Distilled water is created through a purification process involving boiling water into vapor and then condensing the steam back into a liquid. The fundamental question of whether this highly purified substance is classified as an acid or a base is not straightforward, as the answer depends on whether the water is considered in its theoretical, pure state or its practical, real-world condition. Understanding the acidity or basicity of any aqueous solution, including distilled water, requires first examining the scale used to measure this property: the pH scale.
Understanding the pH Scale
The pH scale is a logarithmic measurement that quantifies the concentration of hydrogen ions (\(\text{H}^+\)) in an aqueous solution. This scale ranges from 0 to 14. A substance with a pH below 7 is considered acidic, indicating a relatively high concentration of hydrogen ions. Conversely, a substance with a pH above 7 is classified as basic, or alkaline, which means it has a higher concentration of hydroxide ions (\(\text{OH}^-\)) relative to hydrogen ions.
A neutral substance sits directly in the middle of the scale at pH 7, where the concentration of hydrogen ions is exactly equal to the concentration of hydroxide ions. Because the scale is logarithmic, a change of one pH unit represents a tenfold change in acidity or basicity. For instance, a solution with a pH of 5 is ten times more acidic than a solution with a pH of 6.
The Ideal State of Distilled Water
In a perfect, theoretical setting, distilled water is neither an acid nor a base; it is perfectly neutral with a pH of 7.0. This neutrality arises from the chemical phenomenon known as the self-ionization of water, which is a constant, reversible reaction. In this process, a very small fraction of water molecules spontaneously dissociate, with one water molecule donating a proton (\(\text{H}^+\)) to another.
This reaction results in the formation of equal numbers of hydronium ions (\(\text{H}_3\text{O}^+\), which is often simplified to \(\text{H}^+\)) and hydroxide ions (\(\text{OH}^-\)). The equal concentration of these two ions cancels out any acidic or basic tendency, establishing the defining characteristic of a neutral solution.
This perfect neutrality is strictly dependent on the water being completely isolated from its environment, such as being held in a vacuum. The distillation process removes virtually all dissolved solids and ions, making the water an extremely poor electrical conductor and theoretically perfectly neutral. However, maintaining this precise pH 7 in a real-world setting is practically impossible.
Real-World Acidity: The Role of Carbon Dioxide
Once distilled water is exposed to the atmosphere, it quickly loses its theoretical neutrality and becomes slightly acidic. Because it lacks dissolved minerals, the water acts as a strong solvent and readily absorbs atmospheric carbon dioxide (\(\text{CO}_2\)).
As the carbon dioxide gas dissolves into the water, it reacts with the water molecules to form carbonic acid (\(\text{H}_2\text{CO}_3\)). This weak acid then partially dissociates, releasing a small number of hydrogen ions into the solution. This slight increase in \(\text{H}^+\) concentration is enough to shift the pH value below 7.
Consequently, the pH of distilled water measured in a typical environment is usually found in a slightly acidic range, often between 5.5 and 6.5. This slight acidity is a consistent chemical reality for distilled water stored outside of a controlled, air-tight system.
Distilled Water Compared to Tap Water
The contrast between distilled water and ordinary tap water highlights the role of dissolved substances in maintaining pH stability. Tap water, sourced from municipal supplies, contains a variety of dissolved minerals, salts, and ions, such as calcium, magnesium, and bicarbonates. These dissolved solids give tap water a much higher electrical conductivity compared to distilled water.
Crucially, the minerals in tap water often act as natural chemical buffers that resist changes in pH. For example, bicarbonates in tap water can neutralize the carbonic acid formed from dissolved \(\text{CO}_2\), thereby stabilizing the water’s pH.
While the pH of tap water can vary widely depending on the source, generally falling between 6.5 and 9.5, this value tends to remain stable even upon exposure to air. The buffering capacity of the dissolved compounds prevents the \(\text{CO}_2\) from having a significant acidic effect. In contrast, distilled water’s lack of these stabilizing buffers allows its pH to drop rapidly toward the acidic range when it absorbs carbon dioxide.