Interaction of oxides and hydroxy compounds with water. Advances of modern natural science Reaction of sulfur and iron compounds

§ 1 Oxide and its characteristics

While studying the chemical properties of oxygen, we became familiar with oxidation reactions and oxides. Oxides, for example, include substances with the following formulas: Na2O, CuO, Al2O3, SiO2, P2O5, SO3, Mn2O7.

So, all oxides are characterized in composition by three common features: any oxide is a complex substance, consisting of atoms of two chemical elements, one of the elements is oxygen.

All these characteristics can be expressed by the general formula ExOy, in which E are the atoms of the chemical element that formed the oxide, O are the oxygen atoms; x, y are indices indicating the number of atoms of the elements forming the oxide.

There are a lot of oxides. Almost all simple substances form oxides when oxidized. Atoms of many elements, exhibiting different valence values, participate in the formation of several oxides, for example, nitrogen corresponds to five oxides: nitrogen oxide (I) N2O, nitrogen oxide (II) NO, nitrogen oxide (III) N2O3, nitrogen oxide (IV) NO2, Nitric oxide (V) N2O5.

§ 2 Properties of oxides and their classification

Let's get acquainted with the properties of some oxides.

Carbon monoxide (IV) is a colorless, odorless gas with a slightly sour taste, turning into a solid white snow-like substance, bypassing the liquid state at - 780C, soluble in water.

Hydrogen oxide is water, under normal conditions it is a colorless liquid with a boiling point of 1000C.

Calcium oxide is a white solid whose melting point is 26270C; when mixed with water, it actively interacts with it.

Iron (III) oxide is a red-brown solid that melts at 15620C and is insoluble in water.

Let's pass carbon monoxide (IV) through water and add a few drops of litmus to the resulting solution. Litmus will change color from blue to red, therefore, when carbon monoxide (IV) reacts with water, an acid is formed. The reaction equation is as follows: CO2 + H2O → H2CO3. As a result of the reaction, carbonic acid was formed. Similarly, oxides of other nonmetals interact with water to form acids. Therefore, non-metal oxides are called acidic. Metal oxides with a valency greater than IV are also considered acidic, for example, vanadium (V) oxide V2O5, chromium (VI) oxide CrO3, manganese (VII) oxide Mn2O7.

Place some white calcium oxide powder in a test tube with water and add a few drops of phenolphthalein to the resulting slightly cloudy solution. Phenolphthalein changes color from colorless to crimson, which indicates the appearance of a base in the test tube. CaO + H2O → Ca(OH)2. As a result of the reaction, a base was formed - calcium hydroxide. Metal oxides whose valency is not more than III are called basic.

Metals exhibiting valence III and IV, and sometimes II, form amphoteric oxides. These oxides differ from others in their chemical properties. We will get to know them in more detail later, but for now we will focus our attention on acidic and basic oxides.

§ 3 Dissolution of oxides in water

Many acids and bases can be prepared by dissolving the corresponding oxides in water.

The dissolution of oxides in water is a chemical process accompanied by the formation of new chemical compounds - acids and bases.

For example, when sulfur oxide (VI) is dissolved in water, sulfuric acid is formed: SO3 + H2O → H2SO4. And when phosphorus oxide (V) is dissolved, phosphoric acid is formed: P2O5 + 3H2O → 2H3PO4. When sodium oxide is dissolved, a base is formed - sodium hydroxide: Na2O + H2O → 2NaOH, when barium oxide is dissolved - barium hydroxide: BaO + H2O → Ba(OH)2.

The names of oxide groups reflect their relationship with other classes of inorganic compounds: most acidic oxides correspond to acids, and almost all basic oxides correspond to bases.

However, not all oxides are soluble. Thus, most basic oxides are insoluble, and the only exceptions are oxides formed by elements of the main subgroups of the first and second groups of the periodic table of elements.

