How to determine the oxidation state of an atom of a chemical element. Sulfur compounds

The subgroup of chalcogens includes sulfur - this is the second of the elements that can form a large number of ore deposits. Sulfates, sulfides, oxides and other sulfur compounds are very widespread, important in industry and nature. Therefore, in this article we will consider what they are, what sulfur itself is, its simple substance.

Sulfur and its characteristics

This element has the following position in the periodic table.

1. The sixth group, the main subgroup.
2. Third minor period.
3. Atomic mass - 32.064.
4. The serial number is 16, there are the same number of protons and electrons, and there are also 16 neutrons.
5. Refers to non-metal elements.
6. In the formulas it is read as "es", the name of the element sulfur, Latin sulfur.

There are four stable isotopes found in nature. mass numbers 32,33,34 and 36. This element is the sixth most common in nature. Refers to biogenic elements, as it is part of important organic molecules.

The electronic structure of the atom

Sulfur compounds owe their diversity to the features of the electronic structure of the atom. It is expressed by the following configuration formula: 1s 2 2s 2 2p 6 3s 2 3p 4 .

The given order reflects only the stationary state of the element. However, it is known that if additional energy is imparted to an atom, then electrons can be depaired at the 3p and 3s sublevels, followed by another transition to 3d, which remains free. As a result, not only the valency of the atom changes, but also all possible oxidation states. Their number increases significantly, as well as the number of various substances with the participation of sulfur.

The oxidation states of sulfur in compounds

There are several main options for this indicator. For sulfur it is:

Of these, S +2 is the most rare, the rest are dispersed everywhere. The chemical activity and oxidizing ability of the entire substance depends on the degree of oxidation of sulfur in compounds. So, for example, compounds with -2 are sulfides. In them, the element we are considering is a typical oxidizing agent.

The higher the value of the oxidation state in the compound, the more pronounced the oxidizing abilities of the substance will be. This is easy to verify if we recall the two main acids that sulfur forms:

• H 2 SO 3 - sulfurous;
• H 2 SO 4 - sulfuric.

It is known that the latter is a much more stable, strong compound, which in high concentration has a very serious ability to oxidize.

simple substance

As a simple substance, sulfur is yellow beautiful crystals of even, regular, elongated shape. Although this is only one of its forms, because there are two main of this substance. The first, monoclinic or rhombic - this is the yellow that cannot be dissolved in water, but only in organic solvents. It is brittle and beautiful shape structure represented as a crown. The melting point is about 110 0 С.

If, however, an intermediate moment is not missed when such a modification is heated, then another state can be detected in time - plastic sulfur. It is a rubbery viscous solution Brown, which, upon further heating or rapid cooling, again turns into a rhombic shape.

If we talk about chemically pure sulfur obtained by repeated filtration, then it is a bright yellow small crystals, fragile and completely insoluble in water. Able to ignite on contact with moisture and oxygen in the air. Differ in rather high chemical activity.

Being in nature

In nature, there are natural deposits from which sulfur compounds are extracted and sulfur itself as a simple substance. In addition, it contains:

• in minerals, ores and rocks;
• in the body of animals, plants and humans, as it is part of many organic molecules;
• in natural gases, oil and coal;
• in oil shale and natural waters.

You can name some of the richest minerals in sulfur:

• cinnabar;
• pyrite;
• sphalerite;
• antimonite;
• galena and others.

Most of the sulfur produced today goes to sulfate production. Another part is used for medical purposes, Agriculture, industrial processes production of substances.

Physical Properties

They can be described in several points.

1. It is insoluble in water, in carbon disulfide or turpentine - it dissolves well.
2. With prolonged friction accumulates a negative charge.
3. The melting point is 110 0 C.
4. Boiling point 190 0 С.
5. Upon reaching 300 0 C, it passes into a liquid, easily mobile.
6. A pure substance is capable of spontaneous combustion, combustible properties are very good.
7. By itself, it has practically no smell, however, hydrogen sulfur compounds emit a pungent smell of rotten eggs. Just like some gaseous binary representatives.

