The transformation of one substance into another with the formation of new compounds is called a chemical reaction. Understanding this process is of great importance for the life of people, because with its help you can get a huge amount of necessary and useful substances that are found in nature in small quantities or do not exist at all in their natural form. Among the most important varieties are redox reactions (abbreviated OVR or redox). They are characterized by a change in the oxidation states of atoms or ions.
Processes occurring during the reaction
During the reaction, two processes take place - oxidation and reduction. The first of them is characterized by the donation of electrons by reducing agents (donors) with an increase in their oxidation state, the second by the addition of electrons by oxidizing agents (acceptors) with a decrease in their oxidation state. The most common reducing agents are metals and non-metal compounds in the lowest oxidation state (hydrogen sulfide, ammonia). typicaloxidizing agents are halogens, nitrogen, oxygen, as well as substances that contain an element in the highest oxidation state (nitric or sulfuric acid). Atoms, ions, molecules can donate or gain electrons.
Before 1777, it was hypothesized that oxidation resulted in the loss of an invisible combustible substance called phlogiston. However, the combustion theory created by A. Lavoisier convinced scientists that oxidation occurs when interacting with oxygen, and reduction occurs under the action of hydrogen. Only after a while it became clear that not only hydrogen and oxygen can affect redox reactions.
Oxidation
The process of oxidation can occur in the liquid and gaseous phases, as well as on the surface of solids. A special role is played by electrochemical oxidation occurring in solutions or melts at the anode (an electrode connected to the positive pole of the power source). For example, when fluorides are melted by electrolysis (the decomposition of a substance into its constituent elements on electrodes), the strongest inorganic oxidizing agent, fluorine, is obtained.
Another classic example of oxidation is combustion in air and pure oxygen. Various substances are capable of this process: metals and non-metals, organic and inorganic compounds. Of practical importance is the combustion of fuel, which is mainly a complex mixture of hydrocarbons with small amounts of oxygen, sulfur, nitrogen and other elements.
Classic oxidizer –oxygen
A simple substance or chemical compound in which atoms attach electrons is called an oxidizing agent. A classic example of such a substance is oxygen, which turns into oxides after the reaction. But also an oxidizing agent in redox reactions is ozone, which is reduced to organic substances (for example, ketones and aldehydes), peroxides, hypochlorites, chlorates, nitric and sulfuric acids, manganese oxide and permanganate. It is easy to see that all these substances contain oxygen.
Other common oxidizers
However, the redox reaction is not only a process involving oxygen. Instead, halogens, chromium, and even metal cations and a hydrogen ion (if it turns into a simple substance as a result of the reaction) can act as an oxidizing agent.
How many electrons will be accepted depends largely on the concentration of the oxidizing agent, as well as on the activity of the metal interacting with it. For example, in the reaction of concentrated nitric acid with a metal (zinc), 3 electrons can be accepted, and in the interaction of the same substances, provided that the acid is in a very dilute form, already 8 electrons.
The strongest oxidizers
All oxidizing agents differ in the strength of their properties. So, the hydrogen ion has a low oxidizing ability, while atomic chlorine, formed in aqua regia (a mixture of nitric and hydrochloric acids in a ratio of 1:3), can oxidize even gold and platinum.
Concentrated selenic acid has similar properties. This makes it unique among other organic acids. When diluted, it is not able to interact with gold, but it is still stronger than sulfuric acid, and can even oxidize other acids, such as hydrochloric acid.
Another example of a strong oxidizing agent is potassium permanganate. It successfully interacts with organic compounds and is able to break strong carbon bonds. Copper oxide, cesium ozonide, cesium superoxide, as well as xenon difluoride, tetrafluoride and xenon hexafluoride also have high activity. Their oxidizing ability is due to the high electrode potential when reacting in a dilute aqueous solution.
However, there are substances in which this potential is even higher. Among inorganic molecules, fluorine is the strongest oxidizing agent, but it is not able to act on the inert gas xenon without additional heat and pressure. But this is successfully coped with by platinum hexafluoride, difluorodioxide, krypton difluoride, silver difluoride, divalent silver s alts and some other substances. For their unique ability to redox reactions, they are classified as very strong oxidizers.
