2 3 or Na 2 SO 3 S, a salt of sodium and thiosulfuric acid, forms the crystalline hydrate Na 2 S 2 O 3 5H 2 O.

Receipt

• oxidation of polysulfides;
• boiling excess sulfur with Na 2 SO 3:

$\backslash mathsf(Na\_2SO\_3\; +\; S\; \backslash rightarrow\; Na\_2S\_2O\_3)$

• interaction of H 2 S and SO 2 with NaOH (by-product in the production of NaHSO 3, sulfur dyes, during the purification of industrial gases from):

$\backslash mathsf(4SO\_2\; +\; 2H\_2S\; +\; 6NaOH\; \backslash rightarrow\; 3Na\_2S\_2O\_3\; +\; 5H\_2O)$

• boiling excess sulfur with sodium hydroxide:

$\backslash mathsf(4S\; +\; 6NaOH\; \backslash rightarrow\; 2Na\_2S\; +\; Na\_2S\_2O\_3\; +\; 3H\_2O)$

Then, in the above reaction, sodium sulfite adds sulfur to form sodium thiosulfate.

At the same time, during this reaction, sodium polysulfides are formed (they give the solution a yellow color). To destroy them, SO 2 is passed into the solution.

• pure anhydrous sodium thiosulfate can be prepared by reacting sulfur with sodium nitrite in formamide. This reaction proceeds quantitatively (at 80 °C for 30 minutes) according to the equation:

$\backslash mathsf(2NaNO\_2\; +\; 2S\; \backslash rightarrow\; Na\_2S\_2O\_3\; +\; N\_2O)$

• dissolution of sodium sulfide in water in the presence of atmospheric oxygen:

$\backslash mathsf(2Na\_2S\; +\; 2O\_2\; +\; H\_2O\; \backslash rightarrow\; Na\_2S\_2O\_3\; +\; 2NaOH)$

Physical and chemical properties

Colorless monoclinic crystals. Molar mass 248.17 g/mol (pentahydrate).

Soluble in water (41.2% at 20 o C, 69.86% at 80 o C).

At 48.5 °C, the crystalline hydrate dissolves in its water of crystallization, forming a supersaturated solution; dehydrates at about 100 o C.

When heated to 220 °C, it decomposes according to the following scheme:

$\backslash mathsf(4Na\_2S\_2O\_3\; \backslash rightarrow\; 3Na\_2SO\_4\; +\; Na\_2S\; +\; 4S)$

Sodium thiosulfate is a strong reducing agent:

With strong oxidizing agents, such as free chlorine, it is oxidized to sulfates or sulfuric acid:

$\backslash mathsf(Na\_2S\_2O\_3\; +\; 4Cl\_2\; +\; 5H\_2O\; \backslash rightarrow\; 2H\_2SO\_4\; +\; 2NaCl\; +\; 6HCl)$

With weaker or slow-acting oxidizing agents, for example, iodine, it is converted into salts of tetrathionic acid:

$\backslash mathsf(2Na\_2S\_2O\_3\; +\; I\_2\; \backslash rightarrow\; Na\_2S\_4O\_6\; +\; 2NaI)$

The above reaction is very important, as it serves as the basis for iodometry. It should be noted that in an alkaline environment, the oxidation of sodium thiosulfate with iodine can proceed to sulfate.

It is impossible to isolate thiosulfuric acid (hydrogen thiosulfate) by the reaction of sodium thiosulfate with a strong acid, since it is unstable and immediately decomposes into water, sulfur and sulfur dioxide:

$\backslash mathsf(Na\_2S\_2O\_3\; +\; H\_2SO\_4\; \backslash rightarrow\; Na\_2SO\_4\; +\; H\_2O\; +\; S\; +\; SO\_2)$

Molten crystalline hydrate Na 2 S 2 O 3 ·5H 2 O is very prone to supercooling.

Application

• for removing traces of chlorine after bleaching fabrics
• for extracting silver from ores;
• fixer in photography;
• reagent in iodometry
• antidote for poisoning: , , and other heavy metals, cyanides (translates them into thiocyanates), etc.
• for intestinal disinfection;
• for the treatment of scabies (together with hydrochloric acid);
• anti-inflammatory and anti-burn agent;
• can be used as a medium for determining molecular weights by lowering the freezing point (cryoscopic constant 4.26°)
• registered in the food industry as a food additive E539.
• additives for concrete.
• for cleansing tissues of iodine
• gauze dressings soaked in a sodium thiosulfate solution were used to protect the respiratory system from the poisonous agent chlorine in World War I.
• antidote for lidocaine overdose.

see also

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• // Encyclopedic Dictionary of Brockhaus and Efron: in 86 volumes (82 volumes and 4 additional). - St. Petersburg. , 1890-1907.

