Sulfuric acid. Sulfuric acid: chemical properties, characteristics, production of sulfuric acid in production

When sulfur dioxide (SO2) is dissolved in water, it produces a chemical compound known as sulfurous acid. The formula of this substance is written as follows: H 2 SO 3. In truth, this connection is extremely unstable, with a certain assumption it can even be argued that it does not actually exist. Nevertheless, this formula is often used for the convenience of writing equations of chemical reactions.

Sulfurous acid: basic properties

An aqueous solution of sulfur dioxide is characterized by an acidic environment. It itself has all the properties that are inherent in acids, including the neutralization reaction. Sulfurous acid is capable of forming two types of salts: hydrosulfites and ordinary sulfites. Both belong to the group of reducing agents. The first type is usually obtained when sulfurous acid is present in quite large quantities: H 2 SO 3 + KOH -> KHSO 3 + H 2 O. Otherwise, ordinary sulfite is obtained: H 2 SO 3 + 2KOH -> K 2 SO 3 + 2H 2 O. A qualitative reaction to these salts is their interaction with strong acid. As a result, SO 2 gas is released, which is easily distinguished by its characteristic pungent odor.

Sulfurous acid can have a bleaching effect. It is no secret that a similar effect also gives chlorine water. However, the compound in question has one important advantage: unlike chlorine, sulfurous acid does not lead to the destruction of dyes; sulfur dioxide forms colorless chemical compounds with them. This property It is often used for bleaching fabrics made of silk, wool, plant material, as well as anything that is destroyed by oxidizing agents containing Cl. In the old days, this compound was even used to restore ladies' straw hats to their original appearance. H 2 SO 3 is a fairly strong reducing agent. With the access of oxygen, its solutions gradually turn into sulfuric acid. In those cases when it interacts with a stronger reducing agent (for example, hydrogen sulfide), sulfuric acid, on the contrary, exhibits oxidizing properties. The dissociation of this substance occurs in two stages. First, the hydrosulfite anion is formed, and then the second step occurs, and it turns into the sulfite anion.

Where is sulfurous acid used?

The production of this substance plays a big role in the production of all kinds of wine materials as an antiseptic; in particular, with its help it is possible to prevent the process of fermentation of the product in barrels and thereby ensure its safety. It is also used to prevent grain fermentation during the extraction of starch from it. Sulfurous acid and preparations based on it have broad antimicrobial properties, and therefore they are often used in the fruit and vegetable industry for canning. Calcium hydrosulfite, also called sulfite liquor, is used to process wood into sulfite pulp, from which paper is subsequently made. It remains to add that this compound is poisonous for humans, and therefore any laboratory work and experiments with it require caution and increased attention.

In redox processes, sulfur dioxide can be both an oxidizing agent and a reducing agent because the atom in this compound has an intermediate oxidation state of +4.

How SO 2 reacts with stronger reducing agents, such as:

SO 2 + 2H 2 S = 3S↓ + 2H 2 O

How does the reducing agent SO 2 react with stronger oxidizing agents, for example with in the presence of a catalyst, with, etc.:

2SO2 + O2 = 2SO3

SO 2 + Cl 2 + 2H 2 O = H 2 SO 3 + 2HCl

Receipt

1) Sulfur dioxide is formed when sulfur burns:

2) In industry it is obtained by roasting pyrite:

3) In the laboratory, sulfur dioxide can be obtained:

Cu + 2H 2 SO 4 = CuSO 4 + SO 2 + 2H 2 O

Application

Sulfur dioxide is widely used in the textile industry for bleaching various products. In addition, it is used in agriculture for the destruction of harmful microorganisms in greenhouses and cellars. Large quantities of SO 2 are used to produce sulfuric acid.

Sulfur oxide (VI) – SO 3 (sulfuric anhydride)

Sulfuric anhydride SO 3 is a colorless liquid, which at temperatures below 17 o C turns into a white crystalline mass. Absorbs moisture very well (hygroscopic).

Chemical properties

Acid-base properties

How a typical acid oxide, sulfuric anhydride, reacts:

SO 3 + CaO = CaSO 4

c) with water:

SO 3 + H 2 O = H 2 SO 4

A special property of SO 3 is its ability to dissolve well in sulfuric acid. A solution of SO 3 in sulfuric acid is called oleum.

Formation of oleum: H 2 SO 4 + n SO 3 = H 2 SO 4 ∙ n SO 3

Redox properties

Sulfur oxide (VI) is characterized by strong oxidizing properties (usually reduced to SO 2):

3SO 3 + H 2 S = 4SO 2 + H 2 O

Receipt and use

Sulfuric anhydride is formed by the oxidation of sulfur dioxide:

2SO2 + O2 = 2SO3

In its pure form, sulfuric anhydride has no practical significance. It is obtained as an intermediate product in the production of sulfuric acid.

H2SO4

Mention of sulfuric acid is first found among Arab and European alchemists. It was obtained by calcining iron sulfate (FeSO 4 ∙ 7H 2 O) in air: 2FeSO 4 = Fe 2 O 3 + SO 3 + SO 2 or a mixture with: 6KNO 3 + 5S = 3K 2 SO 4 + 2SO 3 + 3N 2, and the released sulfuric anhydride vapors condensed. Absorbing moisture, they turned into oleum. Depending on the method of preparation, H 2 SO 4 was called oil of vitriol or sulfur oil. In 1595, the alchemist Andreas Libavius established the identity of both substances.

For a long time, oil of vitriol was not widely used. Interest in it increased greatly after in the 18th century. The process of obtaining indigo carmine, a stable blue dye, from indigo was discovered. The first factory for the production of sulfuric acid was founded near London in 1736. The process was carried out in lead chambers, at the bottom of which water was poured. A molten mixture of saltpeter and sulfur was burned in the upper part of the chamber, then air was introduced into it. The procedure was repeated until an acid of the required concentration was formed at the bottom of the container.

In the 19th century the method was improved: instead of saltpeter, they began to use nitric acid (it gives when decomposed in the chamber). To return nitrous gases to the system, special towers were constructed, which gave the name to the whole process - the tower process. Factories operating using the tower method still exist today.

Sulfuric acid– it is a heavy oily liquid, colorless and odorless, hygroscopic; dissolves well in water. When concentrated sulfuric acid is dissolved in water, a large amount of heat is released, so it must be carefully poured into the water (and not vice versa!) and the solution must be mixed.

A solution of sulfuric acid in water with a H 2 SO 4 content of less than 70% is usually called dilute sulfuric acid, and a solution of more than 70% is concentrated sulfuric acid.

Chemical properties

Acid-base properties

Dilute sulfuric acid reveals everything characteristic properties strong acids. She reacts:

H 2 SO 4 + NaOH = Na 2 SO 4 + 2H 2 O

H 2 SO 4 + BaCl 2 = BaSO 4 ↓ + 2HCl

The process of interaction of Ba 2+ ions with SO 4 2+ sulfate ions leads to the formation of a white insoluble precipitate BaSO 4 . This qualitative reaction to sulfate ion.

Redox properties

In dilute H 2 SO 4 the oxidizing agents are H + ions, and in concentrated H 2 SO 4 the oxidizing agents are SO 4 2+ sulfate ions. SO 4 2+ ions are stronger oxidizing agents than H + ions (see diagram).

IN dilute sulfuric acid metals that are in the electrochemical voltage series are dissolved to hydrogen. In this case, metal sulfates are formed and the following is released:

Zn + H 2 SO 4 = ZnSO 4 + H 2

Metals that are located after hydrogen in the electrochemical voltage series do not react with dilute sulfuric acid:

Cu + H 2 SO 4 ≠

Concentrated sulfuric acid is a strong oxidizing agent, especially when heated. It oxidizes many and some organic substances.

