Electronegativity. Oxidation state and valence of chemical elements

A chemical reaction is the process by which the starting materials are converted into reaction products. Substances obtained after the end of the reaction are called products. They may differ from the original ones in structure, composition, or both.

By changing the composition, the following types of chemical reactions are distinguished:

with a change in composition (such are the majority);
without changing the composition (isomerization and transformation of one allotropic modification into another).

If the composition of a substance does not change as a result of the reaction, then its structure necessarily changes, for example: Graphite Diamond

Let us consider in more detail the classification of chemical reactions proceeding with a change in composition.

I. By the number and composition of substances

Compound reactions

As a result of such chemical processes, one substance is formed from several substances: A + B + ... = C

Can connect:

simple substances: 2Na + S = Na2S;
simple with complex: 2SO2 + O2 = 2SO3;
two complex ones: CaO + H2O = Ca (OH) 2.
more than two substances: 4Fe + 3O2 + 6H2O = 4Fe (OH) 3

Decomposition reactions

One substance in such reactions decomposes into several others: A = B + C + ...

Products in this case can be:

simple substances: 2NaCl = 2Na + Cl2
simple and complex: 2KNO3 = 2KNO2 + O2
two complex ones: CaCO3 = CaO + CO2
more than two products: 2AgNO3 = 2Ag + O2 + 2NO2

Substitution reactions

Such reactions in which simple and complex substances react with each other, and the atoms of the simple substance replace the atoms of one of the elements in the complex, and are called substitution reactions. The process of atomic substitution can be shown schematically as follows: A + BC = B + AC.

For example, CuSO4 + Fe = FeSO4 + Cu

Exchange reactions

This group includes reactions in which two complex substances change their parts: AB + CD = AD + CB. According to Berthollet's rule, the irreversible course of such reactions is possible if at least one of the products:

precipitate (insoluble matter): 2NaOH + CuSO4 = Cu (OH) 2 + Na2SO4;
slightly dissociating substance: NaOH + HCl = NaCl + H2O;
gas: NaOH + NH4Cl = NaCl + NH3 + H2O (first, ammonia hydrate NH3 H2O is formed, which, upon receipt, immediately decomposes into ammonia and water).

II. Thermal effect

Exothermic - processes proceeding with the release of heat:
C + O2 = CO2 + Q
Endothermic - reactions in which heat is absorbed:
Cu (OH) 2 = CuO + H2O - Q

III. Types of chemical reactions in the direction

Reversible are called reactions that occur at the same time, both in the forward and in the opposite direction: N2 + O2 ↔ 2NO
Irreversible the processes proceed to the end, that is, until at least one of the reacting substances is completely consumed. Examples of irreversible exchange reactions were discussed above.

IV. By the presence of a catalyst

V. According to the state of aggregation of substances

If all reactants are in the same state of aggregation, the reaction is called homogeneous... Such processes take place in their entirety. For example: NaOH + HCl = NaCl + H2O
Heterogeneous called the reactions between substances in different states of aggregation, occurring at the interface. For example: Zn + 2HCl = ZnCl2 + H2

Vi. Types of chemical reactions to change the oxidation state of reactants

Redox (ОВР) - reactions in which the oxidation states of the reacting substances change.
Reactions proceeding without changing oxidation states reagents (BISO).

The combustion and substitution processes are always redox processes. Metabolic reactions proceed without changing the oxidation states of substances. All other processes can be either OVR or BISO.

OXIDATION-REDUCTION REACTIONS

Oxidation state

The oxidation state is the conditional charge of an atom in a molecule, calculated on the assumption that the molecule consists of ions and is generally electrically neutral.

The most electronegative elements in the compound have negative oxidation states, while the atoms of elements with less electronegativity are positive.

The oxidation state is a formal concept; in some cases, the oxidation state does not coincide with the valence.

For example:

N2H4 (hydrazine)

oxidation state of nitrogen - -2; nitrogen valence - 3.

Calculation of the oxidation state

To calculate the oxidation state of an element, the following points should be considered:

1. The oxidation states of atoms in simple substances are zero (Na 0; H2 0).

2. The algebraic sum of the oxidation states of all atoms that make up a molecule is always zero, and in a complex ion this sum is equal to the charge of the ion.

3.Atoms have a constant oxidation state: alkali metals (+1), alkaline earth metals (+2), hydrogen (+1) (except for hydrides NaH, CaH2, etc., where the oxidation state of hydrogen is -1), oxygen (-2) (except for F 2 -1 O +2 and peroxides containing the –O – O– group, in which the oxidation state of oxygen is -1).