Nystatin ( ) Natamycin ( ) Hachimitsin ( ) Pecilocin ( ) Mepartricin ( ) Pyrrolnitrine ( ) Griseofulvin ( ) Combinations ( ) Clotrimazole ( ) Miconazole ( ) Econazole ( ) Chlormidazole ( ) Isoconazole ( ) Thiabendazole ( ) Tioconazole ( ) Ketoconazole ( ) Sulconazole ( ) Bifonazole ( ) Oxiconazole ( ) Fenticonazole ( ) Omoconazole ( ) Sertaconazole ( ) Fluconazole ( ) Flutrimazole ( ) Combinations ( ) Miconazole in combination ( ) Bifonazole in combination ( ) Bromochlorosalicylanilide ( ) Methylrosaniline ( ) Tribromometacresol ( ) Undecylenic acid ( ) Polynoxylin ( ) 2-(4-chlorophenoxy)-ethanol ( ) Chlorphenesin ( ) Tiklaton ( ) Sulbenthine ( ) Ethyl hydroxybenzoate ( ) Haloprogin ( ) Salicylic acid ( ) Selenium sulfide ( ) Ciclopirox ( ) Terbinafine ( ) Amorolfine ( ) Dimazole ( ) Tolnaftat ( ) Tolcyclate ( ) Combinations ( )