The physical properties of the substance in question have been known to people since antiquity. It is for its combustibility that sulfur got its name. In wars, asphyxiating and poisonous fumes, which are formed during the combustion of this compound, were used as a weapon against enemies. In addition, acids containing sulfur have also always been of great industrial importance.

Chemical properties

Topic: "Sulfur and its compounds" on school course Chemistry takes not one lesson, but several. After all, there are a lot of them. This is due to the chemical activity of this substance. It can exhibit both oxidizing properties with stronger reducing agents (metals, boron, and others), and reducing properties with most non-metals.

However, despite such activity, the interaction occurs only with fluorine at normal conditions. All others require heating. There are several categories of substances with which sulfur can interact:

• metals;
• non-metals;
• alkalis;
• strong oxidizing acids - sulfuric and nitric.

Sulfur compounds: varieties

Their diversity will be explained by the unequal value of the oxidation state of the main element - sulfur. So, we can distinguish several main types of substances on this basis:

• compounds with an oxidation state of -2;

If we consider classes, and not the valency index, then this element forms molecules such as:

• acids;
• oxides;
• salt;
• binary compounds with non-metals (carbon disulfide, chlorides);
• organic substances.

Now consider the main ones and give examples.

Substances with an oxidation state of -2

Sulfur compounds 2 are its conformations with metals, as well as with:

• carbon;
• hydrogen;
• phosphorus;
• silicon;
• arsenic;
• boron.

In these cases, it acts as an oxidizing agent, since all of the listed elements are more electropositive. Let's take a look at some of the more important ones.

1. Carbon disulfide - CS 2 . Transparent liquid with a characteristic pleasant aroma of ether. It is toxic, flammable and explosive. It is used as a solvent for most types of oils, fats, non-metals, silver nitrate, resins and rubbers. It is also an important part in the production of artificial silk - viscose. In industry, it is synthesized in large quantities.
2. Hydrogen sulfide or hydrogen sulfide - H 2 S. A colorless gas with a sweet taste. The smell is pungent, extremely unpleasant, reminiscent of rotten egg. Poisonous, depresses the respiratory center, as it binds copper ions. Therefore, when poisoned by them, suffocation and death occur. It is widely used in medicine, organic synthesis, production of sulfuric acid, and also as an energy-efficient raw material.
3. Metal sulfides are found wide application in medicine, in sulfate production, in the production of paints, in the manufacture of phosphors and in other places. The general formula is Me x S y .

Compounds with an oxidation state of +4

Sulfur compounds 4 are predominantly an oxide and its corresponding salts and an acid. All of them are fairly common compounds that have certain value in industry. They can also act as oxidizing agents, but more often they exhibit reducing properties.

The formulas for a sulfur compound with an oxidation state of +4 are as follows:

• oxide - sulfur dioxide SO 2 ;
• acid - sulfurous H 2 SO 3;
• salts have general formula Mex(SO3)y.

One of the most common is or anhydride. It is a colorless substance with the smell of a burnt match. In large clusters, it is formed during volcanic eruptions; at this moment it is easy to identify it by smell.

It dissolves in water with the formation of easily decomposing acid - sulfurous. It behaves like a typical salt forms, which enters in the form of a sulfite ion SO 3 2-. This anhydride is the main gas that affects the pollution of the surrounding atmosphere. It is he who affects education. In industry, it is used in sulfate production.

Compounds in which sulfur has an oxidation state of +6

These include, first of all, sulfuric anhydride and sulfuric acid with their salts:

• sulfates;
• hydrosulfates.

Since the sulfur atom in them is in the highest degree of oxidation, the properties of these compounds are quite understandable. They are strong oxidizing agents.

Sulfur oxide (VI) - sulfuric anhydride - is a volatile colorless liquid. Characteristic- strong moisture absorption capacity. Smokes outdoors. When dissolved in water, it gives one of the strongest mineral acids - sulfuric. Its concentrated solution is a heavy oily slightly yellowish liquid. If the anhydride is dissolved in sulfuric acid, then a special compound called oleum will be obtained. It is used industrially in the production of acid.

Among the salts - sulfates - great importance has connections like:

• gypsum CaSO 4 2H 2 O;
• barite BaSO 4 ;
• mirabilite;
• lead sulfate and others.

