# Oxidation potential: Wikis

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Reduction potential (also known as redox potential, oxidation / reduction potential or ORP) is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. Reduction potential is measured in volts (V), millivolts (mV), or Eh (1 Eh = 1 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.

## Measurement and Interpretation

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). 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.

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. 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).

## Standard reduction potential, E0

The standard reduction potential (E0) 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, 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 $E_{0}^{r}$ and $E_{0}^{o}$.

## Converting potentials between different types reference electrodes

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.

Eobs | ref = EhEref

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
Ag/AgCl, 1 M KCl + 236
Ag/AgCl, 4 M KCl + 200
Ag/AgCl, sat. KCl +197

For example, if one measured 300 mV using a saturated KCl Ag/AgCl reference and wanted to refer it to the standard reduction potential (E0) measured using an SHE reference electrode, then 197 mV should be added to the 300 mV to obtain 497 mV, since

300mV = Eh − 197mV

it follows that

Eh = 300mV + 197mV = 497mV

and therefore

Eobs | SHE = E0 = 497mV − 0mV = 497mV

Likewise, if one measured 300 mV using a saturated KCl Ag/AgCl reference and wanted to determine the corresponding measurement using an SCE reference, then given

300mV = Eh − 197mV

it follows that

Eh = 300mV + 197mV = 497mV

and therefore

Eobs | SCE = 497mV − 241mV = 256mV

## Half cells

The relative reactivities of different half cells can be compared to predict the direction of electron flow. A higher E0 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. Eh is measured in millivolts (mV). A high positive Eh indicates an environment that favors oxidation reaction such as free oxygen. A low negative Eh 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

The Eh and pH of a solution are related. For a half cell equation (conventionally written as reduction, or with electrons on the right side):

aA + bB + n[e ] + h[H + ] = cC + dD

The half cell standard potential E0 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 Eh:

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

where curly brackets indicate activities and exponents are shown in the conventional manner. This equation is the equation of a straight line for Eh as a function of pH with a slope of − 0.05916h / n volt (pH has no units.) This equation predicts lower Eh 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 Eh at higher pH). An example of that would be the formation of magnetite (Fe3O4) from HFeO2-(aq):[1]

3 HFeO2- + H+ = Fe3O4 + 2 H2O + 2 [[e-]]

where Eh = -1.1819 - 0.0885 log[HFeO2-] + 0.0296 pH. Note that the slope of the line is -1/2 the -0.05916 value above, since h / n = -1/2.

## In 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 (Eh).

Strictly aerobic microorganisms can be active only at positive Eh values, whereas strict anaerobes can be active only at negative Eh 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.

## In geology

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.

## References

1. ^ Garrels, R.M.; Christ, C.L. (1990). Minerals, Solutions, and Equilibria. London: Jones and Bartlett.