The conversion of one substance to 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 it you can get a huge amount of useful and useful substances that are found in nature in small quantities or do not exist in kind. Among the most important varieties, redox reactions (abbreviated as OVR or redox) are distinguished. They are characterized by a change in the oxidation state of atoms or ions.
Processes during the reaction
During the reaction, two processes take place - oxidation and reduction. The first of them is characterized by electron donation by reducing agents (donors) with an increase in their oxidation state, and the second by electron attachment by oxidizing agents (acceptors) with a decrease in oxidation state. The most commonly used reducing agents are metals and non-metal compounds in the lowest oxidation state (hydrogen sulfide, ammonia). Typical oxidizing agents are halogens, nitrogen, oxygen, as well as substances that contain an element of the highest oxidation state (nitric or sulfuric acid). Electrons can be donated or attached by atoms, ions, molecules.
Until 1777, the hypothesis was that oxidation results in the loss of an invisible combustible substance called phlogiston. However, the theory of combustion created by A. Lavoisier convinced scientists that oxidation occurs upon interaction with oxygen, and reduction under the influence of hydrogen. Only after a while did it become clear that not only hydrogen and oxygen can influence redox reactions.
Oxidation
The oxidation process can occur in the liquid and gaseous phases, as well as on the surface of solids. A special role is played by electrochemical oxidation, which occurs in solutions or melts at the anode (an electrode connected to the positive pole of the power source). For example, when fluoride is melted by electrolysis (decomposition of a substance into constituent elements on electrodes), the strongest inorganic oxidizer, fluorine, is obtained.
Another classic example of oxidation is combustion in an atmosphere of air and pure oxygen. Various substances are capable of this process: metals and non-metals, organic and inorganic compounds. Fuel combustion is of practical importance, which is mainly a complex mixture of hydrocarbons with small amounts of oxygen, sulfur, nitrogen and other elements.
The classic oxidizer is 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 notice that all of these substances contain oxygen.
Other common oxidizing agents
However, the redox reaction is not only a process proceeding with the participation of oxygen. Instead, halogens, chromium and even metal cations and a hydrogen ion can act as an oxidizing agent (if the reaction turns into a simple substance).
How many electrons will be taken 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 taken, and in the interaction of the same substances, provided that the acid is in a very dilute form, already 8 electrons.
The strongest oxidizing agents
All oxidizing agents differ in the strength of their properties. So, a hydrogen ion has a low oxidizing ability, while atomic chlorine formed in a solution of aqua regia (a mixture of nitric and hydrochloric acids in a ratio of 1: 3) is capable of oxidizing even gold and platinum.
Concentrated selenic acid has similar properties . This makes it unique among other organic acids. In a diluted form, it is not able to interact with gold, but it still remains stronger than sulfuric acid, and can even oxidize other acids, for example, hydrochloric.
Another example of a strong oxidizing agent is potassium permanganate. It successfully interacts with organic compounds and is capable of breaking strong carbon bonds. Copper oxide, cesium ozonide, cesium superoxide, and also difluoride, tetrafluoride and xenon hexafluoride also have high activity. Their oxidizing ability is due to the high electrode potential during the reaction 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 capable of acting on an inert gas, xenon, without additional supply of heat and pressure. But platinum hexafluoride, difluorodioxide, krypton difluoride, silver difluoride, salts of divalent silver and some other substances successfully cope with this. Due to their unique ability to redox reactions, they are considered to be very strong oxidizing agents.
Recovery
Initially, the term "recovery" was synonymous with deoxidation, that is, deprivation 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 which the electronegative part of a substance is replaced by a positively charged element, for example, 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. Usually, affinity is weaker in endothermic compounds (heat is absorbed during their formation). Their recovery is quite simple. A vivid example of this is explosives.
In order for a reaction to occur with the participation of exothermic compounds (formed with the release of heat), you need to apply a strong source of energy, for example, electric current.
Standard reducing agents
The oldest and most common reducing agent is coal. It mixes with ore oxides, when heated, oxygen is released from the mixture, combining with carbon. As a result, a powder, granules or metal alloy is formed.
Another common reducing agent is hydrogen. It can also be used to mine metals. To do this, the oxides are driven into a tube through which a stream of hydrogen is passed. Basically, this method is applied to copper, lead, tin, nickel or cobalt. You can apply it to iron, but the recovery will be incomplete and water will form. The same problem is observed when trying to expose zinc oxides to hydrogen, moreover, it is exacerbated by the volatility of the metal. Potassium and some other elements cannot be reduced with hydrogen at all.
Features of reactions in organic chemistry
In the process of reduction, the particle receives electrons and thereby reduces the oxidative number of one of its atoms. However, it is convenient to determine the nature of the reaction by changing the oxidation state with the participation of inorganic compounds, whereas in organic chemistry the oxidation number is difficult to calculate, it often has a fractional value.
To navigate in redox reactions involving organic substances, you need to remember the following rule: reduction occurs when a compound gives off 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 the catalytic hydrogenation used for laboratory or industrial purposes, in particular, the purification of substances and systems from impurities of hydrocarbons and oxygen.
The reaction can proceed both at low temperatures and pressures (up to 100 degrees Celsius and 1-4 atmospheres, respectively), and at high (up to 400 degrees and several hundred atmospheres). The production of organic substances requires sophisticated instruments that provide suitable conditions.
In this case, the active metals of the platinum group or base metal nickel, copper, molybdenum and cobalt are used as catalysts. The latter option is more economical. Recovery occurs due to the simultaneous sorption of the substrate and hydrogen with facilitating the reaction between them.