Excerpt characterizing sodium thiosulfate

In April, Rostov was on duty. At 8 o'clock in the morning, returning home after a sleepless night, he ordered the heat to be brought, changed his rain-wet clothes, prayed to God, drank tea, warmed up, put things in order in his corner and on the table, and with a weather-beaten, burning face, wearing only a shirt, he lay on his back with his hands under his head. He pleasantly thought that one of these days he should receive his next rank for the last reconnaissance, and expected Denisov to go somewhere. Rostov wanted to talk to him.
Behind the hut, Denisov’s rolling cry was heard, obviously getting excited. Rostov moved to the window to see who he was dealing with and saw Sergeant Topcheenko.
“I told you not to let them burn this fire, some kind of machine!” Denisov shouted. “After all, I saw it myself, Lazag” was dragging the chuk from the field.
“I ordered, your honor, they didn’t listen,” answered the sergeant.
Rostov lay down on his bed again and thought with pleasure: “Let him fuss and fuss now, I’ve finished my job and I’m lying down - great!” From behind the wall he heard that, in addition to the sergeant, Lavrushka, that lively rogue lackey of Denisov, was also speaking. Lavrushka told something about some carts, crackers and bulls, which he saw while going for provisions.
Behind the booth, Denisov’s scream was heard again, retreating, and the words: “Saddle up! Second platoon!
“Where are they going?” thought Rostov.
Five minutes later, Denisov entered the booth, climbed onto the bed with dirty feet, angrily smoked a pipe, scattered all his things, put on a whip and a saber and began to leave the dugout. To Rostov’s question, where? he answered angrily and vaguely that there was a matter.
- God and the great sovereign judge me there! - Denisov said, leaving; and Rostov heard the feet of several horses splashing in the mud behind the booth. Rostov didn’t even bother to find out where Denisov went. Having warmed himself up in his coal, he fell asleep and just left the booth in the evening. Denisov has not returned yet. The evening cleared up; Near the neighboring dugout, two officers and a cadet were playing pile, laughingly planting radishes in the loose, dirty soil. Rostov joined them. In the middle of the game, the officers saw carts approaching them: about 15 hussars on thin horses followed them. The carts, escorted by the hussars, drove up to the hitching posts, and a crowd of hussars surrounded them.
“Well, Denisov kept grieving,” said Rostov, “and now the provisions have arrived.”
- And then! - said the officers. - Those are very welcome soldiers! - Denisov rode a little behind the hussars, accompanied by two infantry officers with whom he was talking about something. Rostov went to meet him halfway.
“I’m warning you, captain,” said one of the officers, thin, small in stature and apparently embittered.
“After all, I said that I wouldn’t give it back,” Denisov answered.
- You will answer, captain, this is a riot - take away the transports from your own! We didn't eat for two days.
“But mine didn’t eat for two weeks,” answered Denisov.
- This is robbery, answer me, my dear sir! – the infantry officer repeated, raising his voice.
- Why are you pestering me? A? - Denisov shouted, suddenly getting excited, - I will answer, not you, and you don’t buzz around here while you’re still alive. March! – he shouted at the officers.
- Good! - without timidity and without moving away, the little officer shouted, - to rob, so I tell you...
“To chog” that march at a fast pace, while he’s still intact.” And Denisov turned his horse towards the officer.
“Okay, okay,” the officer said with a threat, and, turning his horse, he rode away at a trot, shaking in the saddle.
“A dog is in trouble, a living dog is in trouble,” Denisov said after him - the highest mockery of a cavalryman at a mounted infantryman, and, approaching Rostov, he burst out laughing.
– He recaptured the infantry, recaptured the transport by force! - he said. - Well, shouldn’t people die of hunger?
The carts that approached the hussars were assigned to an infantry regiment, but, having been informed through Lavrushka that this transport was coming alone, Denisov and the hussars repulsed it by force. The soldiers were given plenty of crackers, even shared with other squadrons.
The next day, the regimental commander called Denisov to him and told him, covering his eyes with open fingers: “I look at it like this, I don’t know anything and I won’t start anything; but I advise you to go to headquarters and there, in the provisions department, settle this matter, and, if possible, sign that you received so much food; otherwise, the demand is written down on the infantry regiment: the matter will arise and may end badly.”
Denisov went directly from the regimental commander to headquarters, with a sincere desire to carry out his advice. In the evening he returned to his dugout in a position in which Rostov had never seen his friend before. Denisov could not speak and was choking. When Rostov asked him what was wrong with him, he only uttered incomprehensible curses and threats in a hoarse and weak voice...
Frightened by Denisov's situation, Rostov asked him to undress, drink water and sent for a doctor.
- Try me for crime - oh! Give me some more water - let them judge, but I will, I will always beat the scoundrels, and I will tell the sovereign. Give me some ice,” he said.
The regimental doctor who came said that it was necessary to bleed. A deep plate of black blood came out of Denisov’s shaggy hand, and only then was he able to tell everything that happened to him.
“I’m coming,” Denisov said. - “Well, where is your boss here?” Shown. Would you like to wait? “I have work, I came 30 miles away, I don’t have time to wait, report.” Okay, this chief thief comes out: he also decided to teach me: This is robbery! - “Robbery, I say, is committed not by the one who takes provisions to feed his soldiers, but by the one who takes it to put it in his pocket!” So would you like to remain silent? "Fine". Sign, he says, with the commission agent, and your case will be handed over to the command. I come to the commission agent. I enter - at the table... Who?! No, just think!...Who is starving us, - Denisov shouted, hitting the table with the fist of his sore hand, so hard that the table almost fell and the glasses jumped on it, - Telyanin! “What, are you starving us?!” Once, once in the face, deftly it was necessary... “Ah... with this and that and... began to roll. But I was amused, I can say,” Denisov shouted, baring his white teeth joyfully and angrily from under his black mustache. “I would have killed him if they hadn’t taken him away.”

Introduction

One of the fairly well-known chemicals is sodium thiosulfate. Previously, every photographer and amateur photographer knew about him. But even today, sodium thiosulfate is quite widely used in the mining industry, in veterinary medicine and medicine, and in photography.

But, despite its fairly widespread use, its properties and the properties of thiosulfates are poorly known. We are more familiar with the properties of sulfates and sulfides, sulfites. Although in terms of production tonnage, sodium thiosulfate is slightly inferior to sodium sulfite or sodium sulfide.

The task of this course work will be to consider the properties and applications of sodium thiosulfate. We will also try to consider the production of sodium thiosulfate and highlight the most promising methods for its production, taking into account the environmental friendliness and cost-effectiveness of the method.

Properties of sodium thiosulfate

In this chapter we will look at the general properties of sodium thiosulfate, the history of its discovery and, most importantly, the structure of its molecule, since it is the structure of the molecule that significantly affects the chemical and physical properties of the substance.

General properties of sodium thiosulfate

Sodium thiosulfate (sodium hyposulfite) is the disodium salt of thiosulfuric (sulphurous) acid.

