Redox potential
The redox potential (correct designation according to DIN 38404-6 "redox voltage") is a measure of the chemistry of redox reactions . It is the reduction - / oxidation - standard potential of a substance as measured under standard conditions against a reference standard hydrogen half cell . In biological systems the standard redox potential is defined at pH 7.0 against a standard hydrogen electrode and at a partial pressure of hydrogen of 1 bar .
Basics
There react two partners: one is red uziert, the other ox idiert. In contrast to acid-base reactions , where H + ions (protons) change from one partner to the other, electrons change in redox reactions. The partner that accepts electrons is reduced, the other is oxidized. Thus, the redox reaction can be divided into two half-reactions ( redox pairs ). One of them is oxidized with the oxidation potential as the driving force, the other is reduced with the reduction potential as the driving force. The redox potential of two partners is the sum of the oxidation potential and reduction potential. The "prefer" one partner to be oxidized and the "prefer" the other to be reduced, the greater their common redox potential. The reducing power of a substance is described by its redox potential: the willingness to give up electrons and thus to change into the oxidized form.
- The more negative a redox potential, the stronger the reducing power
- Electrons flow from the redox couple with a more negative potential to the more positive redox couple
Standard potentials
Since redox potentials are dependent on external conditions such as pressure , temperature or the pH value , a state was defined in which the half-elements in the electrochemical series are located for better comparability . The standard conditions prevail in this state : the pressure is 101.325 kPa, the temperature 298.15 K (25 ° C), the activity is one.
Example: Potassium permanganate is a powerful oxidizing agent; the oxidizing power and thus the redox potential, however, depend considerably on the pH value. If a reducing agent is added to potassium permanganate, manganese (II) cations are formed at pH = 1 , manganese (IV) oxide (manganese dioxide) at pH = 7 and manganate (VI) ions at pH = 14 .
The conversion from the standard state to any other state is possible using the Nernst equation .
Measurement and quantification
In addition to the above-mentioned calculation using the Nernst equation , various measurement methods are available for determining redox potentials:
The standard redox potential of a system can be determined by setting up a galvanic element with the hydrogen electrode and measuring the electrical voltage . However, both systems must be in their standard state for this.
Redox potentials can also be accessed by determining the voltage when interconnecting them with other half-elements whose redox potential is already known. For this reason, other half-elements are often used as reference elements in practice. The calomel electrode , for example, is common , since with its potential fluctuations due to temperature changes lead to fewer measurement errors than with the hydrogen electrode.
| temperature | Potential difference |
|---|---|
| + 18 ° C | + 0.2511 V. |
| + 20 ° C | + 0.2496 V |
| + 22 ° C | + 0.2481 V |
As can be seen from the table, the redox potentials fluctuate by only around 0.6% when increasing or decreasing by 2 K.
Redox potentials in biochemistry
For biochemical processes one calculates with the potentials E o ' related to pH 7 . For reactions in which protons are involved, there is thus a potential difference of 0.413 V, as indicated in the table below.
- Please note: If redox potentials are given as E o or E o ' (table), they formally denote the potential relative to the normal hydrogen electrode . The redox potential of every other reaction, ΔE o or ΔE o ' , is then obtained by calculating the difference between the applicable E o and E o' values . The n denotes the number of electrons that are transferred during the redox reaction. According to general convention, the formal potential E 0 or E 0 ' relates to the reduction potential. Therefore the reduction reaction is shown in the table on the left.
| Oxidized form / reduced form | n | E o in V at pH 0 | E o ' in V at pH 7 |
|---|---|---|---|
| Ferredoxin Fe 3+ / Fe2 + | 1 | −0.43 | −0.43 |
| 2 H + / H 2 | 2 | 0 | −0.413 |
| NAD + , 2H + / NADH , H + | 2 | +0.09 | −0.32 |
| Lipoic Acid : Lipons., 2 H + /Lipons.-H 2 | 2 | +0.21 | −0.29 |
| Acetaldehyde 2 H + / ethanol | 2 | +0.21 | −0.20 |
| Flavin nucleotides ( FAD , FMN ): F, 2H + / F-H 2 | 2 | +0.22 | −0.19 1) |
| Glutathione : (GS) 2 , 2 H + / 2GSH | 2 | +0.31 | −0.10 |
| Fumarate , 2 H + / succinate | 2 | +0.38 | +0.03 |
| Dehydroascorbate / ascorbate, 2 H + | 2 | +0.35 | +0.08 |
| Ubiquinone , 2 H + / hydroquinone | 2 | +0.51 | +0.10 |
| ½O 2 , 2 H + / H 2 O | 2 | +1.23 | +0.82 |
| Heme iron proteins | |||
| Catalase Fe 3+ / Fe 2+ | 1 | −0.5 | −0.5 |
| Peroxidase Fe 3+ / Fe 2+ | 1 | −0.2 | −0.2 |
| Cytochrome b 562 Fe 3+ / Fe 2+ | 1 | −0.1 | −0.1 |
| Cytochrome b Fe 3+ / Fe 2+ ( mitochondria ) | 1 | +0.077 | +0.077 |
| Cytochrome b 5 Fe 3+ / Fe 2+ | 1 | +0 | +0 |
| Hemoglobin , myoglobin Fe 3+ / Fe 2+ | 1 | +0.1 | +0.1 2) |
| Cytochrome c 1 Fe 3+ / Fe 2+ | 1 | +0.22 | +0.22 |
| Cytochrome c Fe 3+ / Fe 2+ | 1 | +0.235 | +0.235 |
| Cytochrome a Fe 3+ / Fe 2+ | 1 | +0.29 | +0.29 |
| Cytochrome a 3 Fe 3+ / Fe 2+ | 1 | +0.385 | +0.385 |
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1)Flavin nucleotides are firmly bound prosthetic groups , the exact redox potential of which depends on the protein partner.2)What is remarkable is the reluctance of hemoglobin to donate electrons : this would lead to a loss of function.
Web links
Individual evidence
- ↑ a b Entry on redox potential . In: IUPAC Compendium of Chemical Terminology (the “Gold Book”) . doi : 10.1351 / goldbook.RT06783 .
- ↑ Rolf Dolder: Stabilization of oxidation-sensitive drugs as redox systems. Dissertation Zurich 1950 (PDF file)