4. For elements, the positive oxidation state cannot exceed a value equal to the group number of the periodic system.

Examples:

V 2 +5 O 5 -2 ; Na 2 +1 B 4 +3 O 7 -2 ; K +1 Cl +7 O 4 -2 ; N -3 H 3 +1 ; K2 +1 H +1 P +5 O 4 -2 ; Na 2 +1 Cr 2 +6 O 7 -2

Reactions without and with a change in the oxidation state

There are two types of chemical reactions:

A Reactions in which the oxidation state of the elements does not change:

Addition reactions

SO 2 + Na 2 O → Na 2 SO 3

Decomposition reactions

Cu (OH) 2 → CuO + H 2 O

Exchange reactions

AgNO 3 + KCl → AgCl + KNO 3

NaOH + HNO 3 → NaNO 3 + H 2 O

B Reactions in which there is a change in the oxidation states of the atoms of the elements that make up the reacting compounds:

2Mg 0 + O 2 0 → 2Mg +2 O -2

2KCl +5 O 3 -2 → 2KCl -1 + 3O 2 0

2KI -1 + Cl 2 0 → 2KCl -1 + I 2 0

Mn +4 O 2 + 4HCl -1 ® Mn +2 Cl 2 + Cl +1 2 0 + 2H 2 O

Such reactions are called redox reactions.

Redox reactions are reactions in which there is a change in the oxidation state of atoms. Redox reactions are very common. All combustion reactions are redox.
Redox reaction consists of two processes that cannot occur separately from each other. The process of increasing the oxidation state is called oxidation. Simultaneously with oxidation, reduction occurs, that is, the process of lowering the oxidation state.

Oxidation, reduction

Accordingly, two main participants are distinguished in redox reactions: an oxidizing agent and a reducing agent. The process of donating electrons is oxidation. Oxidation increases the oxidation state. The oxidizing agent in the course of the reaction lowers its oxidation state by being reduced. Here one should distinguish between an oxidizing chemical element and an oxidizing agent.

N +5 - oxidizing agent; HN +5 O 3 and NaN +5 O 3 - oxidizing agents.
If we say that nitric acid and its salts are strong oxidizing agents, then by this we mean that the oxidizing agent is nitrogen atoms with an oxidation state of +5, and not the whole substance as a whole.
The second obligatory participant in the redox reaction is called a reducing agent. The process of electron attachment is reduction. Reduction reduces the oxidation state.

The reducing agent increases its oxidation state by being oxidized during the reaction. As in the case of an oxidizing agent, a distinction should be made between a reducing agent and a reducing chemical element. Carrying out the reaction of reduction of aldehyde to alcohol, we cannot just take hydrogen with an oxidation state of -1, but we take some kind of hydride, preferably lithium aluminum hydride.

H -1 - reducing agent; NaH -1 and LiAlH -1 4 - reducing agents.
In redox reactions, the complete transition of electrons from a reducing agent to an oxidizing agent is extremely rare, since there are few compounds with an ionic bond. But when placing the coefficients, we proceed from the assumption that such a transition does occur. This makes it possible to correctly determine the basic coefficients before the formulas of the oxidizing agent and reducing agent.
5H 2 SO 3 + 2КМnO 4 = 2H 2 SO 4 + 2MnSO 4 + К 2 SO 4 + 3Н 2 О
S +4 - 2e → S +6 5 - reducing agent, oxidation
Mn +7 + 5e → Mn +2 2 - oxidizing agent, reduction

The atoms or ions that attach electrons in this reaction are oxidizing agents, and those that donate electrons are reducing agents.

Redox properties of a substance and the oxidation state of its constituent atoms

Compounds containing atoms of elements with the maximum oxidation state can only be oxidizing agents due to these atoms, because they have already given up all their valence electrons and are only able to accept electrons. The maximum oxidation state of an atom of an element is equal to the number of the group in the periodic table to which this element belongs. Compounds containing atoms of elements with a minimum oxidation state can only serve as reducing agents, since they are only capable of donating electrons, because the external energy level of such atoms is completed with eight electrons. The minimum oxidation state for metal atoms is 0, for non-metals - (n – 8) (where n is the group number in the periodic system). Compounds containing atoms of elements with an intermediate oxidation state can be both oxidizing and reducing agents, depending on the partner with which they interact and on the reaction conditions.

The most important reducing and oxidizing agents

Reducing agents:

Metals,

hydrogen,

coal.

Carbon monoxide (II) (CO).

Hydrogen sulfide (H 2 S);

sulfur oxide (IV) (SO 2);

sulfurous acid H 2 SO 3 and its salts.

Hydrohalic acids and their salts.

Metal cations in the lowest oxidation states: SnCl 2, FeCl 2, MnSO 4, Cr 2 (SO 4) 3.

Nitrous acid HNO 2;

ammonia NH 3;

hydrazine NH 2 NH 2;

nitric oxide (II) (NO).

Electrolysis cathode.

Oxidants

Halogens.

Potassium permanganate (KMnO 4);

potassium manganate (K 2 MnO 4);

manganese (IV) oxide (MnO 2).