In appearance they are colorless crystals. The crystalline form is monoclinic. Sodium thiosulfate is stable in air up to 80°C; when heated in a vacuum at 300°C, it decomposes into sodium sulfite and sulfur. Let's dissolve well in water. At 11 - 48°C it crystallizes from water in the form of pentahydrate. In addition to sodium thiosulfate pentahydrate, we also know sodium thiosulfate decahydrate, which has the formula: . Crystalline hydrates of a different molecular formula for sodium thiosulfate were not found.

Sodium thiosulfate exhibits reducing properties. The molar mass of the substance is: . The molar mass of sodium thiosulfate pentahydrate is 248.17 g/mol.

Density

The solubility in 100 grams of cold water is 66.7 g, and in hot water 266 grams of sodium thiosulfate is soluble in ammonia, aqueous solutions, and slightly soluble in alcohols (ethanol).

At 48.5°C it melts in its water of crystallization and dehydrates at about 100°C.

atria thiosulfate Natrii thiosulfas

Na 2 S 2 0 3 -5H 2 0 M. m. 248.17

Sodium thiosulfate is not a natural product; it is obtained synthetically.

In industry, sodium thiosulfate is obtained from gas production waste. This method, despite its multi-stage nature, is economically profitable, since the raw materials are gas production waste and, in particular, illuminating gas formed during the coking of coal.

Illuminating gas always contains an admixture of hydrogen sulfide, which is captured by absorbers, for example calcium hydroxide. This produces calcium sulfide.

But calcium sulfide undergoes hydrolysis during the production process, so the reaction proceeds somewhat differently - with the formation of calcium hydrosulfide.

When oxidized by atmospheric oxygen, calcium hydrosulfide forms calcium thiosulfate.

When the resulting calcium thiosulfate is fused with sodium sulfate or sodium carbonate, sodium thiosulfate Na 2 S 2 0 3 is obtained.

After evaporation of the solution, sodium thio-sulfate crystallizes, which is a pharmacopoeial drug.

In appearance, sodium thiosulfate (II) is colorless transparent crystals with a salty-bitter taste. Very easily soluble in water. At a temperature of 50 °C it melts in its water of crystallization. Its structure is a salt of thiosulfuric acid (I).

As can be seen from the formula of these compounds, the degree of oxidation of the sulfur atoms in their molecules is different. One sulfur atom has an oxidation state of +6, the other -2. The presence of sulfur atoms in various oxidation states determines their properties.

Thus, having S 2- in the molecule, sodium thiosulfate exhibits reducing ability.

Like thiosulfuric acid itself, its salts are not strong compounds and easily decompose under the influence of acids, even such weak ones as carbonic acid.

This property of sodium thiosulfate to decompose by acids to release sulfur is used to identify the drug. When adding hydrochloric acid to a solution of sodium thiosulfate, cloudiness of the solution is observed due to the release of sulfur.

Very characteristic of sodium thiosulfate is its reaction with a solution of silver nitrate. This produces a white precipitate (silver thiosulfate), which quickly turns yellow. When standing under the influence of air moisture, the sediment turns black due to the release of silver sulfide.

If, when sodium thiosulfate is exposed to silver nitrate, a black precipitate immediately forms, this indicates contamination of the drug with sulfides, which, when interacting with silver nitrate, immediately release a precipitate of silver sulfide.

A pure preparation does not immediately darken when exposed to a solution of silver nitrate.

As an authenticity reaction, the reaction of sodium thiosulfate with a solution of iron (III) chloride can also be used. In this case, iron oxide thiosulfate is formed, colored violet. The color quickly disappears due to the reduction of this salt to colorless ferrous iron salts (FeS 2 0 3 and FeS 4 0 6).

When interacting with sodium iodine, sodium thiosulfate acts as a reducing agent. Taking electrons from S 2-, iodine is reduced to I-, and sodium thiosulfate is oxidized by iodine to sodium tetrathioiate.

Chlorine is similarly reduced to hydrogen chloride.

When there is an excess of chlorine, the released sulfur is oxidized to sulfuric acid.

The use of sodium thiosulfate to absorb chlorine in the first gas masks was based on this reaction.

The preparation is not allowed to contain impurities of arsenic, selenium, carbonates, sulfates, sulfides, sulfites, calcium salts.

GF X allows the presence of impurities of chlorides and heavy metal salts within the standard.

Quantitative determination of sodium thiosulfate is carried out using the iodometric method, which is based on the reaction of its interaction with iodine. GF requires a sodium thio-sulfate content in the preparation of no less than 99% and no more than 102% (due to the permissible limit of weathering of the preparation).

