During electrolysis, the electrons from the cathode are bound to H+ ions.
They don't flow around in water as free electrons.
Now that we understand the “P” in “ORP,” we should also discuss the “O,” which refers to “oxidation,” and the “R,” which refers to “reduction.” Oxidation and reduction reactions are chemical reactions in which electrons are transferred between two different species (ions, atoms, or molecules), with one species gaining electrons (electron deficiency) and another donating electrons (excess electrons). Chemists call these reactions “redox” reactions.
Oxidation and reduction reactions always occur together; when one species loses electrons, another must gain those electrons. The species that loses electrons is oxidized and the species that gains electrons is reduced. A simple way to remember oxidation (electron loss) and reduction (electron gain) in relation to electrons is the mnemonic "OIL-RIG" (English: "Oxidation Is Loss, Reduction Is Gain", translated: "Oxidation is loss, reduction is Profit").
Since electron configurations determine the chemical properties of a substance, the transfer of electrons leads to a change in the electron configurations of those involved in the reaction. This creates products that have different chemical properties than the original substance.
The reactions that produce H2 and O2 gas during the electrolysis of water are examples of oxidation and reduction reactions that occur simultaneously at the anode and cathode immersed in the water. During electrolysis, the cathode supplies electrons that reduce hydrogen ions2 into hydrogen molecules. Equation 1 describes the generation of H2 gas at the negative cathode.
The reactants, hydrogen ions (H+) and electrons (e-, which are supplied by the cathode), are on the left side of the equation.
On the right there is only one product, molecular hydrogen gas (H2).
Two hydrogen ions (2H+) undergo reduction because they accept electrons (2e-) from the cathode and form two hydrogen atoms (2H+), which then “pair” to form a hydrogen molecule (H2).
Figure 3 shows a pictorial representation of the reduction reaction of this equation:
Note that before the reaction, the two H+ ions are without electrons (only protons). The end product is the H2 molecule, which shares two electrons. Because electrons are charged particles, this difference in electrons results in an energy potential between the two species.
When dissolved in water, these two parts represent the working potential of the water and, just like the battery in Figure 2, create a corresponding voltage potential (ORP) on the surface of the meter's platinum electrode. Any change in the concentration 3 of either type leads to a change in the voltage potential generated at the ORP electrode.
Before leaving the topic of oxidation and reduction, let us recall what was mentioned that every reduction reaction must be accompanied by a simultaneous oxidation reaction. Equation 2 shows the accompanying oxidation reaction during electrolysis at the positive anode, where hydroxide ions (OH-) are oxidized to produce oxygen gas (O2), water (H2O), and electrons.
These electrons are returned to the anode, which is connected to the positive terminal of the power supply.
Since the waters produced at the anode and cathode are normally isolated from each other with a membrane so that they cannot mix, a positive ORP is measured in the water from the secondary tube (lower sour water tube).
The reason for this is that instead of the H+/H2 redox couple, this water contains another redox couple (usually an oxygen species) which has a positive reduction potential.
Other types of devices, such as the brown gas generators, which do not use a membrane, produce a water containing both dissolved oxygen and hydrogen gas. This water therefore contains both oxidizing and reducing redox couples at a pH close to neutral. Since our discussion of ORP focuses on the reaction at the cathode that produces hydrogen gas, we will not spend time examining the anode reaction in detail.
Excerpt from the book by Randy Sharpe: “The relationship between dissolved H2, pH and redox potential”
