• What are redox reactions?
• Characteristics of redox reactions
• Types of redox reactions
• Examples of redox reactions
• Industrial applications

We explain what redox reactions are, the types that exist, their applications, characteristics and examples of redox reactions.

In redox reactions, one molecule loses electrons and another takes them.

What are redox reactions?

In chemistry, is known as redox reactions, oxide-reduction reactions or reduction-oxidation reactions to chemical reactions in which an exchange of electrons between the atoms or molecules involved.

This exchange is reflected in the change of state of oxidation of the reagents. The reactant that gives up electrons undergoes oxidation and the one that receives them, reduction.

The oxidation state indicates the amount of electrons that an atom of a chemical element gives up or accepts when it is part of a chemical reaction. It can also be interpreted as the supposed electric charge that a certain atom would have if all its bonds with other atoms were completely ionic. Also called the oxidation number or Valencia.

The oxidation state is expressed in integer numbers, being zero the oxidation state for neutral elements. Thus, it can take positive or negative values ​​depending on the type of atom and the reaction where it participates. On the other hand, some atoms They have variable oxidation states depending on the reaction in which they are involved.

Know how to correctly determine the state or oxidation number of each atom in a chemical compound It is essential to understand and analyze redox reactions. There are certain rules that allow you to calculate their values:

• The oxidation number of neutral elements or molecules is zero. For example: solid metals (Fe, Cu, Zn…), molecules (O2, N2, F2).
• The ions compounds of a single atom have their oxidation number equal to their charge. For example: Na +, Li +, Ca2 +, Mg2 +, Fe2 +, Fe3 +, Cl–.
• Fluorine always has a -1 oxidation state because it is the most electronegative element that exists (F–).
• Hydrogen always has oxidation number +1 (H +), with the exception of metal hydrides (potassium hydride, KH), where it has oxidation number -1 (H–).
• Oxygen has oxidation number -2, with a few exceptions:
  • When it forms compounds with fluorine, it has an oxidation number of 2+. For example: oxygen difluoride (OF2).
  • When it forms peroxides, it has oxidation number -1 (O22-). For example: hydrogen peroxide (H2O2), sodium peroxide (Na2O2).
  • When it forms superoxides, it has the oxidation number -½ (O2–). For example: potassium superoxide (KO2).
• The algebraic sum of the oxidation numbers of the atoms that make up a neutral compound is zero.
• The algebraic sum of the oxidation numbers of the atoms that make up a polyatomic ion is equal to the charge on the ion. For example: the sulfate anion (SO42-) has oxidation number -2, which is equal to the sum of the oxidation numbers of sulfur and oxygen, each multiplied by the amount of each atom in the compound, in this case, it has one sulfur atom and four oxygen atoms.
• The oxidation numbers of some chemical elements they can vary depending on the neutral compound or ion of which they are part. Then, it is possible to calculate the oxidation number of an atom in a compound as follows:

Where no() means oxidation number, and the chemical element is found within the parentheses.

In this way, in every redox reaction there are two types of reactants, one that gives up electrons and the other that accepts them:

• An oxidizing agent. It is the atom that captures the electrons. In this sense, its initial oxidation state decreases, and a reduction is experienced. In this way, it increases its negative electrical charge by gaining electrons.
• A reducing agent. It is the atom that gives up the electrons and increases its initial oxidation state, undergoing oxidation. In this way, it increases its positive electric charge by giving up electrons.

Some chemicals can be oxidized and reduced at the same time. These elements are called ampholytes, and the process in which this happens is called ampholization.

Redox reactions are one of the most common chemical reactions in the universe, as they are part of the processes of photosynthesis in the plants and of the breathing in animals, which allow the continuity of life.

Characteristics of redox reactions

Redox reactions are all around us on a daily basis. The oxidation of metals, the combustion of gas in the kitchen or even the oxidation of glucose to obtain ATP in our body are some examples.

In most cases, redox reactions release a significant amount of Energy.

Generally, each redox reaction is made up of two stages or half-reactions. In one of the half-reactions, oxidation occurs (the reactant is oxidized) and in the other, reduction occurs (the reactant is reduced).

The total redox reaction, which is obtained as a result of combining all the half-reactions algebraically, is often called the “global reaction”. It is important to note that when semi-reactions are combined algebraically, both mass and charge must be adjusted. That is, the number of electrons released during oxidation must be the same as the number of electrons gained during reduction, and the mass of each reactant must equal the mass of each product.