The use of sodium thiosulfate is based on its ability to release sulfur. The drug is used as an antidote for poisoning with halogens, cyanogen and hydrocyanic acid.

The resulting potassium thiocyanate is much less toxic than potassium cyanide. Therefore, in case of poisoning with hydrocyanic acid or its salts, sodium thiosulfate should be used as first aid. The drug can also be used for poisoning with arsenic, mercury, and lead compounds; in this case, non-toxic sulfides are formed.

Sodium thiosulfate is also used for allergic diseases, arthritis, neuralgia - intravenously in the form of a 30% aqueous solution. In this regard, GF X provides a 30% solution of sodium thiosulfate for injection (Solutio Natrii thiosulfatis 30% pro injectionibus).

Available in powders and ampoules of 5, 10, 50 ml of 30% solution.

Sodium thiosulfate contains water of crystallization, which easily evaporates, so it should be stored in a cool place, in well-sealed dark glass bottles, since light promotes its decomposition. Solutions become cloudy when standing due to the sulfur released. This process is accelerated in the presence of carbon dioxide. Therefore, flasks or bottles with sodium thiosulfate solutions are equipped with a calcium chloride tube filled with soda lime, which absorbs it.

Medicines elements VI And IV groups of the periodic table of elements.

ANALYSIS OF SULFUR COMPOUNDS. 6 GROUP PSE.

Sulfur in the human body is found in the epidermis, skin, muscles, pancreas, and hair. It is part of some amino acids (methionine, cysteine), peptides that participate in the processes of tissue respiration and catalyze enzymatic processes.

In medicine, sulfur itself is used in the form of ointments and sodium thiosulfate.

Sodium thiosulfate Natrii thiosulfas (ln)

Na 2 S 2 O 3 5 H 2 OSodiumthiosulfate (MHH)

Sodium salt of thiosulfuric acid

Structural formula:

Sulfur atoms have different oxidation states. Due to S 2 - LB exhibit restorative properties.

Receipt

When heating sodium sulfite and sulfur ( was first obtained in 1799):

Na 2 SO 3 +S→Na 2 S 2 O 3

Oxidation of sodium sulfide with sulfur dioxide:

2Na 2 S+ 3S0 2 → 2Na 2 S 2 0 3 +S↓

Currently it is obtained using gas production waste containing hydrogen sulfide. The method, despite its multi-stage nature, is economically beneficial:

Hydrogen sulfide is captured by an absorber - calcium hydroxide:

Ca(OH) 2 + H 2 S → CaS + 2H 2 S

however, due to the hydrolysis of calcium sulfide, the following reactions occur:

CaS + 2H 2 O → Ca(OH) 2 + H 2 S

2Ca(OH) 2 + 3H 2 S → CaS + Ca(SH) 2 + 4H 2 O

Calcium hydrosulfide is oxidized by atmospheric oxygen to calcium thiosulfate:

Ca(SH) 2 + 2O 2 → CaS 2 O 3 + H 2 O

calcium thiosulfate is fused with calcium carbonate:

CaS 2 O 3 + Na 2 CO 3 → Na 2 S 2 O 3 + CaCO 3 ↓

Description and solubility

Colorless transparent crystals, odorless. In warm, dry air it loses crystallization water (eroses). In humid air it spreads (turns into a liquid state). At a temperature of +50 0 C it melts in water of crystallization.

Very easily soluble in water, practically insoluble in alcohol.

Chemical properties

As can be seen from the formula, the oxidation state of sulfur is different (6+ and 2-). Having S 2- in the molecule, the drug exhibits restorative properties.

Sodium thiosulfate, like thiosulfuric acid, of which it is a salt, is not a strong compound, easily decompose under the influence of acids, even carbonic (air humidity + carbon dioxide):

Na 2 S 2 O 3 + CO 2 + H 2 O → Na 2 CO 3 + H 2 S 2 O 3

H 2 S 2 O 3 → S↓ + SO 2 + H 2 O

yellow smell

sediment (turbidity)

This property is used in authenticity reactions:

Authenticity

Reactions to sodium ion(see catine anions).