Most acid oxides, on the contrary, are soluble in water. The exception here is, for example, silicon (IV) oxide - SiO2. This substance is well known to everyone. Silicon oxide forms the basis of river sand and many minerals, including rare and very beautiful ones: rock crystal, amethyst, citrine, jasper. Many acidic oxides formed by metals are slightly soluble or insoluble.

If the oxides do not dissolve in water, then the corresponding acids and bases are obtained in other ways (indirectly), which we will get acquainted with later.

List of used literature:

1. NOT. Kuznetsova. Chemistry. 8th grade. Textbook for general education institutions. – M. Ventana-Graf, 2012.

Oxides are called complex substances whose molecules include oxygen atoms in oxidation state - 2 and some other element.

can be obtained through the direct interaction of oxygen with another element, or indirectly (for example, during the decomposition of salts, bases, acids). Under normal conditions, oxides come in solid, liquid and gaseous states; this type of compound is very common in nature. Oxides are found in the Earth's crust. Rust, sand, water, carbon dioxide are oxides.

They are either salt-forming or non-salt-forming.

Salt-forming oxides- these are oxides that, as a result, chemical reactions form salts. These are oxides of metals and non-metals, which, when interacting with water, form the corresponding acids, and when interacting with bases, the corresponding acidic and normal salts. For example, Copper oxide (CuO) is a salt-forming oxide, because, for example, when it reacts with hydrochloric acid (HCl), a salt is formed:

CuO + 2HCl → CuCl 2 + H 2 O.

As a result of chemical reactions, other salts can be obtained:

CuO + SO 3 → CuSO 4.

Non-salt-forming oxides These are oxides that do not form salts. Examples include CO, N 2 O, NO.

Salt-forming oxides, in turn, are of 3 types: basic (from the word « base » ), acidic and amphoteric.

Basic oxides These metal oxides are called those that correspond to hydroxides belonging to the class of bases. Basic oxides include, for example, Na 2 O, K 2 O, MgO, CaO, etc.

Chemical properties basic oxides

1. Water-soluble basic oxides react with water to form bases:

Na 2 O + H 2 O → 2NaOH.

2. React with acid oxides, forming the corresponding salts

Na 2 O + SO 3 → Na 2 SO 4.

3. React with acids to form salt and water:

CuO + H 2 SO 4 → CuSO 4 + H 2 O.

4. React with amphoteric oxides:

Li 2 O + Al 2 O 3 → 2LiAlO 2.

If the composition of the oxides contains a non-metal or a metal exhibiting the highest valence (usually from IV to VII) as the second element, then such oxides will be acidic. Acidic oxides (acid anhydrides) are those oxides that correspond to hydroxides belonging to the class of acids. These are, for example, CO 2, SO 3, P 2 O 5, N 2 O 3, Cl 2 O 5, Mn 2 O 7, etc. Acidic oxides dissolve in water and alkalis, forming salt and water.

Chemical properties of acid oxides

1. React with water to form an acid:

SO 3 + H 2 O → H 2 SO 4.

But not all acidic oxides react directly with water (SiO 2, etc.).

2. React with based oxides to form a salt:

CO 2 + CaO → CaCO 3

3. React with alkalis, forming salt and water:

CO 2 + Ba(OH) 2 → BaCO 3 + H 2 O.

Part amphoteric oxide includes an element that has amphoteric properties. Amphotericity refers to the ability of compounds to exhibit acidic and basic properties depending on conditions. For example, zinc oxide ZnO can be either a base or an acid (Zn(OH) 2 and H 2 ZnO 2). Amphotericity is expressed in the fact that, depending on the conditions, amphoteric oxides exhibit either basic or acidic properties.

Chemical properties of amphoteric oxides

1. React with acids to form salt and water:

ZnO + 2HCl → ZnCl 2 + H 2 O.

2. React with solid alkalis (during fusion), forming as a result of the reaction salt - sodium zincate and water:

ZnO + 2NaOH → Na 2 ZnO 2 + H 2 O.

