Testing for pH: What the big pHuss is about

I first heard the axiom, “everything is pH-dependent,” many years ago from a coworker who was an Austrian-trained analytical chemist. She was right, or very close to it. Life processes, cheesemaking, and, yes, wine fermentations are among the things that are pH-dependent. Measuring the pH at various stages of your home wine production is among the most important lab tests you can do. In today’s column, first we’ll discuss what pH is and why it’s important. Then why we want to measure it. Finally we’ll look at the methods as well as the dos and do nots when taking these critical pH measurements.

The origins of the notation “pH” are subject to different interpretations and descriptions. The version I like best is one that is consistent with other chemical notation. That is, the “p” represents “the negative base ten logarithm of” and the “H” represents “the hydrogen ion activity.” The term “activity” in this case is so close to concentration that we may consider them interchangeable here. What it describes is the active, working expression of acidity (or alkalinity) in an aqueous (water) solution. Water itself has a dissociation constant — a natural tendency to separate into acid and base — of 10–14. That very low molar concentration of ionized water in pure water represents the concentration or activity of the two separated components of water, [H]+ and [OH].

In practice, [H]+ combines with another water molecule to form a hydronium ion, [H3O]+, and we can write the product of water dissociation as [H3O]+ x [OH] = 1 x 10-14. Since the concentrations of the two components are equal to one another (each pair representing a dissociated water molecule), [H3O]+ = [OH] = 1 x 10-7. The base ten logarithm of 10-7 is -7, so the negative of that is 7, giving pure water a theoretical pH of 7.0. The scale goes from zero to 14 for practical use, although very strong acids can display pH’s lower than zero and very strong bases can be higher than 14. Nearly all wine has a pH from 3.0 to 4.0.

While the precision of tabletap models is preferred by advanced winemakers, properly calibrated hand-held pH meters can still provide plenty of insight to sulfite and acid adjustments. Photo by Charles A. Parker/Images Plus

Unlike water, strong acids dissociate completely. For instance, hydrochloric acid (HCl) separates completely into H+ and Cl when dissolved in water. If we dissolve 0.1 mole of hydrochloric acid in a liter of water, that 0.1 molar concentration can be expressed as 10-1. The negative base ten logarithm of that figure is 1 and the solution has a theoretical pH of 1.0, with the actual pH very close to that. On the other hand, the acids of wine are weak acids and they do not completely dissociate. Tartaric, malic, lactic, and other organic acids only partially separate. They also interact with each other under changing pH conditions, creating a mixed buffered system in wine that is somewhat resistant to pH change. pH measures the “activity” of the acid and is critical for estimating how stable the wine will be. Since it is a logarithmic scale, it has no units — it’s a dimensionless number.

A related number is titratable acidity (also called total acidity, or TA). That is the sum of the mass of the acids in the wine or juice, independent of their state of dissociation. It is expressed as the matching weight of tartaric acid. Higher TA means lower pH, but they are not locked in a tight theoretical relationship. That means we home winemakers must measure both pH and TA; the latter will be covered in a future column in this series.

pH is critical to wine stability, especially in terms of the effectiveness of sulfur dioxide (SO2, sulfites) to prevent oxidation and spoilage. It is usually considered that it takes about 0.8 ppm (mg/L) of molecular sulfur dioxide to protect white wine and about 0.5 ppm in reds. We don’t measure molecular sulfur dioxide directly, we measure “free” sulfur dioxide. This is where pH comes in. In a low-pH wine, at say pH 3.1, the desired molecular level takes just 11 ppm free in white wine or 7 ppm in red. At a higher pH, say 3.7, white wine needs 63 ppm of free and red needs 39 ppm. The more acid activity in wine, the lower the pH, the more molecular SO2, and better protection for your wine. In terms of wine stability, low pH is good and high pH is bad.

Measuring the pH at various stages of your home wine production is among the most important lab tests you can do.

There are at least four times in a wine’s production cycle that you should consider measuring pH. First, it can support a harvest decision, along with TA and Brix (sugar content). For white grapes, pH at harvest is usually between 3.1 and 3.5 and for reds about 3.3 to 3.6. For readings much lower than the bottom of these ranges, the grapes are probably not fully mature and you should postpone picking. For readings much above these ranges, the grapes are possibly overripe and should be picked right away. That leads to the second opportunity, at crush time. In conjunction with TA readings, a pH measurement may tell you to add tartaric acid (if pH is high) or something like potassium bicarbonate (if pH is low). Next, measure again after fermentation is complete and you are about to begin your sulfite program. The online calculator at winemakermag.com/sulfitecalculator will guide you to a target free SO2 level based on your pH and wine type. Finally, you may want to measure again at bottling to verify stability for cellaring.

The original method for measuring pH is based on natural color changes of certain plant materials called indicators. Indeed, you can make a crude pH indicator by soaking shredded red cabbage in hot water, straining it off, and allowing the water to cool. Under acid conditions at low pH, the cabbage water is reddish-pink. Near neutral, it is purple. At high pH, it turns blue or even greenish-yellow if the pH is high enough. The color changes are from anthocyanin pigments in the cabbage and the behavior is mirrored in red grapes; at low pH the pigments are quite red and at higher pH they are purple or bluish.