They are used in construction, chemical synthesis, medicine, manufacturing optical devices and glass and even the food industry.

Hydrosulfates are widely used in metallurgy, where they are used as a flux. And also they help to convert many complex oxides into soluble sulfate forms, which is used in the corresponding industries.

The study of sulfur in the school chemistry course

When is the best time for students to learn about what sulfur is, what are its properties, what is a sulfur compound? Grade 9 - best period. This is not the very beginning, when everything is new and incomprehensible for children. This is the middle ground in the study of chemical science, when the foundations laid earlier will help to fully understand the topic. Therefore, it is the second half of the graduating class that is allocated for consideration of these issues. At the same time, the whole topic is divided into several blocks, in which there is a separate lesson "Sulfur compounds. Grade 9".

This is due to their abundance. The issue of industrial production of sulfuric acid is also considered separately. In general, on this topic takes an average of 3 hours.

But sulfur is taken out for study only in the 10th grade, when organic issues are considered. They are also affected in biology in high school. After all, sulfur is part of such organic molecules as:

• thioalcohols (thiols);
• proteins (tertiary structure on which the formation of disulfide bridges occurs);
• thioaldehydes;
• thiophenols;
• thioethers;
• sulfonic acids;
• sulfoxides and others.

They are distinguished in special group organosulfur compounds. They are important not only in the biological processes of living beings, but also in industry. For example, sulfonic acids are the basis of many medicines(aspirin, sulfanilamide or streptocide).

In addition, sulfur is a constant component of compounds such as some:

• amino acids;
• enzymes;
• vitamins;
• hormones.

The oxidation state is the conditional charge of an atom in a compound, calculated on the assumption that it consists only of ions. When defining this concept, it is conditionally assumed that the binding (valence) electrons pass to more electronegative atoms (see Electronegativity), and therefore the compounds consist, as it were, of positively and negatively charged ions. The oxidation state can have zero, negative and positive values, which are usually placed above the element symbol at the top: .

The zero value of the oxidation state is assigned to the atoms of the elements in the free state, for example: . Negative meaning oxidation states are those atoms towards which the binding electron cloud (electron pair) is displaced. For fluorine in all its compounds, it is -1. Atoms that donate valence electrons to other atoms have a positive oxidation state. For example, in alkali and alkaline earth metals, it is respectively equal to and In simple ions, like , K, it is equal to the charge of the ion. In most compounds, the oxidation state of hydrogen atoms is equal, but in metal hydrides (their compounds with hydrogen) - and others - it is -1. Oxygen is characterized by an oxidation state of -2, but, for example, in combination with fluorine it will be, and in peroxide compounds, etc.) -1. In some cases, this value can be expressed and fractional number: for iron in iron oxide (II, III) it is equal to .

The algebraic sum of the oxidation states of atoms in a compound is zero, and in a complex ion it is the charge of the ion. Using this rule, we calculate, for example, the oxidation state of phosphorus in orthophosphoric acid. Denoting it through and multiplying the oxidation state for hydrogen and oxygen by the number of their atoms in the compound, we get the equation: whence. Similarly, we calculate the oxidation state of chromium in the ion -.

In compounds, the oxidation state of manganese will be, respectively.

The highest oxidation state is its highest positive value. For most elements, it is equal to the group number in the periodic system and is important quantitative characteristic element in its compounds. Lowest value the oxidation state of an element that occurs in its compounds is commonly called the lowest oxidation state; all others are intermediate. So, for sulfur, the highest oxidation state is equal to, the lowest -2, intermediate.

Change in the oxidation states of elements by groups periodic system reflects the frequency of their change chemical properties with increasing serial number.

The concept of the oxidation state of elements is used in the classification of substances, describing their properties, formulating compounds and their international names. But it is especially widely used in the study of redox reactions. The concept of "oxidation state" is often used in inorganic chemistry instead of the concept of "valency" (see Valency).

Valence is a complex concept. This term has undergone a significant transformation simultaneously with the development of the theory of chemical bonding. Initially, valence was the ability of an atom to attach or replace a certain number of other atoms or atomic groups to form a chemical bond.

The quantitative measure of the valency of an element atom was the number of hydrogen or oxygen atoms (these elements were considered mono- and divalent, respectively), which the element adds to form a hydride of the formula EH x or an oxide of the formula E n O m .