When zinc oxide interacts with an alkali solution (the same NaOH), another reaction occurs:

ZnO + 2 NaOH + H 2 O => Na 2.

Coordination number is a characteristic that determines the number of nearby particles: atoms or ions in a molecule or crystal. Each amphoteric metal has its own coordination number. For Be and Zn it is 4; For and Al it is 4 or 6; For and Cr it is 6 or (very rarely) 4;

Amphoteric oxides are usually insoluble in water and do not react with it.

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The invention relates to methods for dissolving uranium oxides and can be used in the technology of producing fuel cycle materials, in particular for producing enriched uranium. According to the method, uranium oxide powder is placed under a layer of water with a ratio of the height of the water layer and the height of the uranium oxide layer of at least 1.3. Nitric acid is supplied under the layer of uranium oxides at a flow rate of (0.30-0.36) ton HNO 3 per 1 ton of uranium per hour. The invention makes it possible to reduce the volume of gases leaving the solvent reactor and subject to purification before being discharged into the atmosphere, while reducing the nitrogen dioxide content in them. 1 salary files, 1 table.

The invention relates to methods for dissolving uranium oxides and can be used in the technology of producing fuel cycle materials, in particular for producing enriched uranium. As a feedstock for uranium enrichment, its oxides in the form of technical oxide - U 3 O 8 oxide (2UO 3 + UO 2), obtained from natural raw materials, can be used. In this case, before the fluoridation operation, uranium must be further purified from accompanying impurities present in the ore concentrate, including impurities that form volatile fluorides (molybdenum, silicon, iron, vanadium, etc.). In addition, it is necessary to remove impurities that enter uranium during the processing of natural ores into uranium oxide (scale, under-burned residue, graphite, coal, etc.). To purify uranium from impurities, you can use extraction technology for purifying uranium nitrate solutions using tributyl phosphate. Before extraction, uranium oxides must be dissolved. There is a known method for dissolving uranium oxides in a mixture of concentrated nitric and concentrated hydrochloric acids (Uranium and its compounds. USSR industry standard OST 95175-90, p. 5). However, due to the high corrosion of equipment, this method is used only on a laboratory scale. There is a known method for dissolving uranium oxide in nitric acid (V.M. Vdovenko. Modern radiochemistry. - M., 1969, p. 257) (prototype). The method is carried out according to the following reaction: 2U 3 O 8 +14HNO 3 =6UO 2 (NO) 3)2+7H 2 O+NO+NO 2. As a result of the reaction, nitrogen oxide and dioxide are formed, which have a harmful effect on environment and man. In this regard, there is a need to purify waste gases from nitrogen oxides. Nitrogen dioxide (NO 2) is a brown gas, nitrogen oxide (NO) is a colorless gas. Nitrogen oxide (NO) upon contact with atmospheric oxygen is oxidized to NO 2. Nitrogen dioxide is the main component in gas discharges subject to treatment. If a raw material containing over 80% tetravalent uranium oxide is dissolved, then the formation of nitrogen oxides per unit of raw material increases compared to the dissolution of uranium oxide containing approximately 30% tetravalent uranium oxide. The dissolution process of such raw materials is characterized by a significant release of nitrogen dioxide. In oxide raw materials, the uranium (IV) content is 30%: In oxide raw materials, the uranium (IV) content is 80%: When stirring the reaction system, which is used to improve mass transfer in the system, the release of nitrogen oxides from the reaction mixture occurs especially rapidly. The objective of the invention is to reduce the volume of gases (nitrogen oxides) leaving the solvent reactor and subject to purification before being discharged into the atmosphere, while reducing the nitrogen dioxide content in them. The problem is solved by the fact that in the method of dissolving uranium oxides, including their interaction with nitric acid, the uranium oxide powder is placed under a layer of water at a ratio of the height of the water layer and the height of the uranium oxide layer of at least 1.3, and nitric acid is supplied under the layer of uranium oxides at a rate of (0.3-0.36) t HNO 3 per 1 ton of uranium per hour. The reaction mixture is irrigated with water in an amount equal to 10-20% of the aqueous layer. Example. The uranium oxide powder is placed under a layer of water. The acid solution is fed under the oxide layer. The acid solution is supplied under the layer of uranium oxides through a pipe lowered to the bottom of the solvent reactor. Four series of experiments are carried out. In the first series, the ratio of the height of the water layer to the height of the uranium oxide layer is changed. In the second series of experiments, the consumption of HNO 3 per unit time is changed. In the third series of experiments, the reaction mixture is stirred by supplying compressed air. In the fourth series of experiments, water is sprayed over the surface of the water layer to create a water mist in the solvent reactor. In experiment 6 of the first series, there is no layer of water above the layer of uranium oxides. Experiments are carried out without heating the reaction mixture. The experimental results are presented in the table. When nitric acid is supplied under a layer of uranium oxides located under water, the dissolution of uranium oxides occurs evenly throughout the entire volume. Nitrogen dioxide formed during the dissolution of uranium oxides, passing through a layer of water, interacts with the latter to form nitric acid, which, in turn, interacts with uranium oxides; the consumption of nitric acid (in total per experiment) supplied to the solvent reactor is reduced. As can be seen from the table, a decrease in the volume of gases leaving the solvent reactor, with a decrease in the nitrogen dioxide content in them, occurs when the ratio of the height of the water layer to the height of the uranium oxide layer is at least 1.3 and the consumption of nitric acid per unit time is 0.30- 0.36 t НNO 3 / t U per hour (experiments 3-5 of the first series, 1, 2 of the second series). Irrigation of the space above the water layer with water promotes additional capture of nitrogen dioxide and suppression of foam formation (experiments 1, 2 of the fourth series). The absence of a water layer above the uranium oxides during the dissolution process (experiment 6 of the first series) or its insufficient height (the ratio of the height of the water layer to the height of the uranium oxide layer is less than 1, 3, experiments 1, 2 of the first series) leads to an increase in gas release from the solvent reactor, in this case, the gas has a brown color characteristic of nitrogen dioxide. An increase in the consumption of nitric acid per unit time (more than 0.36 t НNO 3 / t U per hour) also leads to strong gas evolution; the gas contains a significant amount of brown nitrogen dioxide (experiments 3, 4 of the second series). Stirring the reaction mixture with air increases the total consumption of nitric acid and leads to strong gas evolution (experiments 1, 2 of the third series). The ratio of the height of the water layer to the height of the powder layer, equal to 1.30-1.36, is optimal from the point of view of obtaining a solution suitable in concentration for the subsequent operation in the technology of fuel cycle materials - extraction.