Potassium dichromate (K 2 Cr 2 O 7);

potassium chromate (K 2 CrO 4).

Nitric acid (HNO 3).

Sulfuric acid (H 2 SO 4) conc.

Copper (II) oxide (CuO);

lead (IV) oxide (PbO 2);

silver oxide (Ag 2 O);

hydrogen peroxide (H 2 O 2).

Iron (III) chloride (FeCl 3).

Berthollet's salt (KClO 3).

Electrolysis anode.

Each such half-reaction is characterized by a standard redox potential E 0 (dimension - volts, V). The more E 0, the stronger the oxidative form as an oxidizing agent and the weaker the reduced form as a reducing agent, and vice versa.

The half-reaction is taken as the reference point for the potentials: 2H + + 2ē ® H 2, for which E 0 = 0

For half-reactions M n + + nē ® M 0, E 0 is called the standard electrode potential. According to the magnitude of this potential, metals are usually arranged in a number of standard electrode potentials (a number of metal voltages):

Li, Rb, K, Ba, Sr, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Cd,

Co, Ni, Sn, Pb, H, Sb, Bi, Cu, Hg, Ag, Pd, Pt, Au

A number of stresses characterize the chemical properties of metals:

1. The more to the left the metal is located in the series of voltages, the stronger its reducing ability and the weaker the oxidizing ability of its ion in solution (ie, the easier it gives up electrons (oxidizes) and the more difficult its ions attach electrons back).

2. Each metal is capable of displacing from salt solutions those metals that are in the series of voltages to the right of it, i.e. restores the ions of subsequent metals into electrically neutral atoms, donating electrons and itself turning into ions.

3. Only metals in the series of voltages to the left of hydrogen (H) are capable of displacing it from acid solutions (for example, Zn, Fe, Pb, but not Cu, Hg, Ag).

Galvanic cells

Every two metals, being immersed in solutions of their salts, which communicate with each other by means of a siphon filled with electrolyte, form a galvanic cell. Plates of metals immersed in solutions are called cell electrodes.

If you connect the outer ends of the electrodes (poles of the element) with a wire, then electrons begin to move from the metal with a lower potential to the metal with a higher potential (for example, from Zn to Pb). The escape of electrons disturbs the equilibrium that exists between the metal and its ions in the solution, and causes a new amount of ions to pass into the solution - the metal gradually dissolves. At the same time, electrons passing to another metal discharge ions in the solution at its surface - the metal is released from the solution. The electrode on which oxidation takes place is called the anode. The electrode on which the reduction takes place is called the cathode. In a lead-zinc cell, the zinc electrode is the anode and the lead electrode is the cathode.

Thus, in a closed galvanic cell, an interaction occurs between a metal and a solution of a salt of another metal, which are not in direct contact with each other. The atoms of the first metal, donating electrons, are converted into ions, and the ions of the second metal, by attaching electrons, are converted into atoms. The first metal displaces the second from its salt solution. For example, during the operation of a galvanic cell composed of zinc and lead immersed in Zn (NO 3) 2 and Pb (NO 3) 2 solutions, the following processes occur at the electrodes:

Zn - 2ē → Zn 2+

Pb 2+ + 2ē → Pb

Summing up both processes, we obtain the equation Zn + Pb 2+ → Pb + Zn 2+, which expresses the reaction in the element in the ionic form. The molecular equation of the same reaction will be:

Zn + Pb (NO 3) 2 → Pb + Zn (NO 3) 2

The electromotive force of a galvanic cell is equal to the potential difference between its two electrodes. When determining it, the smaller one is always subtracted from the larger potential. For example, the electromotive force (EMF) of the considered element is equal to:

E.m.s. =	-0,13		(-0,76)	0.63v
	E Pb		E Zn

It will have such a value provided that the metals are immersed in solutions in which the ion concentration is 1 g-ion / l. At other concentrations of solutions, the values of the electrode potentials will be somewhat different. They can be calculated using the formula:

E = E 0 + (0.058 / n) lgC

where E is the sought-for potential of the metal (in volts)

E 0 - its normal potential

n - valence of metal ions

С - concentration of ions in solution (g-ion / l)

Example

Find the electromotive force of the element (emf) formed by a zinc electrode immersed in a 0.1 M solution of Zn (NO 3) 2 and a lead electrode immersed in a 2 M solution of Pb (NO 3) 2.

Solution

We calculate the potential of the zinc electrode:

E Zn = -0.76 + (0.058 / 2) log 0.1 = -0.76 + 0.029 (-1) = -0.79 v

We calculate the potential of the lead electrode:

E Pb = -0.13 + (0.058 / 2) log 2 = -0.13 + 0.029 0.3010 = -0.12 v

We find the electromotive force of the element:

E. d. With. = -0.12 - (-0.79) = 0.67 v

Electrolysis

Electrolysis the process of decomposition of a substance by an electric current is called.