Reactions to thiosulfate ion:

Decomposition reaction with diluted hydrochloric acid when dilute hydrochloric acid is added to a solution of the drug, the solution gradually becomes cloudy - free sulfur is released (unlike salts of sulfurous acid), then the specific smell of sulfur dioxide SO 2 appears:

Na 2 S 2 O 3 + 2HCl → 2NaCl + SO 2 + S↓+ H 2 O

smell yellow

sediment (turbidity)

S 2 O 3 2- + H 2 O - 2ē → 2SO 2 + 2H +

S 2 O 3 2- + 6H + + 4ē → 2S↓ + 3H 2 O

Reaction with silver nitrate solution.

When an excess of silver nitrate is added, a white precipitate is released, which quickly turns yellow, when standing it turns brown and finally turns black due to the formation of silver sulfide

First, a white precipitate of silver thiosulfate is formed:

Na 2 S 2 O 3 + 2AgN0 3 → Ag 2 S 2 O 3 ↓ + 2NaN0 3

Silver thiosulfate quickly decomposes (intramolecular redox reaction), silver sulfite and sulfur are formed (yellow precipitate):

Ag 2 S 2 O 3 → Ag 2 SO 3 ↓ + S↓

When standing, a black precipitate of silver sulfide is formed:

Ag 2 SO 3 + S + H 2 O → Ag 2 S↓ + H 2 SO 4

If the reaction procedure is changed - adding sodium thiosulfate to a solution of silver nitrate, then the white precipitate of silver thiosulfate dissolves in excess sodium thiosulfate:

Ag 2 S 2 O 3 + 3Na 2 S 2 O 3 → 2Na 3

Thermally very unstable:

In the presence of sulfuric acid it decomposes:

Reacts with alkalis:

Reacts with halogens:

Thiosulfuric acid

If you boil an aqueous solution of sodium sulfite with sulfur and, after filtering out the excess sulfur, leave to cool, then colorless transparent crystals of a new substance are released from the solution, the composition of which is expressed by the formula. This substance is the sodium salt of thiosulfuric acid.

Thiosulfuric acid is unstable. Already at room temperature it disintegrates. Its salts, thiosulfates, are much more stable. Of these, the most commonly used is sodium thiosulfate, also known incorrectly as hyposulfite.

When some acid, such as hydrochloric acid, is added to a solution of sodium thiosulfate, the smell of sulfur dioxide appears and after a while the liquid becomes cloudy from the released sulfur.

The study of the properties of sodium thiosulfate leads to the conclusion that the sulfur atoms included in its composition have different oxidation levels: one of them has an oxidation state of +4, the other has 0 . Sodium thiosulfate - reducing agent . Chlorine, bromine and other strong oxidizing agents oxidize it to sulfuric acid or its salt.

Thiosulf? are you- salts and esters of thiosulfuric acid, H2S2O3. Thiosulfates are unstable and therefore do not occur in nature. The most widely used are sodium thiosulfate and ammonium thiosulfate.

Structure. Structure of the thiosulfate ion

Thiosulfate ion is close in structure to sulfate ion. In the 2− tetrahedron, the S-S bond (1.97A) is longer than the S-O bonds

Sodium thiosulfate can be classified as rather unstable substances. Sodium thiosulfate decomposes when heated to 220°C: In the reaction of thermal decomposition of sodium thiosulfate, we obtain sodium polysulfide, which also further decomposes into sodium sulfide and elemental sulfur. Interaction with acids: it is impossible to isolate thiosulfuric acid (hydrogen thiosulfate) by the reaction of sodium thiosulfate with a strong acid, since it is unstable and immediately decomposes: Hydrochloric and nitric acids will also undergo the same reaction. Decomposition is accompanied by a discharge that has an unpleasant odor.

Redox properties of sodium thiosulfate: due to the presence of sulfur atoms with an oxidation state of 0, the thiosulfate ion has reducing properties, for example, with weak oxidizing agents (I2, Fe3+), thiosulfate is oxidized to the tetrathionate ion: In an alkaline environment, the oxidation of sodium thiosulfate with iodine can proceed to sulfate.

And stronger oxidizing agents oxidize it to sulfate ion :

Strong reducing agents ion is reduced to S2- derivatives: Depending on conditions, sodium thiosulfate can exhibit both oxidizing and reducing properties.

Complexing properties of thiosulfates:

Thiosulfate ion is a strong complexing agent , used in photography to remove unreduced silver bromide from photographic film: The S2O32 ion is coordinated by metals through a sulfur atom, so thiosulfate complexes are easily converted into the corresponding sulfides.

Applications of sodium thiosulfate

Sodium thiosulfate is quite widely used both in everyday life and in industry. The main areas of application of sodium thiosulfate will be medicine, textile and mining industries, photography.

