What is a redox process


A Redox reaction (spoken: redóx) (actually: Reduction-oxidation reaction) is a chemical reaction in which one reaction partner transfers electrons to the other. In such an electron transfer reaction, electrons are released (oxidation) by a substance and electrons are accepted. Oxidation almost always takes place with the release of light and / or heat energy (exothermic). If oxidation and reduction also take place separately from one another, such a system is called a galvanic element - the flow of electrons can be used here to generate energy (example: battery, accumulator). Redox reactions that are forced from the outside are called electrolysis.

In every redox reaction, a substance A that releases electrons (reducing agent, called donor) reacts with at least one substance B that accepts these electrons (oxidizing agent, acceptor):

Oxidation: Substance A releases an electron as a reducing agent.
Reduction: The electron is absorbed by oxidizing agent B.
Redox reaction: substance A releases an electron to substance B.

Redox reactions are of fundamental importance in chemistry: every metabolic and combustion process, many technical production processes and many detection reactions are based on such electron transfer reactions.

Oxidation and oxidizing agents

Oxidation (donation of electrons)

The term oxidation has been defined differently in the history of chemistry:

Oxidation: Substance A releases one or more electrons.

The oxidation is not necessarily associated with a complete release of electrons and thus the ionization of the substances involved - especially in organic chemistry and biochemistry.

Especially here, in biology and biochemistry, oxidation is NOT defined as the "release of hydrogen".

In many biochemical processes in the cell, e.g. B. in glycolysis, organic compounds hydrogen atoms are "torn away" by certain coenzymes (NAD, NADP, FAD).

The oxidizing agent

Oxygen as an oxidizing agent

An oxidizing agent is a substance that oxidizes the reaction partner: it snatches electrons from it. Oxidation in the original sense was not initially referred to as the release of electrons, but rather the chemical reaction of a substance with oxygen. Even today, this term is often only associated with the reaction with oxygen or the formation of oxides: here, oxygen is the oxidizing agent. A copper pipe through which pure oxygen flows can be ignited with a welding torch, for example. The redox reaction started in this way supplies such intense heat of combustion that a thick lump of stone or boulder can effortlessly melt and penetrate with this oxygen lance (illustration). If a metal atom reacts with an oxygen atom, the oxidation of the metal (here: copper) and thus the metal oxide formation can be understood using the following reaction equations:

Oxidation: The metal M gives off two electrons.
Reduction: Oxygen (O) accepts two electrons.
Redox reaction: Oxygen oxidizes the metal and is itself reduced in the process.

In this case, oxygen tends to build up a stable valence electron shell with a total of eight electrons by absorbing two electrons (octet rule). The metal, in turn, can dissolve partially occupied shells by releasing the electrons and thus reach the next lower stable electron configuration. As an oxidizing agent, oxygen always reacts by absorbing 2 electrons per atom:

  • Iron rusts (corrodes) under the influence of oxygen and forms various iron oxides (rust: Fe2O3, Fe3O4, FeO).
  • The reaction of hydrogen with oxygen (oxyhydrogen) produces dihydrogen monoxide, better known as water (H.2O):
Hydrogen + oxygen → hydrogen oxide

A 2: 1 hydrogen-oxygen gas mixture reacts explosively when ignited and is therefore called oxyhydrogen. Johann W. Döbereiner discovered that the redox reaction can also be set in motion by a platinum wire as a catalyst (as early as 1816 he succeeded in oxidizing alcohol to acetic acid with the help of Platinmoor. A few years later he succeeded in igniting an oxyhydrogen gas mixture under the influence of platinum sponge. This discovery led to the invention of the first lighter (the Döbereiner platinum lighter).

  • Another classic example of oxidation by the oxidizing agent oxygen is any type of combustion of carbonaceous substances under atmospheric oxygen, e.g. burning coal, forest fires, gasoline in the engine, candles, etc. Starting from coal (pure carbon), every carbon atom is produced four electrons to two oxygen atoms to form two double bonds. Carbon dioxide (CO2).
Carbon + oxygen → carbon dioxide
  • In the body, food is also converted into the body's own substances, carbon dioxide (CO2) and water oxidizes. Not only in vivo, also in vitro organic substances can react with oxygen in a variety of ways: A primary alcohol (alkanol) is gently oxidized. In the process, an aldehyde (alkanal) is formed first, and when it is oxidized again, a carboxylic acid (alkanoic acid).

