Compounds with an oxidation state of –2. The most important sulfur compounds in oxidation state -2 are hydrogen sulfide and sulfides. Hydrogen sulfide - H 2 S - a colorless gas with a characteristic smell of rotting protein, toxic. The hydrogen sulfide molecule has an angular shape, the bond angle is 92º. It is formed by the direct interaction of hydrogen with sulfur vapor. In the laboratory, hydrogen sulfide is produced by the action of strong acids for metal sulfides:

Na 2 S + 2HCl \u003d 2NaCl + H 2 S

Hydrogen sulfide is a strong reducing agent, oxidized even by sulfur oxide (IV).

2H 2 S -2 + S +4 O 2 \u003d 3S 0 + 2H 2 O

Depending on the conditions, the products of sulfide oxidation can be S, SO 2 or H 2 SO 4:

2KMnO 4 + 5H 2 S -2 + 3H 2 SO 4 ® 2MnSO 4 + 5S + K 2 SO 4 + 8H 2 O;

H 2 S -2 + 4Br 2 + 4H 2 O = H 2 S +4 O 4 + 8HBr

In air and in an oxygen atmosphere, hydrogen sulfide burns, forming sulfur or SO 2, depending on the conditions.

Hydrogen sulfide is slightly soluble in water (2.5 volumes of H 2 S per 1 volume of water) and behaves like a weak dibasic acid.

H 2 S H + + HS - ; K 1 \u003d 1 × 10 -7

HS - H + + S 2-; K 2 \u003d 2.5 × 10 -13

As a dibasic acid, hydrogen sulfide forms two series of salts: hydrosulfides ( acid salts) and sulfides (medium salts). For example, NaHS is hydrosulfide and Na 2 S is sodium sulfide.

Sulfides of most metals in water are sparingly soluble, painted in characteristic colors and differ in solubility in acids: ZnS - white, CdS - yellow-orange, MnS - flesh-colored, HgS, CuS, PbS, FeS - black, SnS - brown, SnS 2 - yellow. Alkaline sulfides are readily soluble in water. alkaline earth metals and also ammonium sulfide. Soluble sulfides are highly hydrolysed.

Na 2 S + H 2 O NaHS + NaOH

Sulfides, like oxides, are basic, acidic, and amphoteric. The main properties are sulfides of alkali and alkaline earth metals, acid properties- non-metal sulfides. The difference in the chemical nature of sulfides is manifested in hydrolysis reactions and in the interaction of sulfides of different nature with each other. During hydrolysis, basic sulfides form an alkaline medium, acidic sulfides are irreversibly hydrolyzed with the formation of the corresponding acids:

SiS 2 + 3H 2 O \u003d H 2 SiO 3 + 2H 2 S

Amphoteric sulfides are insoluble in water, some of them, for example, aluminum, iron (III), chromium (III) sulfides, are completely hydrolyzed:

Al 2 S 3 + 3H 2 O \u003d 2Al (OH) 3 + 3H 2 S

When basic and acid sulfides interact, thiosalts are formed. The thioacids corresponding to them are usually unstable, their decomposition is similar to the decomposition of oxygen-containing acids.

CS 2 + Na 2 S \u003d Na 2 CS 3; Na 2 CS 3 + H 2 SO 4 \u003d H 2 CS 3 + Na 2 SO 4;

sodium thiocarbonate thiocarbonic acid

H 2 CS 3 = H 2 S + CS 2

persulfide compounds. The tendency of sulfur to form homochains is realized in persulfides (polysulfides), which are formed by heating solutions of sulfides with sulfur:

Na 2 S + (n-1) S \u003d Na 2 S n

Persulfides are found in nature, for example, the widespread mineral pyrite FeS 2 is iron(II) persulfide. Under the action of mineral acids on solutions of polysulfides, polysulfanes were isolated - unstable oil-like substances of the composition H 2 S n, where n varies from 2 to 23.

Persulfides, like peroxides, exhibit both oxidizing and reducing properties, and also easily disproportionate.

Na 2 S 2 + SnS \u003d SnS 2 + Na 2 S; 4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2;

Na 2 S 2 -1 \u003d S 0 + Na 2 S -2

Compounds with an oxidation state of +4. The most important is sulfur oxide (IV) - a colorless gas with a sharp unpleasant smell of burning sulfur. The SO 2 molecule has an angular structure (OSO angle is 119.5 °):

In industry, SO 2 is obtained by roasting pyrite or burning sulfur. Laboratory method for obtaining sulfur dioxide - the action of strong mineral acids on sulfites.