So, the valence of the nitrogen atom in the NH 3 ammonia molecule is three, and the sulfur atom in the H 2 S molecule is two, since the valency of the hydrogen atom is one.

In the compounds Na 2 O, BaO, Al 2 O 3, SiO 2, the valences of sodium, barium and silicon are 1, 2, 3 and 4, respectively.

The concept of valence was introduced into chemistry before the structure of the atom became known, namely in 1853 by the English chemist Frankland. It has now been established that the valency of an element is closely related to the number of outer electrons of atoms, since the electrons of the inner shells of atoms do not participate in the formation of chemical bonds.

In the electronic theory of covalent bonding, it is believed that atom valency is determined by the number of its unpaired electrons in the ground or excited state, participating in the formation of common electron pairs with electrons of other atoms.

For some elements, valency is a constant value. So, sodium or potassium in all compounds is monovalent, calcium, magnesium and zinc are divalent, aluminum is trivalent, etc. But most chemical elements exhibit variable valence, which depends on the nature of the partner element and the conditions of the process. So, iron can form two compounds with chlorine - FeCl 2 and FeCl 3, in which the valency of iron is 2 and 3, respectively.

Oxidation state- a concept that characterizes the state of an element in a chemical compound and its behavior in redox reactions; numerically, the oxidation state is equal to the formal charge that can be attributed to the element, based on the assumption that all the electrons of each of its bonds have passed to the more electronegative atom.

Electronegativity- a measure of the ability of an atom to acquire a negative charge during the formation of a chemical bond, or the ability of an atom in a molecule to attract valence electrons involved in the formation of a chemical bond. Electronegativity is not an absolute value and is calculated various methods. Therefore, the electronegativity values ​​given in different textbooks and reference books may differ.

Table 2 shows the electronegativity of some chemical elements on the Sanderson scale, and Table 3 shows the electronegativity of the elements on the Pauling scale.

The value of electronegativity is given under the symbol of the corresponding element. The greater the numerical value of the electronegativity of an atom, the more electronegative the element is. The most electronegative is the fluorine atom, the least electronegative is the rubidium atom. In a molecule formed by atoms of two different chemical elements, the formal negative charge will be on the atom whose numerical value of electronegativity will be higher. So, in a sulfur dioxide molecule SO 2, the electronegativity of the sulfur atom is 2.5, and the value of the electronegativity of the oxygen atom is greater - 3.5. Therefore, the negative charge will be on the oxygen atom, and the positive charge on the sulfur atom.

In the ammonia molecule NH 3, the electronegativity value of the nitrogen atom is 3.0, and that of hydrogen is 2.1. Therefore, the nitrogen atom will have a negative charge, and the hydrogen atom will have a positive charge.

You should clearly know the general trends in electronegativity. Since an atom of any chemical element tends to acquire a stable configuration of the outer electron layer - an octet shell of an inert gas, then the electronegativity of the elements in the period increases, and in the group, the electronegativity generally decreases with an increase in the atomic number of the element. Therefore, for example, sulfur is more electronegative than phosphorus and silicon, and carbon is more electronegative than silicon.

When compiling formulas for compounds consisting of two non-metals, the more electronegative of them is always placed to the right: PCl 3, NO 2. There are some historical exceptions to this rule, such as NH 3 , PH 3 , etc.

The oxidation state is usually indicated by an Arabic numeral (with a sign in front of the digit) located above the element symbol, for example:

To determine the oxidation state of atoms in chemical compounds, the following rules are followed:

1. The oxidation state of the elements in simple substances equals zero.
2. The algebraic sum of the oxidation states of atoms in a molecule is zero.
3. Oxygen in compounds exhibits mainly an oxidation state of –2 (in oxygen fluoride OF 2 + 2, in metal peroxides such as M 2 O 2 –1).
4. Hydrogen in compounds exhibits an oxidation state of +1, with the exception of hydrides active metals, for example, alkaline or alkaline earth, in which the oxidation state of hydrogen is -1.
5. For monatomic ions, the oxidation state is equal to the charge of the ion, for example: K + - +1, Ba 2+ - +2, Br - - -1, S 2- - -2, etc.
6. In compounds with a covalent polar bond, the oxidation state of a more electronegative atom has a minus sign, and a less electronegative atom has a plus sign.
7. AT organic compounds the oxidation state of hydrogen is +1.

