[1] The sensing electrode acts as a platform for electron transfer to or from the reference half cell; it is typically made of platinum, although gold and graphite can be used as well.
Therefore, other more stable reference electrodes such as silver chloride and saturated calomel (SCE) are commonly used because of their more reliable performance.
Although measurement of the redox potential in aqueous solutions is relatively straightforward, many factors limit its interpretation, such as effects of solution temperature and pH, irreversible reactions, slow electrode kinetics, non-equilibrium, presence of multiple redox couples, electrode poisoning, small exchange currents, and inert redox couples.
Nevertheless, reduction potential measurement has proven useful as an analytical tool in monitoring changes in a system rather than determining their absolute value (e.g. process control and titrations).
Similar to how the concentration of hydrogen ions determines the acidity or pH of an aqueous solution, the tendency of electron transfer between a chemical species and an electrode determines the redox potential of an electrode couple.
Like pH, redox potential represents how easily electrons are transferred to or from species in solution.
is defined as the negative logarithm of the free electron concentration in solution, and is directly proportional to the redox potential.
Both pH and redox potentials are properties of solutions, not of elements or chemical compounds themselves, and depend on concentrations, temperature etc.
The more positive the reduction potential the greater the species' affinity for electrons and tendency to be reduced.
Hydrogen (whose reduction potential is 0.0) acts as an oxidizing agent because it accepts an electron donation from the reducing agent lithium (whose reduction potential is -3.04), which causes Li to be oxidized and Hydrogen to be reduced.
The relative reactivities of different half cells can be compared to predict the direction of electron flow.
Any system or environment that accepts electrons from a normal hydrogen electrode is a half cell that is defined as having a positive redox potential; any system donating electrons to the hydrogen electrode is defined as having a negative redox potential.
Surface polarization interferes with measurements, but various sources[citation needed] give an estimated potential for the standard hydrogen electrode of 4.4 V to 4.6 V (the electrolyte being positive).
and pH of a solution are related by the Nernst equation as commonly represented by a Pourbaix diagram (
is the standard Gibbs free energy change, z is the number of electrons involved, and F is Faraday's constant.
The acid-base neutralization of each oxide ion consumes 2 H+ or one H2O molecule as follows: This is why protons are always engaged as reagent on the left side of the reduction reactions as can be generally observed in the table of standard reduction potential (data page).
The ability of an organism to carry out oxidation–reduction reactions depends on the oxidation–reduction state of the environment, or its reduction potential (
) determined under standard conditions (T = 298.15 K = 25 °C = 77 °F; Pgas = 1 atm = 1.013 bar) with the concentration of each dissolved species being taken as 1 M, and thus [ H+] = 1 M and pH = 0.
In the field of environmental chemistry, the reduction potential is used to determine if oxidizing or reducing conditions are prevalent in water or soil, and to predict the states of different chemical species in the water, such as dissolved metals.
[1] The reduction potentials in natural systems often lie comparatively near one of the boundaries of the stability region of water.
In places with limitations in air supply, such as submerged soils, swamps and marine sediments, reducing conditions (negative potentials) are the norm.
[1] In environmental situations, it is common to have complex non-equilibrium conditions between a large number of species, meaning that it is often not possible to make accurate and precise measurements of the reduction potential.
However, it is usually possible to obtain an approximate value and define the conditions as being in the oxidizing or reducing regime.
[1] In the soil there are two main redox constituents: 1) anorganic redox systems (mainly ox/red compounds of Fe and Mn) and measurement in water extracts; 2) natural soil samples with all microbial and root components and measurement by direct method.
[7] A study was conducted comparing traditional parts per million (ppm) chlorination reading and ORP in Hennepin County, Minnesota.
The results of this study presents arguments in favor of the inclusion of ORP above 650 mV in the local health regulation codes.
[8] Eh–pH (Pourbaix) diagrams are commonly used in mining and geology for assessment of the stability fields of minerals and dissolved species.
Although the formation of a mineral or its dissolution may be predicted to occur under a set of conditions, the process may practically be negligible because its rate is too slow.
Nevertheless, the equilibrium conditions can be used to evaluate the direction of spontaneous changes and the magnitude of the driving force behind them.
"Preliminary report on the oxidation-reduction potential obtained on surfaces of gingiva and tongue and in interdental space".