Na 2 SO 3 + 2HCl \u003d 2NaCl + SO 2 + H 2 O

Sulfur(IV) oxide is an energetic reducing agent

S +4 O 2 + Cl 2 \u003d S +6 O 2 Cl 2,

but, interacting with strong reducing agents, it can act as an oxidizing agent:

2H 2 S + S + 4 O 2 \u003d 3S 0 + 2H 2 O

Sulfur dioxide is highly soluble in water (40 volumes per 1 volume of water). In an aqueous solution, hydrated SO 2 molecules partially dissociate to form a hydrogen cation:

SO 2 × H 2 O H + + HSO 3 - 2H + + SO 3 2-

For this reason, an aqueous solution of sulfur dioxide is often considered as a solution of sulfurous acid - H 2 SO 3, although this compound does not seem to exist in reality. However, salts of sulfurous acid are stable and can be isolated individually:

SO 2 + NaOH \u003d NaHSO 3; SO 2 + 2NaOH \u003d Na 2 SO 3

sodium hydrosulfite sodium sulfite

The sulfite anion has the structure of a trigonal pyramid with a sulfur atom at the top. The lone pair of the sulfur atom is spatially directed; therefore, the anion, an active donor of an electron pair, easily transforms into tetrahedral HSO 3 - and exists in the form of two tautomeric forms:

Alkali metal sulfites are highly soluble in water, largely hydrolyzed:

SO 3 2- + H 2 O HSO 3 - + OH -

Strong reducing agents, during storage of their solutions, are gradually oxidized by atmospheric oxygen, when heated, they disproportionate:

2Na 2 S +4 O 3 + O 2 \u003d 2Na 2 S +6 O 4; 4Na 2 S +4 O 3 \u003d Na 2 S -2 + 3Na 2 S +6 O 4

The +4 oxidation state appears in halides and oxohalides:

SF 4 SOF 2 SOCl 2 SOBr 2

Sulfur(IV) fluoride Sulfur(IV) oxofluoride Sulfur(IV) oxochloride Sulfur(IV) oxobromide

In all the above molecules, a lone electron pair is localized on the sulfur atom, SF 4 has the shape of a distorted tetrahedron (bisphenoid), SOHal 2 is a trigonal pyramid.

Sulfur(IV) fluoride is a colorless gas. Sulfur(IV) oxochloride (thionyl chloride, thionyl chloride) is a colorless liquid with a pungent odor. These substances are widely used in organic synthesis to obtain organofluorine and chlorine compounds.

Compounds of this type are acidic, as evidenced by their relationship to water:

SF 4 + 3H 2 O \u003d H 2 SO 3 + 4HF; SOCl 2 + 2H 2 O \u003d H 2 SO 3 + 2HCl.

Compounds with an oxidation state of +6:

SF 6 SO 2 Cl 2 SO 3 H 2 SO 4 2-

sulfur(VI) fluoride, sulfur(VI) dioxodichloride, sulfur(VI) oxide sulphuric acid sulfate anion

Sulfur hexafluoride is a colorless inert gas used as a gaseous dielectric. The SF 6 molecule is highly symmetrical and has the geometry of an octahedron. SO 2 Cl 2 (sulphuryl chloride, sulfuryl chloride) is a colorless liquid fuming in air due to hydrolysis, used in organic synthesis as a chlorinating reagent:

SO 2 Cl 2 + 2H 2 O \u003d H 2 SO 4 + 2HCl

Sulfur(VI) oxide is a colorless liquid (bp 44.8 °C, mp 16.8 °C). In the gaseous state, SO 3 has a monomeric structure; in the liquid state, it predominantly exists in the form of cyclic trimeric molecules; in the solid state, it is a polymer.

In industry, sulfur trioxide is obtained by catalytic oxidation of its dioxide:

2SO 2 + O 2 ¾® 2SO 3

In the laboratory, SO 3 can be obtained by distillation of oleum - a solution of sulfur trioxide in sulfuric acid.

SO 3 is a typical acidic oxide that vigorously attaches water and other proton-containing reagents:

SO 3 + H 2 O \u003d H 2 SO 4; SO 3 + HF = HOSO 2 F

fluorosulfuric (fluorosulfonic)

acid

Sulphuric acid- H 2 SO 4 - colorless oily liquid, so pl. 10.4 °C, b.p. 340 °C (with decomposition). Freely soluble in water, strong dibasic acid. Concentrated sulfuric acid is a vigorous oxidizing agent, especially when heated. It oxidizes non-metals and metals that are in the series of standard electrode potentials to the right of hydrogen:

C + 2H 2 SO 4 \u003d CO 2 + 2SO 2 + 2H 2 O; Cu + 2H 2 SO 4 \u003d CuSO 4 + SO 2 + 2H 2 O

Interacting with more active metals, sulfuric acid can be reduced to sulfur or hydrogen sulfide, for example,

4Zn + 5H 2 SO 4 (conc.) = 4ZnSO 4 + H 2 S + 4H 2 O

Cold concentrated sulfuric acid passivates many metals (iron, lead, aluminum, chromium) due to the formation of a dense oxide or salt film on their surface.