Let's illustrate the above rules with several examples.

Example 1 Determine the degree of oxidation of elements in oxides of potassium K 2 O, selenium SeO 3 and iron Fe 3 O 4.

Potassium oxide K 2 O. The algebraic sum of the oxidation states of atoms in a molecule is zero. The oxidation state of oxygen in oxides is –2. Let us denote the oxidation state of potassium in its oxide as n, then 2n + (–2) = 0 or 2n = 2, hence n = +1, i.e. the oxidation state of potassium is +1.

Selenium oxide SeO 3 . The SeO 3 molecule is electrically neutral. The total negative charge of the three oxygen atoms is –2 × 3 = –6. Therefore, in order to equalize this negative charge to zero, the oxidation state of selenium must be +6.

Fe 3 O 4 molecule electrically neutral. The total negative charge of the four oxygen atoms is –2 × 4 = –8. To equalize this negative charge, the total positive charge on the three iron atoms must be +8. Therefore, one iron atom should have a charge of 8/3 = +8/3.

It should be emphasized that the oxidation state of an element in a compound can be a fractional number. Such fractional oxidation states do not make sense in explaining the bond in a chemical compound, but can be used to formulate equations for redox reactions.

Example 2 Determine the degree of oxidation of elements in the compounds NaClO 3, K 2 Cr 2 O 7.

The NaClO 3 molecule is electrically neutral. The oxidation state of sodium is +1, the oxidation state of oxygen is -2. Let us denote the oxidation state of chlorine as n, then +1 + n + 3 × (–2) = 0, or +1 + n – 6 = 0, or n – 5 = 0, hence n = +5. Thus, the oxidation state of chlorine is +5.

The K 2 Cr 2 O 7 molecule is electrically neutral. The oxidation state of potassium is +1, the oxidation state of oxygen is -2. Let us denote the oxidation state of chromium as n, then 2 × 1 + 2n + 7 × (–2) = 0, or +2 + 2n – 14 = 0, or 2n – 12 = 0, 2n = 12, hence n = +6. Thus, the oxidation state of chromium is +6.

Example 3 Let us determine the oxidation states of sulfur in the sulfate ion SO 4 2– . The SO 4 2– ion has a charge of –2. The oxidation state of oxygen is –2. Let us denote the oxidation state of sulfur as n, then n + 4 × (–2) = –2, or n – 8 = –2, or n = –2 – (–8), hence n = +6. Thus, the oxidation state of sulfur is +6.

It should be remembered that the oxidation state is sometimes not equal to the valence of a given element.

For example, the oxidation states of the nitrogen atom in the ammonia molecule NH 3 or in the hydrazine molecule N 2 H 4 are -3 and -2, respectively, while the nitrogen valence in these compounds is three.

The maximum positive oxidation state for elements of the main subgroups, as a rule, is equal to the group number (exceptions: oxygen, fluorine and some other elements).

The maximum negative oxidation state is 8 - the group number.

Training tasks

1. In which compound is the oxidation state of phosphorus +5?

1) HPO 3
2) H3PO3
3) Li 3 P
4) AlP

2. Which compound has the oxidation state of phosphorus -3?

1) HPO 3
2) H3PO3
3) Li3PO4
4) AlP

3. In what compound is the oxidation state of nitrogen equal to +4?

1) HNO2
2) N 2 O 4
3) N 2 O
4) HNO3

4. In which compound is the oxidation number of nitrogen equal to -2?

1) NH3
2) N 2 H 4
3) N 2 O 5
4) HNO2

5. In what compound is the oxidation state of sulfur equal to +2?

1) Na 2 SO 3
2) SO2
3) SCl2
4) H2SO4

6. In what compound is the oxidation state of sulfur equal to +6?

1) Na 2 SO 3
2) SO3
3) SCl2
4) H2SO3

7. In substances whose formulas are CrBr 2, K 2 Cr 2 O 7, Na 2 CrO 4, the oxidation state of chromium, respectively, is

1) +2, +3, +6
2) +3, +6, +6
3) +2, +6, +5
4) +2, +6, +6

8. The minimum negative oxidation state of a chemical element is usually equal to

1) period number
3) the number of electrons missing before the completion of the outer electron layer

