Reduction potential

Reduction potential (also known as redox potential, oxidation / reduction potential, ORP, pE, ε, or $E_{{h}}$ ) is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. Reduction potential is measured in volts (V), or millivolts (mV). Each species has its own intrinsic reduction potential; the more positive the potential, the greater the species' affinity for electrons and tendency to be reduced. ORP is a common measurement for water quality.^[1]

Measurement and interpretation

In aqueous solutions, reduction potential is a measure of the tendency of the solution to either gain or lose electrons when it is subject to change by introduction of a new species. A solution with a higher (more positive) reduction potential than the new species will have a tendency to gain electrons from the new species (i.e. to be reduced by oxidizing the new species) and a solution with a lower (more negative) reduction potential will have a tendency to lose electrons to the new species (i.e. to be oxidized by reducing the new species). Because the absolute potentials are difficult to accurately measure, reduction potentials are defined relative to a reference electrode. Reduction potentials of aqueous solutions are determined by measuring the potential difference between an inert sensing electrode in contact with the solution and a stable reference electrode connected to the solution by a salt bridge.^[2]

The sensing electrode acts as a platform for electron transfer to or from the reference half cell. It is typically platinum, although gold and graphite can be used as well. The reference half cell consists of a redox standard of known potential. 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 mV. However, it is fragile and impractical for routine laboratory use. 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 reduction 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. Consequently, practical measurements seldom correlate with calculated values. 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).

Explanation

Just as the transfer of hydrogen ions between chemical species determines the pH of an aqueous solution, the transfer of electrons between chemical species determines the reduction potential of an aqueous solution. Like pH, the reduction potential represents how strongly electrons are transferred to or from species in solution. It does not characterize the amount of electrons available for oxidation or reduction, in much the same way that pH does not characterize the buffering capacity.

In fact, it is possible to define pE, the negative logarithm of electron concentration (-log[e]) in a solution, which will be directly proportional to the redox potential.^[2]^[3] Sometimes pE is used as a unit of reduction potential instead of Eh, for example in environmental chemistry.^[2] If we normalize pE of hydrogen to zero, we will have the relation pE=16.9 Eh at room temperature. This point of view 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. Theoretically, however, the two approaches are equivalent.

Conversely, one could define a potential corresponding to pH as a potential difference between a solute and pH neutral water, separated by porous membrane (that is permeable to hydrogen ions). Such potential differences actually do occur from differences in acidity on biological membranes. This potential (where pH neutral water is set to 0V) is analogous with redox potential (where standardized hydrogen solution is set to 0V), but instead of hydrogen ions, electrons are transferred across in the redox case. Both pH and redox potentials are properties of solutions, not of elements or chemical compounds per se, and depend on concentrations, temperature etc.

Standard reduction potential

Converting potentials between different types of reference electrodes

Reference electrode	Electrode potential with respect to SHE (mV)
Standard hydrogen electrode (SHE)	0
Saturated calomel electrode (SCE)	+241
Ag/AgCl, 1 M KCl	+192
Ag/AgCl, 4 M KCl	+228
Ag/AgCl, sat. KCl	+236

Half cells

The relative reactivities of different half cells can be compared to predict the direction of electron flow. A higher $E_{0}$ means there is a greater tendency for reduction to occur, while a lower one means there is a greater tendency for oxidation to occur.

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. $E_{{h}}$ is measured in millivolts (mV). A high positive $E_{{h}}$ indicates an environment that favors oxidation reaction such as free oxygen. A low negative $E_{{h}}$ indicates a strong reducing environment, such as free metals.

Sometimes when electrolysis is carried out in an aqueous solution, water, rather than the solute, is oxidized or reduced. For example, if an aqueous solution of NaCl is electrolyzed, water may be reduced at the cathode to produce H_2(g) and OH⁻ ions, instead of Na⁺ being reduced to Na_(s), as occurs in the absence of water. It is the reduction potential of each species present that will determine which species will be oxidized or reduced.

Absolute reduction potentials can be determined if we find the actual potential between electrode and electrolyte for any one reaction. Surface polarization interferes with measurements, but various sources give an estimated potential for the standard hydrogen electrode of 4.4 V to 4.6 V (the electrolyte being positive.)

Half-cell equations can be combined if one is reversed to an oxidation in a manner that cancels out the electrons to obtain an equation without electrons in it.