When concentrated sulfuric acid interacts with metals that are located after hydrogen in the electrochemical voltage series (Cu, Ag, Hg), metal sulfates are formed, as well as the reduction product of sulfuric acid - SO 2.

Reaction of sulfuric acid with zinc

With more active metals (Zn, Al, Mg), concentrated sulfuric acid can be reduced to free sulfuric acid. For example, when sulfuric acid reacts with, depending on the concentration of the acid, various reduction products of sulfuric acid - SO 2, S, H 2 S - can be formed simultaneously:

Zn + 2H 2 SO 4 = ZnSO 4 + SO 2 + 2H 2 O

3Zn + 4H 2 SO 4 = 3ZnSO 4 + S↓ + 4H 2 O

4Zn + 5H 2 SO 4 = 4ZnSO 4 + H 2 S + 4H 2 O

In the cold, concentrated sulfuric acid passivates some metals, for example and, so it is transported in iron tanks:

Fe + H 2 SO 4 ≠

Concentrated sulfuric acid oxidizes some non-metals (, etc.), reducing to sulfur oxide (IV) SO 2:

S + 2H 2 SO 4 = 3SO 2 + 2H 2 O

C + 2H 2 SO 4 = 2SO 2 + CO 2 + 2H 2 O

Receipt and use

In industry, sulfuric acid is produced by contact method. The obtaining process occurs in three stages:

Obtaining SO 2 by roasting pyrite:

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

Oxidation of SO 2 to SO 3 in the presence of a catalyst – vanadium (V) oxide:

2SO2 + O2 = 2SO3

Dissolution of SO 3 in sulfuric acid:

H2SO4+ n SO 3 = H 2 SO 4 ∙ n SO 3

The resulting oleum is transported in iron tanks. Sulfuric acid of the required concentration is obtained from oleum by adding it to water. This can be expressed with a diagram:

H2SO4∙ n SO 3 + H 2 O = H 2 SO 4

Sulfuric acid has a variety of uses in a wide variety of applications. National economy. It is used for drying gases, in the production of other acids, for the production of fertilizers, various dyes and medicines.

Sulfuric acid salts

Most sulfates are highly soluble in water (CaSO 4 is slightly soluble, PbSO 4 is even less soluble and BaSO 4 is practically insoluble). Some sulfates containing water of crystallization are called vitriols:

CuSO 4 ∙ 5H 2 O copper sulfate

FeSO 4 ∙ 7H 2 O iron sulfate

Everyone has salts of sulfuric acid. Their relationship to heat is special.

Sulfates of active metals (,) do not decompose even at 1000 o C, while others (Cu, Al, Fe) decompose with slight heating into metal oxide and SO 3:

CuSO 4 = CuO + SO 3

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DEFINITION

Anhydrous sulfuric acid is a heavy, viscous liquid that is easily miscible with water in any proportion: the interaction is characterized by an extremely large exothermic effect (~880 kJ/mol at infinite dilution) and can lead to explosive boiling and splashing of the mixture if water is added to the acid; This is why it is so important to always reverse the order in preparing solutions and add the acid to the water, slowly and with stirring.

Some physical properties of sulfuric acid are given in the table.

Anhydrous H 2 SO 4 is a remarkable compound with unusually high dielectric constant and very high electrical conductivity, which is due to ionic autodissociation (autoprotolysis) of the compound, as well as the relay mechanism of conductivity with proton transfer, which ensures the occurrence of electric current through a viscous liquid with a large number of hydrogen bonds.

Table 1. Physical properties of sulfuric acid.

Preparation of sulfuric acid

Sulfuric acid is the most important industrial chemical and the cheapest acid produced in large volume anywhere in the world.

Concentrated sulfuric acid (“oil of vitriol”) was first obtained by heating “green vitriol” FeSO 4 × nH 2 O and was consumed in large quantities to produce Na 2 SO 4 and NaCl.

The modern process for producing sulfuric acid uses a catalyst consisting of vanadium(V) oxide with the addition of potassium sulfate on a silica or kieselguhr support. Sulfur dioxide SO2 is produced by burning pure sulfur or by roasting sulfide ore (primarily pyrite or ores of Cu, Ni and Zn) in the process of extracting these metals. SO2 is then oxidized to trioxide, and then sulfuric acid is obtained by dissolving in water:

S + O 2 → SO 2 (ΔH 0 - 297 kJ/mol);

SO 2 + ½ O 2 → SO 3 (ΔH 0 - 9.8 kJ/mol);

SO 3 + H 2 O → H 2 SO 4 (ΔH 0 - 130 kJ/mol).

Chemical properties of sulfuric acid

Sulfuric acid is a strong dibasic acid. In the first step, in solutions of low concentration, it dissociates almost completely:

H 2 SO 4 ↔H + + HSO 4 - .

Second stage dissociation

HSO 4 — ↔H + + SO 4 2-

occurs to a lesser extent. The dissociation constant of sulfuric acid in the second stage, expressed in terms of ion activity, K 2 = 10 -2.

As a dibasic acid, sulfuric acid forms two series of salts: medium and acidic. Average salts of sulfuric acid are called sulfates, and acid salts are called hydrosulfates.

Sulfuric acid greedily absorbs water vapor and is therefore often used to dry gases. The ability to absorb water also explains the charring of many organic substances, especially those belonging to the class of carbohydrates (fiber, sugar, etc.), when exposed to concentrated sulfuric acid. Sulfuric acid removes hydrogen and oxygen from carbohydrates, which form water, and carbon is released in the form of coal.

Concentrated sulfuric acid, especially hot, is a vigorous oxidizing agent. It oxidizes HI and HBr (but not HCl) to free halogens, coal to CO 2, sulfur to SO 2. These reactions are expressed by the equations:

8HI + H 2 SO 4 = 4I 2 + H 2 S + 4H 2 O;

2HBr + H 2 SO 4 = Br 2 + SO 2 + 2H 2 O;

C + 2H 2 SO 4 = CO 2 + 2SO 2 + 2H 2 O;

S + 2H 2 SO 4 = 3SO 2 + 2H 2 O.

The interaction of sulfuric acid with metals occurs differently depending on its concentration. Dilute sulfuric acid oxidizes with its hydrogen ion. Therefore, it interacts only with those metals that are in the voltage series only up to hydrogen, for example:

Zn + H 2 SO 4 = ZnSO 4 + H 2.

However, lead does not dissolve in dilute acid because the resulting salt, PbSO 4, is insoluble.

Concentrated sulfuric acid is an oxidizing agent due to sulfur (VI). It oxidizes metals in the voltage range up to and including silver. The products of its reduction may vary depending on the activity of the metal and the conditions (acid concentration, temperature). When interacting with low-active metals, such as copper, the acid is reduced to SO 2:

Cu + 2H 2 SO 4 = CuSO 4 + SO 2 + 2H 2 O.

When interacting with more active metals, reduction products can be both dioxide and free sulfur and hydrogen sulfide. For example, when interacting with zinc, the following reactions can occur:

Zn + 2H 2 SO 4 = ZnSO 4 + SO 2 + 2H 2 O;

3Zn + 4H 2 SO 4 = 3ZnSO 4 + S↓ + 4H 2 O;

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

Application of sulfuric acid

The use of sulfuric acid varies from country to country and from decade to decade. For example, in the USA, the main area of consumption of H 2 SO 4 is currently the production of fertilizers (70%), followed by chemical production, metallurgy, oil refining (~5% in each area). In the UK, the distribution of consumption by industry is different: only 30% of H2SO4 produced is used in the production of fertilizers, but 18% goes to paints, pigments and semi-products of dye production, 16% to chemical production, 12% to the production of soaps and detergents, 10 % for the production of natural and artificial fibers and 2.5% is used in metallurgy.