Sulfuric acid forms two series of salts: containing sulfate anion - SO 4 2- (medium salts) and containing hydrosulfate anion - HSO 4 - (acid salts). Sulfates are generally well soluble in water, poorly soluble BaSO 4 , SrSO 4 , PbSO 4 , Cu 2 SO 4 . The formation of a white finely crystalline precipitate of barium sulfate when exposed to a solution of barium chloride is a qualitative reaction to the sulfate anion. This reaction is also used for the quantitative determination of sulfur.

Ba 2+ + SO 4 2- \u003d BaSO 4 ¯

The most important salts of sulfuric acid are: Na 2 SO 4 × 10H 2 O - mirabilite, Glauber's salt - used in the production of soda and glass; MgSO 4 × 7H 2 O - bitter Epsom salt - used in medicine as a laxative, for finishing fabrics, for tanning leather; CaSO 4 × 2H 2 O - gypsum - used in medicine and construction; CaSO 4 ×1 / 2H 2 O - alabaster - used as a building material; CuSO 4 × 5H 2 O - copper sulfate - used in agriculture to protect plants from fungal diseases; FeSO 4 × 7H 2 O - iron sulphate - is used in agriculture as a microfertilizer and in water treatment as a coagulator; K 2 SO 4 × Al 2 (SO 4) 3 × 24H 2 O - potassium alum - used for tanning leather.

The synthesis of sulfuric acid in industry is carried out by the contact method, the first stage of which is pyrite roasting:

4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2

2SO 2 + O 2 \u003d 2SO 3

When SO 3 is dissolved in concentrated sulfuric acid, a whole series of polysulfuric acids is formed. A mixture of H 2 SO 4, H 2 S 2 O 7, H 2 S 3 O 10, H 2 S 4 O 13 is a thick oily liquid fuming in air - oleum. When oleum is diluted with water S-O-S connections break and polysulfuric acids are converted into sulfuric acid of the required concentration.

Pyrosulfuric (two-sulfuric) acid- H 2 S 2 O 7:

Colorless, fusible crystals released from oleum.

SO 3 + H 2 SO 4 \u003d H 2 S 2 O 7

Salts of pyrosulfuric acid - pyrosulfates (disulfates) - are obtained thermal decomposition hydrosulfates:

KHSO 4 \u003d K 2 S 2 O 7 + H 2 O

Thiosulfuric acid- H 2 S 2 O 3 - exists in two tautomeric forms:

In aqueous solutions, it is unstable and decomposes with the release of sulfur and SO 2:

H 2 S 2 O 3 \u003d S¯ + SO 2 + H 2 O

Salts of thiosulfuric acid - thiosulfates - are stable and can be obtained by boiling sulfur with aqueous solutions of sulfites:

Na 2 SO 3 + S \u003d Na 2 S 2 O 3

The properties of thiosulfates are determined by the presence of sulfur atoms in two different oxidation states -2 and +6. So the presence of sulfur in the oxidation state -2 determines the reducing properties:

Na 2 SO 3 S -2 + Cl 2 + H 2 O \u003d Na 2 S +6 O 4 + S 0 + 2HCl

Sodium thiosulfate is widely used in photography as a fixative and in analytical chemistry for the quantitative determination of iodine and iodine-releasing substances (iodometric analysis).

Polythionic acids. Tetrahedral structural units in polysulfuric acids can be combined through sulfur atoms, and compounds of the general formula H 2 S x O 6 are obtained, in which x \u003d 2 - 6.

Polythionic acids are unstable, but form stable salts - polythionates. For example. sodium tetrathionate is formed by the action of iodine on an aqueous solution of sodium thiosulfate:

Na 2 S 2 O 3 + I 2 = Na 2 S 4 O 6 + 2NaI

Peroxosulfuric (persulfuric) acids. The role of a bridge connecting the structural units of polysulfuric acids can be played by a peroxide group. The same group is part of monopersulfuric acid:

H 2 SO 5 - monopersulfuric acid H 2 S 2 O 8 - peroxodisulfuric acid

(caro acid)

Peroxosulfuric acids are hydrolyzed to form hydrogen peroxide:

H 2 SO 5 + H 2 O \u003d H 2 SO 4 + H 2 O 2; H 2 S 2 O 8 + 2H 2 O \u003d 2H 2 SO 4 + H 2 O 2.

