Reduction potential

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Redox potential (also known as oxidation / reduction potential, ORP, pe, ε, or ${\displaystyle E_{h}}$) 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. Redox potential is measured in volts (V), or millivolts (mV). 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. ORP can reflect the antimicrobial potential of the water.

Measurement and interpretation

In aqueous solutions, redox potential is a measure of the tendency of the solution to either gain or lose electrons when it is subjected 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 next to impossible 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.

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 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. 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

Similar to the concentration of hydrogen ion 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. 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. 

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. Sometimes pe is used as a unit of reduction potential instead of ${\displaystyle E_{h}}$, for example in environmental chemistry. If we normalize pe of hydrogen to zero, we will have the relation pe=16.9 ${\displaystyle E_{h}}$ 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 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. 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

The standard reduction potential (${\displaystyle E_{0}}$) is measured under standard conditions: 25 °C, a 1 activity for each ion participating in the reaction, a partial pressure of 1 bar 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 V. 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 ${\displaystyle E_{0}^{r}}$ and ${\displaystyle E_{0}^{o}}$.

Half cells

The relative reactivities of different half cells can be compared to predict the direction of electron flow. A higher ${\displaystyle E_{0}}$ 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. ${\displaystyle E_{h}}$ is measured in millivolts (mV). A high positive ${\displaystyle E_{h}}$ indicates an environment that favors oxidation reaction such as free oxygen. A low negative ${\displaystyle E_{h}}$ 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

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

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

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

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

where ${\displaystyle \Delta G^{\ominus }}$ 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 ${\displaystyle E_{h}}$:

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

where curly brackets indicate activities and exponents are shown in the conventional manner. This equation is the equation of a straight line for ${\displaystyle E_{h}}$ as a function of pH with a slope of ${\displaystyle -0.05916h/n}$ volt (pH has no units.) This equation predicts lower ${\displaystyle E_{h}}$ 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 ${\displaystyle E_{h}}$ at higher pH). An example of that would be the formation of magnetite (Fe₃O₄) from HFeO⁻
_2 (aq):

3 HFeO⁻
₂ + H⁺ = Fe₃O₄ + 2 H₂O + 2 [[e⁻]]

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 (${\displaystyle E_{h}}$). 

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

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.

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. pe values in water range from -12 to 25; the levels where the water itself becomes reduced or oxidized, respectively.

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.

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.

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 [Husson O. et al.: Practical improvements in soil redox potential (${\displaystyle E_{h}}$) measurement for characterisation of soil properties. Application for comparison of conventional and conservation agriculture cropping systems. Anal. Chim. Acta 906 (2016): 98-109].

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. 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.

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.

Geology

Eh-pH (Pourbaix) diagrams are commonly used in mining and geology for assessment of the stability fields of minerals and dissolved species. Under the conditions where a mineral (solid) phase is predicted to be the most stable form of an element, these diagrams show that mineral. As the predicted results are all from thermodynamic (at equilibrium state) 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 practically be negligible because its rate is too slow. Consequently, kinetic evaluations at the same time are necessary. Nevertheless, the equilibrium conditions can be used to evaluate the direction of spontaneous changes and the magnitude of the driving force behind them.