produce astringency in the finished beer.

Most polyphenols are extracted during the final stages of sparging, when the pH of the wort being run off from the mash rises. It is therefore important to stop collecting wort when the pH of the last runnings climbs to 5.8–6.0. (Note that it’s the pH of the wort being run off that matters, not the pH of the sparge water.)

The optimal pH for many aspects of mashing actually varies due to temperature, mash thickness and other factors, including whether an infusion or decoction mash is employed. As such, optimal pH ranges cited in the brewing literature sometimes vary by quite a bit.

For the homebrewer, getting your mash pH in the right ballpark will greatly improve your beer if you have previously missed the  mark. However, making small adjustments within the acceptable range will likely not result in major changes in your beer. And generally, once you learn how to control your pH for a given beer, you will not need to monitor the pH each brewing session. In many cases, the first time a brewer checks his pH, he will find that everything has been OK all along.

Controlling Mash pH

If the pH of your mash does not naturally fall into the acceptable range, there are a variety of ways to manipulate it. The most common problem for brewers, especially those with lots of carbonate ions in their water, is a mash pH that is too high. To lower pH, brewers often add calcium ions, from gypsum (calcium sulfate) or calcium chloride. In a 5-gallon (19-L) batch, one or two teaspoons of either of these will often solve the problem. Likewise, organic acids — such as lactic acid or phosphoric acid — can be added to directly lower mash pH. Adding sour malt, up to about 5% of the grist, is a “natural” way to add lactic acid to the mash.

If the brewer’s water has a lot of carbonates, and this is what is keeping his or her pH level too high, the carbonate level can be greatly reduced by boiling the water and racking it off the precipitate. It is usually easier, however, to simply treat carbonate-rich water with acid (to neutralize the carbonates) or dilute it with distilled water or water prepared by reverse osmosis (RO).

In some cases, especially if a brewer is using very soft water and making a dark beer, the pH of the mash may be too low. In these cases, adding a little bit of chalk (calcium carbonate) or baking soda (sodium bicarbonate) will help.

The Importance of Boil pH

After mashing, the wort is run off to the kettle and boiled. Just as pH is important in the mash, it affects many different processes in the boil as well. During the boil, calcium phosphate will continue to be precipitated — just as it was during the mash — as long as sufficient calcium is still present in the wort. As such, the pH decreases and continue to decrease as the boil progresses.

Optimally, a post-boil wort pH of 5.0–5.2 shuld be achieved. Landing in the right range will help you get the best character extracted from your hops, maximize the amount of hop break that is formed and keep color pickup during the boil to a minimum. Usually, establishing the proper mash pH will allow you to hit the right boil pH without any manipulation, but this isn’t always the case.

Isomerization of alpha acids into iso-alpha acids during wort boiling is influenced by pH. This isomerization reaction is favored by higher pH. Within a pH range of 8–10, the conversion to iso-alpha acids can approach 90%. (This is why hop extracts are produced at very high pH levels.) At typical wort pH ranges (5.2–5.4). the conversion is limited to a theoretical maximum of about 60%, with a final utilization value of about 35%. This does not mean that a high boil pH is a good thing; although high-pH boils extract more bitterness from the hops, the character of the bitterness is more “coarse” and the beer will likely suffer from many other pH-related problems.

Coagulating the hot break — a complex of proteins and polyphenols — is another important function of the boil. The pH of your boil has a very visible affect on this. The optimal pH for break formation is 5.2. If, at the beginning of your boil, you see big, fluffy bits of break material in your wort, you will have visual confirmation that your pH is in the right range.

Wort color generally increases during the wort boiling process due to Maillard reactions, reactions between amino acids and sugars. Maillard reactions are not favored at lower pH values, so having a wort of lower pH is important if a beer of lighter color is to be produced.

If your kettle pH needs to be lowered, adding a little bit of calcium usually helps. For five gallons (19 L) of wort, 1⁄4–1⁄2 tsp. of gypsum or calcium chloride should do the trick. You can also add acid.

And Finally, Fermentation

During fermentation, the pH continues to drop for a variety of reasons. Yeast cells take in ammonium ions (which are strongly basic) and excrete organic acids (including lactic acid). The yeast strain chosen can affect the final beer pH. Most lager beers finish at 4.2–4.6, with some ales ending as low as 3.8. (Sour beers may have pH values around 3.0.)

Achieving an optimal pH, less than 4.4, favors faster beer maturation (including uptake of diacetyl), better beer clarity, better biological stability and a “more refined” beer taste.

Brewers rarely adjust final beer pH with acid. To reach a suitable final pH, all that is needed is to conduct a good, vigorous fermentation. As pH decreases with attenuation, drier beers tend to have slightly lower pH values. One interesting tidbit about fermentation is that some molecules in the fermenting beer become decolorized as the pH lowers and so the color of beer actually lightens slighty during fermentation.

Summary

pH affects almost all of the physical, chemical and biochemical reactions that occur within the brewing process. Brewers who understand the factors that affect pH and how to manage them during the brewing process will be more able to consistently produce good beer. Although pH is clearly an important variable in the brewing process, it rarely requires a great deal of attention from the homebrewer. Usually matching the grain bill with suitable brewing water should be all you need to ensure a successful brew day.

Measuring pH (by Chris Colby)

The best way to measure pH in a homebrewery is with an inexpensive pH meter. There are many adequate models that cost less than $100.

When you first get your pH meter, begin soaking the electrode in electrode storage solution. Whenever the meter is not in use, it will need to be stored in this solution. Ideally, the electrode should never be allowed to dry out.

Calibrate the meter according to the meter’s instructions, using a pH 7.01 buffer and a pH 4.01 buffer.

Take your wort sample in a clean glass. If the sample is from the mash, cool it down to room temperature, even if your pH meter has automatic temperature control. Taking readings of hot samples will decrease the life of the electrode. Rinse the electrode with distilled water then dry the electrode with a tissue. Don’t let the tissue touch the electrode, just bring it close enough to wick the liquid away.

Place the electrode in the sample and give the sample a quick swirl. Make sure there are no bubbles attached to the electrode. Turn on the power to the electrode. The power to the electrode should never be on unless the electrode is submerged.

With the power on the electrode, the meter will take the reading. Note it in your lab notebook, then turn the power to the electrode off before pulling it out of solution. Rinse the electrode with distilled water again, dry and return to the storage solution.

Since pH changes with temperature, you need to compensate for this change. At room temperature, the pH of the cooled sample will be around 0.35 units higher than the pH at mash temperature. Thus, if you get a reading of pH 5.60 for your cooled sample, your corrected reading would be pH 5.25. (You need to do this even if your meter has temperature correction.)