Claim

1. A method for dissolving uranium oxides, including their interaction with nitric acid, characterized in that the uranium oxide powder is placed under a layer of water when the ratio of the height of the water layer and the height of the uranium oxide layer is at least 1.3, and nitric acid is supplied under the layer of uranium oxides at a flow rate (0.300.36) t HNO 3 per 1 ton of uranium per hour. 2. The method according to claim 1, characterized in that the reaction mixture is irrigated with water in an amount equal to 10-20% of the aqueous layer.

Modern chemical science represents many different branches, and each of them, in addition to its theoretical basis, has great applied and practical significance. Whatever you touch, everything around you is a chemical product. The main sections are inorganic and organic chemistry. Let's consider what main classes of substances are classified as inorganic and what properties they have.

Main categories of inorganic compounds

These include the following:

1. Oxides.
2. Salt.
3. Grounds.
4. Acids.

Each of the classes is represented by a wide variety of compounds of inorganic nature and is important in almost any structure of human economic and industrial activity. All the main properties characteristic of these compounds, their occurrence in nature and their production are studied in a school chemistry course without fail, in grades 8-11.

There is a general table of oxides, salts, bases, acids, which presents examples of each substance and their state of aggregation and occurrence in nature. Interactions that describe chemical properties are also shown. However, we will look at each of the classes separately and in more detail.