Recovery reactions occur 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 to primary amines, which, among other useful functions, are protein substances, which are the building material of tissues. At the same time, when using products stained with aniline, toxic compounds are formed.
Types of reactions
What kind of redox reactions, it becomes clear if you look at the presence of changes in oxidation states. But inside this type of chemical transformation there are varieties.
So, if the interaction involves molecules of different substances, one of which includes an oxidizing atom, and the other a reducing agent, the reaction is considered intermolecular. The equation of the redox reaction can be as follows:
Fe + 2HCl = FeCl 2 + H 2 .
It can be seen from the equation that the degrees of oxidation change for iron and hydrogen, while they are part of different substances.
But there are also intramolecular redox reactions in which one atom is oxidized in a chemical compound and the other is reduced, and new substances are obtained:
2H 2 O = 2H 2 + O 2 .
A more complex process occurs when one and the same element acts as a donor and acceptor of electrons and forms several new compounds, which enter into different degrees of oxidation. This process is called dismutation or disproportionation. His example is the following transformation:
4KClO 3 = KCl + 3KClO 4 .
From the above equation of the redox reaction, it can be seen that the Bertolet salt, in which chlorine is in the oxidation state of +5, decomposes into two components - potassium chloride with an oxidation state of chlorine of -1 and perchlorate with an oxidation number of +7. It turns out that one and the same element simultaneously increased and decreased its oxidation state.
The reverse process of dismutation is the reaction of disproportionation or reproportionation. It contains two compounds, which contain the same element in different oxidation states, react with each other with the formation of a new substance with a single oxidizing number:
SO 2 + 2H 2 S = 3S + 2H 2 O.
As can be seen from the above examples, in some equations, the numbers are in front of the substance. They show the number of molecules involved in the process and are called the stoichiometric coefficients of redox reactions. In order for the compiled equation to be correct, you need to know how to arrange them.
Electronic balance method
The balance of redox reactions is always maintained. This means that the oxidizing agent receives exactly as many electrons as were given by the reducing agent. To correctly formulate the equation of the redox reaction, you need to follow this algorithm:
- Determine the oxidation state of the elements before and after the reaction. For example, in the reaction between nitric acid and phosphorus in the presence of water, orthophosphoric acid and nitric oxide are obtained: HNO 3 + P + H 2 O = H 3 PO 4 + NO. In all compounds, hydrogen has an oxidation state of +1, and oxygen has a -2. In nitrogen, before the start of the reaction, the oxidation number is +5, and after its occurrence +2, in phosphorus - 0 and +5, respectively.
- Mark the elements in which the oxidation number has changed (nitrogen and phosphorus).
- Create electronic equations: N +5 + 3e = N +2 ; P 0 - 5e = P +5 .
- Equalize the number of electrons received by selecting the smallest total multiple and calculating the factor (numbers 3 and 5 are divisors for number 15, respectively, the factor for nitrogen is 5 and for phosphorus 3): 5N +5 + (3 x 5) e = 5N + 2 ; 3P 0 - 15e = 3P +5 .
- Add the resulting half-reactions in accordance with the left and right sides: 5N +5 + 3P 0 = 5N + 2 - 15e = 3P +5 . If everything is done correctly at this stage, the electrons will contract.
- Rewrite the equation completely, putting down the coefficients according to the electronic balance of the redox reaction: 5HNO 3 + 3P + H 2 O = 3H 3 PO 4 + 5NO.
- Check whether everywhere the number of elements before and after the reaction remains the same, and if necessary, add coefficients to other substances (in this example, the amount of hydrogen and oxygen was not equalized, so that the reaction equation looks right, you need to add the coefficient before water): 5HNO 3 + 3P + 2H 2 O = 3H 3 PO 4 + 5NO.
Such a simple method allows you to correctly arrange the coefficients and prevent confusion.
Reaction Examples
A good example of the redox reaction is the interaction of manganese with concentrated sulfuric acid, proceeding according to the following scheme:
Mn + 2H 2 SO 4 = MnSO 4 + SO 2 + 2 H2O.
The redox reaction proceeds with a change in the oxidation state of manganese and sulfur. Before the process began, manganese was in an unbound state and had a zero oxidation state. But when interacting with sulfur, which is part of the acid, he increased the oxidation state to +2, thus acting as an electron donor. Sulfur, in contrast, played the role of an acceptor, lowering the oxidation state from +6 to +4.
However, there are also such reactions in which manganese acts as an electron acceptor. For example, this is the interaction of its oxide with hydrochloric acid, proceeding by the reaction:
MnO 2 + 4HCl = MnCl 2 + Cl 2 + 2 H2O.
In this case, the redox reaction proceeds with a decrease in the degree of oxidation of manganese from +4 to +2 and an increase in the degree of oxidation of chlorine from -1 to 0.
Earlier, the oxidation of sulfur oxide with nitric oxide in the presence of water, which produced 75% sulfuric acid, was of great practical importance:
SO 2 + NO 2 + H 2 O = NO + H2So4.
The redox reaction used to be carried out in special towers, and the final product was called a 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 is not only industrial, but also of historical importance, since it was 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 nitric oxide in the atmosphere. As a result of this phenomenon, mortality of infants, the elderly and those suffering from respiratory diseases has increased. The event was given the name Great Smog.
Thus, redox reactions are of great practical importance. Understanding their mechanism makes it possible to better understand natural processes and to obtain new substances in the laboratory.