Experiments in chemistry. Sulfur and its compounds

Being in nature

Copper(I) sulfide occurs naturally as the mineral chalcocite (obsolete names: chalcocite, redrutite, copper luster). Also, the composition of this compound is similar to the mineral jurleite, sometimes called jarleite (for more details, see Copper sulfides). The main properties are presented in the table.

"Mysterious Copper" (page 2)

In an excess of a concentrated solution of strong alkali, copper(II) hydroxide is soluble due to the formation of blue cuprites (NaHCuO2, NaCuO2, etc.). However, the latter are very unstable and, when the solution is diluted, decompose with the release of Cu(OH)2.

It is used as a pigment called "brem blue".

Copper(II) chloride CuCl2.

2H2O
.
Forms dark green crystals, easily soluble in water. Very concentrated solutions of copper (II) chloride are green, diluted solutions are blue-blue, because in this case the formation of aqua complexes [Cu(H2O)6]2+ occurs. Dissolves in water, ammonia, pyridine, ether, alcohol. CuCl2 crystal hydrates are known .

nH2O (n=1, 2, 3, 4). It is obtained by the action of chlorine on copper, the reaction of copper(II) oxide or hydroxide with hydrochloric acid.

Copper (II) bromide CuBr2

.
Black crystals. Dissolves in water and alcohol, pyridine, acetone. Crystalline hydrate CuBr2 is known .
2H2O. It is obtained by the action of liquid bromine on copper, the reaction of copper (II) oxide or hydroxide with hydrobromic acid.

Copper(II) fluoride CuF2

. White small crystals. Dissolves in ammonia, hydrofluoric, hydrochloric and nitric acids, pyridine, ethyl acetate. Reduced to copper by hydrogen. Decomposes with water vapor. Obtained by heating copper with fluorine.

Copper(II) nitrate Cu(NO3)2.

In solid form it is blue crystals of Cu(NO3)2
.
H2O. When heated, copper(II) nitrate crystals first lose water and then easily decompose with the release of oxygen and brown nitrogen dioxide, turning into copper(II) oxide. It is obtained by dissolving copper in nitric acid.

Copper(II) sulfate CuSO4

In an anhydrous state, it is a white powder that turns blue when absorbing water.
Therefore, it is used to detect traces of moisture in organic liquids. An aqueous solution of copper sulfate has a characteristic blue-blue color. This color is characteristic of hydrated ions [Cu(H2O)4]2+ or [Cu(H2O)6]2+, therefore all dilute solutions of copper(II) salts have the same color, unless they contain any colored anions. From aqueous solutions, copper sulfate crystallizes with five molecules of water, forming transparent blue crystals of copper sulfate. It is obtained by dehydration of crystalline hydrates and treatment of copper with concentrated sulfuric acid. Copper sulfate is used for electrolytic coating of metals with copper, for the preparation of mineral paints, and also as a starting material in the preparation of other copper compounds. In agriculture, a diluted solution of copper sulfate is used to spray plants and treat grain before sowing to destroy spores of harmful fungi. It occurs in nature in the form of the mineral chalcanthite (CuSO4 .
5H2O).

Copper(II) hydroxycarbonate (CuOH)2CO3

. It occurs naturally in the form of the mineral malachite, which has a beautiful emerald green color. It is obtained artificially by the action of Na2CO3 on solutions of copper(II) salts.

2CuSO4 + 2Na2CO3 + H2O = (CuOH)2CO3 + 2Na2SO4 + CO2↑

It is used for the production of copper(II) chloride, for the preparation of blue and green mineral paints, as well as in pyrotechnics.

Copper(II) acetate Cu(CH3COO)2.

H2O
.
It is obtained by treating copper(II) hydroxide or oxide with acetic acid. Usually it is a mixture of basic salts of various compositions and colors (green and blue-green). Under the name verdigris, it is used to prepare oil paint. Most cupric salts are soluble in water, and solutions are often acidic. With weak acids, copper forms basic salts.

Volatile copper compounds color the non-luminous flame of a gas burner blue-green.

