The acidity or basicity of an aqueous solution is quantified by its pH, a fundamental measurement in chemistry, biology, and environmental science. pH measures the concentration of hydrogen ions (H+) present in the liquid. A higher concentration of H+ indicates an acidic substance, while a lower concentration indicates a basic substance. Accurately determining pH is important for applications ranging from food safety to monitoring water quality. Experimental determination of pH uses methods from simple color-changing agents to sophisticated electronic instrumentation.
Measuring pH Using Chemical Indicators
The simplest approach to determining pH relies on colorimetric methods, which use chemical compounds called indicators. These indicators are typically weak organic acids or bases that undergo a reversible structural change when they interact with hydrogen ions. This molecular rearrangement alters the way the compound absorbs light, resulting in a visible color change across a specific pH range.
One common application involves pH test strips, which are small pieces of paper infused with indicator dyes. When dipped into a solution, the paper quickly changes color, and the resulting hue is compared against a provided color chart to estimate the pH value. While incredibly fast and inexpensive, this method only provides a rough, semi-quantitative reading, often accurate only to the nearest whole number.
Liquid chemical indicators are also widely used, either added directly to a solution or employed in procedures like titrations. For example, phenolphthalein is colorless in acidic solutions but turns pink in basic conditions, typically above a pH of 8.2. Universal indicators, which are a blend of several different dyes, are designed to display a spectrum of colors across the entire 0–14 pH range. These colorimetric techniques serve as a convenient first check when high precision is not required.
Precision Measurement with pH Meters
For applications demanding accuracy, the potentiometric method using a pH meter is the standard technique. A pH meter measures the electrical potential difference, or voltage, that develops across a specialized glass membrane when submerged in a solution. This glass electrode is selectively sensitive to hydrogen ions, generating a voltage directly proportional to the H+ concentration of the sample.
Modern pH systems typically use a combination electrode, housing both the glass measuring electrode and a stable reference electrode within a single probe body. The reference electrode provides a constant, known voltage for comparison against the variable voltage generated by the glass electrode. The meter converts this measured voltage difference into a corresponding pH reading based on the Nernst equation.
To ensure accurate readings, the pH meter must be calibrated regularly using standardized buffer solutions of known, stable pH values, commonly pH 4.01, 7.00, and 10.00. Calibration adjusts the meter’s internal electronic zero point and slope to match the electrode’s current response characteristics. Proper technique requires maintaining the glass bulb of the electrode by keeping it hydrated in a storage solution when not in use. A dried-out membrane will yield unstable and inaccurate results.
Interpreting the Logarithmic pH Scale
Results from experimental measurements are expressed on the pH scale, which conventionally spans from 0 to 14. A reading of 7.0 represents a neutral solution, such as pure water, where the concentrations of hydrogen ions (H+) and hydroxide ions (OH-) are equal. Solutions below pH 7.0 are acidic, with acidity increasing as the number approaches zero.
Conversely, solutions above pH 7.0 are considered basic, or alkaline, with basicity increasing as the number approaches 14. The structure of this scale is important because it is logarithmic, not linear.
A change of one whole pH unit represents a tenfold change in the hydrogen ion concentration. For instance, a solution with a pH of 5.0 is ten times more acidic than a solution with a pH of 6.0, and one hundred times more acidic than a solution with a pH of 7.0.
The logarithmic nature of the scale allows scientists to manage and represent the vast range of H+ concentrations found in nature and industry. This structure highlights the significant chemical difference that even a small numerical change in a pH reading represents.