Temperature impact on pH

I think the best way to explain this is to start with a brief discussion about pH and why temperature affects it. pH is a measure of hydrogen ion concentration using a logarithmic scale, where pH = -log [H⁺]. The pH of pure water is 7.0 at 77 °F (25 °C) because the concentration of hydrogen ions is 10-7 moles per liter — noted as [10^-7] using standard chemistry shorthand — because of the equilibrium of water with its dissociated ions as shown below:

H²O  H⁺+ OH^–

The equilibrium of molecules is governed by an equilibrium constant at a specific temperature. The equilibrium constant of water (written as K^w) is 10^-14 at 77 °F (25 °C). As temperature increases above 77 °F (25 °C), or standard temperature used in chemistry, dissociation increases as does the concentration of hydrogen ions. Because pH is defined as the -log [H⁺], an increase in [H⁺] corresponds to a lower pH. Acidic solutions have a higher concentration of hydrogen ions than pure water and bases have lower hydrogen ion concentrations compared to pure water.

The graph shown in Figure 2 illustrates that water pH ranges from 7.5 to 6.1 over the temperature range from 32–212 °F (0–100 °C). 

Using water as the topic of discussion, Figure 2 shows that water with pH 6.5 measured at 140 °F (60 °C) will increase to pH 7.0 when cooled to 77 °F (25 °C). However, this assumption becomes invalid if there is anything in the water that acts as a pH buffer. Buffers are systems of organic acids that can bind hydrogen ions through their own equilibria. For example, carbon dioxide readily dissolves in water and exists in three forms — carbon dioxide, bicarbonate, and carbonate, as shown in the following equation:

H²O + CO²  HCO^3– + H⁺   CO^3–2+ 2H⁺

Back to the assumption that water at 140 °F (60 °C) with pH 6.5 has a pH of 7.0 at 77 °F (25 °C). This is a poor assumption because the atmosphere contains about 0.04% carbon dioxide. Mashes contain much more buffering compounds compared to the small amount of carbon dioxide contributed by the atmosphere. These buffers include proteins, amino acids, phosphates, and nucleic acids from malt, plus carbonate from brewing water. To further complicate things, calcium and magnesium from brewing water both cause a reduction in mash pH because they react with malt compounds. In practical terms, this means that the mash system is heavily buffered and that changes in mash pH as a function of temperature are not as big as changes in pure water pH.

Life is full of approximations. The typical thumb is about an inch (2.5 cm) wide. A stone fetched from a pile of standard stones weighs 14 pounds (6.4 kg). A hand is 4 inches (10 cm) measured from thumb to opposite side of palm. And mash pH drops by about 0.30 pH units when cooled from mash to room temperature. One thing we know about these approximations is that they are indeed approximate!

The best way to consistently monitor mash pH is to either cool it to 68 °F (20 °C) — not 77 °F (25 °C) because biochemists use a different set of rules than physical chemists — or measure mash pH hot. If you prefer measuring mash pH at 68 °F (20 °C), you should use published pH ranges that are associated with cooled samples for your target range. Although the ranges vary by source, 5.45–5.65 at 68 °F (20 °C)  agrees with textbook information. Some references, most notably Malting & Brewing Science by Hough, Briggs, Stevens, and Young, provide mash pH at mash temperature and at room temperature. The true confusion with this subject comes from the lack of temperature reference in nearly all published data about mash pH. Given the well-known effect that temperature has on pH, it’s appalling that brewing scientists and academics have omitted this important detail.

Hopefully the background about pH and temperature is useful. Now let’s apply this information to your specific questions. You correctly understand what John is saying. The mash pH is lower than the pH measured in a cooled sample. Is it 0.3 pH units lower? The only way to know is to measure the pH at two temperatures because the mash buffering systems are too numerous and variable to predict the temperature effect.

Yes, if you are targeting pH 5.2 for your mash pH, then you want your reading to be higher when measuring a cooled sample. I will come back to this in a moment.

If you cool your sample to 63 °F (17 °C) instead of 68 °F (20 °C), you cannot use the same approximation for the offset. Instead of the difference being ~0.3 pH units, it may be closer to 0.32 pH units. Is this difference going to change your beer? Probably not, unless you are brewing the same beer many times a year on a commercial scale.

Now that I have answered your questions, let’s muddy things up a bit! pH 5.2 likely became a target for mash pH because of the following excerpt from Malting & Brewing Science:“ An infusion mash is best carried out at pH 5.2–5.4. Consequently, the pH in the cooled wort will be 5.5–5.8.” I think the first sentence became part of the homebrewing zeitgeist while the values in the second sentence were forgotten! I suggest changing your target pH at 68 °F (20 °C) to be in the 5.5–5.8 range.