The essence of electrolysis lies in the fact that when current is passed through an electrolyte solution (or molten electrolyte), positively charged ions move to the cathode, and negatively charged ions move to the anode. Upon reaching the electrodes, the ions are discharged, as a result of which the constituents of the dissolved electrolyte or hydrogen and oxygen from the water are released from the electrodes.

To convert different ions into neutral atoms or groups of atoms, different voltages of electric current are required. Some ions lose their charges more easily, others more difficult. The ease with which metal ions are discharged (electrons are attached) is determined by the position of the metals in the series of voltages. The farther to the left the metal is in the series of voltages, the greater its negative potential (or less positive potential), the more difficult, other things being equal, its ions are discharged (the ions Au 3+, Ag + are discharged most easily; the most difficult are Li +, Rb +, K +).

If there are ions of several metals in the solution at the same time, then first of all, the ions of the metal with a lower negative potential (or more positive) are discharged. For example, from a solution containing Zn 2+ and Cu 2+ ions, metallic copper is first precipitated. But the magnitude of the potential of the metal also depends on the concentration of its ions in the solution; the ease of discharge of ions of each metal changes in the same way, depending on their concentration: an increase in concentration facilitates the discharge of ions, while a decrease makes it difficult. Therefore, during the electrolysis of a solution containing ions of several metals, it can happen that the release of a more active metal will occur earlier than the release of a less active metal (if the concentration of ions of the first metal is significant, and the second is very low).

In aqueous solutions of salts, in addition to salt ions, there are always water ions (H + and OH -). Of these, hydrogen ions will be discharged more easily than ions of all metals preceding hydrogen in a series of voltages. However, due to the negligible concentration of hydrogen ions during the electrolysis of all salts, except for the salts of the most active metals, a metal is released at the cathode, and not hydrogen. Only during the electrolysis of salts of sodium, calcium and other metals up to and including aluminum, hydrogen ions are discharged and hydrogen is released.

At the anode, either ions of acid residues or hydroxyl ions of water can be discharged. If the ions of acidic residues do not contain oxygen (Cl -, S 2-, CN -, etc.), then these ions are usually discharged, and not hydroxyl ions, which lose their charge much more difficult, and Cl 2, S, etc. are released at the anode. .d. On the contrary, if the salt of an oxygen-containing acid or the acid itself is subjected to electrolysis, then the hydroxyl ions are discharged, and not the ions of the oxygen residues. The neutral OH groups formed during the discharge of hydroxyl ions immediately decompose according to the equation:

4OH → 2H 2 O + O 2

As a result, oxygen is released at the anode.

Electrolysis of nickel chloride solution NiCl 2

The solution contains ions Ni 2+ and Cl -, as well as ions H + and OH - in negligible concentration. When a current is passed, Ni 2+ ions move to the cathode, and Cl ions - to the anode. Taking two electrons from the cathode, Ni 2+ ions turn into neutral atoms released from the solution. The cathode is gradually covered with nickel.

Chlorine ions, reaching the anode, donate electrons to it and turn into chlorine atoms, which, when combined in pairs, form chlorine molecules. Chlorine is released at the anode.

Thus, at the cathode, recovery process, at the anode - oxidation process.

Electrolysis of potassium iodide solution KI

Potassium iodide is in solution in the form of K + and I - ions. When a current is passed, the K + ions move to the cathode, and the I - ions to the anode. But since potassium is in the series of voltages much more to the left of hydrogen, then not potassium ions are discharged at the cathode, but hydrogen ions of water. The resulting hydrogen atoms combine to form H2 molecules, and thus hydrogen is released at the cathode.

As the hydrogen ions discharge, more and more water molecules dissociate, as a result of which hydroxyl ions (released from the water molecule) accumulate at the cathode, as well as K + ions, which continuously move to the cathode. A KOH solution is formed.

At the anode, iodine is released, since the I - ions are discharged more easily than the hydroxyl ions of water.

Potassium sulfate solution electrolysis

The solution contains ions K +, SO 4 2- and ions H + and OH - from water. Since K + ions are discharged more difficult than H + ions, and SO 4 2- ions than OH - ions, then when an electric current is passed, hydrogen ions will be discharged at the cathode, hydroxyl groups at the anode, that is, electrolysis of water... At the same time, due to the discharge of hydrogen and hydroxyl ions of water and the continuous movement of K + ions to the cathode, and SO 4 2- ions to the anode, an alkali solution (KOH) is formed at the cathode, and a sulfuric acid solution at the anode.

Electrolysis of copper sulfate solution at a copper anode

Electrolysis proceeds in a special way when the anode is made of the same metal, the salt of which is in solution. In this case, no ions are discharged at the anode, but the anode itself gradually dissolves, sending ions into the solution and donating electrons to the current source.

The whole process is reduced to the release of copper at the cathode and the gradual dissolution of the anode. The amount of CuSO 4 in the solution remains unchanged.