Recovery
Originally, the term "recovery" was synonymous with deoxidation, that is, the deprivation of a substance of oxygen. However, over time, the word acquired a new meaning, it meant the extraction of metals from compounds containing them, as well as any chemical transformations in whichthe electronegative part of a substance is replaced by a positively charged element, such as hydrogen.
The complexity of the process largely depends on the chemical affinity of the elements in the compound. The weaker it is, the easier the reaction is carried out. Typically, the affinity is weaker in endothermic compounds (heat is absorbed during their formation). Their recovery is quite simple. A striking example of this is explosives.
In order for a reaction involving exothermic compounds (formed with the release of heat), a strong source of energy, such as an electric current, must be applied.
Standard reducing agents
The most ancient and common reducing agent is coal. It mixes with ore oxides, when heated, oxygen is released from the mixture, which combines with carbon. The result is a powder, granules or metal alloy.
Another common reducing agent is hydrogen. It can also be used to mine metals. To do this, the oxides are clogged into a tube through which a stream of hydrogen is passed. Basically, this method is applied to copper, lead, tin, nickel or cob alt. You can apply it to iron, but the reduction will be incomplete and water is formed. The same problem is observed when trying to treat zinc oxides with hydrogen, and it is aggravated by the volatility of the metal. Potassium and some other elements are not reduced by hydrogen at all.
Features of reactions in organic chemistry
In progressthe reduction particle accepts electrons and thereby lowers the oxidation number of one of its atoms. However, it is convenient to determine the essence of the reaction by changing the oxidation state with the participation of inorganic compounds, while in organic chemistry it is difficult to calculate the oxidation number, it often has a fractional value.
To navigate the redox reactions involving organic substances, you need to remember the following rule: reduction occurs when a compound gives up oxygen atoms and acquires hydrogen atoms, and vice versa, oxidation is characterized by the addition of oxygen.
The reduction process is of great practical importance for organic chemistry. It is he who underlies catalytic hydrogenation used for laboratory or industrial purposes, in particular, the purification of substances and systems from hydrocarbon and oxygen impurities.
The reaction can proceed both at low temperatures and pressures (up to 100 degrees Celsius and 1-4 atmospheres, respectively), and at high temperatures (up to 400 degrees and several hundred atmospheres). The production of organic substances requires complex instruments to provide the right conditions.
Active platinum group metals or non-precious nickel, copper, molybdenum and cob alt are used as catalysts. The latter option is more economical. Restoration occurs due to the simultaneous sorption of the substrate and hydrogen with the facilitation of the reaction between them.
Reduction reactions proceedand inside the human body. In some cases, they can be useful and even vital, in others they can lead to serious negative consequences. For example, nitrogen-containing compounds in the body are converted into primary amines, which, among other useful functions, constitute protein substances that are the building material of tissues. At the same time, aniline-dyed foods produce toxic compounds.
Types of reactions
What kind of redox reactions, it becomes clear if you look at the presence of changes in oxidation states. But within this type of chemical transformation, there are variations.
So, if molecules of different substances participate in the interaction, one of which includes an oxidizing atom, and the other a reducing agent, the reaction is considered intermolecular. In this case, the redox reaction equation can be as follows:
Fe + 2HCl=FeCl2 + H2.
The equation shows that the oxidation states of iron and hydrogen change, while they are part of different substances.
But there are also intramolecular redox reactions, in which one atom in a chemical compound is oxidized and another is reduced, and new substances are obtained:
2H2O=2H2 + O2.
A more complex process occurs when the same element acts as an electron donor and acceptor and forms several new compounds, which are included in different oxidation states. Such a process is calleddismutation or disproportionation. An example of this is the following transformation:
4KClO3=KCl + 3KClO4.
From the above equation of the redox reaction, it can be seen that Bertolet s alt, in which chlorine is in the oxidation state of +5, decomposes into two components - potassium chloride with the oxidation state of chlorine -1 and perchlorate with an oxidation number of +7. It turns out that the same element simultaneously increased and lowered its oxidation state.
The reverse of the dismutation process is the reaction of coproportionation or reproportionation. In it, two compounds, which contain the same element in different oxidation states, react with each other to form a new substance with a single oxidation number:
SO2 +2H2S=3S + 2H2O.