Other oxidizing agents

In addition to oxygen, other non-metallic elements and many chemical compounds can also react as oxidizing agents and accept electrons.

The term oxidation was later expanded to include reactions that follow the same chemical principle, even if no oxygen is involved. In addition to oxygen, strong oxidizing agents are therefore also:


Like oxygen, halogens are so powerful oxidizing agents that they can attack metals (corrosion). For example, when the light metal sodium in the poison gas chlorine ignites, the sodium atom releases an electron to the chlorine atom, so sodium is oxidized to sodium chloride by chlorine.

In return, the oxidizing agent chlorine is reduced:

Oxidation partial reaction: Sodium releases an electron.
Partial reaction reduction: In return, chlorine is reduced by absorbing an electron.
Overall reaction: Sodium and chlorine react with each other in a redox reaction.

Since chlorine is only molecularly known as Cl2 goes into the reaction, one writes more precisely:

Atoms with a high oxidation number, like the halogens, are particularly good oxidizing agents. One such example is potassium permanganate, because manganese (Mn) here has the oxidation number + VII. It therefore "effortlessly" oxidizes sulphite to sulphate ions, iron-II-ions to iron-III-ions and hydrochloric acid to chlorine gas.

The oxidation number

The Oxidation number indicates how many electrons a neutral atom has formally taken up (negative values) or given up (positive values) within a compound. The Oxidation number of an atom In a chemical compound, there is a formal measure for specifying the proportions of the electron density around this atom. A positive oxidation number indicates that the electron density is lower than its normal state, a negative one indicates that the electron density has increased by the atom.

The oxidation numbers are used in redox reactions to better recognize the processes: The transfer of electrons from one atom to another is shown by the fact that the oxidation number of one (which emits electrons) increases, that of the other (which accepts electrons) decreases. It is often only by determining the oxidation number of individual atoms that it becomes clear whether an oxidation (release of electrons) or a reduction (acceptance of electrons by the oxidizing agent) is taking place. In the past, the oxidation numbers in compounds were written in Roman numerals over the atomic symbols (e.g. O-II). Today this is no longer permissible according to the IUPAC nomenclature rules; therefore they are noted in Arabic numerals, with the sign of the number in contrast to the designation of the ionic charge in front the number stands.

Determination of the oxidation number

The oxidation number can easily be derived using the following rules:

  1. Elements have an oxidation number of zero
  2. If several identical atoms are linked to one another in a chain-like manner, the binding partners also have the oxidation number zero.
  3. In the case of ionic compounds, the oxidation number of an element is identical to its ionic charge.
  4. The sum of the oxidation numbers of all atoms in a polyatomic compound is equal to the charge of the compound.
  5. In the case of covalently formulated compounds (so-called line formulas, Lewis formulas), the connection is formally divided into ions. It is assumed that the electrons involved in a bond are completely taken over by the more electronegative atom.
  6. Most elements occur in several oxidation states.

In practice it has proven to be helpful to formulate a few auxiliary rules for determining the oxidation numbers:

  1. In compounds and as ions, metal atoms always have a positive oxidation number.
  2. Alkali metals always have + I and alkaline earth metals always + II as the oxidation number.
  3. The fluorine atom (F) always has the oxidation number -I in compounds except with itself.
  4. A hydrogen atom generally has the oxidation number + I. (Does not apply if hydrogen is directly linked to more "electropositive" atoms such as metals (hydrides) or to itself.)
  5. An oxygen atom generally has the oxidation number -II; in peroxides it has the oxidation number -I. (Does not apply to compounds in which oxygen is directly connected to itself or to the more electronegative fluorine, see auxiliary rule 3.)
  6. Halogen atoms (chlorine, bromine, iodine) generally have the oxidation number -I, except in connection with oxygen or another halogen that is higher in the periodic table.
  7. covalent bond: binding electrons are assigned to the more electronegative binding partner, the same binding partner each receive half of the binding electrons; the oxidation number corresponds to the assigned binding electrons compared to the number of external electrons normally present.

