Given zinc copper zinc oxide. Zinc - general characteristics of an element, chemical properties of zinc and its compounds
I.V. TRIGUBCHAK
Chemistry Tutor Benefit
Continuation. For the beginning, see No. 22/2005; 1, 2, 3, 5, 6, 8, 9, 11, 13, 15, 16, 18, 22/2006;
3, 4, 7, 10, 11, 21/2007;
2, 7, 11/2008
LESSON 24
10th grade(first year of study)
Zinc and its compounds
1. Position in the table of DI Mendeleev, the structure of the atom.
2. Origin of the name.
3. Physical properties.
4. Chemical properties.
5. Being in nature.
6. Basic methods of obtaining.
7. Zinc oxide and hydroxide - properties and methods of production.
Zinc is located in a secondary subgroup of group II of Mendeleev's table. His electronic formula 1s 2 2s 2 p 6 3s 2 p 6 d 10 4s 2. Zinc is d-element, shows in compounds the only oxidation state +2 (since the third energy level in the zinc atom is completely filled with electrons). Being an amphoteric element with a predominance of metallic properties, in compounds zinc is more often included in the cation, less often in the anion. For instance,
It is believed that the name of zinc comes from the ancient Germanic word "zinc" (white, thorn). In turn, this word goes back to the Arabic "harasin" (metal from China), which indicates the place of production of zinc, which was brought to Europe from China in the Middle Ages.
PHYSICAL PROPERTIES
Zinc is a white metal; in air it becomes covered with an oxide film, and its surface fades. In the cold, it is a rather brittle metal, but at a temperature of 100–150 ° C, zinc is easily processed and forms alloys with other metals.
Chemical properties
Zinc is a metal of average chemical activity, but it is more active than iron. After the destruction of the oxide film, zinc exhibits the following chemical properties.
Zn + H 2 ZnH 2.
2Zn + O 2 2ZnO.
Metals (-).
Non-metals (+):
Zn + Cl 2 ZnCl 2,
3Zn + 2P Zn 3 P 2.
Zn + 2H 2 O Zn (OH) 2 + H 2.
Basic oxides (-).
Acidic oxides (-).
Reasons (+):
Zn + 2NaOH + 2H 2 O = Na 2 + H 2,
Zn + 2NaOH (melt) = Na 2 ZnO 2 + H 2.
Non-oxidizing acids (+):
Zn + 2HCl = ZnCl 2 + H 2.
Oxidizing acids (+):
3Zn + 4H 2 SO 4 (conc.) = 3ZnSO 4 + S + 4H 2 O.
4Zn + 5H 2 SO 4 (conc.) = 4ZnSO 4 + H 2 S + 4H 2 O,
4Zn + 10HNO 3 (very dil.) = 4Zn (NO 3) 2 + NH 4 NO 3 + 3H 2 O.
Salts (+/–): *
Zn + CuCl 2 = Cu + ZnCl 2,
Zn + NaCl no reaction.
In general, zinc is found in the form of compounds, the most important of which are sphalerite, or zinc blende (ZnS), smithsonite, or zinc spar (ZnCO 3), red zinc ore (ZnO).
In industry, for the production of zinc, zinc ore is roasted in order to obtain zinc oxide, which is then reduced with carbon:
2ZnS + 3O 2 2ZnO + 2SO 2,
2ZnO + C2Zn + CO 2.
The most important zinc compounds are its o to s and d (ZnO) and g and dro to c and d (Zn (OH) 2). These are white crystalline substances that exhibit amphoteric properties:
ZnO + H 2 SO 4 = ZnSO 4 + H 2 O,
ZnO + 2NaOH + H 2 O = Na 2,
Zn (OH) 2 + 2HCl = ZnCl 2 + 2H 2 O,
Zn (OH) 2 + 2NaOH = Na 2.
Zinc oxide can be obtained by oxidizing zinc, decomposing zinc hydroxide, or burning zinc blende:
Zn (OH) 2 ZnO + H 2 O,
2ZnS + 3O 2 2ZnO + 3SO 2.