Electrolysis laws (M. Faraday)

1. The weight amount of the substance released during electrolysis is proportional to the amount of electricity flowing through the solution and practically does not depend on other factors.

2. Equal amounts of electricity release equivalent amounts of substances from various chemical compounds during electrolysis.

3. To isolate one gram-equivalent of any substance from the electrolyte solution, 96500 coulombs of electricity must be passed through the solution.

m (x) = ((I t) / F) (M (x) / n)

where m (x) is the amount of reduced or oxidized substance (g);

I is the strength of the transmitted current (a);

t is the electrolysis time (s);

M (x) - molar mass;

n is the number of electrons acquired or donated in redox reactions;

F is the Faraday constant (96500 coul / mol).

Based on this formula, you can make a number of calculations related to the electrolysis process, for example:

1. Calculate the amount of substances released or decomposed by a certain amount of electricity;

2. Find the current strength by the amount of released substance and the time spent on its release;

3. Establish how long it will take to release a certain amount of a substance at a given current strength.

Example 1

How many grams of copper will be released at the cathode when a current of 5 amperes is passed through a solution of copper sulfate CuSO 4 for 10 minutes?

Solution

Determine the amount of electricity flowed through the solution:

Q = I t,

where I is the current strength in amperes;

t is the time in seconds.

Q = 5A 600 c = 3000 coulombs

The equivalent of copper (atomic mass 63.54) is 63.54: 2 = 31.77. Therefore, 96,500 coulombs emit 31.77 g of copper. The required amount of copper:

m = (31.77 3000) / 96500 "0.98 g

Example 2

How long does it take to pass a current of 10 amperes through the acid solution to obtain 5.6 liters of hydrogen (at normal)?

Solution

We find the amount of electricity that must pass through the solution so that 5.6 liters of hydrogen are released from it. Since 1 g-eq. hydrogen occupies at n. at. volume of 11.2 liters, then the required amount of electricity

Q = (96500 5.6) / 11.2 = 48250 coulombs

Let us determine the time of passage of the current:

t = Q / I = 48250/10 = 4825 s = 1 h 20 min 25 s

Example 3

When a current was passed through a solution of a silver salt at the cathode, it precipitated in 10 min. 1 g of silver. Determine the amperage.

Solution

1 g-eq. silver is equal to 107.9 g. To isolate 1 g of silver, 96500 must pass through the solution: 107.9 = 894 coulombs. Hence the current strength

I = 894 / (10 60) »1.5A

Example 4

Find the equivalent of tin, if at a current of 2.5 amperes from a SnCl 2 solution in 30 min. 2.77 g of tin is released.

Solution

The amount of electricity passed through the solution in 30 minutes.

Q = 2.5 30 60 = 4500 coulombs

Since for the selection of 1 g-eq. requires 96,500 coulombs, then the equivalent of tin.

E Sn = (2.77 96500) / 4500 = 59.4

Corrosion

Before ending the discussion of electrochemistry, let us apply the knowledge we have gained to the study of one very important problem - corrosion metals. Corrosion is caused by redox reactions, in which a metal, as a result of interaction with a substance from its environment, turns an undesirable compound.

One of the most widely known corrosion processes is iron rusting. From an economic point of view, this is a very important process. It is estimated that 20% of the iron produced annually in the United States is used to replace iron items that have deteriorated due to rusting.

It is known that oxygen is involved in the rusting of iron; iron is not oxidized in water in the absence of oxygen. Water also takes part in the rusting process; iron will not corrode in oxygenated oil unless it contains traces of water. Rust is accelerated under the influence of a number of factors, such as the pH of the environment, the presence of salts in it, the contact of iron with metal, which is more difficult to oxidize than iron, and also under the influence of mechanical stress.

Iron corrosion is in principle an electrochemical process. Some areas of the surface of iron serve as an anode on which its oxidation occurs:

Fe (solid) → Fe 2+ (aq.) + 2е - Еº oxid = 0.44 V

The resulting electrons move through the metal to other areas of the surface, which play the role of a cathode. Oxygen is restored on them:

О 2 (g.) + 4Н + (aq.) + 4е - → 2Н 2 О (l.) E ° recovery = 1.23 V

Note that Н + ions participate in the process of О 2 reduction. If the H + concentration decreases (ie, as the pH rises), O 2 reduction becomes more difficult. It is noticed that iron in contact with a solution with a pH higher than 9-10 does not corrode. In the process of corrosion, the Fe 2+ ions formed at the anode are oxidized to Fe 3+. Fe 3+ ions form hydrated iron (III) oxide, which is called rust:

4Fe 2+ (aq.) + O 2 (g.) + 4H 2 O (l.) +2 NS H 2 O (f.) → 2Fe 2 O 3. x H 2 O ( tv.) + 8H + (aq.)