9. The maximum positive oxidation state of chemical elements located in the main subgroups is usually equal to

1) period number
2) the serial number of the chemical element
3) group number
4) total number electrons in the element

10. Phosphorus exhibits the maximum positive oxidation state in the compound

1) HPO 3
2) H3PO3
3) Na 3 P
4) Ca 3 P 2

11. Phosphorus exhibits the lowest oxidation state in the compound

1) HPO 3
2) H3PO3
3) Na3PO4
4) Ca 3 P 2

12. Nitrogen atoms in ammonium nitrite, which are part of the cation and anion, exhibit oxidation states, respectively

1) –3, +3
2) –3, +5
3) +3, –3
4) +3, +5

13. The valency and oxidation state of oxygen in hydrogen peroxide, respectively, are

1) II, -2
2) II, -1
3) I, +4
4) III, -2

14. The valency and oxidation state of sulfur in pyrite FeS2 are, respectively,

1) IV, +5
2) II, -1
3) II, +6
4) III, +4

15. The valency and oxidation state of the nitrogen atom in ammonium bromide, respectively, are

1) IV, -3
2) III, +3
3) IV, -2
4) III, +4

16. The carbon atom shows negative degree oxidation in conjunction with

1) oxygen
2) sodium
3) fluorine
4) chlorine

17. A constant degree of oxidation in its compounds exhibits

1) strontium
2) iron
3) sulfur
4) chlorine

18. +3 oxidation state in their compounds can exhibit

1) chlorine and fluorine
2) phosphorus and chlorine
3) carbon and sulfur
4) oxygen and hydrogen

19. +4 oxidation state in their compounds can exhibit

1) carbon and hydrogen
2) carbon and phosphorus
3) carbon and calcium
4) nitrogen and sulfur

20. The oxidation state, equal to the group number, in its compounds exhibits

1) chlorine
2) iron
3) oxygen
4) fluorine

Electronegativity, like other properties of atoms of chemical elements, changes periodically with an increase in the ordinal number of the element:

The graph above shows the periodicity of the change in the electronegativity of the elements of the main subgroups, depending on the ordinal number of the element.

When moving down the subgroup of the periodic table, the electronegativity of chemical elements decreases, when moving to the right along the period, it increases.

Electronegativity reflects the non-metallicity of elements: the higher the value of electronegativity, the more non-metallic properties of the element are expressed.

Oxidation state

How to calculate the oxidation state of an element in a compound?

1) The oxidation state of chemical elements in simple substances is always zero.

2) There are elements that exhibit a constant oxidation state in complex substances:

3) There are chemical elements that exhibit a constant oxidation state in the vast majority of compounds. These elements include:

4) The algebraic sum of the oxidation states of all atoms in a molecule is always zero. The algebraic sum of the oxidation states of all atoms in an ion is equal to the charge of the ion.

5) The highest (maximum) oxidation state is equal to the group number. Exceptions that do not fall under this rule are elements of the secondary subgroup of group I, elements of the secondary subgroup of group VIII, as well as oxygen and fluorine.

Chemical elements whose group number does not match their the highest degree oxidation (required to remember)

6) The lowest oxidation state of metals is always zero, and the lowest oxidation state of non-metals is calculated by the formula:

lowest oxidation state of a non-metal = group number - 8

Based on the rules presented above, it is possible to establish the degree of oxidation of a chemical element in any substance.

Finding the oxidation states of elements in various compounds

Example 1

Determine the oxidation states of all elements in sulfuric acid.

Decision:

Let's write the formula for sulfuric acid:

The oxidation state of hydrogen in all complex substances is +1 (except for metal hydrides).

The oxidation state of oxygen in all complex substances is -2 (except for peroxides and oxygen fluoride OF 2). Let's arrange the known oxidation states:

Let us denote the oxidation state of sulfur as x:

The sulfuric acid molecule, like the molecule of any substance, is generally electrically neutral, because. the sum of the oxidation states of all atoms in a molecule is zero. Schematically, this can be depicted as follows:

Those. we got the following equation:

Let's solve it:

Thus, the oxidation state of sulfur in sulfuric acid is +6.