Zinc hydroxide is produced by an exchange reaction between a solution of a zinc salt and an alkali:
ZnCl 2 + 2NaOH (deficiency) = Zn (OH) 2 + 2NaCl.
These compounds should be remembered: zinc blende (ZnS), zinc sulfate (ZnSO 4 7H 2 O).
Test on the topic "Zinc and its compounds"
1. The sum of the coefficients in the equation for the reaction of zinc with very dilute nitric acid:
a) 20; b) 22; c) 24; d) 29.
2. Zinc from a concentrated sodium carbonate solution displaces:
a) hydrogen; b) carbon monoxide;
c) carbon dioxide; d) methane.
3. Alkaline solutions can react with the following substances (several correct answers are possible):
a) copper sulfate and chlorine;
b) calcium oxide and copper;
c) sodium hydrogen sulfate and zinc;
d) zinc hydroxide and copper hydroxide.
4. The density of a 27.4% sodium hydroxide solution is 1.3 g / ml. The molar concentration of alkali in this solution is:
a) 0.0089 mol / ml; b) 0.0089 mol / l;
c) 4 mol / l; d) 8.905 mol / l.
5. To obtain zinc hydroxide, you must:
a) add sodium hydroxide solution dropwise to the zinc chloride solution;
b) add the zinc chloride solution dropwise to the sodium hydroxide solution;
c) add an excess of sodium hydroxide solution to the zinc chloride solution;
d) add the sodium hydroxide solution dropwise to the zinc carbonate solution;
6. Eliminate the "extra" connection:
a) H 2 ZnO 2; b) ZnCl 2; c) ZnO; d) Zn (OH) 2.
7. An alloy of copper and zinc weighing 24.12 g was treated with an excess of dilute sulfuric acid. At the same time, 3.36 liters of gas (n.u.) were released. The mass fraction of zinc in this alloy is (in%):
a) 59.58; b) 40.42; c) 68.66; d) 70.4.
8. Zinc granules will interact with an aqueous solution (several correct answers are possible):
a) hydrochloric acid; b) nitric acid;
c) potassium hydroxide; d) aluminum sulfate.
9. Carbon dioxide with a volume of 16.8 liters (NU) was absorbed by 400 g of a 28% potassium hydroxide solution. The mass fraction of a substance in solution is (in%):
a) 34.5; b) 31.9; c) 69; d) 63.7.
10. The mass of a zinc carbonate sample, which contains 4.816 10 24 oxygen atoms, is (in g):
a) 1000; b) 33.3; c) 100; d) 333.3.
The key to the test
| 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 |
| b | a | a, in | G | a | b | b | a B C D | b | G |
Tasks and Exercises for Amphoteric Metals
Chains of transformation
1. Zinc -> zinc oxide -> zinc hydroxide -> zinc sulfate -> zinc chloride -> zinc nitrate -> zinc sulfide -> zinc oxide -> potassium zincate.
2. Aluminum oxide -> potassium tetrahydroxoaluminate -> aluminum chloride -> aluminum hydroxide -> potassium tetrahydroxoaluminate.
3. Sodium -> sodium hydroxide -> sodium bicarbonate -> sodium carbonate -> sodium hydroxide -> sodium hexahydroxochromate (III).
4. Chromium -> chromium (II) chloride -> chromium (III) chloride -> potassium hexahydroxochromate (III) + bromine + potassium hydroxide -> potassium chromate -> potassium dichromate -> chromium (VI) oxide.
5. Iron (II) sulfide -> X 1 -> iron (III) oxide -> X 2 -> iron (II) sulfide.
6. Iron (II) chloride -> A -> B -> C -> D -> E -> iron (II) chloride (all substances contain iron; in the scheme there are only three redox reactions in a row).
7. Chromium -> X 1 -> chromium (III) sulfate -> X 2 -> potassium dichromate -> X 3 -> chromium.
LEVEL A
1. To dissolve 1.26 g of an alloy of magnesium with aluminum, 35 ml of a 19.6% solution of sulfuric acid (density - 1.14 g / ml) were used. The excess acid reacted with 28.6 ml of a 1.4 mol / L potassium hydrogen carbonate solution. Determine the composition of the starting alloy and the volume of gas (n.o.) released during the dissolution of the alloy.