Group of compounds - oxides

4. Reactions as a result of which elements change CO

Me +n O + C = Me 0 + CO

1. Reagent water: formation of acids (SiO 2 exception)

CO + water = acid

2. Reactions with bases:

CO 2 + 2CsOH = Cs 2 CO 3 + H 2 O

3. Reactions with basic oxides: salt formation

P 2 O 5 + 3MnO = Mn 3 (PO 3) 2

4. OVR reactions:

CO 2 + 2Ca = C + 2CaO,

They exhibit dual properties and interact according to the principle of the acid-base method (with acids, alkalis, basic oxides, acid oxides). They do not interact with water.

1. With acids: formation of salts and water

AO + acid = salt + H 2 O

2. With bases (alkalis): formation of hydroxo complexes

Al 2 O 3 + LiOH + water = Li

3. Reactions with acid oxides: obtaining salts

FeO + SO 2 = FeSO 3

4. Reactions with OO: formation of salts, fusion

MnO + Rb 2 O = double salt Rb 2 MnO 2

5. Fusion reactions with alkalis and alkali metal carbonates: formation of salts

Al 2 O 3 + 2LiOH = 2LiAlO 2 + H 2 O

They do not form either acids or alkalis. They exhibit highly specific properties.

Each higher oxide, formed either by a metal or a non-metal, when dissolved in water, gives a strong acid or alkali.

Organic and inorganic acids

In classical sound (based on the positions of ED - electrolytic dissociation - acids are compounds that in an aqueous environment dissociate into cations H + and anions of acid residues An -. However, today acids have been carefully studied in anhydrous conditions, so there are many different theories for hydroxides.

Empirical formulas of oxides, bases, acids, salts consist only of symbols, elements and indices indicating their quantity in the substance. For example, inorganic acids are expressed by the formula H + acid residue n- . Organic substances have a different theoretical representation. In addition to the empirical one, you can write down a full and abbreviated structural formula for them, which will reflect not only the composition and quantity of the molecule, but also the order of the atoms, their connection with each other and the main functional group for carboxylic acids -COOH.

In inorganics, all acids are divided into two groups:

• oxygen-free - HBr, HCN, HCL and others;
• oxygen-containing (oxoacids) - HClO 3 and everything where there is oxygen.

Inorganic acids are also classified by stability (stable or stable - everything except carbonic and sulfurous, unstable or unstable - carbonic and sulfurous). In terms of strength, acids can be strong: sulfuric, hydrochloric, nitric, perchloric and others, as well as weak: hydrogen sulfide, hypochlorous and others.

Organic chemistry offers not the same variety. Acids that are organic in nature are classified as carboxylic acids. Their common feature is the presence of the -COOH functional group. For example, HCOOH (formic), CH 3 COOH (acetic), C 17 H 35 COOH (stearic) and others.

There are a number of acids that are especially carefully emphasized when considering this topic in a school chemistry course.

1. Solyanaya.
2. Nitrogen.
3. Orthophosphoric.
4. Hydrobromic.
5. Coal.
6. Hydrogen iodide.
7. Sulfuric.
8. Acetic or ethane.
9. Butane or oil.
10. Benzoin.

These 10 acids in chemistry are fundamental substances of the corresponding class both in the school course and in general in industry and syntheses.

Properties of inorganic acids

The main physical properties include, first of all, the different state of aggregation. After all, there are a number of acids that have the form of crystals or powders (boric, orthophosphoric) under normal conditions. The vast majority of well-known inorganic acids represents different liquids. Boiling and melting points also vary.

Acids can cause severe burns because they have the power to destroy organic tissue and skin covering. Indicators are used to detect acids:

• methyl orange (in normal environment - orange, in acids - red),
• litmus (in neutral - violet, in acids - red) or some others.

The most important chemical properties include the ability to interact with both simple and complex substances.