In dilute solutions, the color of divalent copper salts with colorless anions is blue-blue (the color of the Cu2+ ion). On the contrary, the color of solid salts is different. In particular, the most practically important of them is copper sulfate (CuSO4 .

5H2O) – has a blue color.

Complexation is very typical for divalent copper. This is already evident from the fact that almost all Cu2+ salts are released from solutions in the form of crystalline hydrates. With the corresponding alkali metal salts, Cu2+ salts give double compounds containing copper as part of complex anions (for example, [CuCl4]2-). However, most of the latter are unstable and disintegrate into their component parts in solution. Much more stable is the dark blue complex cation [Cu(NH3)4] 2+, very characteristic of divalent copper, which is formed by adding excess ammonia to Cu2+ solutions according to the reaction, for example:

CuSO4 + 4NH 3.

H2O = [Cu(NH3)4]SO4 + 4H2O

Due to the intense color of this complex, ammonia can be used as a reagent for copper.

Amino complexes [Cu(NH3)2]+ are also characteristic of monovalent copper. Therefore, Cu(I) compounds can dissolve in the presence of ammonia:

CuClt + 2NH3.

H2Ozh = [Cu(NH3)2]Clр

Using copper salts, mineral paints of various colors are obtained: green, blue, brown, purple and black.

All copper salts are poisonous. For example, mixed copper(II) acetate-arsenite—Cu(CH3COO)2•3Cu(AsO2)2—is used under the name “Paris green” to kill plant pests.

In everyday life, to avoid the formation of toxic copper salts, copper utensils are coated on the inside with a layer of tin (tinned).

After analyzing all the information received, we discovered that our ideas about the properties of copper and its compounds differed from reality. These differences are summarized in Table 1.

Table 1

To verify new information about copper and its compounds, the following experiments were performed in the experimental part of the work:

· Dissolving copper in ammonia solution.

· Interaction of copper with concentrated hydrochloric and

dilute sulfuric acids.

· Dissolution of copper in a solution of iron(III.) chloride

· Interaction of copper with hydrobromic acid.

· Study of amphoteric properties of copper(I) oxide. Dissolution in saline

acid, ammonia and ammonium hydroxide solutions.

· Study of amphoteric properties of copper(II) oxide. Dissolution in saline

acid, ammonia and ammonium hydroxide solutions.

· Study of amphoteric properties of copper(II) hydroxide. Dissolution in

hydrochloric acid, ammonia and ammonium hydroxide solutions.

The experiments used freshly deposited copper, copper(I) oxide, copper(II) oxide, copper(II) hydroxide and hydrobromic acid, which were obtained immediately before the experiments.

A saturated solution of copper(II) sulfate was poured into a porcelain or glass cup and zinc granules were added. The released copper was deposited on the zinc in the form of a loose mass.

CuSO4 + Zn = ZnSO4 + Cu

When stirring the solution, sediment accumulated at the bottom of the cup

(Fig. 1). The precipitate was washed, granules of unreacted zinc were removed; The resulting copper, without drying, was used for experiments.

Fig.1. Preparation of freshly deposited copper

5 ml of a ten percent solution of copper (II) sulfate was poured into the test tube and then a slightly larger volume of potassium hydroxide. An excess of copper(II) hydroxide was added to the resulting blue precipitate

glucose and stirred (Figure 2, bright blue solution). The mixture was heated with water at a temperature of ~100 °C. The formation of a yellow copper(I) precipitate and an orange-colored mixture of copper(I) hydroxide and oxide precipitates were observed. After completion of the reaction, a red precipitate of copper(I) oxide formed in the test tube (Fig. 2).

Rice. 2. Preparation of copper(I) oxide.

2C6H12O6 + 2Cu(OH)2 → 2C6H12O7 + 2CuOH + H2O

2CuOH → Cu2O + H2O

A solution of sodium hydroxide was placed in a test tube and heated. A solution of copper (II) sulfate was added to the hot solution until a precipitate formed. The resulting mixture was heated. As a result, a black precipitate of copper(II) oxide was obtained.

