Hydrogen Ion Concentration (pH)
In simple terms, pH refers to the concentration of hydrogen ions (H⁺) in an aqueous solution. Although the term “hydrogen ion” is commonly used, free hydrogen ions are highly reactive in water and do not exist in isolation. Instead, they bind with water molecules to form hydronium ions (H₃O⁺). Thus, pH more accurately represents the concentration of hydronium ions.
The concentration of hydrogen ions in water typically ranges from 10⁻¹ to 10⁻¹⁴ mol/L, a span that is inconvenient for direct numerical manipulation. To address this, the pH scale employs a negative logarithmic transformation of hydrogen ion concentration. The pH scale generally ranges from 1 to 14, with lower values indicating acidity and higher values indicating alkalinity. A pH of 7 is considered neutral. Due to its logarithmic nature, each unit change in pH represents a tenfold change in hydrogen ion concentration. For instance, a solution with pH 4 is ten times more acidic than one with pH 5 and one hundred times more acidic than one with pH 6.
Common acids encountered in coffee and food science include acetic acid (CH₃COOH), citric acid, and malic acid. These compounds dissociate in water to release hydrogen ions. For example:
CH₃COOH ⇌ CH₃COO⁻ + H⁺
HCl ⇌ H⁺ + Cl⁻
It is a common misconception that dissolving equal masses of different acids yields identical pH values or comparable sensory acidity. However, the critical variables are the dissociation constant of each acid and the proportion of molecules that actually release hydrogen ions in solution. Furthermore, two solutions with the same pH may exhibit radically different sensory perceptions of acidity.
For instance, dark-roasted coffee may register a pH around 5.0 to 5.5, indicating moderate acidity, yet its perceived acidity is minimal. Conversely, lighter roasts with similar pH values can taste markedly brighter. Given such inconsistencies and measurement errors—including those introduced by temperature, probe calibration, and variable acid content—direct comparative tasting may offer a more accurate assessment than pH readings.
In assessing perceived acidity, the concept of titratable acidity (TA) often provides a more comprehensive and reliable metric than pH. While pH measures the activity of free hydrogen ions, TA quantifies the total potential acidity, including weakly dissociated acids. If pH represents the tip of the iceberg, TA represents the entire mass below the surface. Strong acids like hydrochloric acid dissociate completely in water, while weak acids such as acetic acid dissociate only partially. For example, out of 100 molecules of hydrochloric acid, nearly all release hydrogen ions, whereas only 1–4 molecules of acetic acid may do so. This disparity explains why two solutions may contain the same number of acidic molecules but exhibit different pH values. Conversely, titratable acidity would remain the same in both, as it accounts for the total amount of acid present, regardless of dissociation.
Titratable acidity is typically determined by titrating the sample with a strong base such as sodium hydroxide (NaOH) until the pH reaches 8.2.
In contexts outside of coffee extraction, pH can significantly influence chemical behavior. For instance, a higher pH can increase the dissociation of organic acids and modify the binding affinity of calcium and magnesium ions to those acids. However, in coffee extraction, these effects are largely negligible. This is because coffee extraction occurs in a slurry—a suspension of coffee particles in water—rather than in pure water. Coffee itself is inherently acidic. Regardless of the initial pH of the brewing water (whether 8.0 or 7.0), contact with ground coffee causes rapid acidification, driving the final pH toward 5.0–5.5. In this context, the water’s starting pH becomes irrelevant almost immediately upon brewing.
One might presume that water with a high initial pH could neutralize coffee acids. However, pH alone does not confer buffering capacity. Once additional hydrogen ions are introduced, the pH will plummet accordingly. Therefore, when evaluating a solution’s ability to resist changes in pH due to acid influx, alkalinity—not pH—is the relevant metric. Alkalinity represents the actual buffering capacity that stabilizes pH during extraction.
Although the pH of the brewing water prior to extraction is relatively unimportant, the pH after extraction can be quite significant. Lower post-extraction pH values help preserve volatile compounds such as esters and lactones, which are prone to hydrolysis at elevated pH levels. A lower pH environment thus retards the degradation of aromatic compounds, contributing to greater flavor longevity. Additionally, polyphenols such as chlorogenic acids—which contribute to bitterness and astringency—oxidize more rapidly under alkaline conditions. Their oxidation leads to unpleasant flavors and a dull sensory profile. Maintaining a lower pH slows this oxidative process and helps preserve the desirable attributes of the coffee.