Answer. 57.6% Mg; 42.4% Al; 1.34 L H 2.
2. A mixture of calcium and aluminum weighing 18.8 g was calcined in the absence of air with an excess of graphite powder. The reaction product was treated with diluted hydrochloric acid, while 11.2 liters of gas (n.u.) were released. Determine the composition of the original mixture.
Solution
Reaction equations:
Let (Ca) = x mol, (Al) = 4 y mole.
Then: 40 x + 4 27y = 18,8.
By the condition of the problem:
v (C 2 H 2 + CH 4) = 11.2 l.
Hence,
(C 2 H 2 + CH 4) = 11.2 / 22.4 = 0.5 mol.
According to the reaction equation:
(C 2 H 2) = (CaC 2) = (Ca) = X mole,
(CH 4) = 3/4 (Al) = 3 y mole,
x + 3y = 0,5.
We solve the system:
x = 0,2, y = 0,1.
Hence,
(Ca) = 0.2 mol,
(Al) = 4 0.1 = 0.4 mol.
In the original mixture:
m(Ca) = 0.2 40 = 8 g,
(Ca) = 8 / 18.8 = 0.4255, or 42.6%;
m(Al) = 0.4 27 = 10.8 g,
(Al) = 10.8 / 18.8 = 0.5744, or 57.4%.
Answer... 42.6% Ca; 57.4% Al.
3. When 11.2 g of metal of group VIII of the periodic system interacted with chlorine, 32.5 g of chloride were formed. Identify the metal.
Answer... Iron.
4. Firing pyrite produced 25 m 3 of sulfur dioxide (temperature 25 ° C and pressure 101 kPa). Calculate the mass of the resulting solid.
Answer. 40.8 kg Fe 2 O 3.
5. On calcining 69.5 g of the crystalline hydrate of iron (II) sulfate, 38 g of anhydrous salt are formed. Determine the formula for the crystalline hydrate.
Answer. Heptahydrate FeSO 4 7H 2 O.
6. Under the action of an excess of hydrochloric acid on 20 g of a mixture containing copper and iron, a gas with a volume of 3.36 L (NU) was released. Determine the composition of the original mixture.
Answer. 58% Cu; 42% Fe.
Level B
1. What volume of a 40% solution of potassium hydroxide (density - 1.4 g / ml) should be added to 50 g of a 10% solution of aluminum chloride in order to completely dissolve the initially precipitated precipitate?
Answer. 15 ml
2. The metal was burned in oxygen with the formation of 2.32 g of oxide, for the reduction of which to metal it is necessary to spend 0.896 L (NU) of carbon monoxide. The reduced metal was dissolved in dilute sulfuric acid, the resulting solution gives a blue precipitate with red blood salt. Determine the oxide formula.
Answer: Fe 3 O 4.
3. What volume of a 5.6 M potassium hydroxide solution is required to completely dissolve 5 g of a mixture of chromium (III) and aluminum hydroxides, if the mass fraction of oxygen in this mixture is 50%?
Answer. 9.3 ml.
4. Sodium sulfide was added to a 14% solution of chromium (III) nitrate, the resulting solution was filtered and boiled (without loss of water), while the mass fraction of the chromium salt decreased to 10%. Determine the mass fractions of the remaining substances in the resulting solution.
Answer. 4.38% NaNO 3.
5. A mixture of iron (II) chloride with potassium dichromate was dissolved in water and the solution was acidified with hydrochloric acid. After some time, an excess of potassium hydroxide solution was added dropwise to the solution, the precipitate formed was filtered off and calcined to constant weight. The mass of the dry residue is 4.8 g. Find the mass of the initial mixture of salts, taking into account that the mass fractions of iron (II) chloride and potassium dichromate in it are in the ratio of 3: 2.