In the case of a molecule, the binding electron pair (s) is / are assigned to the atom whose electronegativity is higher. Numbers of similar atomic compounds are divided equally among each other. Ions have their charge number as the sum of their oxidation numbers, molecules the sum 0. The oxidation numbers of those atoms that are part of a functional group are important.

Reduction and reducing agent

The reduction (electron uptake)

The word reduction is from Latin. reductio (= Tracing back) and means in sophisticated everyday language a reduction in the extent of a property. This can refer to measurable quantities (e.g. number, circumference, weight) as well as abstract quantities. In chemistry is reduction a type of chemical reaction that absorbs electrons and thus reduces oxidation numbers.

The term reduction has also been defined differently in the history of chemistry:

  • By “reduction”, Antoine Laurent de Lavoisier understood the breakdown of oxides into an element and oxygen: reduction as the release of oxygen. If z. B. red mercury oxide is split, it is a reduction: 2HgO → 2Hg + O2 So even if you heat copper oxide in a stream of hydrogen, the copper oxide is reduced: Metallic copper is formed because the hydrogen and the oxygen in the CuO form water; the CuO is deprived of oxygen; and if you heat CuO together with C in a test tube under a gas burner, you get Cu + CO2. The CuO has converted to Cu and the oxygen is now in the CO2. (The CO released during the experiment2 can be detected with lime water)
  • Reduction as the uptake of hydrogen: At a higher level, reduction is defined as the uptake of hydrogen (and oxidation as the release of hydrogen). For example, if acetaldehyde absorbs hydrogen, it creates ethanol: CH3CHO + H2 → CH3CH2OH (This term of reduction is especially important for biological processes. In many metabolic pathways of a cell a reduction takes place through the transfer of hydrogen. The cell has even developed special coenzymes for this, e.g. NAD, NADP or FAD, which are capable Transferring hydrogen from one compound to another).
  • Reduction as electron uptake: Reduction means lowering the oxidation number and thus the uptake of electrons; Oxidation means the release of electrons and thus an increase in the oxidation number. If, for example, an iron nail is placed in a copper sulphate solution, a red-brown deposit of metallic copper forms on the nail because the iron atoms donate electrons to the copper ions ("cementation"). The copper is reduced and the iron is oxidized. In this general Definition occurs the reduction always together with the oxidation on:

Oxidation:Fe → Fe2+ + 2e-

Reduction:Cu2+ + 2e- → Cu
Redox equation:Fe + Cu2+ → Cu + Fe2+

The reducing agent

The Reducing agent is accordingly the substance that provides electrons to the reaction partner: it lowers its oxidation number. It is a substance that can reduce other substances and is itself oxidized (electron donor) - i.e. the opposite of the oxidizing agent (electron acceptor).

Quality Reducing agent (Electron donors) are:

  • Hydrogen (H2)
  • base metals (alkali, alkaline earth metals, aluminum, zinc, etc.)
  • Carbon,
  • Methane, ethane and other combustibles as well
  • Hydrides such as lithium aluminum hydride and sodium borohydride,
  • Hydrogen sulfide and sodium sulfide, sodium sulfite, sodium dithionite, sodium thiosulfate,

The redox reaction: electron transfer through oxidation and reduction

If an electron transfer takes place during a material conversion, this means that an electron release (oxidation) and an electron uptake (reduction) take place: the substance that gives off electrons is the reducing agent (electron donor, its oxidation number increases) - the oxidising agent takes it up (electron acceptor, its oxidation number is lower)

This electron exchange - for example between ammonium and nitrate ions in the fertilizer ammonium nitrate - can be uncontrolled and explosive if handled incorrectly (see figure):

Examples of different types of electron exchange

Salt formation: reaction of metal and non-metal

The Elements magnesium and Iodine form the connection Magnesium iodide

Magnesium releases electrons to the oxidizing agent iodine.

Cementation: reaction of noble metal cation and metal

The conversion of aluminum and copper sulfate to copper and aluminum sulfate takes place as a redox reaction, since aluminum, as a reducing agent, releases electrons to the copper (II) cations of the blue salt: 2 Al (s) + 3 CuSO4 Al2(SO4)3 (aq) + 3 Cu (s)

Elemental copper is also deposited on a zinc rod (as a reducing agent), which is immersed in a solution with copper (II) ions (oxidizing agent) (cementation). The solution is enriched with zinc (II) ions.

