The electrodes used in potentiometric analysis are referred to as galvanic cells. A galvanic cell is a system that converts chemical reaction energy into electrical energy. The voltage of such a cell is known as the electromotive force (EMF). This EMF is composed of two half-cells: one is the measuring electrode, whose potential is related to the activity of a specific ion; the other is the reference half-cell, commonly called the reference electrode, which is typically in contact with the solution being measured and connected to the measuring instrument.

For example, an electrode consists of a silver wire immersed in a salt solution containing silver ions. At the interface between the metal wire and the solution, the difference in the activity of silver ions in the two phases—metal and solution—gives rise to an ion-exchange charging process, resulting in a finite potential difference. Silver ions that have lost electrons enter the solution. If no external current is applied for countercharging—that is, if no current flows—this process will eventually reach equilibrium. The voltage that exists in this equilibrium state is referred to as the half-cell potential or the electrode potential.

Such electrodes, as described above, consisting of a metal and a solution containing ions of that metal, are referred to as first-class electrodes.

The measurement of this potential is carried out relative to a reference electrode whose potential is independent of the composition of the salt solution. Such a reference electrode with an intrinsic, stable potential is also referred to as a secondary electrode. In these electrodes, the metallic conductor is coated with a sparingly soluble salt of that metal (e.g., Ag/AgCl) and is immersed in an electrolyte solution containing the anion of that metal salt. Under these conditions, the value of the half-cell potential or electrode potential depends on the activity of the relevant anion.

The voltage between these two electrodes obeys the Nernst equation:

E=	E0+	R·T	·In a+
		n·F

In the formula:

E—Potential

E ₀ — Standard electrode voltage

R —— gas constant (8.31439 joules per mole per degree Celsius)

T —— Absolute temperature in Kelvin (e.g., 20°C = 273 + 293 Kelvin)

F — Faraday constant (96,493 coulombs per equivalent)

n — the oxidation state of the ion being measured (silver = 1, hydrogen = 1)

a ⁺ — Activity of ions

The standard hydrogen electrode serves as the reference point for all potential measurements. It consists of a platinum wire electroplated (or coated) with platinum chloride and surrounded by hydrogen gas at a fixed pressure of 1013 hPa.

Immerse this electrode in H at 25°C. ₃ O ⁺ In a solution with an ion concentration of 1 mol/L, the half-cell potential or electrode potential that serves as the reference for all electrochemical potential measurements is established. Since the hydrogen electrode is difficult to implement as a reference electrode in practice, second-type electrodes are used instead. The most commonly employed of these is the silver/silver chloride electrode, which responds to changes in chloride-ion concentration via the dissolution of AgCl.

The electrode potential of this reference electrode is maintained constant by a saturated KCl reservoir (e.g., 3 mol/L KCl). The electrolyte solution, in liquid or gel form, is connected to the solution under measurement via a porous separator.

Using the aforementioned electrode combination—silver electrode and Ag/AgCl reference electrode—it is possible to measure the silver ion concentration in film developing solutions. Alternatively, the silver electrode can be replaced with a platinum or gold electrode to measure redox potential, for example, during the oxidation of a particular metal ion.

The most familiar and widely used pH indicator electrode is the glass electrode. It consists of a glass tube with a pH-sensitive glass membrane blown onto one end. The interior of the tube is filled with a 3 mol/L KCl buffer solution saturated with AgCl, which has a pH of 7. The potential difference across the two surfaces of the glass membrane, which reflects the pH value, is transmitted via an Ag/AgCl reference system, functioning as the second electrode.

This potential difference also obeys the Nernst equation:

E=	E ₀+	R·T	·In a H3O+
		n·F

E = 59.16 mV/25°C per pH, where R and F are constants and n is the valence; each ion has a fixed value for n. For the hydrogen ion, n = 1. Temperature “T” is a variable that plays a significant role in the Nernst equation: as temperature increases, the potential value also increases. Specifically, for every 1°C rise in temperature, the potential changes by 0.2 mV per pH unit. Expressed in terms of pH, this corresponds to a change of 0.0033 pH units per 1°C. In other words, for measurements conducted between 20°C and 30°C at around pH 7, temperature compensation is not required; however, for applications where the temperature exceeds 30°C or falls below 20°C, or where the pH is greater than 8 or less than 6, temperature compensation is essential.

Figure 1: Relationship among pH, potential, and ion concentration