Peroxodisulfuric acid is obtained by electrolysis of an aqueous solution of sulfuric acid:

2HSO 4 - - 2e - \u003d H 2 S 2 O 8

Forms salts - persulfates. Ammonium persulfate - (NH 4) 2 S 2 O 8 - is used in the laboratory as an oxidizing agent.

The formal charge of an atom in compounds is an auxiliary quantity, it is usually used in descriptions of the properties of elements in chemistry. This conditional electric charge is the degree of oxidation. Its meaning changes as a result of many chemical processes. Although the charge is formal, it vividly characterizes the properties and behavior of atoms in redox reactions (ORDs).

Oxidation and reduction

In the past, chemists used the term "oxidation" to describe the interaction of oxygen with other elements. The name of the reactions comes from the Latin name for oxygen - Oxygenium. Later it turned out that other elements also oxidize. In this case, they are restored - they attach electrons. Each atom during the formation of a molecule changes the structure of its valence electron shell. In this case, a formal charge appears, the value of which depends on the number of conditionally given or received electrons. To characterize this value, the English chemical term "oxidation number" was previously used, which means "oxidation number" in translation. Its use is based on the assumption that the bonding electrons in molecules or ions belong to the atom with the higher electronegativity (EO). The ability to retain their electrons and attract them from other atoms is well expressed in strong non-metals (halogens, oxygen). Strong metals (sodium, potassium, lithium, calcium, other alkali and alkaline earth elements) have opposite properties.

Determination of the degree of oxidation

The oxidation state is the charge that an atom would acquire if the electrons involved in the formation of the bond were completely shifted to a more electronegative element. There are substances that do not have a molecular structure (alkali metal halides and other compounds). In these cases, the oxidation state coincides with the charge of the ion. The conditional or real charge shows what process took place before the atoms acquired their current state. A positive oxidation number is total electrons that have been removed from atoms. The negative value of the oxidation state is equal to the number of acquired electrons. By change in oxidation state chemical element judge what happens to its atoms during the reaction (and vice versa). The color of the substance determines what changes in the state of oxidation have occurred. Compounds of chromium, iron and a number of other elements in which they exhibit different valences are colored differently.

Negative, zero and positive oxidation state values

Simple substances are formed by chemical elements with the same EO value. In this case, the bonding electrons belong to all structural particles equally. Therefore, in simple substances the elements do not have an oxidation state (H 0 2, O 0 2, C 0). When atoms accept electrons or the general cloud shifts in their direction, it is customary to write charges with a minus sign. For example, F -1, O -2, C -4. By donating electrons, atoms acquire a real or formal positive charge. In OF 2 oxide, the oxygen atom donates one electron each to two fluorine atoms and is in the O +2 oxidation state. It is believed that in a molecule or a polyatomic ion, the more electronegative atoms receive all the binding electrons.

Sulfur is an element that exhibits different valencies and oxidation states.

Chemical elements of the main subgroups often exhibit a lower valence equal to VIII. For example, the valency of sulfur in hydrogen sulfide and metal sulfides is II. The element is characterized by intermediate and higher valencies in the excited state, when the atom gives up one, two, four or all six electrons and exhibits valences I, II, IV, VI, respectively. The same values, only with a minus or plus sign, have the oxidation states of sulfur:

• in fluorine sulfide gives one electron: -1;
• in hydrogen sulfide, the lowest value: -2;
• in dioxide intermediate state: +4;
• in trioxide, sulfuric acid and sulfates: +6.

In its highest oxidation state, sulfur only accepts electrons; in its lowest state, it exhibits strong reducing properties. The S +4 atoms can act as reducing or oxidizing agents in compounds, depending on the conditions.

Transfer of electrons in chemical reactions

In the formation of a sodium chloride crystal, sodium donates electrons to the more electronegative chlorine. The oxidation states of the elements coincide with the charges of the ions: Na +1 Cl -1 . For molecules created by the socialization and displacement of electron pairs to a more electronegative atom, only the concept of a formal charge is applicable. But it can be assumed that all compounds are composed of ions. Then the atoms, by attracting electrons, acquire a conditional negative charge, and by giving away, they acquire a positive one. In reactions, indicate how many electrons are displaced. For example, in the carbon dioxide molecule C +4 O - 2 2, the index indicated in the upper right corner of the chemical symbol for carbon displays the number of electrons removed from the atom. Oxygen in this substance has an oxidation state of -2. The corresponding index when chemical sign O is the number of added electrons in the atom.