Examples of problem solving

EXAMPLE 1

Exercise

Determine the mass of sulfuric acid that can be obtained from one ton of pyrite if the yield of sulfur (IV) oxide in the roasting reaction is 90%, and sulfur (VI) oxide in the catalytic oxidation of sulfur (IV) is 95% of theoretical.

Solution

Let us write the equation for the pyrite firing reaction:

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

Let's calculate the amount of pyrite substance:

n(FeS 2) = m(FeS 2) / M(FeS 2);

M(FeS 2) = Ar(Fe) + 2×Ar(S) = 56 + 2×32 = 120g/mol;

n(FeS 2) = 1000 kg / 120 = 8.33 kmol.

Since in the reaction equation the coefficient for sulfur dioxide is twice as large as the coefficient for FeS 2, then the theoretically possible amount of sulfur oxide (IV) substance is equal to:

n(SO 2) theor = 2 ×n(FeS 2) = 2 ×8.33 = 16.66 kmol.

And the practically obtained amount of moles of sulfur oxide (IV) is:

n(SO 2) pract = η × n(SO 2) theor = 0.9 × 16.66 = 15 kmol.

Let us write the reaction equation for the oxidation of sulfur oxide (IV) to sulfur oxide (VI):

2SO 2 + O 2 = 2SO 3.

The theoretically possible amount of sulfur oxide (VI) is equal to:

n(SO 3) theor = n(SO 2) pract = 15 kmol.

And the practically obtained amount of moles of sulfur oxide (VI) is:

n(SO 3) pract = η × n(SO 3) theor = 0.5 × 15 = 14.25 kmol.

Let us write the reaction equation for the production of sulfuric acid:

SO 3 + H 2 O = H 2 SO 4.

Let's find the amount of sulfuric acid:

n(H 2 SO 4) = n(SO 3) pract = 14.25 kmol.

The reaction yield is 100%. The mass of sulfuric acid is equal to:

m(H 2 SO 4) = n(H 2 SO 4) × M(H 2 SO 4);

M(H 2 SO 4) = 2×Ar(H) + Ar(S) + 4×Ar(O) = 2×1 + 32 + 4×16 = 98 g/mol;

m(H 2 SO 4) = 14.25 × 98 = 1397 kg.

Answer

The mass of sulfuric acid is 1397 kg

Russian Peoples' Friendship University

Faculty of Foreign Languages and General Education Disciplines

Sulfur. Its use in medicine.

Performed

student of group SV-53

Head of chemistry seminars

Departments of Chemistry

Professor V.F. Zakharov

Moscow, 2002

Finding sulfur in nature.

Physical properties of sulfur.

Chemical properties of sulfur and its compounds.

1) Properties of a simple substance.

Properties of oxides:

sulfur(IV) oxide;

sulfur(VI) oxide.

Properties of acids and their salts:

sulfurous acid and its salts;

hydrogen sulfide and sulfides;

sulfuric acid and its salts.

Use of sulfur in medicine.

General characteristics of the oxygen subgroup

The oxygen subgroup includes five elements: oxygen, sulfur, selenium, tellurium and polonium (polonium is a radioactive element). These are p-elements of the VI group of the periodic system of D.I. Mendeleev. They have a group name - chalcogens, which means “ore-forming”.

Properties of oxygen subgroup elements

Properties
Serial number
Valence electrons
Atom ionization energy, eV
Relative electronegativity
Oxidation state in compounds
Atomic radius, nm

Chalcogen atoms have the same structure of the external energy level - ns 2 np 4. This explains the similarity of their chemical properties. All chalcogens in compounds with hydrogen and metals exhibit an oxidation state of –2, and in compounds with oxygen and other active non-metals – usually +4 and +6. For oxygen, as for fluorine, an oxidation state equal to the group number is not typical. It exhibits an oxidation state of usually –2 and in compounds with fluorine +2.

Hydrogen compounds of elements of the oxygen subgroup correspond to the formula H 2 R(R– element symbol ): H 2 O, H 2 S, H 2 Se, H 2 Te. They are called chalcohydrogens. When they are dissolved in water, acids are formed (the formulas are the same). The strength of these acids increases with increasing serial number element, which is explained by a decrease in binding energy in a series of compounds H 2 R. Water dissociating into ions H + And HE - , is an amphoteric electrolyte.

Sulfur, selenium and tellurium form the same forms of compounds with oxygen type R.O. 2 And R.O. 3 . They correspond to acids of the type H 2 R.O. 3 And H 2 R.O. 4 . As the atomic number of an element increases, the strength of these acids decreases. All of them exhibit oxidizing properties, and acids like H 2 R.O. 3 also restorative.

The properties of simple substances naturally change: with an increase in the charge of the nucleus, non-metallic properties weaken and metallic properties increase. Thus, oxygen and tellurium are non-metals, but the latter has a metallic luster and conducts electricity.

Finding sulfur in nature

Sulfur is widely distributed in nature. It makes up 0.05% of the mass of the earth's crust. In a free state (native sulfur) it is found in large quantities in Italy (the island of Sicily) and the USA. Deposits of native sulfur are available in the Kuibyshev region (Volga region), in the states of Central Asia, in the Crimea and other areas.

Sulfur often occurs in compounds with other elements. Its most important natural compounds are metal sulfides: FeS 2 – iron pyrite, or pyrite; HgS – cinnabar, etc., as well as sulfuric acid salts (crystalline hydrates): CaSO 4 ּ 2 H 2 O – plaster, Na 2 SO 4 ּ 10 H 2 O- Glauber's salt, MgSO 4 ּ 7 H 2 O– bitter salt, etc.

Physical properties of sulfur

Natural sulfur consists of a mixture of four stable isotopes: ,
,
,
.

Sulfur forms several allotropic modifications. Stable at room temperature rhombic sulfur It is a yellow powder, poorly soluble in water, but highly soluble in carbon disulfide, aniline and some other solvents. Conducts heat and electricity poorly. When crystallized from chloroform CHCl 3 or from carbon disulfide C.S. 2 it stands out in the form of transparent crystals of octahedral shape. Orthorhombic sulfur consists of cyclic molecules S 8 shaped like a crown. At 113 0 Sona melts, turning into a yellow, easily mobile liquid. With further heating, the melt thickens, as long polymer chains are formed in it. And if you heat sulfur to 444.6 0 C, it boils. Pouring boiling sulfur in a thin stream into cold water, available plastic sulfur – rubber-like modification consisting of polymer chains. As the melt cools slowly, dark yellow needle-shaped crystals form monoclinic sulfur.(t pl =119 0 C). Like rhombic sulfur, this modification consists of molecules S 8 . At room temperature plastic and monoclinic sulfur are unstable and spontaneously transform into orthorhombic sulfur powder.

Chemical properties of sulfur and its compounds

Properties of a simple substance.

The sulfur atom, having an incomplete external energy level, can add two electrons and exhibit an oxidation state of –2. Sulfur exhibits this degree of oxidation in compounds with metals and hydrogen (for example, Na 2 S And H 2 S). When electrons are given up or withdrawn to an atom of a more electronegative element, the oxidation state of sulfur can be +2, +4 and +6.