We still use pH indicators, like phenolphthalein for the TA titration. The indicator is colorless at low pH and turns pink at the endpoint near pH 8.2. pH sticks or strips use similar color changes, sometimes on more than one pad, to show the pH of a sample. While suitable for endpoint detection in a titration, they are not generally accurate enough to guide decisions like harvest or sulfite additions. Most pH sticks or papers will give you a result that is plus or minus about 0.5 pH. Without even seeing it, I can tell you that your wine is probably pH 3.5 ± 0.5, just like a pH stick might, because almost all wine is in that range.

Your other choice is the correct one all winemakers should take: Purchase a pH meter. These are available in many different models at various price points, but all use the same principle: Electronically measuring the voltage between a glass pH electrode and a reference electrode. When I began studying chemistry in the twentieth century, there were often two separate electrodes for these functions. Now, you will almost always encounter a configuration where the two are combined in a single electrode body called a combination pH electrode. The bulb of the pH electrode develops a molecular gel layer of wet glass when it is kept in aqueous solution and that layer is critical to its sensitivity — that’s why pH electrodes are kept wet. When measuring, the reference electrode must be in electrical contact with the sample, so somewhere on that part of the probe you will see a fritted orifice or a small fiber that allows ions and current to pass back and forth from the reference electrode to the sample.

As the glass electrode comes to equilibrium with the solution, sodium, hydrogen, potassium, and possibly other positive ions adopt the same concentration in the gel layer as in the surrounding liquid. The ratios of those ions change the electrical potential or voltage as measured by a voltmeter between the glass electrode and the reference electrode. That voltmeter, marked as pH, is your pH meter. While a theoretical pH response for an electrode can be calculated, the actual output of any given electrode is subject to variation. Because of that, every pH meter and electrode must first be calibrated against solutions of known pH: pH buffers. For winemaking, we usually use buffers near pH 7 and pH 4, which are widely commercially available. This is because the meters need to “sense” neutral pH and then “sense” pH that more resembles the wine being tested. To use a pH meter, first soak the electrode in tap water for at least a half hour to saturate the glass gel. Do not use distilled or deionized water for this soak, as that will deplete the gel of ions and cause a slow response when you try to calibrate.

After soaking, rinse the electrode (now use distilled water) and blot dry with a tissue. Immerse the electrode tip (past any visible reference junction!) in pH 7.0 (or 7.01) buffer and follow the instructions that came with your meter to set the first calibration point. Rinse the electrode and blot dry again, immerse in the pH 4.0 (or 4.01) buffer and follow the rest of the instructions to complete the calibration. For the technically minded, the first buffer calibrated the response and the second buffer calibrated the slope. Best accuracy is near the second buffer, which is why we do 7 first and 4 last when measuring wine. Users can also test their meters using the pH 7 buffer to find potential problems with the electrode. If the offset deviates significantly, the meter’s readings may not be accurate and the electrode may need to be replaced. Rinse and blot the electrode again, and measure your sample the same way. Since you need to present fresh buffer or wine to the glass electrode throughout the test, swirl or stir buffers and samples.

When you are finished with your samples for the day, rinse the electrode and then store it in electrode storage solution or pH 4 buffer. That will keep the glass saturated and ready to use again next time. The most common failure with a pH meter is a dried-out electrode. If yours has dried out, soaking in tap water as much as overnight may bring it back. Another failure is a clogged reference junction. Some ribbon-type junctions can be pulled out a small amount and clipped off, exposing a new unclogged junction. A fritted glass junction (a little white spot) may be cleaned by swirling in water with some dishwashing detergent, or by alternating between solutions of mild acid and base (like a solution of citric acid and another of potassium bicarbonate). Generally speaking, if your pH meter will calibrate at 4 and 7, it is working and you may proceed. If it won’t, even after cleaning and soaking, you probably need a new electrode.

Effect of temperature

Temperature also affects the pH of a solution and inconsistent temperature contributes to inaccurate readings. This can be troublesome for winemakers when calibrating with buffers that are, for example, near room temperature and then measuring the pH of a must or wine that is cold (especially in the winter months where cold stabilization may be performed). To address this, ensure the sample of wine and buffer have had time to come to the same temperature before taking any pH readings.
A quality pH meter will have a temperature compensation feature that will take into consideration the temperature of the solution before displaying a reading. Check the manufacturer’s recommendations for purchasing the right probes to take advantage of this feature. They are commonly labeled “ATC” for Automatic Temperature Compensation.

Buffer storage and shelf life

  • Buffers have a shelf life of a couple years, and this is reduced to just a few months once the buffer comes in contact with air. It is important to use buffers that are not expired and have not been open for too long.
  • To avoid contaminating the buffer, always pour a small amount of buffer from the bottle into a smaller vessel to perform a calibration and then discard the portion. Never calibrate or store the electrode in the original bottle.
  • As an alternative to purchasing buffer in liquid form (which has a limited shelf life) consider using single use sachets of buffer powder. They can be opened and reconstituted in distilled water on an as-needed basis.
  • Before throwing out a costly electrode, calibrate it using brand new buffers just to make sure that any problems in measuring pH are due to the equipment and not a buffer that is past its prime.