Reduction potential

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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 reduction potential will have a tendency to gain electrons from new species (i.e. oxidize them) and a solution with a lower reduction potential will have a tendency to lose electrons to new species (i.e. reduce them). 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 characterise the capacity of the system for oxidation or reduction, in much the same way that pH does not characterise the buffering capacity.

Reduction potential is measured in volts (V), millivolts (mV), or Eh (1Eh = 1mV). Because the true or 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 (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. Historically, many countries, including the United States, 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_r⁰ and E_o⁰.

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. 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 4.4V to 4.6V (the electrolyte being positive.)

1 Reduction potential in biochemistry
2 Practical measurement of reduction potential
3 See also
4 External links

[edit] Reduction potential 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. Oxygen strongly affects redox potential.

[edit] Practical measurement of reduction potential

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, Ag/AgCl and saturated calomel (SCE) reference electrodes are commonly used. The voltage relationships for several different reference electrodes at 25 degrees C can be interrelated as follows:

Reference electrode	Electrode potential with respect to SHE (mV)
Standard hydrogen electrode (SHE)	0
Saturated calomel electrode (SCE)	+ 245
Ag/AgCl, 1 M KCl	+ 236
Ag/AgCl, 4 M KCl	+ 200
Ag/AgCl, sat. KCl	+199

For example: If you had a reading of 100mV using a saturated KCl Ag/AgCl reference and wanted to refer it back to an SHE you would add 199mV to obtain 299mV. Alternatively, if you took a reading in the same solution using an SCE, you would obtain 54mV (subtract 245mV from 299mV).