Nernst equation

Main article: Nernst equation

The $E_{{h}}$ and pH of a solution are related. For a half cell equation, conventionally written as reduction (electrons on the left side):

$aA+bB+n[e^{{-}}]+h[H^{{+}}]=cC+dD$

The half cell standard potential $E_{0}$ is given by:

$E_{{0}}({\textrm {volts}})=-{\frac {\Delta G^{\ominus }}{nF}}$

where $\Delta G^{\ominus }$ is the standard Gibbs free energy change, $n$ is the number of electrons involved, and $F$ is Faraday's constant. The Nernst equation relates pH and $E_{{h}}$ :

$E_{{h}}=E_{{0}}+{\frac {0.05916}{n}}\log \left({\frac {\{A\}^{{a}}\{B\}^{{b}}}{\{C\}^{{c}}\{D\}^{{d}}}}\right)-{\frac {0.05916h}{n}}{\text{pH}}$

where curly brackets indicate activities and exponents are shown in the conventional manner. This equation is the equation of a straight line for $E_{{h}}$ as a function of pH with a slope of $-0.05916h/n$ volt (pH has no units.) This equation predicts lower $E_{{h}}$ at higher pH values. This is observed for reduction of O₂ to OH⁻ and for reduction of H⁺ to H₂. If H⁺ were on the opposite side of the equation from H⁺, the slope of the line would be reversed (higher $E_{{h}}$ at higher pH). An example of that would be the formation of magnetite (Fe₃O₄) from HFeO−
2 (aq):^[4]

3 HFeO−
2 + H⁺ = Fe₃O₄ + 2 H₂O + 2 [[e<sup>−</sup>]]

where $E_{{h}}$ = −1.1819 − 0.0885 log([HFeO−
2]³) + 0.0296 pH. Note that the slope of the line is −1/2 the −0.05916 value above, since $h/n$ = −1/2.

Biochemistry

Many enzymatic reactions are oxidation-reduction reactions in which one compound is oxidized and another compound is reduced. The ability of an organism to carry out oxidation-reduction reactions depends on the oxidation-reduction state of the environment, or its reduction potential ( $E_{{h}}$ ).

Strictly aerobic microorganisms are generally active at positive $E_{{h}}$ values, whereas strict anaerobes are generally active at negative $E_{{h}}$ values. Redox affects the solubility of nutrients, especially metal ions.^[5]

There are organisms that can adjust their metabolism to their environment, such as facultative anaerobes. Facultative anaerobes can be active at positive Eh values, and at negative Eh values in the presence of oxygen bearing inorganic compounds, such as nitrates and sulfates.

Direct redox potential (Eh) recording in plant tissue. Measurement can be done in aerobic conditions or in hypoxia. Analytical Pt plate electrode 5x5 mm in form of lancet or needle shaped electrode and saturated calomel standard electrode (+244 mV) were used. For aerobic conditions the Pt electrode was inserted into the plant tissue and calomel electrode beside it on moist filter paper The Eh reading began to sink to a bottom turn point which was declared as Eh of a particular place. For measurement in hypoxia a potato tuber slice was put in a glass beaker and overflooded with deionised water. Full text can be found on www.vukrom.cz/contacts/Benada

Environmental chemistry

Water quality

Oxidation reduction potential (ORP) can be used for water system monitoring with the benefit of a single-value measure of the disinfection potential, showing the activity of the disinfectant rather than the applied dose.^[1] For example, E. coli, Salmonella, Listeria and other pathogens have survival times of under 30 s when the ORP is above 665 mV, compared against >300 s when it is below 485 mV.^[1]

A study was conducted comparing traditional parts per million chlorination reading and ORP in Hennepin County, Minnesota. The results of this study argue for the inclusion of ORP above 650mV in local health codes.^[6]

Geology

References

1 2 3 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
1 2 3 4 5 6 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.
↑ 1981 Stumm, W. and Morgan, J. J. (1981): Aquatic chemistry, 2nd Ed.; John Wiley & Sons, New York
↑ Garrels, R.M.; Christ, C.L. (1990). Minerals, Solutions, and Equilibria. London: Jones and Bartlett.
↑ 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. doi:10.1007/BF00282668.
↑ Bastian, Tiana; Brondum, Jack. "Do Traditional Measures of Water Quality in Swimming Pools and Spas Correspond with Beneficial Oxidation Reduction Potential?". Retrieved 21 January 2014.

Additional notes

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.

External links

Redox potential definition
Large table of potentials (Site broken. Archived version on the Internet Archive.)

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