Pure water is neither acidic nor basic; it is neutral, defined by a pH value of exactly 7.0 at a standard temperature of 25 degrees Celsius. This neutrality is a fundamental chemical property of the water molecule itself, meaning it is not dependent on external dissolved substances. The concept of pure water serves as the theoretical benchmark against which all other aqueous solutions are measured.
Defining the pH Scale
The pH scale is a logarithmic measure used to specify the acidity or basicity of an aqueous solution, corresponding directly to the concentration of hydrogen ions ($\text{H}^+$). This scale typically ranges from 0 to 14. A value below 7 indicates an acidic solution, and a value above 7 indicates a basic, or alkaline, solution. The scale is inversely related to the hydrogen ion concentration, meaning a lower pH signifies a higher concentration of $\text{H}^+$ ions.
Solutions with a pH near the extremes exhibit highly reactive properties. For example, lemon juice is acidic (pH 2.35), while a baking soda solution is alkaline (pH 8.3). A neutral solution, such as pure water, sits squarely in the middle at pH 7, indicating an equal balance between acidic and basic components.
The Chemical Reason for Water’s Neutrality
The neutrality of pure water stems from a process called auto-ionization, where water molecules spontaneously react with each other to produce ions. In this reaction, one water molecule ($\text{H}_2\text{O}$) donates a proton to another water molecule, forming a hydronium ion ($\text{H}_3\text{O}^+$) and a hydroxide ion ($\text{OH}^-$). The hydronium ion is the acidic component, while the hydroxide ion is the basic component, and their presence dictates the water’s $\text{pH}$.
In pure water, this dissociation occurs on an extremely small but measurable scale. At 25 degrees Celsius, the concentration of both the hydronium ions ($\text{H}_3\text{O}^+$) and the hydroxide ions ($\text{OH}^-$) are exactly equal, measuring $1.0 \times 10^{-7}$ M (moles per liter). The $\text{pH}$ value is mathematically derived from the negative logarithm of the hydronium ion concentration, and the negative logarithm of $1.0 \times 10^{-7}$ is 7.
The concentrations of $\text{H}_3\text{O}^+$ and $\text{OH}^-$ are linked by the ion product constant for water ($\text{K}_w$), which is $1.0 \times 10^{-14}$ at 25°C. Since the two ions are produced in equal amounts, their concentrations must be the square root of the $\text{K}_w$ value, confirming they are both $1.0 \times 10^{-7}$ M. This perfect balance of acidic and basic ions is the precise chemical definition of neutrality.
Distinguishing Pure Water from Natural Water Sources
The concept of chemically pure water, with its precise $\text{pH}$ of 7.0, is largely a laboratory ideal. Natural water sources, including tap water and rainwater, are never truly pure because they dissolve gases, minerals, and other contaminants from their environment, causing the $\text{pH}$ to deviate from the neutral 7.0 benchmark. Rainwater, for example, is naturally slightly acidic (typically $\text{pH}$ 5.6) because atmospheric carbon dioxide ($\text{CO}_2$) dissolves into the water droplets, forming carbonic acid.
Tap water and groundwater $\text{pH}$ fluctuate depending on local geology and treatment processes. Water passing through mineral-rich soil, such as limestone, often becomes slightly basic due to dissolved compounds. Conversely, the moment pure water is exposed to the air, it begins to absorb atmospheric $\text{CO}_2$, causing its $\text{pH}$ to drop almost immediately.
Maintaining a $\text{pH}$ of 7.0 is a challenge outside of a controlled, sealed environment. Most water encountered day-to-day is a solution containing various dissolved substances, meaning its $\text{pH}$ will be slightly acidic or basic. This deviation, however, is not a reflection of pure water’s fundamental neutral nature.
Measuring Water’s pH
Determining the $\text{pH}$ of water can be accomplished using several methods that vary in accuracy and convenience. For quick, general estimates, color-changing litmus paper or $\text{pH}$ strips are often used, providing a rough indication of acidity or alkalinity based on a color chart. For highly accurate measurements, an electronic $\text{pH}$ meter is employed, which uses a specialized electrode to measure the electrical potential difference in the solution and provides a precise numerical reading.
Temperature is an important consideration during measurement, as the $\text{pH}$ of neutrality is defined as 7.0 only at 25 degrees Celsius. Since the auto-ionization of water is an endothermic process, increasing the temperature causes more ions to form. This increase in ion concentration slightly lowers the $\text{pH}$ value.