2.5.1. How Indicators Work

Consider a generic indicator HIn (a weak acid): HIn(aq) ⇌ H plus + In minus

• HIn has one color (for example, red).
• In minus has another color (for example, yellow).
• The apparent color of the solution depends on the ratio [In minus] ÷ [HIn].

Indicators are substances that change color depending on the acidity or basicity of a solution. For example, HIn represents the acid form of the indicator. When we add it to a solution, it can either remain in its acid form (HIn) or dissociate into its base form (In minus). The color seen in the solution will depend on the balance of these two forms. The ratio between the colored forms dictates the visible color. If the concentration of the hydrogen ions ([H plus]) in the solution is higher than a certain value (related to the indicator's dissociation constant, Ki), then the equilibrium will shift to the left, favoring the HIn form, and we will see its color (e.g., red). Conversely, if the [H plus] is lower than Ki, the equilibrium shifts to the right, and we see the other color (e.g., yellow).

Imagine you're trying to determine how ripe a tomato is by its color. A perfectly ripe tomato might be red, while an unripe tomato could be green. Similarly, indicators operate on a color shift based on their environment—just like the tomato varies in color based on ripeness, the indicator changes color based on the acidity of the solution.

At pH values where [H plus] is much greater than the indicator’s dissociation constant Ki, the equilibrium lies far left (dominant form is HIn, color = red). At pH values where [H plus] is much smaller than Ki, equilibrium lies far right (dominant form is In minus, color = yellow).

The pH of a solution affects the ionization of indicators. When there is a high concentration of hydrogen ions, this favors the left side of the equilibrium equation, meaning that the acidic form HIn will dominate and thus reveal its color (red). Conversely, at low pH (or high pH in this case), the reaction will shift to the right, favoring the In minus form which presents its unique color (yellow). This dynamic shows how indicators provide a visual representation of changes in pH through color change.

Think of preparing a drink that changes color based on temperature. For example, if it’s cold, it might be a specific shade of blue, but when it warms up, it could turn pink. This serves a purpose like our pH indicator: it visually signals a change—just as our indicator changes color with pH levels.

The pH at which the indicator is half-dissociated ([HIn] = [In minus]) satisfies: Ki = [H plus] Taking negative logarithms: pKi = pH

Typically, the color change range is pKi ± 1 pH unit (because at [HIn] : [In minus] ratios around 10:1 or 1:10, one color is just overtaken by the other).

The point where an indicator transitions between its two colors occurs at a specific pH called the pKi. At this pH, the concentrations of the acid form and the base form of the indicator are equal. Since the logarithmic nature of pH reflects hydrogen ion concentration, the equation reveals that for a given indicator at pKi, you can expect it to show significant color change at values within approximately one unit of pH on either side of this point. This informs chemists about the effective use of indicators for determining the pH during titrations.

Consider a traffic light system: the transitions between red (stop) and green (go) are most critical at the yellow phase, which serves as the point of caution. In chemistry, pKi functions similarly; at values near this point, we see the most significant visual change in the indicator, helping us make critical decisions based on color.

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