Example 2

Determine the oxidation state of all elements in ammonium dichromate.

Decision:

Let's write the formula of ammonium dichromate:

As in the previous case, we can arrange the oxidation states of hydrogen and oxygen:

However, we see that the oxidation states of two chemical elements at once, nitrogen and chromium, are unknown. Therefore, we cannot find the oxidation states in the same way as in the previous example (one equation with two variables does not have a unique solution).

Let us pay attention to the fact that the indicated substance belongs to the class of salts and, accordingly, has an ionic structure. Then we can rightly say that the composition of ammonium dichromate includes NH 4 + cations (the charge of this cation can be seen in the solubility table). Therefore, since there are two positive singly charged NH 4 + cations in the formula unit of ammonium dichromate, the charge of the dichromate ion is -2, since the substance as a whole is electrically neutral. Those. the substance is formed by NH 4 + cations and Cr 2 O 7 2- anions.

We know the oxidation states of hydrogen and oxygen. Knowing that the sum of the oxidation states of the atoms of all elements in the ion is equal to the charge, and denoting the oxidation states of nitrogen and chromium as x and y accordingly, we can write:

Those. we get two independent equations:

Solving which, we find x and y:

Thus, in ammonium dichromate, the oxidation states of nitrogen are -3, hydrogen +1, chromium +6, and oxygen -2.

How to determine the oxidation state of elements in organic substances can be read.

Valence

The valency of atoms is indicated by Roman numerals: I, II, III, etc.

The valence possibilities of an atom depend on the quantity:

1) unpaired electrons

2) unshared electron pairs in the orbitals of valence levels

3) empty electron orbitals of the valence level

Valence possibilities of the hydrogen atom

Let's depict the electronic graphic formula of the hydrogen atom:

It was said that three factors can affect the valence possibilities - the presence of unpaired electrons, the presence of unshared electron pairs at the outer level, and the presence of vacant (empty) orbitals of the outer level. We see one unpaired electron in the outer (and only) energy level. Based on this, hydrogen can exactly have a valency equal to I. However, at the first energy level there is only one sublevel - s, those. the hydrogen atom at the outer level does not have either unshared electron pairs or empty orbitals.

Thus, the only valency that a hydrogen atom can exhibit is I.

Valence possibilities of a carbon atom

Consider the electronic structure of the carbon atom. In the ground state, the electronic configuration of its outer level is as follows:

Those. In the ground state, the outer energy level of an unexcited carbon atom contains 2 unpaired electrons. In this state, it can exhibit a valency equal to II. However, the carbon atom very easily goes into an excited state when energy is imparted to it, and the electronic configuration of the outer layer in this case takes the form:

Although a certain amount of energy is spent on the process of excitation of the carbon atom, the expenditure is more than compensated for by the formation of four covalent bonds. For this reason, valence IV is much more characteristic of the carbon atom. So, for example, carbon has valency IV in the molecules of carbon dioxide, carbonic acid and absolutely all organic substances.

In addition to unpaired electrons and lone electron pairs, the presence of vacant () orbitals of the valence level also affects the valence possibilities. The presence of such orbitals in the filled level leads to the fact that the atom can act as an electron pair acceptor, i.e. form additional covalent bonds by the donor-acceptor mechanism. So, for example, contrary to expectations, in the molecule carbon monoxide CO bond is not double, but triple, which is clearly shown in the following illustration:

Valence possibilities of the nitrogen atom

Let's write down the electron-graphic formula of the external energy level of the nitrogen atom:

As can be seen from the illustration above, the nitrogen atom in its normal state has 3 unpaired electrons, and therefore it is logical to assume that it can exhibit a valency equal to III. Indeed, a valence equal to three is observed in ammonia molecules (NH 3), nitrous acid(HNO 2), nitrogen trichloride (NCl 3), etc.

