Indicator Method for pH Determination: Advantages and Limitations

Water is a weak electrolyte and dissociates according to the equation: H₂O ↔ H⁺ + OH^–. In pure water at 25 °C, the concentrations of [H⁺] and [OH^–] are equal. In this case, the solution has a neutral reaction. When an acid is added to water, the concentration of hydrogen ions increases, while the concentration of hydroxide ions decreases, i.e., [H⁺] > [OH^–]. Such a solution has an acidic reaction. When a base is added, the opposite occurs-the hydroxide ion concentration rises, and the hydrogen ion concentration decreases, i.e., [OH^–] > [H⁺], and such a solution has an alkaline reaction.

For convenience in expressing the reaction medium of aqueous solutions, a special value-pH-was introduced. The hydrogen index (pH) is defined as the negative decimal logarithm of the hydrogen ion concentration:

pH = –lg[H⁺],
where [H⁺] is the hydrogen ion concentration, mol/L.

The nature of the medium can be determined qualitatively using acid-base indicators, semi-quantitatively using indicator strips, and quantitatively via potentiometric methods.

Acid-base indicators are organic compounds, weak electrolytes, exhibiting either acidic or basic properties:
HInd ↔ H⁺ + Ind^–,
where HInd is the molecular form of the indicator,
Ind^– is the ionic form of the indicator.

The higher the concentration of H⁺ ions, the lower the degree of indicator ionization. The molecular (HInd) and ionic (Ind^–) forms of the indicator have different colors. If the equilibrium shifts to the right, the solution predominantly exhibits the color of the ionic form. If the equilibrium shifts to the left, the color of the molecular form dominates. For example, methyl orange at pH = 3.1 exists in its molecular form, which is red. When the pH changes to 4.4, the ionic form (yellow) prevails. At a certain ratio of molecular and ionic forms, the solution may display an intermediate (transitional) color due to the mixing of the two forms. Thus, at pH = 3.8, methyl orange appears orange-red. This is why indicators change color gradually rather than abruptly, within a specific pH range. Studies show that the pH range over which an individual indicator changes color is approximately 2 logarithmic units.

In addition to individual indicators, mixed and universal indicators are used, allowing pH measurements over broader logarithmic ranges.

Delta-CT LLC manufactures and sells indicator strips for determining the pH of aqueous solutions across various logarithmic ranges. The strips consist of polymer strips with test zones containing a mixture of acid-base indicators. When the indicator zone is immersed in a solution of unknown pH, it develops a specific color. By comparing this color to a reference scale, the pH of the solution is determined.

Delta-CT’s indicator strips are a simple, rapid means of assessing aqueous pH. They offer higher accuracy compared to paper strips, are cost-effective, require no expensive instruments or specialized equipment, and do not demand highly skilled operators. Additionally, their hydrophobic plastic holder design allows safe testing of aggressive and toxic solutions.

When measuring pH with indicator methods, potential errors leading to inaccurate results must be considered.

Measuring non-buffered solutions may introduce “acid” or “alkaline” errors, as the indicator itself (a weak acid or base) alters the solution’s reaction. These errors can be significant-e.g., in pure water, the acid error may exceed 1 pH unit. Buffered solutions do not exhibit such errors.

Measurement accuracy also depends on salt concentration. At ~2M salt concentrations, a “salt error” arises due to the influence of ionic strength on indicator solubility and dissociation. Typically, salt errors do not exceed 0.2 pH units.

In protein-containing solutions (e.g., physiological fluids), a “protein error” may distort pH readings due to indicator adsorption by proteins. At low protein levels, this error is minor, usually ≤0.2–0.3 pH units.

Alcohols in the solution can cause an “alcohol error” in colorimetric pH measurements. Alcohol significantly affects indicator dissociation constants, shifting their transition intervals. For instance, phenolphthalein in 70% alcohol exhibits an error of up to 2.2 pH units.

Temperature fluctuations introduce “temperature errors” by altering indicator dissociation constants. For example, p-nitrophenol has pK = 7.3 at 0°C but 6.81 at 50°C.

Technical errors may also occur due to poor lighting, subjective color perception, etc.

Despite these limitations, the colorimetric pH method using indicator strips offers simplicity, accessibility, and eliminates the need for specialized equipment.