and, especially, to minimize color leaching from the red grape varieties, Pinot Noir and Pinot Meunier, used in Champagne. At this stage in the ripening process, the grape skins are much less fragile than they are when completely ripe. The last method for controlling the total acidity of must is by taking great care in pressing the grapes and keeping the juice from each pressing separate (Volume 1, Section 14.3.2). In the Champagne region, the cuvée corresponds to juice from the mid‐part of the flesh (furthest from the skin and seeds), where it has the highest sugar and acidity levels.
Once the grapes have been pressed, winemakers have other means of raising or lowering the acidity of a must or wine. It may be necessary to acidify “flat” white wines by adding tartaric acid after malolactic fermentation in years when the grapes have a high malic acid content. This is mainly the case in cool‐climate vineyards, where the malic acid is not consumed during ripening. The disadvantage is that it causes an imbalance in the remaining total acidity, which then consists exclusively of a diacid, tartaric acid, and its monopotassium salt.
One method that is little‐known, or at least rarely used to avoid this total acidity imbalance, consists in partially or completely eliminating the malic acid by chemical means using a mixture of calcium tartrate and calcium carbonate. This method precipitates the double salt, calcium tartromalate (Section 1.4.4, Figure 1.9), and is a very flexible process. When the malic acid is partially eliminated, the wine has a buffer capacity based on those of both tartaric and malic acids, and not just on that of the former. Tartrate buffer capacity is less stable over time, as it decreases due to the precipitation of monopotassium and calcium salts during aging, whereas the malic acid salts are much more soluble.
Another advantage of partial elimination of malic acid over malolactic fermentation, followed by the addition of tartrate, is that, due to the low acidification rate, it does not produce wines with too low a pH. Low pH can be responsible for difficult or stuck secondary fermentation in the bottle during sparkling winemaking via the traditional method (méthode champenoise), leaving residual sugar in the wine.
Traditional acidification and deacidification methods are aimed solely at changing total acidity levels, with no concern for the impact on pH and even less for the buffer capacity of the wine and with all the unfortunate consequences this may have on flavor and aging potential.
This is certainly due to the lack of awareness of the importance of the acid–base buffer capacity in winemaking. Changes in the acid–base characteristics of a wine require knowledge of not only its total acidity and pH but also of its buffer capacity. These three parameters may be measured using a pH meter. Few articles in the literature deal with the buffer capacity of wine (Genevois and Ribéreau‐Gayon, 1935; Vergnes, 1940; Hochli, 1997; Dartiguenave et al., 2000a). This lack of knowledge is probably related to the fact that buffer capacity cannot be measured directly but rather requires readings of four or five points on a neutralization curve (Figure 1.3), and this is not one of the regular analyses carried out by winemakers.
It is now possible to automate the plotting of a neutralization curve, based on the wine's initial pH and total acidity, and thus measuring buffer capacity at the main stages in winemaking should become a routine.
Mathematically and geometrically, buffer capacity, β, and buffer range are deduced from the Henderson–Hasselbalch equation (Section 1.4.2, Equation (1.2)). Buffer capacity is defined by Equation (1.3):
where ΔB is the number of strong base equivalents that cause an increase in pH equal to ΔpH. Buffer range is a way of assessing buffer capacity. For an organic acid alone, with its salt in solution, it may be defined as the pH interval in which the buffer effect is optimum (Equation (1.4)):
Buffer capacity is normally defined in relation to a strong base, but it could clearly be defined in the same way in relation to a strong acid. In this case, the pH = f (strong acid) function decreases, and its β differential is negative, i.e.:
Strictly speaking, buffer capacity is obtained from the differential of the Henderson–Hasselbalch expression, i.e. from the following derived formula:
as only the Napierian logarithm is geometrically significant and provides access to the slope of the titration curve around its pKa (Figure 1.4).
The differential of the equation is as follows:
FIGURE 1.4 Determining the buffer capacity β from the titration curves of two model buffer solutions.
Making the assumption that the quantity of strong base added, d[B], generates the same variation in acidity in salt form, d[A−], and leads to an equal decrease in free acidity d[HA], per unit, i.e.:
the differential equation for pH is then
or,
Dividing both sides of the equation by d[B] gives the inverse of Equation (1.3), defining the buffer capacity. Theoretically, variations ΔB and ΔpH must be infinitely small, as the value of the ΔB/ΔpH ratio at a fixed pH corresponds geometrically to the tangent on each point on the titration curve (Figure 1.4). More practically, buffer capacity can be defined as the number of strong base equivalents required to cause an increase in pH of 1 unit per liter of must or wine. It is even more practical to calculate smaller pH variations in much smaller samples (e.g. 30 ml). Figure 1.4 clearly shows the difference in buffer capacity of a model solution between pH 3 and 4, as well as between pH 4 and 5.
This raises the issue of the pH and pKa at which buffer range should be assessed. Champagnol (1986) suggested that pH should be taken as the mean of the pKa