Answer. 4.5 g
6. 139 g of ferrous sulfate was dissolved in water at a temperature of 20 ° C and received a saturated solution. When this solution was cooled to 10 ° C, a precipitate of ferrous sulfate precipitated. Find the mass of the precipitate and mass fraction iron (II) sulfate in the remaining solution (the solubility of iron (II) sulfate at 20 ° C is 26 g, and at 10 ° C - 20 g).
Answer. 38.45 g FeSO 4 7H 2 O; 16.67%.
Qualitative tasks
1. A silvery-white light simple substance A, which has good thermal and electrical conductivity, reacts when heated with another simple substance B. The resulting solid dissolves in acids with the release of gas C, when passed through a solution of sulfurous acid, a precipitate of substance B precipitates. substances, write the reaction equations.
Answer. Substances: A - Al, B - S, C - H 2 S.
2. There are two gases, A and B, whose molecules are triatomic. When each of them is added to the potassium aluminate solution, a precipitate forms. Suggest possible formulas for gases A and B, considering that these gases are binary. Write down the reaction equations. How can these gases be distinguished chemically?
Solution
Gas A - CO 2; gas B - H 2 S.
2KAlO 2 + CO 2 + 3H 2 O = 2Al (OH) 3 + K 2 CO 3,
2KAlO 2 + H 2 S + 2H 2 O = 2Al (OH) 3 + K 2 S.
3. Brown compound A, insoluble in water, decomposes on heating to form two oxides, one of which is water. Another oxide, B, is reduced by carbon to form metal C, the second most common metal in nature. Identify substances, write down the reaction equations.
Answer. Substances: A - Fe (OH) 3,
B - Fe 2 O 3, C - Fe.
4. Salt A is formed by two elements; when it is fired in air, two oxides are formed: B - solid, brown, and gaseous. Oxide B enters into a substitution reaction with the silvery white metal C (when heated). Identify substances, write down the reaction equations.
Answer. Substances: A - FeS 2, B - Fe 2 O 3, C - Al.
* The +/– sign means that this reaction does not take place with all reagents or under specific conditions.
To be continued
Write the reaction equations according to the Pozhaaaluist schemes 1) calcium phosphate + barium chloride = barium phosphate + calcium chloride 2) Sodium carbonate + potassium nitrate = carbonatecalcium + sodium nitrate 3) Sulfuric acid + magnesium hydroxide = magnesium sulfate + foda 4) Lithium oxide + hydrochloric acid = lithium chloride + water 5) Sulfur oxide (V1) + sodium hydroxide = sodium sulfate + water 6) Aluminum + hydrobromic acid = aluminum bromide + hydrogen 7) Lead nitrate (11) + sodium sulfide = lead sulfide (11) + silicic acid 8) Potassium silicate + phosphoric acid = potassium phosphate + silicic acid 9) zinc hydroxide-hydroiodic acid = zinc iodide + water 10) Nitric oxide (V) + sodium hydroxide = potassium netrate + water 11) Barium nitrate + sulfuric acid = barium sulfate + nitric acid 12) Carbon monoxide (1V) -calcium hydroxide = calcium carbonate + water 13) Sulfur oxide (1V) + oxide potassium = potassium sulfate 14) Magnesium oxide + phosphorus (V) oxide = magnesium phosphate 15) Nitric acid + chromium godroxide (111) = chromium nitrate (111) + water 16) Hydrogen sulfide acid + silver netrate = silver sulfide + nitric acid 17) Iron oxide (111) + hydrogen = iron + water 18) Copper nitrate (11) + aluminum = copper + aluminum nitrate 19) Aluminum hydroxide = aluminum oxide + water
