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This article relies largely or entirely upon a single source. Relevant discussion may be found on the talk page. Please help improve this article by introducing citations to additional sources. (October 2011) |
Reduction potential (also known as redox potential, oxidation / reduction potential, ORP, pE, ε, or ) 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]
Contents
- 1 Measurement and interpretation
- 2 Explanation
- 3 Standard reduction potential
- 4 Converting potentials between different types of reference electrodes
- 5 Half cells
- 6 Nernst equation
- 7 Biochemistry
- 8 Environmental chemistry
- 9 Water quality
- 10 Geology
- 11 See also
- 12 References
- 13 Additional notes
- 14 External links
Measurement and interpretation[edit]
In aqueous solutions, the 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. 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[edit]
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 an intensity factor. It does not characterize the capacity of the system for oxidation or reduction, in much the same way that pH does not characterize the buffering capacity.[clarification needed]
In fact, it is possible to define pE, logarithm of electron concentration 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]
The actual value of electron concentration is not measurable as it is of order of 10^-55 M, yet it is thermodynamically a well defined quantity. What matters is the logarithm of the quotient of concentration of electrons in one solute with respect to the other, and that is, via Nernst equation, proportional to the redox potential. 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[edit]
See also: Standard hydrogen electrode
The standard reduction potential () is measured under standard conditions: 25°C, a 1 M concentration for each ion participating in the reaction, a partial pressure of 1 atm for each gas that is part of the reaction, and metals in their pure state. The standard reduction potential is defined relative to a standard hydrogen electrode (SHE) reference electrode, which is arbitrarily given a potential of 0.00 volts. Historically, many countries, including the United States and Canada[citation needed], used standard oxidation potentials rather than reduction potentials in their calculations. These are simply the negative of standard reduction potentials, so it is not a major problem in practice. However, because these can also be referred to as "redox potentials", the terms "reduction potentials" and "oxidation potentials" are preferred by the IUPAC. The two may be explicitly distinguished in symbols as and .
Converting potentials between different types of reference electrodes[edit]
See also: Saturated calomel electrode
Often a reduction potential is quoted as measured against a different reference electrode than the one desired and it becomes necessary to convert to the desired reference potential. Alternatively, it may be necessary to convert measurements to the standard reduction potential for reporting purposes. This is easily done by recognizing that the observed potential represents the difference between the potential at the sensing electrode and the potential at the reference electrode, i.e.
Where is the observed reaction, is the reference used in experiment, and is the reference that is desired. The voltage relationships for several different reference electrodes at 25 °C can be interrelated as follows:
| Reference electrode |
Electrode potential with respect to SHE (mV) |
| Standard hydrogen electrode (SHE) |
0 |
| Saturated calomel electrode (SCE) |
+241[citation needed] |
| Ag/AgCl, 1 M KCl |
+192[citation needed] |
| Ag/AgCl, 4 M KCl |
+228[citation needed] |
| Ag/AgCl, sat. KCl |
+236[citation needed] |
For example, if one measured 300 mV using a saturated KCl Ag/AgCl reference(ref2) and wanted to refer it to the standard reduction potential () measured using a SHE reference electrode (ref1), then 197 mV should be added to the 300 mV to obtain 497 mV, since
it follows that
and therefore
Likewise, if one measured 300 mV using a saturated KCl Ag/AgCl reference (ref2) and wanted to determine the corresponding measurement using an SCE reference (ref1), then given
it follows that
and therefore
Half cells[edit]
The relative reactivities of different half cells can be compared to predict the direction of electron flow. 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.
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. is measured in millivolts (mV). A high positive indicates an environment that favors oxidation reaction such as free oxygen. A low negative 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 H2(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[edit]
Main article: Nernst equation
The and pH of a solution are related. For a half cell equation, conventionally written as reduction (electrons on the left side):
The half cell standard potential is given by:
where is the standard Gibbs free energy change, is the number of electrons involved, and is Faraday's constant. The Nernst equation relates pH and :
where curly brackets indicate activities and exponents are shown in the conventional manner. This equation is the equation of a straight line for as a function of pH with a slope of volt (pH has no units.) This equation predicts lower 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 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−]]
where = −1.1819 − 0.0885 log[HFeO−
2] + 0.0296 pH. Note that the slope of the line is −1/2 the −0.05916 value above, since = −1/2.
Biochemistry[edit]
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 ().
Strictly aerobic microorganisms are generally active at positive values, whereas strict anaerobes are generally active at negative values. Redox affects the solubility of nutrients, especially metal ions.[citation needed]
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 to.[citation needed]
Environmental chemistry[edit]
See also: Pourbaix diagram
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. pE values in water ranges from -12 to 25; the levels where the water itself becomes reduced or oxidized, respectively.[2]
The reduction potentials in natural systems often lie comparatively near one of the boundaries of the stability region of water. Aerated surface water, rivers, lakes, oceans, rainwater and acid mine water, usually have oxidizing conditions (positive potentials). In places with limitations in air supply, such as submerged soils, swamps and marine sediments, reducing conditions (negative potentials) are the norm. Intermediate values are rare and usually a temporary condition found in systems moving to higher or lower pE values.[2]
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.[2]
Water quality[edit]
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]
Geology[edit]
See also: Pourbaix diagram
Eh-pH (Pourbaix) diagrams are commonly used in mining and geology for assessment of the stability fields of minerals and dissolved species. Under conditions where a mineral (solid) phase is the most stable form of an element, these diagrams show that mineral. As with results from all thermodynamic (equilibrium) evaluations, these diagrams should be used with caution. Although the formation of a mineral or its dissolution may be predicted to occur under a set of conditions, the process may be negligible because its rate is so slow. Under those circumstances, kinetic evaluations are necessary. However, the equilibrium conditions can be used to evaluate the direction of spontaneous changes and the magnitude of the driving force behind them.
See also[edit]
- Galvanic cell
- Electrolytic cell
- Electromotive force
- Electrochemical potential
- Standard electrode potential
- Table of standard electrode potentials
- Oxygen radical absorbance capacity
- Redox
References[edit]
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This article needs additional citations for verification. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. (May 2008) |
- ^ a b c 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
- ^ a b c d e f vanLoon, Gary; Duffy, Stephen (2011). Environmental Chemistry - 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.
Additional notes[edit]
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[edit]
- Redox potential definition
- Large table of potentials (Site broken. Archived version on the Internet Archive.)
