Acid Base Indicators
In our study of acids and bases, we use many methods to determine the pH of a solution, and to determine the point of equivalence when mixing acids and bases. These methods range from the use of litmus paper, indicator paper Figure 2, specifically designed electrodes, and the use of colored molecules in solution. Other than the electrodes, all of the methods are visual, and rely on some fundamental changes that occur in a molecule when the pH of its environment changes.
In general, a molecule that changes color with the pH of the environment it is in can be used as an indicator. In the equation
the weak acid HIn is shown in equilibrium with its ionized anion In- -- in the case of the indicator methyl orange, the HIn is colored red and the ionized In- form is yellow. The structures are shown in Figure 3.
In this example,
For methyl orange, Ka = 1.6 X 10-4 and pKa = 3.8. The neutral (red) and dissociated (yellow) forms of the indicator are present at equal concentrations when the pH = 3.8. The eye is sensitive to color changes over a range of concentration ratios of approximately 100 or over two pH units. Below pH 2.8 is a solution containing methyl orange is red, and above approximately 4.8 it is clearly yellow.
pH indicators are frequently employed in titrations in analytical chemistry and biology to determine the extent of a chemical reaction. Because of the subjective choice (determination) of color, pH indicators are susceptible to imprecise readings. For applications requiring precise measurement of pH, a pH meter is frequently used. Sometimes a blend of different indicators is used to achieve several smooth color changes over a wide range of pH values. These commercial indicators (e.g., universal indicator and Hydrion papers) are used when only rough knowledge of pH is necessary. Tabulated below Figure 1 are several common laboratory pH indicators. Indicators usually exhibit intermediate colors at pH values inside the listed transition range. For example, phenol red exhibits an orange color between pH 6.8 and pH 8.4. The transition range may shift slightly depending on the concentration of the indicator in the solution and on the temperature at which it is used.
When viewed on the pH scale itself, the color transitions as determined by their transition ranges becomes clearer, and the context of the indicator sensitivity over ranges of pH is laid out more informatively.