The pH scale provides a measure of how acidic or basic a substance is by quantifying the concentration of hydrogen ions (\(\text{H}^+\)) in an aqueous solution. Ranging from 0 to 14, values below 7 indicate acidity, values above 7 indicate basicity (alkalinity), and 7 represents neutrality. The scale is logarithmic, meaning each whole number change represents a tenfold difference in hydrogen ion concentration. Monitoring pH is important across many fields, including health (for optimal enzyme function) and agriculture (where soil pH affects nutrient availability).
What pH Indicators Are
A pH indicator is a chemical compound added to a solution to determine its acidity or basicity. These substances are typically weak acids or weak bases that react reversibly to changes in hydrogen ion concentration. The indicator molecule exists in two forms—one acidic and one basic—with each form exhibiting a different color. The color change occurs over a narrow range of pH values, known as the transition range, where the indicator shifts from the color of its acidic form to the color of its basic form.
The Chemistry Behind the Color Change
The mechanism of a pH indicator’s color change is rooted in a reversible chemical equilibrium involving the gain or loss of a proton (\(\text{H}^+\)). In an acidic solution, the high concentration of hydrogen ions causes the molecule to gain a proton, entering its protonated form. In a basic solution, the molecule loses a proton, resulting in its deprotonated form. These two structural forms have different arrangements of electrons, which fundamentally changes how they absorb light.
The color change becomes visible when the concentration of one colored form significantly outweighs the other. The point where the indicator’s acidic and basic forms are present in equal concentrations, resulting in a blended color, is close to the indicator’s specific pKa value. This shift in molecular structure allows a tiny amount of indicator to provide a clear, visual signal of the solution’s environmental change.
Common Examples of pH Indicators
Indicators are selected based on their specific transition range. Litmus, a natural indicator derived from lichens, is widely recognized, turning red in acidic solutions and blue in basic solutions. It is used for quick, general testing, though its wide transition range (pH 5 to 8) makes it less suitable for precise measurements. Phenolphthalein is a common synthetic example, colorless below pH 8.2 but turning bright pink in basic solutions above pH 10.
Universal Indicators
For estimating pH across the entire scale, a universal indicator is often employed. This substance is a mixture of several different indicators, allowing it to exhibit a full rainbow of colors from pH 1 to 14. Universal indicator paper, or pH paper, uses this mixture to provide a color that can be matched to a reference chart, yielding a quick, approximate pH value. Natural indicators, such as the anthocyanins found in red cabbage juice, also demonstrate this principle, shifting from red in strong acids to green or blue in bases.