How to calculate oxidation states

Counting the number of electrons donated and added by atoms can be time consuming. The following rules make this task easier:

1. In simple substances, the oxidation states are zero.
2. The sum of the oxidation of all atoms or ions in a neutral substance is zero.
3. In a complex ion, the sum of the oxidation states of all elements must correspond to the charge of the entire particle.
4. A more electronegative atom acquires a negative oxidation state, which is written with a minus sign.
5. Less electronegative elements receive positive oxidation states, they are written with a plus sign.
6. Oxygen generally exhibits an oxidation state of -2.
7. For hydrogen, the characteristic value is: +1, in metal hydrides it occurs: H-1.
8. Fluorine is the most electronegative of all elements, its oxidation state is always -4.
9. For most metals, oxidation numbers and valences are the same.

Oxidation state and valence

Most compounds are formed as a result of redox processes. The transition or displacement of electrons from one element to another leads to a change in their oxidation state and valency. Often these values ​​coincide. As a synonym for the term "oxidation state", the phrase "electrochemical valency" can be used. But there are exceptions, for example, in the ammonium ion, nitrogen is tetravalent. At the same time, the atom of this element is in the oxidation state -3. In organic substances, carbon is always tetravalent, but the oxidation states of the C atom in methane CH 4, formic alcohol CH 3 OH and acid HCOOH have different values: -4, -2 and +2.

Redox reactions

Redox processes include many of the most important processes in industry, technology, animate and inanimate nature: combustion, corrosion, fermentation, intracellular respiration, photosynthesis, and other phenomena.

When compiling the OVR equations, the coefficients are selected using the electronic balance method, in which the following categories are operated:

• oxidation states;
• the reducing agent donates electrons and is oxidized;
• the oxidizing agent accepts electrons and is reduced;
• the number of given electrons must be equal to the number of attached ones.

The acquisition of electrons by an atom leads to a decrease in its oxidation state (reduction). The loss of one or more electrons by an atom is accompanied by an increase in the oxidation number of the element as a result of reactions. For OVR, flowing between ions of strong electrolytes in aqueous solutions, not the electronic balance, but the method of half-reactions is more often used.

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: . The negative value of the degree of oxidation have 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 an important quantitative characteristic of the element in its compounds. The lowest value of 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).

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 are expressed in the element.

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 manifest in complex substances constant oxidation state:

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

Element	The oxidation state in almost all compounds	Exceptions
hydrogen H	+1	Alkali and alkaline earth metal hydrides, for example:
oxygen O	-2	Hydrogen and metal peroxides: Oxygen fluoride -

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 highest oxidation state (mandatory to memorize)

6) The lowest oxidation state of metals is always zero, and lowest degree oxidation 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.

Solution:

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.

Solution:

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 states of elements in organic matter 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 offset 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 valence IV in molecules 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 valence equal to III. Indeed, a valency of three is observed in the molecules of ammonia (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 valency 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 valence 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.

The subgroup of chalcogens includes sulfur - this is the second of the elements that is able to form big number 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 system.

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 in nature with mass numbers 32,33,34 and 36. This element is the sixth most abundant 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 features electronic structure atom. It is expressed by the following configuration formula: 1s 2 2s 2 2p 6 3s 2 3p 4 .

The above order reflects only steady state 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 is increasing significantly, as well as the number various substances with 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, is the yellow that cannot be dissolved in water, but only in organic solvents. Differs in fragility and a beautiful form of the structure presented in the form of 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 brown viscous solution, which, upon further heating or sudden 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 for the 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 almost no smell, however hydrogen compounds sulfur gives off a strong 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" in the school chemistry course takes not one lesson, but several. After all, there are a lot of them. This is due to the chemical activity of this substance. She can appear as oxidizing properties with stronger reducing agents (metals, boron, etc.), and reducing ones with most non-metals.

However, despite such activity, only interaction with fluorine occurs under 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 sharp, extremely unpleasant, reminiscent of a 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 widely used 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 a 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 oxidation, then the properties of these compounds are quite understandable. They are strong oxidizing agents.

Sulfur oxide (VI) - sulfuric anhydride - is a volatile colorless liquid. Feature- 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, the manufacture of optical instruments and glasses, 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? 9th grade is the best period. This is not the very beginning, when everything is new and incomprehensible for children. This is the middle in learning 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 isolated in a special group of 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 drugs (aspirin, sulfanilamide or streptocide).

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

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