Sodium thiosulfate is used in the textile and paper industries to remove traces of chlorine after bleaching fabrics and paper, and in leather production it is used as a chromic acid reducer.

In the mining industry, sodium thiosulfate is used to extract silver from ores with low silver concentrations. Complex compounds of silver with thiosulfates are quite stable, at least more stable than complex compounds with fluorine, chlorine, bromides, and thiocyanates. Therefore, the isolation of silver in the form of a soluble complex compound of the composition or is industrially profitable. Work is underway on its use in gold extraction. But in this case, the instability constant of the complex compound is much higher and the complexes are less stable compared to silver ones.

The first use of sodium thiosulfate was in medicine. And to this day it has not lost its importance in medicine. True, other, more effective drugs have already been found for the treatment of many diseases, so sodium thiosulfate has begun to be used more widely in veterinary medicine. Sodium thiosulfate is used in medicine as an antidote for poisoning with arsenic, mercury and other heavy metals, cyanides (translates them into thiocyanates):

As mentioned above, the thiosulfate ion creates stable complex compounds with many metals, including many toxic heavy metals. The created complex compounds are low-toxic and are excreted from the body. This feature of sodium thiosulfate is the basis for its use in toxicology and treatment of poisoning.

Sodium thiosulfate is also used to disinfect the intestines in case of food poisoning, to treat scabies (together with hydrochloric acid), as an anti-inflammatory and anti-burn agent.

Sodium thiosulfate is widely used in analytical chemistry because it is a reagent in iodometry. Iodometry is one of the methods for quantitatively determining the concentrations of substances and to determine the concentration of iodine, a redox reaction with sodium thiosulfate is used:

A final fairly common use of sodium thiosulfate is its use as a fixative in photography. And although ordinary black and white photography has already given way to color and ordinary photographic film is used quite rarely, in many ways inferior to digital image capture, there are quite a few places where photographic plates and photographic film are still used. Examples include X-ray machines, both medical and industrial, scientific equipment, and phototelescopes.

In order for us to obtain a photographic image, it is enough for about 25% of the silver bromide in the photographic film to develop. And the rest of it remains in the photographic film and retains its photosensitivity. If the photographic film is taken out into the light after development, then the undeveloped halogen silver that remains in it will be developed by the developer and the negative will darken. Even if all the developer is washed out, the negative will somehow darken in the light due to the decomposition of the silver halide.

To preserve the image on the film, the undeveloped silver halogen must be removed from it. To do this, an image fixation process is used, during which silver halides are converted into soluble compounds and washed out of the film or photograph. Sodium thiosulfate is used to fix the image.

Depending on the concentration of sodium thiosulfate in the solution, various compounds are formed. If the fixer solution contains a small amount of thiosulfate, then the reaction proceeds according to the equation:

The resulting silver thiosulfate is insoluble in water, so it is difficult to isolate it from the photo layer; it is quite unstable and decomposes with the release of sulfuric acid:

Silver sulfide blackens the image and cannot be removed from the photo layer.

If there is excess sodium thiosulfate in the solution, complex silver salts will form:

The resulting complex salt, sodium thiosulfate argentate, is quite stable, but poorly soluble in water.

When there is a large excess of thiosulfates in the solution, complex silver complex salts that are highly soluble in water are formed:

These properties of sodium thiosulfate are the basis for its use as a fixative in photography.

Tetratnopic acid belongs to the group of polynoid acids. These are dibasic acids of a general formula, where they can take values ​​from 2 to 6, and possibly more. Polnithionic acids unstable and known only in aqueous solutions. Salts of polythioic acids—polythionates—are more stable; some of them are obtained in the form of crystals.

Polythionic acids - sulfur compounds with the general formula H2SnO6, where n>=2. Their salts are called polythionates.

Tetrathionate ion can be obtained by oxidation of the thiosulfate ion with iodine (the reaction is used in iodometry):

Pentationate ion obtained by the action of SCl2 on the thiosulfate ion and from Wackenroder's liquid by adding potassium acetate to it. First, prismatic crystals of potassium tetrathionate fall out, then plate-like crystals of potassium pentathionate, from which an aqueous solution of pentathionic acid is obtained by the action of tartaric acid.

Potassium hexathionate K2S6O6 best synthesized by the action of KNO2 on K2S2O3 in concentrated HCl at low temperatures.