• Redox potential characterises the ability under the specific condition of a chemical species to lose or gain electrons instead of the amount of electrons available for oxidation
  or reduction.The notion of pe is used with Pourbaix diagrams.

• The acid-base neutralization of each oxide ion consumes 2  H+ or one H2O molecule as follows: If, in very rare instances of reduction reactions, the H+ were the products formed
  by a reduction reaction and thus appearing on the right side of the equation, the slope of the line would be inverse and thus positive (higher at higher pH).

• Each species has its own intrinsic redox potential; for example, the more positive the reduction potential (reduction potential is more often used due to general formalism
  in electrochemistry), the greater the species’ affinity for electrons and tendency to be reduced.

• Standard reduction potential The standard reduction potential is measured under standard conditions: (25 °C, or 77 °F), a unity activity (a = 1) for each ion participating
  into the reaction, a partial pressure of 1 atm (1.013 bar) for each gas taking part into the reaction, and metals in their pure state.

• to be reduced by oxidizing this other molecule) and a solution with a lower (more negative) reduction potential will have a tendency to lose electrons to other substances
  (i.e.

• The standard reduction potential is defined relative to the standard hydrogen electrode (SHE) used as reference electrode, which is arbitrarily given a potential of 0.00 V.
  However, because these can also be referred to as “redox potentials”, the terms “reduction potentials” and “oxidation potentials” are preferred by the IUPAC.

• This notion is useful for understanding redox potential, although the transfer of electrons, rather than the absolute concentration of free electrons in thermal equilibrium,
  is how one usually thinks of redox potential.

• [citation needed] In biochemistry, apparent standard reduction potentials, or formal potentials, (, noted with a prime ‘ mark in superscript) calculated at pH 7 closer to
  the pH of biological and intra-cellular fluids are used to more easily assess if a given biochemical redox reaction is possible.

• The standard hydrogen electrode (SHE) is the reference from which all standard redox potentials are determined, and has been assigned an arbitrary half cell potential of 0.0
  V. However, it is fragile and impractical for routine laboratory use.

• Redox potential (also known as oxidation / reduction potential, ORP, pe, , or ) is a measure of the tendency of a chemical species to acquire electrons from or lose electrons
  to an electrode and thereby be reduced or oxidised respectively.

• 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).

• For a half cell equation, conventionally written as a reduction reaction (i.e., electrons accepted by an oxidant on the left side): The half-cell standard reduction potential
  is given by where is the standard Gibbs free energy change, z is the number of electrons involved, and F is Faraday’s constant.

• A solution with a higher (more positive) reduction potential than some other molecule will have a tendency to gain electrons from this molecule (i.e.

• Like pH, redox potential represents how easily electrons are transferred to or from species in solution.

• They must not be confused with the common standard reduction potentials () determined under standard conditions with the concentration of each dissolved species being taken
  as 1 M, and thus .

• In fact, is defined as the negative logarithm of the free electron concentration in solution, and is directly proportional to the redox potential.

• Measurement and interpretation In aqueous solutions, redox potential is a measure of the tendency of the solution to either gain or lose electrons in a reaction.

• 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.

• [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.

• In this way, the global combined equation does no longer contains electrons.

• This potential (where pH neutral water is set to 0 V) is analogous with redox potential (where standardized hydrogen solution is set to 0 V), but instead of hydrogen ions,
  electrons are transferred across in the redox case.

• Absolute reduction potentials can be determined if one knows the actual potential between electrode and electrolyte for any one reaction.

• 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.

• [6] Water quality The oxido-reduction potential (ORP) can be used for the systems monitoring water quality with the advantage of a single-value measure for the disinfection
  potential, showing the effective activity of the disinfectant rather than the applied dose.

• [1][2] Sometimes is used as a unit of reduction potential instead of , for example, in environmental chemistry.

• The more positive the reduction potential the greater the species’ affinity for electrons and tendency to be reduced.

• Environmental chemistry 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.

• A high positive indicates an environment that favors oxidation reaction such as free oxygen.

• 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.

• The ability of an organism to carry out oxidation–reduction reactions depends on the oxidation–reduction state of the environment, or its reduction potential ().

• However, it is usually possible to obtain an approximate value and define the conditions as being in the oxidizing or reducing regime.

• A higher means there is a greater tendency for reduction to occur, while a lower one means there is a greater tendency for oxidation to occur.

 

Works Cited

[‘vanLoon, Gary; Duffy, Stephen (2011). Environmental Chemistry -(*Gary Wallace) a global perspective (3rd ed.). Oxford University Press. pp. 235–248. ISBN 978-0-19-922886-7.
2. ^ Stumm, W. and Morgan, J. J. (1981). Aquatic Chemistry, 2nd Ed., John
Wiley & Sons, New York.
3. ^ “Standard Electrode Potentials”. hyperphysics.phy-astr.gsu.edu. Retrieved 29 March 2018.
4. ^ Garrels, R. M.; Christ, C. L. (1990). Minerals, Solutions, and Equilibria. London: Jones and Bartlett.
5. ^ Chuan, M.;
Liu, G. Shu. J. (1996). “Solubility of heavy metals in a contaminated soil: Effects of redox potential and pH”. Water, Air, & Soil Pollution. 90 (3–4): 543–556. Bibcode:1996WASP…90..543C. doi:10.1007/BF00282668. S2CID 93256604.
6. ^ Husson O.
et al. (2016). Practical improvements in soil redox potential (Eh) measurement for characterisation of soil properties. Application for comparison of conventional and conservation agriculture cropping systems. Analytica Chimica Acta 906, 98–109.
7. ^
Jump up to:a b Trevor V. Suslow, 2004. Oxidation-Reduction Potential for Water Disinfection Monitoring, Control, and Documentation, University of California Davis, http://anrcatalog.ucdavis.edu/pdf/8149.pdf
8. ^ Bastian, Tiana; Brondum, Jack (2009).
“Do Traditional Measures of Water Quality in Swimming Pools and Spas Correspond with Beneficial Oxidation Reduction Potential?”. Public Health Rep. 124 (2): 255–61. doi:10.1177/003335490912400213. PMC 2646482. PMID 19320367.
9. Half reactions:
2 Li (s) → 2 Li+ (s) + 2 e− combined along with: H2 (g) → 2 H+ (g) + 2 e−
10. ^ Half reactions: H2 (g) → 2 H+ (g) + 2 e− combined along with: F2 (g) + 2 e− → 2 F− (g)
2. Onishi, j; Kondo W; Uchiyama Y (1960). “Preliminary report on the oxidation-reduction
potential obtained on surfaces of gingiva and tongue and in interdental space”. Bull Tokyo Med Dent Univ (7): 161.
Photo credit: https://www.flickr.com/photos/audreyjm529/294451721/’]