For example:

• Reduction half-reaction. Reduction of copper by capturing two electrons. Decreases its oxidation state.
  
• Oxidation semi-reaction. Iron oxidation by losing two electrons. Increases its oxidation state.
  
  Global reaction:
  

Types of redox reactions

Combustion reactions (redox reactions) release energy that can create movement.

There are different types of redox reactions, endowed with different characteristics. The most common types are:

• Combustion. Combustion are redox chemical reactions that release a significant amount of energy in the form of heat Y light. These reactions are fast oxidations that give off a lot of energy. The released energy can be used in a controlled way to generate movement in car engines. An element called oxidizer (which is reduced and oxidized to the fuel) and a fuel element (which is oxidized and reduced to the oxidizer). Some examples of fuels are gasoline and the gas that we use in our kitchens, while the best known oxidizer is gaseous oxygen (O2).
• Oxidation of metals. They are reactions slower than combustion. They are commonly described as the degradation of certain materials, especially metallic ones, by the action of oxygen on them. It is a worldwide known and everyday phenomenon, especially in coastal populations, where salts from the environment accelerate (catalyze) the reaction. That is why a car, after taking us to the beach, must be cleaned of all traces of salt water.
• Disproportion. Also known as dismutation reactions, they present a single reagent that is reduced and oxidized at the same time. A typical case of this is the decomposition of hydrogen peroxide (H2O2).
• Simple scrolling. Also called "simple substitution reactions", they occur when two elements exchange their respective places within the same compound. That is, one element replaces another in its exact place in the formula, balancing their respective electric charges with other atoms as appropriate. An example is what happens when a metal displaces hydrogen in an acid and salts are formed, as occurs when batteries of an appliance break down.

Examples of redox reactions

Examples of redox reactions are very abundant. We will try to give an example of each of the types previously described:

• The combustion of octane. Octane is a hydrocarbon component of gasoline used to run the engine of our cars. When octane reacts with oxygen, octane is oxidized and oxygen is reduced, releasing a large amount of energy as a result of this reaction. This released energy is used to generate work in the engine, also producing carbon dioxide and water vapor in the process. The equation that represents this reaction is:
  
• The decomposition of hydrogen peroxide. It is a dismutation reaction in which hydrogen peroxide breaks down into its constituent elements, water and oxygen. In this reaction, oxygen is reduced by decreasing its oxidation number from -1 (H2O2) to -2 (H2O), and it is oxidized by increasing its oxidation number from -1 (H2O2) to 0 (O2).
  
• Displacement of silver by copper. It is a reaction of displacement simple in which you can see how by immersing a fragment of metallic copper in a solution of silver nitrate, the color of the solution turns blue and a thin layer of metallic silver is deposited on the copper fragment. In this case, part of the metallic copper (Cu) is transformed into the Cu2 + ion, as part of the copper (II) nitrate (Cu (NO3) 2), whose solution has a beautiful blue color. On the other hand, part of the Ag + cation, which is part of silver nitrate (AgNO3), is transformed into metallic silver (Ag) which is deposited.
  
• Reaction of zinc with dilute hydrochloric acid. It is a simple displacement reaction in which hydrogen in HCl (aq) is displaced by zinc to form a salt.
  
• Iron oxidation. Metallic iron oxidizes when it comes into contact with oxygen from the air. This is seen in everyday life when iron objects form a layer of brown rust when exposed to air for long periods of time. In this reaction, metallic iron (Fe), which has an oxidation state of 0, is transformed into Fe3 +, that is, its oxidation state increases (it oxidizes). For this reason, it is said intuitively or colloquially: iron rusts.
  

Industrial applications

In power plants, redox reactions are able to move large motors.

The industrial applications of redox reactions are endless. For example, combustion reactions are ideal for producing job that serves to generate movement in the large motors used in power plants to produce electricity.

The process consists of burning fossil fuels to get heat and produce water steam in a boiler, then this steam is used to drive large engines or turbines. On the other hand, combustion reactions are also used to run the engine of motor vehicles that use fossil fuels, such as our cars.

On the other hand, substitution and displacement redox reactions are useful to obtain certain elements in a state of purity that is not often seen in the nature. For example, silver is highly reactive. Although it is rare to find it pure in the mineral subsoil, a high degree of purity can be obtained through a redox reaction. The same happens when it comes to obtaining salts and other compounds.