Reduction potential

Reduction potential (also known as redox potential, oxidation / reduction potential or ORP) is the tendency of a chemical species to acquire electrons and thereby be reduced. 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.

In aqueous solutions, the reduction potential is 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.

Standard hydrogen electrode

Reduction potential is measured in volts (V), millivolts (mV), or Eh (1 Eh = 1 mV). Because the absolute potentials are difficult to accurately measure, reduction potentials are defined relative to the standard hydrogen electrode (SHE) which is arbitrarily given a potential of 0.00 volts. Standard reduction potential ( $E^{0}$ ), 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. 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 $E_{r}^{0}$ and $E_{o}^{0}$ .

Saturated calomel electrode

Although measurement of the reduction potential in aqueous samples is relatively straightforward, many factors limit its interpretation, such as 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).

Reduction potentials of aqueous solutions are determined by measuring the potential difference between an inert indicator electrode in contact with the solution and a stable reference electrode connected to the solution by a salt bridge. The indicator 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, silver chloride and saturated calomel (SCE) reference electrodes are commonly used.

Converting potentials between reference electrodes

Often a reduction potential is quoted against a different reference electrode than the one utilized and it becomes necessary to convert the potential. This is easily done by recognizing that the measured potential represents the difference between the potential measured at the indicator electrode and the reference electrode, i.e. $E_{Msrdvs.Ref}=E_{Indctr}-E_{Ref}$ 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 back to the Standard Reduction Potential that would be measured using an SHE reference, then 197 mV should be added to the 300 mV to obtain 497 mV, since $300mV=E_{Indctr}-197mV$ , then $E_{Indctr}=300mV+197mV=497mV$ , and $E_{MsrdvsSHE}=497mV-0mV=497mV$

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 $300mV=E_{Indctr}-197mV$ , then $E_{Indctr}=300mV+197mV=497mV$ and $E_{MsrdvsSCE}=497mV-241mV=256mV$ Jjotter (talk) 03:25, 24 October 2009 (UTC)

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

The $E_{h}$ 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 $E_{0}$ is given by:

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

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

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

where square 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 HFeO₂^-_(aq):^[1]

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

where $E_{h}$ = -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 ( $E_{h}$ ).

Strictly aerobic microorganisms can be active only at positive $E_{h}$ values, whereas strict anaerobes can be active only at negative $E_{h}$ 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 postive 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

^ Garrels, R.M.; Christ, C.L. (1990). Minerals, Solutions, and Equilibria. London: Jones and Bartlett.{{cite book}}: CS1 maint: multiple names: authors list (link)

External links

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

[garrels-1] Garrels, R.M.; Christ, C.L. (1990). Minerals, Solutions, and Equilibria. London: Jones and Bartlett.{{cite book}}: CS1 maint: multiple names: authors list (link)

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