Chemical properties of inorganic acids

What do they interact with?	Example reaction
1. With simple substances - metals. Mandatory condition: the metal must be in the EHRNM before hydrogen, since metals standing after hydrogen are not able to displace it from the composition of acids. The reaction always produces hydrogen gas and salt.
2. With reasons. The result of the reaction is salt and water. Such reactions of strong acids with alkalis are called neutralization reactions.	Any acid (strong) + soluble base = salt and water
3. With amphoteric hydroxides. Bottom line: salt and water.	2HNO 2 + beryllium hydroxide = Be(NO 2) 2 (medium salt) + 2H 2 O
4. With basic oxides. Result: water, salt.	2HCL + FeO = iron (II) chloride + H 2 O
5. With amphoteric oxides. Final effect: salt and water.	2HI + ZnO = ZnI 2 + H 2 O
6. With salts formed by weaker acids. Final effect: salt and weak acid.	2HBr + MgCO 3 = magnesium bromide + H 2 O + CO 2

When interacting with metals, not all acids react equally. Chemistry (9th grade) at school involves a very shallow study of such reactions, however, even at this level the specific properties of concentrated nitric and sulfuric acid when interacting with metals are considered.

Hydroxides: alkalis, amphoteric and insoluble bases

Oxides, salts, bases, acids - all these classes of substances have a common chemical nature, explained by the structure of the crystal lattice, as well as the mutual influence of atoms in the molecules. However, if it was possible to give a very specific definition for oxides, then this is more difficult to do for acids and bases.

Just like acids, bases, according to the ED theory, are substances that can decompose in an aqueous solution into metal cations Me n + and anions of hydroxyl groups OH - .

• Soluble or alkalis (strong bases that change Formed by metals of groups I and II. Example: KOH, NaOH, LiOH (that is, elements of only the main subgroups are taken into account);
• Slightly soluble or insoluble (medium strength, do not change the color of the indicators). Example: magnesium hydroxide, iron (II), (III) and others.
• Molecular (weak bases, in an aqueous environment they reversibly dissociate into ion molecules). Example: N 2 H 4, amines, ammonia.
• Amphoteric hydroxides (show dual basic-acid properties). Example: beryllium, zinc and so on.

Each group presented is studied in the school chemistry course in the “Fundamentals” section. Chemistry in grades 8-9 involves a detailed study of alkalis and poorly soluble compounds.

Main characteristic properties of bases

All alkalis and slightly soluble compounds are found in nature in a solid crystalline state. At the same time, their melting temperatures are usually low, and poorly soluble hydroxides decompose when heated. The color of the bases is different. If alkalis are white, then crystals of poorly soluble and molecular bases can be of very different colors. The solubility of most compounds of this class can be found in the table, which presents the formulas of oxides, bases, acids, salts, and shows their solubility.

Alkalies can change the color of indicators as follows: phenolphthalein - crimson, methyl orange - yellow. This is ensured by the free presence of hydroxo groups in the solution. That is why poorly soluble bases do not give such a reaction.

The chemical properties of each group of bases are different.

Chemical properties
Alkalis	Slightly soluble bases	Amphoteric hydroxides
I. Interact with CO (result - salt and water): 2LiOH + SO 3 = Li 2 SO 4 + water II. Interact with acids (salt and water): ordinary neutralization reactions (see acids) III. They interact with AO to form a hydroxo complex of salt and water: 2NaOH + Me +n O = Na 2 Me +n O 2 + H 2 O, or Na 2 IV. They interact with amphoteric hydroxides to form hydroxo complex salts: The same as with AO, only without water V. React with soluble salts to form insoluble hydroxides and salts: 3CsOH + iron (III) chloride = Fe(OH) 3 + 3CsCl VI. React with zinc and aluminum in an aqueous solution to form salts and hydrogen: 2RbOH + 2Al + water = complex with hydroxide ion 2Rb + 3H 2	I. When heated, they can decompose: insoluble hydroxide = oxide + water II. Reactions with acids (result: salt and water): Fe(OH) 2 + 2HBr = FeBr 2 + water III. Interact with KO: Me +n (OH) n + KO = salt + H 2 O	I. React with acids to form salt and water: (II) + 2HBr = CuBr 2 + water II. React with alkalis: result - salt and water (condition: fusion) Zn(OH) 2 + 2CsOH = salt + 2H 2 O III. React with strong hydroxides: the result is salts if the reaction occurs in an aqueous solution: Cr(OH) 3 + 3RbOH = Rb 3