Sulfur easily forms compounds with many elements. When it burns in air or oxygen, sulfur oxide (IV) is formed. SO 2 and partially sulfur(VI) oxide SO 3 :

S+O 2 =SO 3

2S + 3O 2 = 2SO 3

These are the most important sulfur oxides.

When heated, sulfur combines directly with hydrogen, halogens (except iodine), phosphorus, coal, and all metals except gold, platinum and iridium. For example:

S+H 2 = H 2 S

3S + 2P = P 2 S 3

S+Cl 2 = SCl 2

2S+C=CS 2

S + Fe = FeS

As follows from the examples, in reactions with metals and some non-metals, sulfur is an oxidizing agent, and in reactions with more active non-metals, such as oxygen, chlorine, it is a reducing agent.

Properties of oxides

Sulfur oxide (IV)

Sulphur dioxide SO 2 - a colorless gas with a suffocating, pungent odor. When dissolved in water (at 0 0 C, 1 volume of water dissolves more than 70 volumes SO 2 ) sulfurous acid is formed H 2 SO 3 , which is known only in solutions.

In laboratory conditions to obtain SO 2 act on solid sodium sulfite with concentrated sulfuric acid:

Na 2 SO 3 + 2H 2 SO 4 = 2NaHSO 4 + SO 2 +H 2 O

In industry SO 2 obtained by roasting sulfide ores, such as pyrite:

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

or when burning sulfur. Sulfur dioxide is an intermediate product in the production of sulfuric acid. It is also used (together with sodium hydrosulfites NaHSO 3 and calcium Ca(HSO 3) 2) to separate cellulose from wood. This gas is used to fumigate trees and shrubs to kill agricultural pests.

Chemical reactions characteristic of SO 2 , can be divided into 3 groups:

Reactions that occur without changing the oxidation state, for example:

SO 2 +Ca(OH) 2 = CaSO 3  +H 2 O

2SO 2 + O 2 = 2SO 3

Reactions that occur with a decrease in the oxidation state of sulfur, for example:

SO 2 + 2H 2 S = 3S + 2H 2 O

Thus, SO 2 can exhibit both oxidizing and reducing properties.

Sulfur oxide (VI)

Sulfuric anhydride SO 3 at room temperature it is a colorless, easily volatile liquid (t boil = 44.8 0 C, t pl = 16.8 0 C), which over time turns into an asbestos-like modification consisting of shiny silky crystals. Sulfuric anhydride fibers are stable only in a sealed container. Absorbing moisture from the air, they turn into a thick, colorless liquid - oleum (from Latin oleum - “oil”). Although formally oleum can be considered a solution SO 3 V H 2 SO 4 , in fact, it is a mixture of various pyrosulfuric acids: H 2 S 2 O 7 ,H 2 S 3 O 10 etc. With water SO 3 interacts very energetically: it releases so much heat that the resulting tiny droplets of sulfuric acid create a fog. You need to work with this substance with extreme caution.

Sulfur (VI) oxide is produced by oxidation SO 2 oxygen only in the presence of a catalyst:

2SO 2 + O 2  2SO 3 +Q.

The need to use a catalyst in this reversible reaction is due to the fact that the good yield SO 3 (i.e., a shift of equilibrium to the right) can only be obtained with a decrease in temperature, however, with low temperatures The reaction rate drops very significantly.

Sulfur (VI) oxide combines vigorously with water to form sulfuric acid:

SO 3 + H 2 O = H 2 SO 4

Properties of acids and their salts

Sulfurous acid and its salts

Sulfur (IV) oxide is highly soluble in water (in 1 40 volumes of SO 2 dissolve in a volume of water at 20 0 C). In this case, sulfurous acid, which exists only in aqueous solution, is formed:

SO 2 + N 2 O = N 2 SO 3

Compound reaction SO 2 reversible with water. In an aqueous solution, sulfur oxide (IV) and sulfurous acid are in chemical equilibrium, which can be displaced. When tying N 2 SO 3 with alkali (neutralization of acid), the reaction proceeds towards the formation of sulfurous acid; when deleting SO 2 (blowing through a nitrogen solution or heating) the reaction proceeds towards the starting substances. A solution of sulfurous acid always contains sulfur oxide (IV), which gives it a pungent odor.

Sulfurous acid has all the properties of acids. In solution N 2 SO 3 dissociates stepwise:

N 2 SABOUT 3  H + + HSO 4 –

HSO 3 -  H + + SO 3 2-

As a dibasic acid, it forms two series of salts - sulfites and hydrosulfites. Sulfites are formed when an acid is completely neutralized with an alkali:

N 2 SO 3 + 2 NaOH =NAH.S.ABOUT 4 + 2H 2 ABOUT

Hydrosulfites are obtained when there is a lack of alkali (compared to the amount required to completely neutralize the acid):

N 2 SO 3 + NaOH = NaHSO 3 + N 2 ABOUT

Like sulfur(IV) oxide, sulfurous acid and its salts are strong reducing agents. At the same time, the degree of sulfur oxidation increases. So, N 2 SABOUT 3 easily oxidized into sulfuric acid even by atmospheric oxygen:

2H 2 SO 3 + O 2 = 2H 2 SO 4

Therefore, solutions of sulfurous acid that have been stored for a long time always contain sulfuric acid.

The oxidation of sulfurous acid with bromine and potassium permanganate occurs even more easily:

N 2 SABOUT 3 + Br 2 + N 2 O = N 2 SO 4 + 2НВr

5H 2 S0 3 + 2KmnABOUT 4 = 2H 2 SO 4 + 2MnSO 4 + K 2 SABOUT 4 + 2H 2 ABOUT

Sulfur (IV) oxide and sulfurous acid decolorize many dyes, forming colorless compounds. The latter can decompose again when heated or exposed to light, as a result of which the color is restored. Therefore, the whitening effect SO 2 And N 2 SO 4 differs from the bleaching effect of chlorine. Typically, sulfur (IV) oxide is used to bleach wool, silk and straw (these materials are destroyed by chlorine water).

Calcium hydrosulfite solution has important applications. Ca(HSO 3 ) 2 (sulfite liquor), which is used to treat wood fibers and paper pulp.

Hydrogen sulfide and sulfides

Hydrogen sulfide N 2 S - colorless gas with the smell of rotten eggs. It is highly soluble in water (at 20 °C, 2.5 volumes of hydrogen sulfide are dissolved in 1 volume of water). A solution of hydrogen sulfide in water is called hydrogen sulfide water or hydrosulfide acid (it exhibits the properties of a weak acid).

Hydrogen sulfide is a very poisonous gas that can damage nervous system. Therefore, it is necessary to work with it in fume hoods or with hermetically sealed devices. Permissible content of H 2 Sv production premises is 0.01 mg in 1 liter of air.

Hydrogen sulfide occurs naturally in volcanic gases and in the waters of some mineral springs, for example Pyatigorsk; Matsesta. It is formed during the decay of sulfur-containing organic substances of various plant and animal residues. This explains the characteristic bad smell sewage, cesspools and garbage dumps.

Hydrogen sulfide can be produced by directly combining sulfur with hydrogen when heated:

S+ N 2 = H 2 S

But it is usually obtained by the action of dilute hydrochloric or sulfuric acid on iron (II) sulfide:

2HCl + FeS =FEUl 2 + N 2 S

This reaction is often carried out in a Kipp apparatus.