CuSO4 + 2NaOH = Na2SO4 + Cu(OH)2

t

Cu(OH)2 = CuO + H2O

5 ml of a ten percent solution of copper(II) sulfate and then a solution of sodium hydroxide were poured into the test tube until a blue gelatinous precipitate of copper(II) hydroxide formed.

CuSO4 + 2NaOH = Na2SO4 + Cu(OH)2

A. The interaction of dilute sulfuric acid with bromides

.

When potassium bromide was exposed to a solution obtained from three volumes of concentrated sulfuric acid (ρ = 1.84 g/cm3) and one volume of water [6, 7], hydrogen bromide was released (Fig. 3-1, 4-1):

Physical properties

Copper(I) sulfide is a dark gray solid, insoluble in water and ethanol.

Cu2S exists in three crystalline modifications:

• α-Cu2S, orthorhombic system, space group Ab
  2m, cell parameters
  a
  = 1.1190 nm,
  b
  = 2.728 nm,
  c
  = 1.341 nm,
  Z
  = 96,
  d
  = 5.81 g/cm3;
• β-Cu2S, hexagonal system, space group P
  63/mmc, cell parameters
  a
  = 0.389 nm,
  c
  = 0.668 nm,
  Z
  = 2,
  d
  = 5.78 g/cm3;
• γ-Cu2S, cubic system, space group F
  m3m, cell parameters
  a
  = 0.5735 nm,
  Z
  = 4,
  d
  = 5.60 g/cm3.

Phase transition temperatures: α → β 103 °C, β → γ 437 °C.

Copper(I) sulfide tends to form non-stoichiometric compounds, the properties of which can differ greatly from stoichiometric ones.

Experiments in chemistry. Sulfur and its compounds

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Interaction of sulfur with copper

Copper is an inactive metal. But at elevated temperatures it actively interacts with sulfur. Heat the sulfur in the test tube to boiling. The sulfur melts and gradually boils. The test tube is filled with brown sulfur vapor. Let's lower the copper wire into the test tube. Copper ignites in sulfur vapor. The reaction product is black copper sulfide.

Cu + S = CuS

Equipment: stand, test tube, burner, crucible tongs.

Safety precautions. The experiment should be carried out under traction. Follow the rules for working with flammable substances and heating devices.

Experiment setting and text – Ph.D. Pavel Bespalov.

Interaction of sulfur with sodium

Sodium is an active alkali metal. To avoid oxidation and interaction with water vapor, sodium is stored in kerosene. A small piece of sodium is placed in a mortar with sulfur powder. Carefully grind the sodium. It actively interacts with sulfur. The reaction product is sodium sulfide.

Na + S = Na 2 S

Equipment: porcelain mortar with pestle, spatula, scalpel, tweezers, filter paper.

Safety precautions. The experiment should be carried out under traction. Follow the rules for working with alkali metals. To avoid chemical burns, wrap your hand in a towel.

Experiment setting and text – Ph.D. Pavel Bespalov.

Interaction of sulfur with zinc

Zinc is less active compared to sodium. Heating is necessary for the reaction with sulfur to occur. Set fire to a mixture of sulfur and zinc powders. The mixture burns vigorously to form zinc sulfide.

Zn + S = ZnS

Equipment: spatula, fireproof gasket, burner, torch.

Safety precautions. Follow the rules for working with heating devices. Conduct the experiment under traction.

Experiment setting and text – Ph.D. Pavel Bespalov.

Obtaining plastic sulfur

Sulfur is a crystalline substance. In crystals, molecules are always arranged in a strict order. Powdered sulfur retains its crystalline structure. Let's heat the sulfur powder. Sulfur melts at a temperature of 120 degrees and becomes an easily mobile yellow liquid. With a further increase in temperature, sulfur begins to darken and turns into a dark brown, viscous liquid - plastic sulfur. When rapidly cooled in water, sulfur retains its plastic properties. Plastic sulfur crystallizes and hardens over time.