Redox reactions can be divided into the partial reactions oxidation and reduction disassemble:

  1. Oxidation:
  2. Reduction:

The amalgam sample - a detection reaction for mercury salts with copper sheet - is also a redox reaction. This "cementation" proceeds according to the following scheme:

The electron donor - here the metal copper as a detection agent - is the reducing agent: It has the task of reducing other substances, but is itself oxidized. The electron acceptor - here the mercury salt in solution - is the oxidizing agent because it oxidizes other substances and itself absorbs electrons (is reduced). This example also shows that electrons are given off with every oxidation, so that the oxidation number increases. Conversely, the oxidation number is reduced (reduced) during the reduction.

Reaction of non-metal with non-metal anion

Elemental chlorine releases elemental iodine from sodium iodide when dissolved in water during a halide detection reaction and sodium chloride (table salt) is formed:

The oxidation number of the oxidizing agent chlorine becomes smaller here (from 0 to -1): It is reduced. Iodide, on the other hand, is oxidized to iodine, which then dissolves in hexane in a red-violet color (figure). Chlorine water also reacts in a similar way with bromide (which then results in hexane colored red-orange by bromine).



At a Comproportioning (also: Synproportionation) react connections that an element in lower Oxidation level (oxidizing agent), together with compounds in which the element in question is in higher Oxidation level (reducing agent) is present to a compound with middle Oxidation state.

Example: Lead and lead IV oxide react in the battery acid of a lead accumulator when discharging to lead II sulphate (car battery): Pb + PbO2 + 2 H.2SO4 → 2 PbSO4 + 2 H.2O


One speaks of a disproportionation when, in a chemical reaction, elements with a medium oxidation state change into those with a lower and a higher one.These reactions often take place under the influence of a catalyst (a substance that accelerates a chemical reaction without changing).

Example: When potassium chlorate (chlorine with oxidation number + V) is heated, oxygen, potassium chloride (chlorine with –I) and potassium perchlorate (chlorine with + VII) are produced:

2 KClO3 → KClO4 + KCl + O2 (endothermic reaction)

Reaction schemes for redox reactions

Like any chemical reaction, a redox reaction is also written as a reaction scheme. An example of a simple inorganic redox equation (reaction scheme in partial steps, with oxidation numbers) is the cementation of copper from copper (II) salt solution with iron:

Copper salts oxidize iron: copper ions take up electrons, iron releases them. The oxidation number of copper becomes smaller (from +2 to 0), that of the reducing agent increases (Fe from 0 to +2).


Redox .:

Formulation of redox reaction equations


Such sometimes more complicated reaction schemes are formulated, if necessary, using coefficients or - as a second example - ideally also step-by-step. For the redox reactions of potassium nitrate (nitric) with sulfur and of copper with concentrated nitric acid, the reaction scheme is created as follows, for example:

1st example: About coefficients (correct)

Oxidation number and electron equation are practical tools for formulating complete reaction equations for redox reactions. First, let's take a simple example:

a KNO3 + b S → c K2SO4 + d NO

The coefficients a, b, c and d can be determined here by comparing the two sides of the reaction scheme:

2KNO3 + S → K2SO4 + 2NO

2nd example: In partial steps (with "electron equations") This is considerably more difficult for the equation with which we describe the reaction between copper and undiluted nitric acid:

a Cu + b ENT3 → d Cu (NO3)2 + e H2O + f NO

The following steps are necessary to determine the coefficients you are looking for:

- First of all, the particles that occur during the reaction must be formulated:

Cu + H+ + NO3- → Cu2+ + NO3- + H2O + NO

- Then write the corresponding oxidation number over each element symbol:

0 + I + V-II + II + V-II + II-II + I-II Cu + H+ + NO3- → Cu2+ NO3- + NO + H2O

- The particles of the starting and end materials are compared. The particles in which the oxidation numbers have changed are determined and the two sub-processes are formulated:

0 + II Cu → Cu2+ + 2e- + V + II NO3- + 3e- → NO

- Now this subset is multiplied by the smallest factor, which causes the number of electrons accepted to be equal to the number of electrons released:

Sub-processes without coefficients Factors Sub-processes with coefficients Cu → Cu2+ + 2e- · 3 3Cu → 3Cu2+ + 6e- NO3- + 3e- → NO 2 2 NO3- + 6e- → 2NO ____________________________________________________________________________ The particle equation therefore applies to the two partial processes: 3Cu + 2NO3- → 3Cu2+ + 2NO

