Redox (ORP) Measurement Theory

Redox measurements and pH are both scales of proton H+ activity. Redox potential, on the other hand, reflects electron activity. In both cases, the measured voltage corresponds to the activity of the relevant chemical species; the only difference is that one is expressed in pH units, while the other is expressed in millivolts (mV).

For both types of measurements, the voltage generated by the potential difference is determined using the zero-current method. Although the sensitive noble-metal electrodes used in redox measurements are low-resistance conductors, the Re voltage must still be measured with a high-impedance amplifier.

In addition, the two share the same electrode technology, and the changes in the measured potential both obey the Nernst equation.

The meaning of REDOX or (ORP)

These two acronyms are easy to translate. REDOX is formed from the initial letters of REDuction and OXidation, while ORP is formed from the initial letters of Oxidation Reduction Potential.

The fundamental principles underlying redox measurements are, to varying degrees, not well understood by users.

For example, hydrochloric acid (HCl) contains H+ protons and carries a positive unit charge; accordingly, it is paired with a negatively charged ion, in this case Cl−. The carrier of the negative charge is the electron.

When zinc or iron is placed in hydrochloric acid, hydrogen gas is produced, and the iron is converted into iron oxide. A molecular transformation has thus taken place: the H+ ions, which have combined with Cl–, are reduced to form hydrogen gas. So what changes occur to the hydrochloric acid and the iron?

Each element and each compound has its own electrochemical potential (relative to the standard hydrogen electrode potential). Substances with higher potentials—or more negative polarity—will donate electrons to substances with lower potentials—or more positive polarity. The substance that accepts electrons is reduced, meaning it transitions from its ionic form to the gaseous form (in the case of hydrogen) or to the metallic form (in the case of metal ions), while the substance that loses electrons is converted back into its ionic form.

In our example

Hydrogen (H ⁺ The potential of ) is 0.0 V (in free-state hydrochloric acid), while the potential of iron (Fe) is −0.4 V.

Iron donates two electrons and is converted into an ionic form. The hydrogen ion, in turn, accepts these two electrons and is converted into gaseous hydrogen—hydrogen gas.

The above process can be seen more clearly through the reaction equation.

Fe → Fe ²⁺ + 2e

2e + 2H ⁺ → H ₂

――――――――――――――――

Fe + 2H ⁺ →Fe ²⁺ + H ₂

A substance that accepts electrons is reduced, meaning it transitions from an oxidized state to a reduced state. A substance that loses electrons is oxidized, meaning it transitions from a reduced state to an oxidized state. Below is another example:

The standard electrode potential of zinc is −0.76 V, while that of copper is +0.337 V.

When a zinc rod is immersed in a copper sulfate solution, copper deposits on the zinc rod without the application of an external electric current.

(-0.76 V) Zn → Zn ²⁺ +2e

2e + Cu ²⁺ →Cu (+0.337 V)

――――――――――――――――――――――――――――

Zn + Cu ²⁺ →Zn ²⁺ + Copper

Zinc donates two electrons to copper ions, reducing the copper ions to metallic copper. When a potential difference exists, a current flows from the higher potential to the lower potential.

The potential difference is the electromotive force (EMF).

To illustrate this phenomenon, let us imagine two tanks, one of which is filled with water. The bottoms of the two tanks are connected by a pipe. When the valve is opened, water flows from the full tank into the empty one. From this observation, we can conclude that there is no reduction process without an accompanying oxidation process, nor is there an oxidation process without a corresponding reduction process. In general, the voltage range for oxidation processes spans from +3 V to −3 V. The greater the potential difference, the stronger the reducing power—from high potential to low potential—or, conversely, the stronger the oxidizing power—from low potential to high potential. However, the relative strengths of oxidizing and reducing properties are not absolute; they depend on the following factors:

a) Concentration of the reference medium

b) Magnitude of the potential difference

c) pH value

It is evident from the Nernst equation that the concentration “C”, the number of electrons “n”, and the pH value “E” ₀ “How does it work?”