a) sodium --- sodium hydroxide - sodium sulfide --- sodium chloride --- sodium sulfate b) magnesium --- magnesium sulfate --- magnesium hydroxide --- magnesium oxide - magnesium chloridec) lead - lead (II) oxide - lead (II) nitrate - lead (II) hydroxide - lead (II) oxide - lead (II) sulfate g) sulfur --- hydrogen sulfide --- potassium sulfite - - potassium chloride - potassium chloride - hydrochloric acid e) calcium - calcium hydroxide - calcium carbonate - calcium nitrate - nitric acid f) aluminum - aluminum sulfate - aluminum hydroxide - aluminum oxide - aluminum nitrate g) sulfur - sulfur (IV) oxide - sulfurous acid --- sodium sulfite - sulfurous acid h) oxygen - aluminum oxide - aluminum sulfate - aluminum hydroxide - sodium metaaluminate j) aluminum - chloride aluminum - aluminum nitrate - aluminum hydroxide - aluminum sulfate l) copper - copper (II) chloride - copper - copper (II) oxide - copper (II) nitrate m) iron - iron (II) chloride - iron (II) hydroxide - iron (II) sulfate - iron n) iron - iron (III) chloride - iron (III) nitrate - iron (III) sulfate - iron
1.Reacts with an aqueous solution of sodium carbonate1) potassium sulfate 3) copper (II) sulfide
2) carbon monoxide (IV) 4) silicic acid
2.Reacts with barium chloride solution
1) calcium hydroxide 3) sodium sulfate
2) copper (II) hydroxide 4) Hydrogen
3.Reacts with calcium nitrate solution
1) sodium carbonate 3) silicon
2) zinc 4) hydrobromic acid
4.the interaction of 1 mol and 2 mol of KoH forms
1) medium salt 3) acidic salt
2) basic salt 4) substances do not react
5. As a result of the reaction of sodium silicate with hydrochloric acid,
1) sodium silicide 3) silicic acid
2) Silicon 4) silicon oxide
1. Salt and alkali are formed by the interaction of solutions
1)
2.Reacts with barium nitrate solution
1) sodium chloride 3) potassium carbonate
2) copper 4) calcium carbonate
3.Reacts with barium nitrate solution
1) sodium sulfate 3) iron
2) chloride words 4) copper
4.Reacts with zinc sulfate solution
1) magnesium 3) sulfur
2) silicon oxide 4) aluminum hydroxide
5.A chemical reaction (in solution) is possible between
6) Between what substances does a chemical reaction take place?
1) calcium carbonate and sodium nitrate
2) magnesium silicate and potassium phosphate
3) iron (II) sulfate and lead sulfide
4) barium chloride and zinc sulfate
Alloy of zinc with copper - brass - was known in Ancient Greece, Ancient Egypt, India (VII century), China (XI century). For a long time it was not possible to isolate pure zinc. In 1746, A.S. Marggraf developed a method for obtaining pure zinc by calcining a mixture of its oxide with coal without access to air in clay refractory retorts, followed by condensation of zinc vapor in refrigerators. On an industrial scale, zinc smelting began in the 17th century.
The Latin zincum translates as "white bloom". The origin of this word has not been precisely established. Presumably, it comes from the Persian "cheng", although this name does not refer to zinc, but to stones in general. The word "zinc" is found in the writings of Paracelsus and other researchers of the 16th and 17th centuries. and possibly goes back to the ancient Germanic "zinc" - a raid, an eyesore. The name "zinc" became commonly used only in the 1920s.
Being in nature, getting:
The most common zinc mineral is sphalerite, or zinc blende. The main component of the mineral is zinc sulfide ZnS, and various impurities give this substance all kinds of colors. Apparently, for this, the mineral is called blende. Zinc blende is considered the primary mineral from which other minerals of element 30 were formed: smithsonite ZnCO 3, zincite ZnO, calamine 2ZnO · SiO 2 · H 2 O. In Altai, you can often find striped "chipmunk" ore - a mixture of zinc blende and brown spar. A piece of such ore from a distance really looks like a hidden striped animal.
The separation of zinc begins with the concentration of ore by sedimentation or flotation methods, then it is roasted to form oxides: 2ZnS + 3О 2 = 2ZnО + 2SO 2
Zinc oxide is processed electrolytically or reduced with coke. In the first case, zinc is leached from the crude oxide with a dilute solution of sulfuric acid, the cadmium impurity is precipitated with zinc dust, and the zinc sulfate solution is subjected to electrolysis. Metal of 99.95% purity is deposited on aluminum cathodes.