Since the role of the cathode is usually played by that part of the surface that is best provided with an influx of oxygen, rust most often appears in these areas. If you carefully examine a shovel that has stood for some time in the open humid air with dirt adhering to the blade, you will notice that depressions have formed on the metal surface under the dirt, and rust has appeared everywhere where O 2 could penetrate.

Increased corrosion in the presence of salts is often encountered by motorists in areas where, in winter, roads are abundantly sprinkled with salt to combat icy conditions. The effect of salts is explained by the fact that the ions they form create the electrolyte necessary for the formation of a closed electric circuit.

The presence of sites of anodic and cathodic nature on the surface of iron leads to the creation of two different chemical environments on it. They can arise due to the presence of impurities or crystal lattice defects (apparently caused by stresses inside the metal). In places where there are such impurities or defects, the microscopic environment of a particular iron atom can cause some increase or decrease in its oxidation state compared to normal positions in the crystal lattice. Therefore, such places can play the role of anodes or cathodes. Ultrapure iron, in which the number of such defects is minimized, is much less corrosive than conventional iron.

Iron is often coated with paint or some other metal, such as tin, zinc, or chromium, to protect its surface from corrosion. The so-called "tinplate" is obtained by coating sheet iron with a thin layer of tin. Tin protects the iron only as long as the protective layer remains intact. One has only to damage it, as air and moisture begin to affect the iron; tin even accelerates iron corrosion because it serves as a cathode in the electrochemical corrosion process. Comparison of the oxidation potentials of iron and tin shows that iron is oxidized more easily than tin:

Fe (solid) → Fe 2+ (aq.) + 2е - Еº oxid = 0.44 V

Sn (tv) → Sn 2+ (aq.) + 2е - Еº oxid = 0.14 V

Therefore, in this case, iron serves as an anode and is oxidized.

"Galvanized" (galvanized) iron is produced by coating the iron with a thin layer of zinc. Zinc protects iron from corrosion even after damage to the integrity of the coating. In this case, iron in the corrosion process plays the role of a cathode, because zinc is oxidized more easily than iron:

Zn (TV) → Zn 2+ (aq.) + 2e - Еº oxid = 0.76 V

Therefore, zinc acts as an anode and corrodes instead of iron. Such protection of the metal, in which it plays the role of a cathode in the process of electrochemical corrosion, is called cathodic protection. Pipes buried underground often protect against corrosion by making them the cathode of an electrochemical cell. To do this, blocks of some active metal, most often magnesium, are buried in the ground along the pipeline and connected with wires to pipes. In moist soil, the active metal acts as an anode, and the iron pipe receives cathodic protection.

While our discussion focuses on iron, it is not the only metal that corrodes. At the same time, it may seem strange that an aluminum can, carelessly left in the open air, corrodes immeasurably slower than an iron one. Judging by the standard oxidation potentials of aluminum (Eº oxide = 1.66 V) and iron (Eº oxide = 0.44 V), it should be expected that aluminum corrosion should occur much faster. The slow corrosion of aluminum is explained by the fact that a thin dense oxide film forms on its surface, which protects the metal located under it from further corrosion. Magnesium, which has a high oxidation potential, is protected from corrosion due to the formation of the same oxide film. Unfortunately, the oxide film on the iron surface has a too loose structure and is not able to provide reliable protection. However, a good protective oxide film is formed on the surface of iron-chromium alloys. Such alloys are called stainless steel.

Based on the change in the oxidation states of the atoms that make up the reactants, chemical reactions are divided into two types.

1) Reactions proceeding without changing the oxidation states of atoms.

For example:

2 + 4-2 t +2 -2 +4 -2
CaCO 3 = CaO + CO 2

In this reaction, the oxidation state of each of the atoms remained unchanged.

2) Reactions proceeding with a change in the oxidation states of atoms.

For example:

0 +2 -1 0 +2 -1
Zn + CuCl 2 = Cu + ZnCl 2

In this reaction, the oxidation states of the zinc and copper atoms changed.

Redox reactions are the most common chemical reactions.

In practice, a redox reaction is the addition or donation of electrons. Some atoms (ions, molecules) donate to others or receive electrons from them.

Oxidation.

The process of giving up electrons by an atom, ion or molecule is called oxidation.

With the donation of electrons, the oxidation state of the atom increases.

A substance whose atoms, ions or molecules donate electrons is called reducing agent.

In our example, atoms in the oxidation state of 0 are converted to atoms in the oxidation state +2. That is, an oxidation process has taken place. In this case, the zinc atom that donated two electrons is a reducing agent (it increased the oxidation state from 0 to +2).

The oxidation process is recorded by an electronic equation, which indicates the change in the oxidation state of the atoms and the number of electrons donated by the reducing agent.

For example:

0 +2 0
Zn - 2e - = Zn (oxidation, Zn is a reducing agent).

Recovery.

The process of electron attachment is called rebuilding.