As you can see from the above examples, in some equations, the substance is preceded by numbers. They show the number of molecules involved in the process and are called stoichiometric coefficients of redox reactions. For the equation to be correct, you need to know how to arrange them.
E-balance method
The balance in redox reactions is always preserved. This means that the oxidizing agent accepts exactly as many electrons as were given away by the reducing agent. To correctly compose an equation for a redox reaction, you need to follow this algorithm:
- Determine the oxidation states of the elements before and after the reaction. For example, inreaction between nitric acid and phosphorus in the presence of water produces phosphoric acid and nitric oxide: HNO3 + P + H2O=H 3PO4 + NO. Hydrogen in all compounds has an oxidation state of +1, and oxygen has -2. For nitrogen, before the reaction begins, the oxidation number is +5, and after it proceeds +2, for phosphorus - 0 and +5, respectively.
- Mark the elements in which the oxidation number has changed (nitrogen and phosphorus).
- Compose electronic equations: N+5 + 3e=N+2; R0 - 5e=R+5.
- Equalize the number of received electrons by choosing the least common multiple and calculating the multiplier (the numbers 3 and 5 are divisors for the number 15, respectively, the multiplier for nitrogen is 5, and for phosphorus 3): 5N+5 + (3 x 5)e=5N+2; 3P0 - 15e=3P+5.
- Add the resulting half-reactions according to the left and right parts: 5N+5 + 3P0=5N+ 2 - 15th=3Р+5. If everything is done correctly at this stage, the electrons will shrink.
- Rewrite the equation completely, putting down the coefficients according to the electronic balance of the redox reaction: 5HNO3 + 3P + H2O=3H 3PO4 + 5NO.
- Check whether the number of elements before and after the reaction remains the same everywhere, and if necessary, add coefficients in front of other substances (in this example, the amount of hydrogen and oxygen did not equalize, in order for the reaction equation to look correct, you need to add a coefficient in front ofwater): 5HNO3 + 3P + 2H2O=3H3PO 4 + 5NO.
Such a simple method allows you to correctly place the coefficients and avoid confusion.
Examples of reactions
An illustrative example of a redox reaction is the interaction of manganese with concentrated sulfuric acid, proceeding as follows:
Mn + 2H2SO4=MnSO4 + SO 2 + 2 H2O.
The redox reaction proceeds with a change in the oxidation states of manganese and sulfur. Prior to the start of the process, manganese was in an unbound state and had a zero oxidation state. But when interacting with sulfur, which is part of the acid, it increased the oxidation state to +2, thus acting as an electron donor. Sulfur, on the contrary, played the role of an acceptor, lowering the oxidation state from +6 to +4.
However, there are also reactions in which manganese acts as an electron acceptor. For example, this is the interaction of its oxide with hydrochloric acid, proceeding according to the reaction:
MnO2+4HCl=MnCl2+Cl2+2 H2O.
The redox reaction in this case proceeds with a decrease in the oxidation state of manganese from +4 to +2 and an increase in the oxidation state of chlorine from -1 to 0.
Previously, the oxidation of sulfur oxide with nitrogen oxide in the presence of water, which produced 75% sulfuric acid, was of great practical importance:
SO2 + NO2 + H2O=NO + H2So4.
The redox reaction used to be carried out in special towers, and the final product was called tower. Now this method is far from the only one in the production of acid, since there are other modern methods, for example, contact using solid catalysts. But obtaining acid by the redox reaction method has not only industrial, but also historical significance, since it was precisely such a process that spontaneously occurred in the air of London in December 1952.
The anticyclone then brought unusually cold weather, and the townspeople began to use a lot of coal to heat their homes. Since this resource was of poor quality after the war, a large amount of sulfur dioxide was concentrated in the air, which reacted with moisture and nitrogen oxide in the atmosphere. As a result of this phenomenon, the mortality of infants, the elderly and those suffering from respiratory diseases has increased. The event was given the name of the Great Smog.
Thus, redox reactions are of great practical importance. Understanding their mechanism allows you to better understand natural processes and achieve new substances in the laboratory.