Equipment: test tube, test tube rack, spatula, beaker, alcohol lamp, test tube holder.

Safety precautions . Follow the rules for handling heating devices. Do not allow molten sulfur to come into contact with the table surface. If sulfur catches fire, extinguish the flame with river sand.

Experiment setup – Elena Makhinenko, text – Ph.D. Pavel Bespalov.

Obtaining hydrogen sulfide and studying its properties

Hydrogen sulfide is a toxic, colorless gas with a rotten egg odor. Hydrogen sulfide (H2S) can be produced in the laboratory by heating a mixture of paraffin and sulfur. We place a piece of paraffin and yellow sulfur powder into a test tube, close it with a stopper with a gas outlet tube and heat it on the flame of an alcohol lamp. The released hydrogen sulfide can be detected using wet universal indicator paper: under the influence of hydrogen sulfide it turns red. A solution of hydrogen sulfide in water is a very weak acid. We pass hydrogen sulfide through cold distilled water. Hydrogen sulfide partially dissolves and hydrogen sulfide water is formed. Litmus turns red in hydrogen sulfide water, which is a really weak acid. A qualitative reaction to sulfide ion is a reaction with lead ions; we will carry out a similar reaction with copper ions. When copper sulfate is added to hydrogen sulfide water, a black precipitate of copper sulfide precipitates.

CuSO 4 + H 2 S = CuS ↓+ H 2 SO 4

This is what lead sulfide precipitate looks like. The black precipitate of lead sulfide can reveal the sulfide ion.

Equipment: test tubes, gas outlet tube, stand, stand for test tubes, spatula, tongs, porcelain crucibles, alcohol lamp, test tube holder.

Safety precautions . Rules for handling toxic gases must be observed. The experiment must be carried out in a fume hood. Follow the rules for handling heating devices.

Experiment setup – Elena Makhinenko, text – Ph.D. Pavel Bespalov.

Preparation of sulfurous acid

Sulfurous acid can be produced by dissolving sulfur dioxide in water. To obtain sulfur dioxide you will need concentrated sulfuric acid and copper: place a piece of copper in a test tube with concentrated sulfuric acid and heat the solution.

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

Let's direct a current of sulfur dioxide into a test tube with water. When sulfur dioxide dissolves, it partially reacts with water to form weak sulfurous acid.

SO 2 + H 2 O = H 2 SO 3

By passing sulfur dioxide through water, we get sulfurous acid. Sulfur dioxide is released in huge quantities from the bowels of the Earth at active volcanoes. Sulfur dioxide irritates the mucous membranes of the respiratory tract, causing inflammation, loss of appetite, and, in high concentrations, loss of consciousness. Therefore, sulfur dioxide must be handled with all precautions.

Equipment: test tubes, gas outlet tube, stand, stand for test tubes, tweezers, alcohol lamp.

Safety precautions . The experiment must be carried out under traction. Follow the rules for handling heating devices. Special care is required when working with concentrated sulfuric acid. Follow the rules for working with concentrated acids.

Inhalation of the sulfur dioxide gas released is unacceptable.

Experiment setup – Elena Makhinenko, text – Ph.D. Pavel Bespalov.

Properties of sulfurous acid

A solution of sulfurous acid H2SO3 has reducing properties. Sulfurous acid reacts with iodine solution, discoloring it. In this case, hydroiodic and sulfuric acids are formed.

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

Like all acids, sulfurous acid changes the color of indicator solutions. Methyl orange turns red in acid solution. In the old days, ladies' straw hats were bleached with sulfurous acid. A solution of sulfurous acid bleaches fabrics made from plant material, wool, and silk.

Equipment: test tubes, gas outlet tube, stand, stand for test tubes, tweezers, alcohol lamp.

Safety precautions . The experiment must be carried out under traction. Follow the rules for handling heating devices. Inhalation of the sulfur dioxide gas released is unacceptable.

Experiment setup – Elena Makhinenko, text – Ph.D. Pavel Bespalov.