H 2 S is a less strong compound than water. This is due to the large size of the sulfur atom compared to the oxygen atom. Therefore, the H-0 bond is shorter and stronger than the H-S bond. When heated strongly, hydrogen sulfide almost completely decomposes into sulfur and hydrogen:

N 2 S = S + N 2

Gaseous H 2 S burns in air with a blue flame to form sulfur oxide (IV) and water:

2H 2 S + 3 O 2 = 2 SO 2 + 2H 2 ABOUT

With a lack of oxygen, sulfur and water are formed:

2H 2 S + O 2 = 2 S+ 2H 2 ABOUT

This reaction is used to produce sulfur from hydrogen sulfide on an industrial scale.

Hydrogen sulfide is a fairly strong reducing agent. This important chemical property of it can be explained as follows. In solution N 2 S relatively easily gives electrons to oxygen molecules in the air:

N 2 S - 2e- = S + 2H + 2

O 2 + 4 e- = 2O 2- 1

In this case, H 2 S is oxidized by atmospheric oxygen to sulfur, which makes hydrogen sulfide water cloudy. Overall reaction equation:

2 N 2 S+O 2 = 2S + 2N 2 O

This also explains the fact that hydrogen sulfide does not accumulate in very large quantities in nature during the decay of organic substances - air oxygen oxidizes it into free sulfur.

Hydrogen sulfide reacts vigorously with solutions of halogens. For example:

N 2 S+I 2 = 2HI + S

Sulfur is released and the iodine solution becomes discolored.

Hydrogen sulfide acid, as a dibasic acid, forms two series of salts - medium (sulfides) and acidic (hydrosulfides). For example, Na 2 S - sodium sulfide, NaHS- sodium hydrosulfide. Hydrosulfides are almost all highly soluble in water. Sulfides of alkali and alkaline earth metals are also soluble in water, while other metals are practically insoluble or slightly soluble; some of them do not dissolve in dilute acids. Therefore, such sulfides can be easily obtained by passing hydrogen sulfide through salts of the corresponding metal, for example:

WITHuSO 4 + N 2 S = CuS + H 2 SO 4

Some sulfides have a characteristic color: CuS And RbS - black, WITHdS- yellow, ZnS- white, MnS- pink, SnS- brown, Sb 2 S 3 - orange, etc. Qualitative analysis of cations is based on the different solubility of sulfides and the different colors of many of them.

Sulfuric acid and its salts

Sulfuric acid is a heavy, colorless, oily liquid. Extremely hygroscopic. It absorbs moisture with the release of a large amount of heat, so you cannot add water to concentrated acid - the acid will splash. To dilute, add small amounts of sulfuric acid to water.

Anhydrous sulfuric acid dissolves up to 70% of sulfur (VI) oxide. At ordinary temperatures it is non-volatile and odorless. When heated, it splits off SO 3 until a solution containing 98.3% is formed N 2 SO 4 . Anhydrous H 2 SO 4 almost does not conduct electric current.

Concentrated sulfuric acid chars organic substances - sugar, paper, wood, fibers, etc., removing water elements from them. In this case, sulfuric acid hydrates are formed. The charring of sugar can be expressed by the equation

WITH 12 N 22 ABOUT 11 + nN 2 SO 4 = 12C + H 2 SO 4 ּ nN 2 ABOUT

The resulting carbon partially reacts with the acid:

C + 2H 2 SO 4 = CO 2 + 2 SO 2 + 2H 2 ABOUT

Therefore, the acid that goes on sale has a brown color due to dust and organic substances that have accidentally fallen into it and become charred in it.

Gas drying is based on the absorption (removal) of water by sulfuric acid.

As a strong non-volatile acid N 2 SO 4 displaces other acids from dry salts. For example:

NaNO3 + H 2 SO 4 = NaHSO 4 + NNO 3

However, if N 2 SABOUT 4 is added to salt solutions, then displacement of acids does not occur.

A very important chemical property of sulfuric acid is its relationship to metals. Dilute and concentrated sulfuric acid react with them differently. Diluted sulfuric acid oxidizes only metals located in the voltage series to the left of hydrogen, due to ions H + , For example:

Zn+H 2 SO 4 ( razb ) = ZnSO 4 +H 2

Concentrated Sulfuric acid does not react with many metals at ordinary temperatures. Therefore, anhydrous sulfuric acid can be stored in iron containers and transported in steel tanks. However, when heated, concentrated N 2 SO 4 interacts with almost all metals (except Rt, Au and some others), as well as with non-metals. At the same time, it acts as an oxidizing agent and is usually reduced to SO 2 . In this case, hydrogen is not released, but water is formed. For example:

WITHu+2N 2 SO 4 = WITHuSO 4 + SO 2  + 2 N 2 O

2Ag + 2H 2 SO 4 = Ag 2 SO 4 + SO 2  + 2H 2 O

C+2H 2 SO 4 + = CO 2  +2SO 2  + 2H 2 O

2P+5H 2 SO 4 = 2H 3 P.O. 4 +5SO 2 

Sulfuric acid has all the properties of acids.

Sulfuric acid, being dibasic, forms two series of salts: medium, called sulfates, and acidic, called hydrosulfates. Sulfates are formed when an acid is completely neutralized by an alkali (for 1 mole of acid there are 2 moles of alkali), and hydrosulfates are formed when there is a lack of alkali (for 1 mole of acid there is 1 mole of alkali):

N 2 SO 4 + 2 NAOH= Na 2 SO 4 + 2H 2 ABOUT

N 2 SO 4 + NaOH = NAHSO 4 + N 2 ABOUT

Many salts of sulfuric acid are of great practical importance.

Most sulfuric acid salts are soluble in water. Salts CaSO 4 And RbSO 4 are slightly soluble in water, and VaSO 4 practically insoluble in both water and acids. This property allows the use of any soluble barium salt, for example Youl 2 , as a reagent for sulfuric acid and its salts (more precisely, for the ion SO 4 2- ):

H 2 SO 4 + BaCl 2 = BaSO 4  + 2HCl

NaSO 4 + BaCl 2 = BaSO 4  + 2NaCl

In this case, a white precipitate of barium sulfate, insoluble in water and acids, precipitates.

Sulfuric acid is the most important product of the basic chemical industry, which produces inorganic acids, alkalis, salts, mineral fertilizers and chlorine.

In terms of variety of applications, sulfuric acid ranks first among acids. The largest amount of it is consumed to produce phosphorus and nitrogen fertilizers. Being a non-volatile acid, sulfuric acid is used to produce other acids - hydrochloric, hydrofluoric, phosphoric, acetic, etc. A lot of it is used for purifying petroleum products - gasoline, kerosene and lubricating oils - from harmful impurities. In mechanical engineering, sulfuric acid is used to clean the metal surface from oxides before coating (nickel plating, chrome plating, etc.). Sulfuric acid is used in the production of explosives, artificial fibers, dyes, plastics and many others. It is used to fill batteries. In agriculture it is used to control weeds (herbicide).

This determines the importance of sulfuric acid in our national economy.

Use of sulfur in medicine

Purified sulfur (Sulfurdepuratum) - a fine lemon-yellow powder - is used for enterobiasis as an anthelmintic. It is also a mild laxative and is part of the complex licorice root powder. A sterile 1-2% solution of purified sulfur in peach oil (sulfozine) is sometimes used for pyrogenic therapy for syphilis.

In addition, sulfur compounds, both organic and inorganic, are widely used in medicine. Sulfur atoms are found in many drugs with very different effects. Since it is not possible to cover them all, we will limit ourselves to a few examples.

Sulfuric acid (H2SO4) is one of the most caustic acids and dangerous reagents known to man, especially in concentrated form. Chemically pure sulfuric acid is a heavy toxic liquid of oily consistency, odorless and colorless. It is obtained by contact oxidation of sulfur dioxide (SO2).

