Redox potential (also known as oxidation / reduction potential, ORP, pe, ${\displaystyle E_{red}}$, 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 expressed in volts (V). 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.

Measurement and interpretation

In aqueous solutions, redox potential is a measure of the tendency of the solution to either gain or lose electrons in a reaction. A solution with a higher (more positive) reduction potential than some other molecule will have a tendency to gain electrons from this molecule (i.e. to be reduced by oxidizing this other molecule) and a solution with a lower (more negative) reduction potential will have a tendency to lose electrons to other substances (i.e. to be oxidized by reducing the other substance). 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.^[1]

The sensing electrode acts as a platform for electron transfer to or from the reference half cell; it is typically made of 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 V. 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 how the concentration of hydrogen ions 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.

The notion of pe is used with Pourbaix diagrams. pe is a dimensionless number and can easily be related to E_H by the following relationship:

${\displaystyle pe={\frac {E_{H}}{V_{T}\lambda }}={\frac {E_{H}}{0.05916}}=16.903\,{\text{×}}\,E_{H}}$

where, ${\displaystyle V_{T}={\frac {RT}{F}}}$ is the thermal voltage, with R, the gas constant (8.314 J⋅K⁻¹⋅mol⁻¹), T, the absolute temperature in Kelvin (298.15 K = 25 °C = 77 °F), F, the Faraday constant (96 485 coulomb/mol of  e⁻), and λ = ln(10) ≈ 2.3026.

In fact, ${\displaystyle pe=-\log}$ is defined as the negative logarithm of the free electron concentration in solution, and is directly proportional to the redox potential.^[1][2] Sometimes ${\displaystyle pe}$ is used as a unit of reduction potential instead of ${\displaystyle E_{h}}$, for example, in environmental chemistry.^[1] If one normalizes ${\displaystyle pe}$ of hydrogen to zero, one obtains the relation ${\displaystyle pe=16.9\ E_{h}}$ at room temperature. This notion 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 themselves, and depend on concentrations, temperature etc.

The table below shows a few reduction potentials, which can be changed to oxidation potentials by reversing the sign. Reducers donate electrons to (or "reduce") oxidizing agents, which are said to "be reduced by" the reducer. The reducer is stronger when it has a more negative reduction potential and weaker when it has a more positive reduction potential. The more positive the reduction potential the greater the species' affinity for electrons and tendency to be reduced. The following table provides the reduction potentials of the indicated reducing agent at 25 °C. For example, among sodium (Na) metal, chromium (Cr) metal, cuprous (Cu⁺) ion and chloride (Cl⁻) ion, it is Na metal that is the strongest reducing agent while Cl⁻ ion is the weakest; said differently, Na⁺ ion is the weakest oxidizing agent in this list while Cl₂ molecule is the strongest.

Some elements and compounds can be both reducing or oxidizing agents. Hydrogen gas is a reducing agent when it reacts with non-metals and an oxidizing agent when it reacts with metals.

2 Li (s) + H₂ (g) → 2 LiH (s)^[a]

Hydrogen (whose reduction potential is 0.0) acts as an oxidizing agent because it accepts an electron donation from the reducing agent lithium (whose reduction potential is -3.04), which causes Li to be oxidized and Hydrogen to be reduced.

H₂ (g) + F₂ (g) → 2 HF (g)^[b]

Hydrogen acts as a reducing agent because it donates its electrons to fluorine, which allows fluorine to be reduced.

Standard reduction potential

The standard reduction potential ${\displaystyle E_{red}^{\ominus }}$ is measured under standard conditions: T = 298.15 K (25 °C, or 77 °F), a unity activity (a = 1) for each ion participating into the reaction, a partial pressure of 1 atm (1.013 bar) for each gas taking part into the reaction, and metals in their pure state. The standard reduction potential ${\displaystyle E_{red}^{\ominus }}$ is defined relative to the standard hydrogen electrode (SHE) used as 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 by the symbols ${\displaystyle E_{red}}$ and ${\displaystyle E_{ox}}$, with ${\displaystyle E_{ox}=-E_{red}}$.

Half cells

The relative reactivities of different half cells can be compared to predict the direction of electron flow. A higher ${\displaystyle E_{red}^{\ominus }}$ 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 usually expressed in volts (V) or 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 one knows the actual potential between electrode and electrolyte for any one reaction. Surface polarization interferes with measurements, but various sources^{[citation needed]} 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 the one corresponding to oxidation is reversed so that each electron given by the reductant is accepted by the oxidant. In this way, the global combined equation no longer contains electrons.

Nernst equation

The ${\displaystyle E_{h}}$ and pH of a solution are related by the Nernst equation as commonly represented by a Pourbaix diagram (${\displaystyle E_{h}}$ – pH plot). For a half cell equation, conventionally written as a reduction reaction (i.e., electrons accepted by an oxidant on the left side):

${\displaystyle a\,A+b\,B+h\,{\ce {H+}}+z\,e^{-}\quad {\ce {<=>}}\quad c\,C+d\,D}$

The half-cell standard reduction potential ${\displaystyle E_{\text{red}}^{\ominus }}$ is given by

${\displaystyle E_{\text{red}}^{\ominus }({\text{volts}})=-{\frac {\Delta G^{\ominus }}{zF}}}$

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

${\displaystyle E_h=E_{\text{red}}=E_{\text{red}}^{\ominus <!--\; \ominus \; -->}-<!--\; -\; -->\frac{0.05916}{z}\log <!--\; \; -->(\frac{\{C{\}}^c\{D{\}}^d}{\{A}}$Zdroj:https://en.wikipedia.org?pojem=Standard_reduction_potential
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