- Now one finds the number of oxygen atoms (with the oxidation number -II) on both sides of the equation. 6 oxygen atoms are found on the left and 2 on the right. On the right hand side of the equation, 4 oxygen atoms are missing. You are in the H2O included as a further reaction product. For the formation of water molecules one needs 8 hydrogen atoms with the oxidation number + I, i.e. 8 H.+-Ions:

3Cu + 2NO3- + 8H+ → 3Cu2+ + 2NO + 4H2O

The complete reaction equation is thus:

3Cu + 8HNO3 → 3Cu (NO3)2 + 2NO + 4H2O

Copper (II) nitrate solutions are colored blue-green. The nitrogen monoxide gas reacts further to form brown nitrogen dioxide gas or nitrous tetroxide. This is used in rocket technology as an oxidizing agent in liquid fuel mixtures.

Redox processes and electrochemistry

The spatial separation of reduction and oxidation creates an electrochemical cell. A distinction is made between electrochemical cells in which the electrochemical reaction (the cell reaction) takes place voluntarily, these cells are called galvanic cells. If the cell reaction is forced by the application of a current from the outside so that the electrochemical reaction takes place, one speaks of an electrolysis cell.

The Electrochemistry is the branch of physical chemistry that deals with the relationship between electrical and chemical processes. If, therefore, a redox reaction is forced or supplies such an electric current, this is an electrochemical process.

Galvanic elements, accumulators and electrolysis


If the redox reaction is forced by an externally applied electrical voltage, this process is called (electrolysis) - if a measurable voltage is produced by the chemical reaction of suitable substances, a galvanic element is present. These voltages (redox potentials) are characteristic of the respective reactions and are documented on a scale, the electrochemical series. So here is the Strength an oxidizing or reducing agent measurable. Electrochemical redox reactions take place in a galvanic cell:

  • During electrolysis and the charging of an accumulator, electrical energy is supplied,
  • when discharging a battery or when drawing electricity from a fuel cell, electrical energy is obtained which, in the reversible case, corresponds to the Gibbs energy of the reaction when the electrical current tends towards zero (I = 0).

The processes that are decisive for electrochemistry take place at the electrode-electrolyte phase boundary - electrochemistry is the science of the processes at the phase boundary between an electron conductor (electrode) and an ion conductor (electrolyte).

Applications of electrochemistry

  • Manufacture of chemical substances
    • Reduction of metal salts for the production of base metals, above all by melt electrolysis, e.g. for the production of lithium, sodium, potassium, calcium, magnesium and aluminum
      The electric current acts here as a reducing agent. Since the voltage can be varied, the reducing force can be adjusted. Electric current is the strongest reducing agent in chemistry, with which even the most base metals can be reduced.
      Electrolytic metal deposition is also used in electroplating.
    • Oxidation of anions, e.g. of halides, e.g. for the production of fluorine and chlorine
    • The electric current allows redox reactions without the addition of reducing or oxidizing agents. Many other redox reactions can therefore be carried out particularly elegantly electrolytically or are only made possible. Electrofluorination and Kolbe electrolysis should be mentioned.
  • Electroplating
  • Provision of an electrical voltage, especially for mobile applications, in:
  • Use of electric current to carry out chemical analyzes and investigations: Electroanalysis, especially polarography
  • Studies of thermodynamics and the mechanism of reactions, also important for corrosion research


  • Michael Wächter: Substances, particles, reactions. Verlag Handwerk und Technik, Hamburg 2000, pp.154-169 ISBN 3-582-01235-2
  • Bertram Schmidkonz: Inorganic Analysis Internship. Harri Deutsch publishing house, Frankfurt 2002, ISBN 3-8171-1671-3
  • Peter W. Atkins, Physical chemistry, Wiley-VCH, ISBN 3527302360
  • C.H. Hamann, W. Vielstich, Electrochemistry, Wiley-VCH, 3rd ed. 1998
  • W. Schmickler, Basics of electrochemistry, Springer 1996
  • L. Landen and Dr. Kledt, Basic knowledge of chemistry 1, Velcro 2005

Categories: Detection Reaction | Chemical reaction