It was said above that the valence of an atom of a chemical element depends not only on the number of unpaired electrons, but also on the presence of unshared electron pairs. This is due to the fact that the covalent chemical bond can be formed not only when two atoms provide each other with one electron each, but also when one atom that has an unshared pair of electrons - a donor () provides it to another atom with a vacant () orbital of the valence level (acceptor). Those. for the nitrogen atom, valency IV is also possible due to an additional covalent bond formed by the donor-acceptor mechanism. So, for example, four covalent bonds, one of which is formed by the donor-acceptor mechanism, is observed during the formation of the ammonium cation:

Despite the fact that one of the covalent bonds is formed by the donor-acceptor mechanism, all N-H bonds in the ammonium cation are absolutely identical and do not differ from each other.

A valency equal to V, the nitrogen atom is not able to show. This is due to the fact that the transition to an excited state is impossible for the nitrogen atom, in which the pairing of two electrons occurs with the transition of one of them to a free orbital, which is the closest in energy level. The nitrogen atom has no d-sublevel, and the transition to the 3s-orbital is energetically so expensive that the energy costs are not covered by the formation of new bonds. Many may wonder, what then is the valence of nitrogen, for example, in molecules nitric acid HNO 3 or nitric oxide N 2 O 5? Oddly enough, the valence there is also IV, as can be seen from the following structural formulas:

The dotted line in the illustration shows the so-called delocalized π -connection. For this reason, NO terminal bonds can be called "one and a half". Similar one-and-a-half bonds are also found in the ozone molecule O 3 , benzene C 6 H 6 , etc.

Valence possibilities of phosphorus

Let us depict the electron-graphic formula of the external energy level of the phosphorus atom:

As we can see, the structure of the outer layer of the phosphorus atom in the ground state and the nitrogen atom is the same, and therefore it is logical to expect for the phosphorus atom, as well as for the nitrogen atom, possible valences equal to I, II, III and IV, which is observed in practice.

However, unlike nitrogen, the phosphorus atom also has d-sublevel with 5 vacant orbitals.

In this regard, it is able to pass into an excited state, steaming electrons 3 s-orbitals:

Thus, the valency V for the phosphorus atom, which is inaccessible to nitrogen, is possible. So, for example, a phosphorus atom has a valence of five in the molecules of such compounds as phosphoric acid, phosphorus (V) halides, phosphorus (V) oxide, etc.

Valence possibilities of the oxygen atom

The electron-graphic formula of the external energy level of the oxygen atom has the form:

We see two unpaired electrons at the 2nd level, and therefore valency II is possible for oxygen. It should be noted that this valency of the oxygen atom is observed in almost all compounds. Above, when considering the valence possibilities of the carbon atom, we discussed the formation of the carbon monoxide molecule. The bond in the CO molecule is triple, therefore, oxygen is trivalent there (oxygen is an electron pair donor).

Due to the fact that the oxygen atom does not have an external level d-sublevels, depairing of electrons s and p- orbitals is impossible, which is why the valence capabilities of the oxygen atom are limited compared to other elements of its subgroup, for example, sulfur.

Valence possibilities of the sulfur atom

The external energy level of the sulfur atom in the unexcited state:

The sulfur atom, like the oxygen atom, has two unpaired electrons in its normal state, so we can conclude that a valency of two is possible for sulfur. Indeed, sulfur has valency II, for example, in the hydrogen sulfide molecule H 2 S.

As we can see, the sulfur atom at the outer level has d sublevel with vacant orbitals. For this reason, the sulfur atom is able to expand its valence capabilities, unlike oxygen, due to the transition to excited states. So, when unpairing a lone electron pair 3 p- sublevel, the sulfur atom acquires the electronic configuration of the outer level of the following form:

In this state, the sulfur atom has 4 unpaired electrons, which tells us about the possibility of sulfur atoms showing a valency equal to IV. Indeed, sulfur has valency IV in the molecules SO 2, SF 4, SOCl 2, etc.

When unpairing the second lone electron pair located on 3 s- sublevel, the external energy level acquires the following configuration:

In such a state, the manifestation of valence VI already becomes possible. An example of compounds with VI-valent sulfur are SO 3 , H 2 SO 4 , SO 2 Cl 2 etc.

Similarly, we can consider the valence possibilities of other chemical elements.