At a temperature of + 10.5 °C, sulfuric acid turns into a frozen glassy crystalline mass, greedily, like a sponge, absorbing moisture from environment. In industry and chemistry, sulfuric acid is one of the main chemical compounds and occupies a leading position in terms of production volume in tons. This is why sulfuric acid is called the “blood of chemistry.” Fertilizers are produced using sulfuric acid. medications, other acids, large fertilizers and much more.

Basic physical and chemical properties of sulfuric acid

Sulfuric acid in its pure form (formula H2SO4), at a concentration of 100%, is a colorless, thick liquid. The most important property of H2SO4 is its high hygroscopicity - the ability to remove water from the air. This process is accompanied by a large-scale release of heat.
H2SO4 is a strong acid.
Sulfuric acid is called a monohydrate - it contains 1 mole of H2O (water) per 1 mole of SO3. Due to its impressive hygroscopic properties, it is used to extract moisture from gases.
Boiling point – 330 °C. In this case, the acid decomposes into SO3 and water. Density – 1.84. Melting point – 10.3 °C/.
Concentrated sulfuric acid is a powerful oxidizing agent. To initiate a redox reaction, the acid must be heated. The result of the reaction is SO2. S+2H2SO4=3SO2+2H2O
Depending on the concentration, sulfuric acid reacts with metals differently. In a dilute state, sulfuric acid is capable of oxidizing all metals that are in the voltage series before hydrogen. The exception is the most resistant to oxidation. Dilute sulfuric acid reacts with salts, bases, amphoteric and basic oxides. Concentrated sulfuric acid is capable of oxidizing all metals in the voltage series, including silver.
Sulfuric acid forms two types of salts: acidic (these are hydrosulfates) and intermediate (sulfates)
H2SO4 reacts actively with organic substances and non-metals, some of which it can turn into coal.
Sulfuric anhydrite dissolves well in H2SO4, and in this case oleum is formed - a solution of SO3 in sulfuric acid. Outwardly, it looks like this: fuming sulfuric acid, releasing sulfuric anhydrite.
Sulfuric acid in aqueous solutions is a strong dibasic acid, and when it is added to water, a huge amount of heat is released. When preparing dilute solutions of H2SO4 from concentrated ones, it is necessary to add a heavier acid to the water in a small stream, and not vice versa. This is done to avoid boiling water and splashing acid.

Concentrated and diluted sulfuric acids

Concentrated solutions of sulfuric acid include solutions from 40% that can dissolve silver or palladium.

Dilute sulfuric acid includes solutions whose concentration is less than 40%. These are not such active solutions, but they are capable of reacting with brass and copper.

Preparation of sulfuric acid

Production of sulfuric acid in industrial scale was launched in the 15th century, but at that time it was called “oil of vitriol.” If earlier humanity consumed only a few tens of liters of sulfuric acid, now modern world the calculation is in millions of tons per year.

Sulfuric acid production is carried out industrially, and there are three of them:

Contact method.
Nitrose method
Other methods

Let's talk in detail about each of them.

Contact production method

The contact production method is the most common, and it performs the following tasks:

The result is a product that satisfies the needs of the maximum number of consumers.
During production, environmental damage is reduced.

In the contact method, the following substances are used as raw materials:

pyrite (sulfur pyrite);
sulfur;
vanadium oxide (this substance acts as a catalyst);
hydrogen sulfide;
sulfides of various metals.

Before starting the production process, raw materials are pre-prepared. To begin with, in special crushing plants, the pyrite is crushed, which allows, by increasing the contact area of the active substances, to speed up the reaction. Pyrite undergoes purification: it is lowered into large containers of water, during which waste rock and all kinds of impurities float to the surface. At the end of the process they are removed.

The production part is divided into several stages:

After crushing, the pyrite is cleaned and sent to the furnace, where it is fired at temperatures up to 800 °C. According to the counterflow principle, air is supplied into the chamber from below, and this ensures that the pyrite is in a suspended state. Today, this process takes a few seconds, but previously it took several hours to fire. During the roasting process, waste appears in the form of iron oxide, which is removed and subsequently transferred to the metallurgical industry. During firing, water vapor, O2 and SO2 gases are released. When purification from water vapor and tiny impurities is completed, pure sulfur oxide and oxygen are obtained.
In the second stage, an exothermic reaction occurs under pressure using a vanadium catalyst. The reaction starts when the temperature reaches 420 °C, but it can be increased to 550 °C to increase efficiency. During the reaction, catalytic oxidation occurs and SO2 becomes SO3.
The essence of the third stage of production is as follows: absorption of SO3 in an absorption tower, during which oleum H2SO4 is formed. In this form, H2SO4 is poured into special containers (it does not react with steel) and is ready to meet the end consumer.

During production, as we said above, a lot of thermal energy is generated, which is used for heating purposes. Many sulfuric acid plants install steam turbines, which use the released steam to generate additional electricity.

Nitrous method for producing sulfuric acid

Despite the advantages of the contact production method, which produces more concentrated and pure sulfuric acid and oleum, quite a lot of H2SO4 is produced by the nitrous method. In particular, at superphosphate plants.

For the production of H2SO4, the starting material, both in the contact and nitrose methods, is sulfur dioxide. It is obtained specifically for these purposes by burning sulfur or roasting sulfur metals.

Processing sulfur dioxide into sulfurous acid involves the oxidation of sulfur dioxide and the addition of water. The formula looks like this:
SO2 + 1|2 O2 + H2O = H2SO4

But sulfur dioxide does not react directly with oxygen, therefore, with the nitrous method, sulfur dioxide is oxidized using nitrogen oxides. Higher oxides of nitrogen (we are talking about nitrogen dioxide NO2, nitrogen trioxide NO3) during this process are reduced to nitrogen oxide NO, which is subsequently oxidized again by oxygen to higher oxides.

The production of sulfuric acid by the nitrous method is technically formalized in two ways:

Chamber.
Tower.

The nitrous method has a number of advantages and disadvantages.

Disadvantages of the nitrous method:

The result is 75% sulfuric acid.
Product quality is low.
Incomplete return of nitrogen oxides (addition of HNO3). Their emissions are harmful.
The acid contains iron, nitrogen oxides and other impurities.

Advantages of the nitrous method:

The cost of the process is lower.
Possibility of SO2 recycling at 100%.
Simplicity of hardware design.

Main Russian sulfuric acid plants

The annual production of H2SO4 in our country is in the six-digit range - about 10 million tons. The leading producers of sulfuric acid in Russia are companies that are, in addition, its main consumers. It's about about companies whose field of activity is the production of mineral fertilizers. For example, “Balakovo mineral fertilizers”, “Ammophos”.

Works in Armyansk, Crimea largest producer titanium dioxide in Eastern Europe "Crimean Titan". In addition, the plant produces sulfuric acid, mineral fertilizers, iron sulfate, etc.

Many factories produce various types of sulfuric acid. For example, battery sulfuric acid is produced by: Karabashmed, FKP Biysk Oleum Plant, Svyatogor, Slavia, Severkhimprom, etc.

Oleum is produced by UCC Shchekinoazot, FKP Biysk Oleum Plant, Ural Mining and Metallurgical Company, Kirishinefteorgsintez PA, etc.

Sulfuric acid of special purity is produced by OHC Shchekinoazot, Component-Reaktiv.

Spent sulfuric acid can be purchased at the ZSS and HaloPolymer Kirovo-Chepetsk plants.

Manufacturers of technical sulfuric acid are Promsintez, Khiprom, Svyatogor, Apatit, Karabashmed, Slavia, Lukoil-Permnefteorgsintez, Chelyabinsk Zinc Plant, Electrozinc, etc.