When electrons are attached, the oxidation state of the atom decreases.

A substance whose atoms, ions or molecules attach electrons is called oxidizing agent.

In our example, the transition of copper atoms with an oxidation state of +2 to atoms with an oxidation state of 0 is a reduction process. In this case, a copper atom with an oxidation state of +2, accepting two electrons, lowers the oxidation state from +2 to 0 and is an oxidizing agent.

The oxidation process is also written with an electronic equation:

2 0 0
Cu + 2e - = Cu (reduction, Cu is an oxidizing agent).

The reduction process and the oxidation process are inseparable and proceed simultaneously.

0 +2 0 +2
Zn + CuCl 2 = Cu + ZnCl 2
reducing agent oxidizing agent
oxidized reduced

Calculation of the oxidation state

Summary

1. Building the workforce is one of the most essential areas of work for a HR manager.

2. In order to provide the organization with the necessary human resources, it is important to develop an adequate situation in the external environment and technology of activity, the structure of the company; calculate the need for staff.

3. To develop recruitment programs, it is necessary to analyze the personnel situation in the region, develop procedures for attracting and evaluating candidates, and carry out adaptation measures to include new employees in the organization.

Control questions

What groups of factors should be considered when creating an organizational structure?
What stages of organization design can be highlighted?
Explain the concept of “qualitative assessment of staffing needs”.
Describe the concept of “additional need for personnel”.
What is the purpose of the analysis of the personnel situation in the region?
What is the purpose of the activity analysis?
What stages of activity analysis can be distinguished?
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To calculate the oxidation state of an element, the following points should be taken into account:

1. The oxidation states of atoms in simple substances are zero (Na 0; H 2 0).

2. The algebraic sum of the oxidation states of all atoms that make up a molecule is always zero, and in a complex ion this sum is equal to the charge of the ion.

3. Atoms have a constant oxidation state: alkali metals (+1), alkaline earth metals (+2), hydrogen (+1) (except for hydrides NaH, CaH 2, etc., where the oxidation state of hydrogen is -1), oxygen (-2 ) (except for F 2 -1 O +2 and peroxides containing the –O – O– group, in which the oxidation state of oxygen is -1).

4. For elements, the positive oxidation state cannot exceed a value equal to the group number of the periodic system.

Examples:

V 2 +5 O 5 -2; Na 2 +1 B 4 +3 O 7 -2; K +1 Cl +7 O 4 -2; N -3 H 3 +1; K 2 +1 H +1 P +5 O 4 -2; Na 2 +1 Cr 2 +6 O 7 -2

There are two types of chemical reactions:

A Reactions in which the oxidation state of the elements does not change:

Addition reactions

SO 2 + Na 2 O Na 2 SO 3

Decomposition reactions

Cu (OH) 2 - t CuO + H 2 O

Exchange reactions

AgNO 3 + KCl AgCl + KNO 3

NaOH + HNO 3 NaNO 3 + H 2 O

B Reactions in which there is a change in the oxidation states of the atoms of the elements that make up the reacting compounds:

2Mg 0 + O 2 0 2Mg +2 O -2

2KCl +5 O 3 -2 - t 2KCl -1 + 3O 2 0

2KI -1 + Cl 2 0 2KCl -1 + I 2 0

Mn +4 O 2 + 4HCl -1 Mn +2 Cl 2 + Cl 2 0 + 2H 2 O

Such reactions are called redox.

Redox reactions include those that are accompanied by the movement of electrons from one particle to another. When considering the regularities of the course of redox reactions, the concept of the oxidation state is used.

Oxidation state

Concept oxidation state introduced to characterize the state of elements in connections. The oxidation state is understood as conditional charge of an atom in a compound, calculated on the assumption that the compound consists of ions... The oxidation state is indicated by an Arabic numeral with a plus sign when electrons are displaced from a given atom to another atom, and by a minus sign when electrons are displaced in the opposite direction. A digit with a "+" or "-" sign is placed above the element symbol. The oxidation state indicates the oxidation state of the atom and is just a convenient form for taking into account the electron transfer: it should not be considered either as the effective charge of the atom in the molecule (for example, in the LiF molecule, the effective charges Li and F are respectively + 0.89 and -0, 89, while the oxidation states are +1 and -1), nor as the valence of the element (for example, in the compounds CH 4, CH 3 OH, HCOOH, CO 2, the valency of carbon is 4, and the oxidation states are, respectively, -4, -2, + 2, +4). The numerical values of the valence and oxidation state can coincide in absolute value only when compounds with an ionic structure are formed.

When determining the oxidation state, the following rules are used:

The atoms of elements that are in a free state or in the form of molecules of simple substances have an oxidation state equal to zero, for example, Fe, Cu, H 2, N 2, etc.

