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Redox ( reduction–oxidation, pronunciation: or ) is a type of chemical reaction in which the of Atom are changed. Redox reactions are characterized by the actual or formal transfer of between chemical species, most often with one species (the reducing agent) undergoing oxidation (losing electrons) while another species (the oxidizing agent) undergoes reduction (gains electrons). The chemical species from which the electron is removed is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. In other words:
Many reactions in organic chemistry are redox reactions due to changes in oxidation states but without distinct electron transfer. For example, during the combustion of wood with molecular oxygen, the oxidation state of carbon atoms in the wood increases and that of oxygen atoms decreases as carbon dioxide and water are formed. The oxygen atoms undergo reduction, formally gaining electrons, while the carbon atoms undergo oxidation, losing electrons. Thus oxygen is the oxidizing agent and carbon is the reducing agent in this reaction.
Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, oxygen is not necessarily included in such reactions, as other chemical species can serve the same function.
Redox reactions can occur relatively slowly, as in the formation of rust, or much more rapidly, as in the case of burning fuel. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such as the oxidation of glucose (C6H12O6) in the human body. Analysis of bond energies and ionization energies in water allow calculation of the redox potentials.
The word reduction originally referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was "reduced" to metal. Antoine Lavoisier demonstrated that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of reduction then became generalized to include all processes involving a gain of electrons.
The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes, respectively, when they occur at . These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists worldwide.
The term "hydrogenation" could often be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions, especially in organic chemistry and biochemistry. However, unlike oxidation, which has been generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance (e.g., the hydrogenation of unsaturated fats into saturated fats, R−CH=CH−R + H2 → R−CH2−CH2−R). The word "redox" was first used in 1928.
Though sufficient for many purposes, these general descriptions are not precisely correct. Although oxidation and reduction properly refer to a change in oxidation state, the actual transfer of electrons may never occur. The oxidation state of an atom is the fictitious charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is best defined as an increase in oxidation state, and reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds). As a result, simple half-reactions cannot be written for the individual atoms undergoing a redox process.
Oxidants are usually chemical substances with elements in high oxidation states (e.g., , , , , ), or else highly electronegative elements (Oxygen, Fluorine, Chlorine, Bromine) that can gain extra electrons by oxidizing another substance.
Reductants in chemistry are very diverse. Electropositive elemental , such as lithium, sodium, magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate or give away electrons relatively readily. Hydride transfer reagents, such as NaBH4 and LiAlH4, are widely used in organic chemistry, primarily in the reduction of carbonyl compounds to alcohols. Another method of reduction involves the use of hydrogen gas (H2) with a palladium, platinum, or nickel catalyst. These catalytic reductions are used primarily in the reduction of carbon-carbon double or triple bonds.
The electrode potential of each half-reaction is also known as its reduction potential E, or potential when the half-reaction takes place at a cathode. The reduction potential is a measure of the tendency of the oxidizing agent to be reduced. Its value is zero for H+ + e− → H2 by definition, positive for oxidizing agents stronger than H+ (e.g., +2.866 V for F2) and negative for oxidizing agents that are weaker than H+ (e.g., −0.763 V for Zn2+).Electrode potential values from:
For a redox reaction that takes place in a cell, the potential difference is:
However, the potential of the reaction at the anode is sometimes expressed as an oxidation potential:
This reaction is spontaneous and releases 542 kJ per 2 g of hydrogen because the H-F bond is much stronger than the weak, high-energy F-F bond. We can write this overall reaction as two :
the oxidation reaction:
and the reduction reaction:
Analyzing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above).
Elements, even in molecular form, always have an oxidation state of zero. In the first half-reaction, hydrogen is oxidized from an oxidation state of zero to an oxidation state of +1. In the second half-reaction, fluorine is reduced from an oxidation state of zero to an oxidation state of −1.
When adding the reactions together the electrons are canceled:
| |align=right | → | 2 H+ + 2 e− |
| + 2 e− | → | 2 F− |
| H2 + F2 | → | 2 H+ + 2 F− |
And the ions combine to form hydrogen fluoride:
The overall reaction is:
Zn(s)+ CuSO4(aq) → ZnSO4(aq) + Cu(s)
In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc because relative to zinc, copper metal is lower in energy due to bonding via its partially filled d-orbitals.
The ionic equation for this reaction is:
As two , it is seen that the zinc is oxidized:
And the copper is reduced:
Oxidation is used in a wide variety of industries such as in the production of cleaning products and oxidizing ammonia to produce nitric acid, which is used in most fertilizers.
Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or support electrosynthesis. Metal often contain metals in oxidized states such as oxides or sulfides, from which the pure metals are extracted by smelting at high temperature in the presence of a reducing agent. The process of electroplating uses redox reactions to coat objects with a thin layer of a material, as in Chrome plating automotive parts, silver plating cutlery, galvanization and gold-plated jewelry.
Many important biology processes involve redox reactions.
Cellular respiration, for instance, is the oxidation of glucose (C6H12O6) to carbon dioxide and the reduction of oxygen to water. The summary equation for cell respiration is:
The process of cell respiration also depends heavily on the reduction of NAD+ to NADH and the reverse reaction (the oxidation of NADH to NAD+). Photosynthesis and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cell respiration:
Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD+) to NADH, which then contributes to the creation of a proton gradient, which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, mitochondria perform similar functions. See the Membrane potential article.
Free radical reactions are redox reactions that occur as a part of homeostasis and killing microorganisms, where an electron detaches from a molecule and then reattaches almost instantaneously. Free radicals are a part of redox molecules and can become harmful to the human body if they do not reattach to the redox molecule or an antioxidant. Unsatisfied free radicals can spur the mutation of cells they encounter and are, thus, causes of cancer.
The term redox state is often used to describe the balance of Glutathione, NAD+/NADH and NADP+/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., lactic acid and pyruvate, beta-hydroxybutyrate, and acetoacetate), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis. Redox mechanism also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the CoRR hypothesis for the function of DNA in mitochondria and chloroplasts.
For instance, when manganese(II) reacts with sodium bismuthate:
| Unbalanced reaction: | Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + (aq) |
| Oxidation: | 4 H2O(l) + Mn2+(aq) → (aq) + 8 H+(aq) + 5 e− |
| Reduction: | 2 e− + 6 H+ + (s) → Bi3+(aq) + 3 H2O(l) |
The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa):
Adding these two reactions eliminates the electrons terms and yields the balanced reaction:
For example, in the reaction between potassium permanganate and sodium sulfite:
| Unbalanced reaction: | KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH |
| Reduction: | 3 e− + 2 H2O + → MnO2 + 4 OH− |
| Oxidation: | 2 OH− + → + H2O + 2 e− |
Balancing the number of electrons in the two half-cell reactions gives:
Adding these two half-cell reactions together gives the balanced equation:
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