Due to the fact that pyrite is the main raw material in the production of H2SO4, and this is a waste of enrichment enterprises, its suppliers are the Norilsk and Talnakh enrichment factories.

The world's leading positions in H2SO4 production are occupied by the USA and China, which account for 30 million tons and 60 million tons, respectively.

Scope of application of sulfuric acid

The world consumes about 200 million tons of H2SO4 annually, from which a wide range of products are produced. Sulfuric acid rightfully holds the palm among other acids in terms of the scale of use for industrial purposes.

As you already know, sulfuric acid is one of the most important products chemical industry, therefore the scope of sulfuric acid is quite wide. The main areas of use of H2SO4 are as follows:

Sulfuric acid is used in enormous volumes for the production of mineral fertilizers, and this consumes about 40% of the total tonnage. For this reason, factories that produce H2SO4 are built next to factories that produce fertilizers. These are ammonium sulfate, superphosphate, etc. During their production, sulfuric acid is taken in its pure form (100% concentration). To produce a ton of ammophos or superphosphate you will need 600 liters of H2SO4. These fertilizers are in most cases used in agriculture.
H2SO4 is used to produce explosives.
Purification of petroleum products. To obtain kerosene, gasoline and mineral oils, purification of hydrocarbons is required, which occurs using sulfuric acid. In the process of refining oil to purify hydrocarbons, this industry “takes” as much as 30% of the world’s tonnage of H2SO4. In addition, the octane number of fuel is increased with sulfuric acid and wells are treated during oil production.
In the metallurgical industry. Sulfuric acid in metallurgy is used to remove scale and rust from wire and sheet metal, as well as to restore aluminum in the production of non-ferrous metals. Before coating metal surfaces with copper, chromium or nickel, the surface is etched with sulfuric acid.
In the production of medicines.
In the production of paints.
In the chemical industry. H2SO4 is used in the production of detergents, ethylene, insecticides, etc., and without it these processes are impossible.
For the production of other known acids, organic and inorganic compounds used for industrial purposes.

Salts of sulfuric acid and their use

The most important salts of sulfuric acid:

Glauber's salt Na2SO4 · 10H2O (crystalline sodium sulfate). The scope of its application is quite capacious: the production of glass, soda, in veterinary medicine and medicine.
Barium sulfate BaSO4 is used in the production of rubber, paper, and white mineral paint. In addition, it is indispensable in medicine for fluoroscopy of the stomach. It is used to make “barium porridge” for this procedure.
Calcium sulfate CaSO4. In nature, it can be found in the form of gypsum CaSO4 2H2O and anhydrite CaSO4. Gypsum CaSO4 · 2H2O and calcium sulfate are used in medicine and construction. When gypsum is heated to a temperature of 150 - 170 °C, partial dehydration occurs, resulting in burnt gypsum, known to us as alabaster. By mixing alabaster with water to the consistency of a batter, the mass quickly hardens and turns into a kind of stone. It is this property of alabaster that is actively used in construction work: casts and casting molds are made from it. In plastering work, alabaster is indispensable as a binding material. Patients in trauma departments are given special fixing hard bandages - they are made on the basis of alabaster.
Iron sulfate FeSO4 · 7H2O is used to prepare ink, impregnate wood, and also in agricultural activities to kill pests.
Alum KCr(SO4)2 · 12H2O, KAl(SO4)2 · 12H2O, etc. are used in the production of paints and the leather industry (leather tanning).
Many of you know copper sulfate CuSO4 · 5H2O firsthand. This is an active assistant in agriculture in the fight against plant diseases and pests - grain is treated with an aqueous solution of CuSO4 · 5H2O and sprayed on plants. It is also used to prepare some mineral paints. And in everyday life it is used to remove mold from walls.
Aluminum sulfate – it is used in the pulp and paper industry.

Sulfuric acid in diluted form is used as an electrolyte in lead batteries. In addition, it is used to produce detergents and fertilizers. But in most cases it comes in the form of oleum - this is a solution of SO3 in H2SO4 (you can also find other formulas of oleum).

Amazing fact! Oleum is more chemically active than concentrated sulfuric acid, but despite this, it does not react with steel! It is for this reason that it is easier to transport than sulfuric acid itself.

The scope of use of the “queen of acids” is truly large-scale, and it is difficult to talk about all the ways it is used in industry. It is also used as an emulsifier in Food Industry, for water purification, in the synthesis of explosives and many other purposes.

The history of sulfuric acid

Who among us has not at least once heard of copper sulfate? So, it was studied in ancient times, and in some works of the beginning of the new era, scientists discussed the origin of vitriol and their properties. Vitriol was studied by the Greek physician Dioscorides and the Roman nature explorer Pliny the Elder, and in their works they wrote about the experiments they carried out. For medical purposes, various vitriol substances were used by the ancient physician Ibn Sina. How vitriol was used in metallurgy was discussed in the works of alchemists Ancient Greece Zosima from Panopolis.

The first way to obtain sulfuric acid is the process of heating potassium alum, and there is information about this in the alchemical literature of the 13th century. At that time, the composition of alum and the essence of the process were unknown to alchemists, but already in the 15th century chemical synthesis sulfuric acid began to be studied purposefully. The process was as follows: alchemists treated a mixture of sulfur and antimony (III) sulfide Sb2S3 by heating with nitric acid.

In medieval times in Europe, sulfuric acid was called "oil of vitriol", but then the name changed to vitriol acid.

In the 17th century, Johann Glauber obtained sulfuric acid as a result of burning potassium nitrate and native sulfur in the presence of water vapor. As a result of the oxidation of sulfur with saltpeter, sulfur oxide was obtained, which reacted with water vapor, resulting in a liquid with an oily consistency. This was oil of vitriol, and this name for sulfuric acid still exists today.

In the thirties of the 18th century, a pharmacist from London, Ward Joshua, used this reaction to industrial production sulfuric acid, but in the Middle Ages its consumption was limited to several tens of kilograms. The scope of use was narrow: for alchemical experiments, purification of precious metals and in pharmacy. Concentrated sulfuric acid in small volumes was used in the production of special matches that contained bertholite salt.

Vitriol acid appeared in Rus' only in the 17th century.

In Birmingham, England, John Roebuck adapted the above method for producing sulfuric acid in 1746 and launched production. At the same time, he used durable large leaded chambers, which were cheaper than glass containers.

This method held its position in industry for almost 200 years, and 65% sulfuric acid was obtained in chambers.

After a while, the English Glover and the French chemist Gay-Lussac improved the process itself, and sulfuric acid began to be obtained with a concentration of 78%. But such an acid was not suitable for the production of, for example, dyes.

At the beginning of the 19th century, new methods were discovered for the oxidation of sulfur dioxide into sulfuric anhydride.

Initially this was done using nitrogen oxides, and then platinum was used as a catalyst. These two methods of oxidizing sulfur dioxide have been further improved. The oxidation of sulfur dioxide on platinum and other catalysts became known as the contact method. And the oxidation of this gas with nitrogen oxides is called the nitrous method for producing sulfuric acid.

The British acetic acid merchant Peregrine Philips patented an economical process for the production of sulfur oxide (VI) and concentrated sulfuric acid only in 1831, and it is this method that is familiar to the world today as a contact method for its production.

Superphosphate production began in 1864.

In the eighties of the nineteenth century in Europe, the production of sulfuric acid reached 1 million tons. The main producers were Germany and England, producing 72% of the total volume of sulfuric acid in the world.

Transportation of sulfuric acid is a labor-intensive and responsible undertaking.