These are most of the chemical properties that bases exhibit. The chemistry of bases is quite simple and follows the general laws of all inorganic compounds.

Class of inorganic salts. Classification, physical properties

Based on the provisions of the ED, salts can be called inorganic compounds that dissociate in an aqueous solution into metal cations Me +n and anions of acidic residues An n-. This is how you can imagine salts. Chemistry gives more than one definition, but this is the most accurate.

Moreover, according to their chemical nature, all salts are divided into:

• Acidic (containing a hydrogen cation). Example: NaHSO 4.
• Basic (containing a hydroxo group). Example: MgOHNO 3, FeOHCL 2.
• Medium (consist only of a metal cation and an acid residue). Example: NaCL, CaSO 4.
• Double (include two different metal cations). Example: NaAl(SO 4) 3.
• Complex (hydroxo complexes, aqua complexes and others). Example: K 2.

The formulas of salts reflect their chemical nature, and also indicate the qualitative and quantitative composition of the molecule.

Oxides, salts, bases, acids have different abilities to solubility, which can be viewed in the corresponding table.

If we talk about the state of aggregation of salts, then we need to notice their uniformity. They exist only in solid, crystalline or powdery states. The color range is quite varied. Solutions of complex salts, as a rule, have bright, saturated colors.

Chemical interactions for the class of medium salts

They have similar chemical properties as bases, acids, and salts. Oxides, as we have already examined, are somewhat different from them in this factor.

In total, 4 main types of interactions can be distinguished for medium salts.

I. Interaction with acids (only strong from the point of view of ED) with the formation of another salt and a weak acid:

KCNS + HCL = KCL + HCNS

II. Reactions with soluble hydroxides producing salts and insoluble bases:

CuSO 4 + 2LiOH = 2LiSO 4 soluble salt + Cu(OH) 2 insoluble base

III. Reaction with another soluble salt to form an insoluble salt and a soluble one:

PbCL 2 + Na 2 S = PbS + 2NaCL

IV. Reactions with metals located in the EHRNM to the left of the one that forms the salt. In this case, the reacting metal should not interact with water under normal conditions:

Mg + 2AgCL = MgCL 2 + 2Ag

These are the main types of interactions that are characteristic of medium salts. The formulas of complex, basic, double and acidic salts speak for themselves about the specificity of the chemical properties exhibited.

The formulas of oxides, bases, acids, salts reflect the chemical essence of all representatives of these classes of inorganic compounds, and in addition, give an idea of ​​​​the name of the substance and its physical properties. Therefore, you should pay attention to their writing Special attention. A huge variety of compounds is offered to us by the generally amazing science of chemistry. Oxides, bases, acids, salts - this is only part of the immense diversity.

Sulfur and its compounds.

Equipment, reagents:

Sulfur (small pieces), sulfur (powder), reduced iron, dry sodium sulfite, concentrated sulfuric acid, copper, sodium hydroxide, phenolphthalein, fuchsin, sugar, crystalline potassium permanganate, alcohol, copper (II) oxide.

Large test tubes - 5 pcs., small test tubes - 6 pcs., a test tube rack, a prefabricated rack, a mortar and pestle, a small crucible, a small flask with a gas outlet tube and a dropping funnel, a small glass, glass stirring rods, flasks, cotton wool, porcelain cups, tiles. electric.

Sulfur and its properties

Features of sulfur melting.

