Redox reactions

Oxidation-reduction reactions – key concepts

Oxidation state

The oxidation state of a chemical element is the hypothetical charge that could accumulate on an atom of a particular element contained in a chemical compound if all chemical bonds in this compound were ionic. In practice, this situation does not always occur (decomposition of a compound into individual cations and anions), so the oxidation state should be treated as a conventional concept. The oxidation state is equal to the charge of a particular ion, so it takes either positive or negative values. It is denoted by a Roman numeral placed after the symbol of a chemical element. Elements in different oxidation states have different oxidation-reduction properties.

Oxidation

During oxidation (deelectronation), the reductant increases its oxidation state, i.e. gives electrons to the oxidant. Neither oxidation nor reduction can proceed independently because electrons donated by one chemical entity must be immediately accepted by the other for the opposite reaction to take place.

Reduction

During reduction (electronation), the oxidant lowers its oxidation state by accepting electrons, i.e. the reduction is about taking electrons. Chemical elements that do this are called oxidants.

Disproportionation reaction (dismutation)

The disproportionation reaction is one of the types of redox reactions. In the literature, you can also find the term: dismutation reaction. Its characteristic feature is that during the course of a redox reaction, the same element is simultaneously oxidised and reduced. In order for a disproportionation to take place, the element in question must have at least three different oxidation states. If this condition is met, the compound which is in the intermediate oxidation state is much less stable compared to the other two states. disproportionation reactions proceed spontaneously. Atoms such as sulphur, nitrogen, phosphorus or manganese are susceptible to this type of redox reaction.

Synproportionation reaction

The synproportionation reaction, like disproportionation, is also a type of redox reaction. This process occurs when two different chemical compounds containing the same element in different oxidation states react with each other. As a result of the oxidation-reduction reaction, another compound containing that element in a new oxidation state is formed.

Electron balance

In every oxidation-reduction reaction that takes place, the same number of electrons are exchanged. If in a particular process, the reductant donates, for example, two electrons, the other of the pair, the oxidant, will also accept two electrons into its electron shell. This situation is referred to as the so-called electron balance of the reaction. For an entire redox reaction, this balance should be zero.

How do redox reactions occur?

The basis of any redox reaction is oxidation and reduction. Taking these into account, any process can be written using the so-called half-equations, in which only atoms donating or accepting electrons are specified. Thus, the whole redox reaction is, in a way, about giving and taking electrons. Only those elements that occur on more than one oxidation state in chemical compounds can do so. Knowing its states in individual chemical entities is essential to correctly write and balance redox reactions. While balancing electrons, in addition to correctly writing the half-equations, the oxidation and reduction reactions as well as the oxidant and the reductant, respectively, should be indicated. Oxidants most commonly include highly electronegative elements (groups 16 and 17 of the periodic table), metal ions in higher oxidation states, noble metal ions and oxidising acids (e.g. nitric (V) acid, sulphuric (VI) acid and their mixtures with other non-oxidising acids). The most common oxidants are compounds such as KMnO₄, K₂Cr₂O₇, KClO₃ or K₂S₂O₈. The reductants, on the other hand, are electropositive elements (usually from groups 1 and 2 of the periodic table), metals in zero oxidation state, molecular hydrogen, carbon, carbon monoxide and anions of inorganic acids. The most popular reductants are: Na, Mg, Fe²⁺, Cl^–, Br^–, SCN^–. The redox reaction formula furthermore indicates the number of electrons exchanged in the process. The course of this electron exchange is determined by the redox potential of the reactants involved. In other words, it can be called half-cell potential or electron potential. By definition, the greater the potential difference in the system, the greater the driving force of the entire oxidation-reduction reaction.

Can redox reactions be observed in daily life?

It may seem that redox reactions appear only on the pages of school textbooks and in chemistry lessons. However, nothing could be further from that. These kinds of reactions accompany us every day. It is worth learning more about them to observe the processes and environment around us with greater understanding. The following are examples of everyday redox reactions that each of us has certainly encountered:

• Corrosion of metals – it is the most common deterioration process of metals and their alloys. It results from the contact of the surface of the material in question with the environment and atmospheric conditions. In terms of the mechanisms of corrosion processes, the most common one is electrochemical corrosion, which occurs in an electrolyte environment, in moist gases or in soil with high moisture levels. In the place where corrosion occurs, a so-called corrosion cell is formed, in which electrode oxidation-reduction reactions take place. Metal deterioration always occurs in the anodic area. There, electrons are donated by the metal, which oxidises and in the form of ions passes into the electrolyte solution. The released charges migrate to the cathode. There, they combine with ions or atoms which have the ability to accept electrons. These are most often oxygen atoms from the air (on the cathode, they will be reduced to hydroxide ions) or hydrogen ions (they will be reduced to molecular hydrogen). At the cathode, either or both of these processes can take place at the same time.
• Photosynthesis – it is a process that accompanies us every day. During photosynthesis, cells convert atmospheric carbon dioxide and water into glucose and oxygen with the use of solar energy. Like many biochemical processes occurring in living organisms, photosynthesis also involves changing the oxidation states of the elements that make up the reactants. In this redox reaction, the oxygen atom in the water molecule is oxidised to molecular oxygen. Therefore, the water molecule is the electron donor, or the reductant. The acceptor of the resulting charge, or oxidant, is carbon dioxide. Its constituent carbon atoms in the fourth oxidation state are reduced to zero oxidation state.
• Galvanic cells – cells are arrangements of two electrodes, immersed in the same electrolyte (or different electrolytes), which are connected to each other by means of an external circuit. Each electrode immersed in its own electrolyte (half-cell), exhibits a certain potential. The resulting potential difference, i.e. the flow of current (electrons), is caused by the oxidation-reduction reactions. Half processes take place at each electrode. At the anode, as a result of the oxidation reaction, electrons are donated, which are then accepted at the second electrode – the cathode – in the reduction reaction. The most common devices using galvanic cells are batteries, which are a source of energy for cars, for example. The most common lead–acid is composed of two electrodes. One is pure lead and the other is coated with lead (IV) oxide. Both are immersed in 37% sulphuric (VI) acid. It allows a free exchange of electrons between the cathode and anode. During battery operation, oxidation-reduction reactions begin to take place. In this case, the anode is the lead electrode. Lead starts to oxidise and goes from zero oxidation state to up to the second oxidation state. At the same time, two electrons are released and migrate to the cathode via the electrolyte. There, the process of lead reduction from the fourth oxidation state to lead (II) begins, i.e. lead (IV) oxide is transformed into lead (II) sulphate. In the case of a battery, the redox reaction is a source of energy, which can be used to power a number of devices.