Sulfuric acid belongs to the class of dangerous chemicals, and upon contact with the skin causes severe burns. In addition, it can cause chemical poisoning in humans. If during transportation the certain rules, then sulfuric acid, due to its explosiveness, can cause a lot of harm to both people and the environment.

Sulfuric acid has been assigned a hazard class of 8 and must be transported by specially trained and trained professionals. An important condition for the delivery of sulfuric acid is compliance with specially developed Rules for the Transportation of Dangerous Goods.

Transportation by road is carried out in accordance with the following rules:

For transportation, special containers are made from a special steel alloy that does not react with sulfuric acid or titanium. Such containers do not oxidize. Dangerous sulfuric acid is transported in special sulfuric acid chemical tanks. They differ in design and are selected for transportation depending on the type of sulfuric acid.
When transporting fuming acid, specialized isothermal thermos tanks are taken, in which the required temperature is maintained to preserve the chemical properties of the acid.
If ordinary acid is transported, then a sulfuric acid tank is selected.
Transportation of sulfuric acid by road, such types as fuming, anhydrous, concentrated, for batteries, glover, is carried out in special containers: tanks, barrels, containers.
The transportation of dangerous goods can only be carried out by drivers who have an ADR certificate.
Travel time has no restrictions, since during transportation you must strictly adhere to the permissible speed.
During transportation, a special route is built, which should pass places of large crowds of people and production facilities.
Transport must have special markings and danger signs.

Dangerous properties of sulfuric acid for humans

Sulfuric acid is increased danger for the human body. Its toxic effect occurs not only upon direct contact with the skin, but upon inhalation of its vapors, when sulfur dioxide is released. Hazardous effects include:

Respiratory system;
Skin;
Mucous membranes.

Intoxication of the body can be enhanced by arsenic, which is often included in sulfuric acid.

Important! As you know, severe burns occur when acid comes into contact with the skin. Poisoning by sulfuric acid vapors is no less dangerous. The safe dose of sulfuric acid in the air is only 0.3 mg per 1 square meter.

If sulfuric acid gets on the mucous membranes or skin, a severe burn appears that does not heal well. If the scale of the burn is impressive, the victim develops burn disease, which can even lead to fatal outcome, if qualified medical assistance is not provided in a timely manner.

Important! For an adult lethal dose sulfuric acid is only 0.18 cm per 1 liter.

Of course, “experiencing” the toxic effects of acid in everyday life is problematic. Most often, acid poisoning occurs due to neglect of industrial safety precautions when working with the solution.

Mass poisoning with sulfuric acid vapor may occur due to technical problems at work or negligence, and a massive release into the atmosphere occurs. To prevent such situations they work special services, whose task is to control the functioning of production where dangerous acid is used.

What symptoms are observed during sulfuric acid intoxication?

If the acid was ingested:

Pain in the area of the digestive organs.
Nausea and vomiting.
Abnormal bowel movements as a result of severe intestinal disorders.
Heavy secretion of saliva.
Due to toxic effects on the kidneys, the urine becomes reddish.
Swelling of the larynx and throat. Wheezing and hoarseness occur. This can be fatal from suffocation.
Brown spots appear on the gums.
The skin turns blue.

For a burn skin There may be all the complications inherent in a burn disease.

In case of vapor poisoning, the following picture is observed:

Burn of the mucous membrane of the eyes.
Nose bleed.
Burn of the mucous membranes of the respiratory tract. In this case, the victim experiences severe pain.
Swelling of the larynx with symptoms of suffocation (lack of oxygen, skin turns blue).
If the poisoning is severe, there may be nausea and vomiting.

It is important to know! Acid poisoning after ingestion is much more dangerous than intoxication from inhalation of vapors.

First aid and therapeutic procedures for sulfuric acid injury

Proceed as follows when in contact with sulfuric acid:

First of all, call an ambulance. If liquid gets inside, do gastric lavage warm water. After this, you will need to drink 100 grams of sunflower or olive oil in small sips. In addition, you should swallow a piece of ice, drink milk or burnt magnesia. This must be done to reduce the concentration of sulfuric acid and alleviate the human condition.
If acid gets into your eyes, you need to rinse them with running water and then drip them with a solution of dicaine and novocaine.
If acid gets on the skin, rinse the burned area well under running water and apply a bandage with soda. You need to rinse for about 10-15 minutes.
In case of vapor poisoning, you need to go to Fresh air, and also rinse the affected mucous membranes with water when available.

In a hospital setting, treatment will depend on the area of the burn and the degree of poisoning. Pain relief is carried out only with novocaine. To avoid the development of infection in the affected area, the patient is given a course of antibiotic therapy.

In case of gastric bleeding, plasma or blood transfusion is administered. The source of bleeding can be eliminated surgically.

Sulfuric acid occurs in nature in its 100% pure form. For example, in Italy, Sicily, in the Dead Sea, you can see a unique phenomenon - sulfuric acid seeps straight from the bottom! What happens is this: pyrite from the earth’s crust serves in this case as a raw material for its formation. This place is also called the Lake of Death, and even insects are afraid to fly near it!
After large volcanic eruptions in earth's atmosphere drops of sulfuric acid can often be found, and in such cases the "culprit" may bring Negative consequences to the environment and cause serious climate change.
Sulfuric acid is an active absorbent of water, so it is used as a gas desiccant. In the old days, to prevent indoor windows from fogging up, this acid was poured into jars and placed between the glass of window openings.
It is sulfuric acid that is the main cause of deposition. acid rain. main reason formation of acid rain - air pollution with sulfur dioxide, and when dissolved in water it forms sulfuric acid. Sulfur dioxide, in turn, is released when fossil fuels are burned. In acid rain studied in recent years, the content has increased nitric acid. The reason for this phenomenon is the reduction of sulfur dioxide emissions. Despite this fact, the main cause of acid rain remains sulfuric acid.

We offer you a video selection interesting experiments with sulfuric acid.

Let's consider the reaction of sulfuric acid when it is poured into sugar. In the first seconds of sulfuric acid entering the flask with sugar, the mixture darkens. After a few seconds the substance turns black. Then the most interesting thing happens. The mass begins to grow rapidly and climb outside the flask. The output is a proud substance, similar to porous charcoal, 3-4 times larger than the original volume.

The author of the video suggests comparing the reaction of Coca-Cola with hydrochloric acid and sulfuric acid. When Coca-Cola is mixed with hydrochloric acid, no visual changes are observed, but when mixed with sulfuric acid, Coca-Cola begins to boil.

An interesting interaction can be observed when sulfuric acid comes into contact with toilet paper. Toilet paper consists of cellulose. When acid hits the cellulose molecule, it instantly breaks down releasing free carbon. Similar charring can be observed when acid comes into contact with wood.

Into a flask with concentrated acid I add a small piece of potassium. In the first second, smoke is released, after which the metal instantly flares up, ignites and explodes, breaking into pieces.

In the following experiment, when sulfuric acid hits a match, it ignites. In the second part of the experiment, aluminum foil with acetone and a match inside is immersed. The foil is instantly heated, releasing a huge amount of smoke and completely dissolving it.

An interesting effect is observed when adding baking soda into sulfuric acid. Soda instantly turns colored yellow. The reaction proceeds with rapid boiling and an increase in volume.

We strongly advise against carrying out all of the above experiments at home. Sulfuric acid is very aggressive and toxic substance. Such experiments must be carried out in special rooms equipped with forced ventilation. The gases released in reactions with sulfuric acid are very toxic and can cause damage to the respiratory tract and poisoning of the body. In addition, similar experiments are carried out in facilities personal protection skin and respiratory organs. Take care of yourself!