Small pieces of sulfur are placed in a test tube to fill 1/3 of its volume (sulfur color is less suitable for these purposes, since strong foaming is observed when it melts). The test tube with sulfur is heated until the sulfur melts (119 "C). With further heating, the sulfur darkens and begins to thicken (maximum thickening at 200 "C). At this moment, the test tube is turned upside down for a moment, and the sulfur will not spill out. With even stronger heating, the sulfur liquefies again, and at 445 "C it boils. Boiling sulfur is poured into a glass or crystallizer with water, while making a circular motion with the test tube. Plastic sulfur solidifies in the water. If you remove it from the water (using a glass rod) , then it stretches like rubber.

Reaction between sulfur and iron.

a) The experiment is carried out in a test tube. First, prepare a mixture of substances in a ratio of 7: 4

(Ar(Fe): Ar(S) = 56: 32). For example, it is enough to take 3.5 g of iron and 2 g of sulfur. In the resulting mixture, individual particles of sulfur, iron and the color of these substances are distinguishable. If you throw a little of the mixture into a glass of water, the sulfur floats (not wetted by water), and the iron sinks (wetted by water).

The mixture can be separated using a magnet. To do this, a magnet is brought to the mixture on a watch glass or glass plate covered with paper, which attracts iron, the sulfur remains on the watch

glass The mixture is transferred into a test tube, which is fixed in a tripod leg at a slightly inclined angle and heated. It is enough to achieve the start of the reaction (red-hot) in one place of the mixture - and the reaction continues by itself (the process is exothermic). To extract the resulting iron sulfide, the test tube is broken. So, from two substances, if they were taken in quantities corresponding to the calculations, one substance was obtained, having properties different from the properties of the original substances.

Possible problems during the experiment

1. For the experiment, you need to take only reduced iron. When using ordinary sawdust, the reaction does not occur, since each grain is covered with a thin film of iron oxides, which

prevents the contact of iron with sulfur.

2. The reaction will not proceed or only isolated outbreaks will be observed if the mixture is poorly mixed and there is not sufficient contact of sulfur with iron.

3. The reaction will not work if the grains of iron are very large, therefore, the surface of contact with sulfur is small.

Sulfur (IV) oxide and sulfurous acid.

Preparation of sulfur (IV) oxide.

a) The flask with solid sodium sulfite is closed with a stopper and a dropping funnel. When adding concentrated sulfuric acid (the acid must be added drop by drop. When observed

strong gas evolution, then the addition of acid is stopped) sulfur oxide (IV) is released. The reaction occurs without heating.

b) Concentrated sulfuric acid is added to copper (shavings, sawdust or wire) and heated. Sulfur (IV) oxide is collected by displacing air.

Dissolution of sulfur oxide (IV) in water.

Place the cylinder with the hole facing up and fill it with sulfur (IV) oxide. The completeness of filling is controlled as with carbon dioxide with a burning torch. The cylinder is covered with glass

plate and hole downwards are lowered into a crystallizer with water. When the cylinder is rocked, water gradually enters it. The solubility of sulfur (IV) oxide in water is very high and at room conditions equal to an average of 40 volumes of gas per 1 volume of water, which is approximately 10% by weight. High solubility always allows students to conclude that in this case a chemical reaction occurs between the dissolving gas and the solvent.

reaction.

Chemical properties of sulfurous acid.

100 - 150 ml of water is poured into a flask and sulfur (IV) oxide is passed through for several minutes so that the solution has a strong odor. This bottle is closed with a stopper.

a) 1/3 of the volume of the test tube is filled with water tinted with magenta. Add sulfurous acid to the colored water and stir the solution. Sulfurous acid gives a colorless solution with organic dyes. Heat the solution to boiling. The magenta color is restored again. Why?

Sulfuric acid

Charring of a splinter.

When a splinter is lowered into concentrated sulfuric acid, it becomes charred and free carbon is released. After rinsing in water, the splinter is shown to students, who conclude that sulfuric acid is capable of removing hydrogen and oxygen from complex substances, which explains some of the rules for working with it.