The oxidation state of an element in the form of a monatomic ion in a compound having an ionic structure is equal to the charge of a given ion,

1 -1 +2 -2 +3 -1

for example, NaCl, Cu S, AlF 3.

Hydrogen in most compounds has an oxidation state of +1, with the exception of metal hydrides (NaH, LiH), in which the oxidation state of hydrogen is -1.

The most common oxidation state of oxygen in compounds is -2, with the exception of peroxides (Na 2 O 2, H 2 O 2), in which the oxidation state of oxygen is –1 and F 2 O, in which the oxidation state of oxygen is +2.

For elements with a variable oxidation state, its value can be calculated knowing the formula of the compound and taking into account that the algebraic sum of the oxidation states of all elements in a neutral molecule is zero. In a complex ion, this sum is equal to the charge of the ion. For example, the oxidation state of the chlorine atom in the HClO 4 molecule, calculated based on the total charge of the molecule = 0, where x is the oxidation state of the chlorine atom), is +7. The oxidation state of the sulfur atom in the (SO 4) 2- [x + 4 (-2) = -2] ion is +6.

Redox properties of substances

Any redox reaction consists of oxidation and reduction processes. Oxidation - it is the process of giving up electrons by an atom, ion, or reagent molecule. Substances that give off their electrons during the reaction and oxidize at the same time, they call reducing agents.

Recovery is the process of accepting electrons by an atom, ion or reagent molecule.

Substances that accept electrons and are reduced at the same time are called oxidizing agents.

Oxidation-reduction reactions always proceed as a single process, called redox reaction. For example, when zinc metal interacts with copper ions reducing agent(Zn) donates its electrons oxidizer- copper ions (Cu 2+):

Zn + Cu 2+ Zn 2+ + Cu

Copper is released on the surface of zinc, and zinc ions pass into solution.

The redox properties of elements are associated with the structure of their atoms and are determined by the position in the periodic system of D.I. Mendeleev. The reducing ability of the element is due to the weak bond of valence electrons with the nucleus. Metal atoms containing a small number of electrons at the external energy level are prone to give up, i.e. easily oxidized, playing the role of reducing agents. The most powerful reducing agents are the most active metals.

The criterion for the redox activity of elements can be the value of their relative electronegativity: the higher it is, the more pronounced the oxidizing ability of the element, and the lower, the brighter its reducing activity is. Atoms of non-metals (for example, F, O) have a high value of electron affinity and relative electronegativity; they easily accept electrons, i.e. are oxidizing agents.

The redox properties of an element depend on its oxidation state. The same element is distinguished lower, higher and intermediate oxidation states.

As an example, consider sulfur S and its compounds H 2 S, SO 2 and SO 3. The relationship between the electronic structure of the sulfur atom and its redox properties in these compounds is clearly shown in Table 1.

In the H2S molecule, the sulfur atom has a stable octet configuration of the external energy level 3s 2 3p 6 and therefore can no longer attach electrons, but can donate them.

The state of an atom in which it can no longer accept electrons is called the lowest oxidation state.

In the lowest oxidation state, the atom loses its oxidizing ability and can only be a reducing agent.

Table 1.

Formula of substance	Electronic formula	Redox properties
	1s 2 2s 2 2p 6 3s 2 3p 6	–2 ; - 6 ; - 8 reducing agent
	1s 2 2s 2 2p 6 3s 2 3p 4	+ 2 oxidizing agent	–4 ; - 6 reducing agent
	1s 2 2s 2 2p 6 3s 2 3p o	+ 4 ; + 6 oxidizing agent	-2 reducing agent
	1s 2 2s 2 2p 6 3s o 3p 0	+ 2 ; + 6 ; + 8 oxidizing agent

In the SO 3 molecule, all the outer electrons of the sulfur atom are displaced towards oxygen atoms. Consequently, in this case, the sulfur atom can only accept electrons, exhibiting oxidizing properties.

The state of an atom in which it gave up all its valence electrons is called the highest oxidation state. An atom in the highest oxidation state can only be an oxidizing agent.

In the SO 2 molecule and elemental sulfur S, the sulfur atom is located in intermediate oxidation states, i.e., having valence electrons, an atom can give them away, but, not having a complete R - sublevel, can and accept electrons before its completion.

An atom of an element with an intermediate oxidation state can exhibit both oxidizing and reducing properties, which is determined by its role in a particular reaction.

So, for example, the role of sulfite - anion SO in the following reactions is different:

5Na 2 SO 3 + 2KMnO 4 + 3H 2 SO 4  2MnSO 4 + 5Na 2 SO 4 + K 2 SO 4 + 3H 2 O (1)

H 2 SO 3 + 2 H 2 S  3 S + 3 H 2 O (2)

In reaction (1) sulfite anion SO in the presence of a strong oxidizing agent, KMnO 4 plays the role of a reducing agent; in reaction (2) sulfite anion SO - an oxidizing agent, since H 2 S can only exhibit reducing properties.