Physical properties:
In its pure form, it is a rather ductile silvery-white metal. At room temperature fragile, when the plate is bent, a crackle is heard from the friction of crystallites (usually stronger than the "cry of tin"). Zinc is ductile at 100-150 ° C. Impurities, even insignificant ones, sharply increase the brittleness of zinc. Melting point - 692 ° C, boiling point - 1180 ° C
Chemical properties:
Typical amphoteric metal. The standard electrode potential is -0.76 V, in the series of standard potentials it is located before iron. In air, zinc is covered with a thin film of ZnO oxide. Burns out when heated up. When heated, zinc reacts with halogens, with phosphorus, forming phosphides Zn 3 P 2 and ZnP 2, with sulfur and its analogs, forming various chalcogenides, ZnS, ZnSe, ZnSe 2 and ZnTe. Zinc does not directly react with hydrogen, nitrogen, carbon, silicon and boron. Zn 3 N 2 nitride is produced by the reaction of zinc with ammonia at 550-600 ° C.
Zinc of ordinary purity reacts actively with solutions of acids and alkalis, forming in the latter case hydroxozincates: Zn + 2NaOH + 2H 2 O = Na 2 + H 2
Very pure zinc does not react with solutions of acids and alkalis.
Zinc is characterized by compounds with the oxidation state: +2.
The most important connections:
Zinc oxide- ZnO, white, amphoteric, reacts with both acid solutions and alkalis:
ZnO + 2NaOH = Na 2 ZnO 2 + H 2 O (fusion).
Zinc hydroxide- is formed in the form of a gelatinous white precipitate when alkali is added to aqueous solutions of zinc salts. Amphoteric hydroxide
Zinc salts... Colorless crystalline substances. In aqueous solutions, zinc ions Zn 2+ form aqua complexes 2+ and 2+ and undergo strong hydrolysis.
Zincats are formed by the interaction of zinc oxide or hydroxide with alkalis. Upon fusion, metazincates are formed (for example, Na 2 ZnO 2), which, dissolving in water, transform into tetrahydroxozincates: Na 2 ZnO 2 + 2H 2 O = Na 2. When solutions are acidified, zinc hydroxide precipitates.
Application:
Production of anti-corrosion coatings. - Metallic zinc in the form of bars is used to protect steel products from corrosion in contact with sea water. About half of all zinc produced is used for the production of galvanized steel, one third for hot-dip galvanizing of finished products, and the rest for strip and wire.
- Alloys of zinc - brass (copper plus 20-50% zinc) are of great practical importance. For die casting, in addition to brass, a rapidly growing number of special zinc alloys are used.
- Another area of application is the production of dry batteries, although in last years it has dropped significantly.
- Zinc telluride ZnTe is used as a material for photoresistors, infrared detectors, dosimeters and radiation counters. - Zinc acetate Zn (CH 3 COO) 2 is used as a fixative for dyeing fabrics, wood preservative, antifungal agent in medicine, catalyst in organic synthesis. Zinc acetate is a component of dental cements and is used in the production of glazes and porcelain.
Zinc is one of the most important biologically active elements and is essential for all forms of life. Its role is mainly due to the fact that it is part of more than 40 important enzymes. The function of zinc in proteins responsible for the recognition of the base sequence in DNA and, therefore, regulating the transfer of genetic information during DNA replication has been established. Zinc is involved in carbohydrate metabolism with the help of a zinc-containing hormone - insulin. Vitamin A acts only in the presence of zinc. Zinc is also needed for the formation of bones.
At the same time, zinc ions are toxic.
Bespomestnykh S., Shtanova I.
KhF Tyumen State University, group 571.
Sources: Wikipedia:
Copper (Cu) belongs to d-elements and is located in group IB of Mendeleev's periodic table. The electronic configuration of the copper atom in the ground state is written in the form 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 instead of the assumed formula 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2. In other words, in the case of a copper atom, the so-called “electron slip” is observed from the 4s sublevel to the 3d sublevel. For copper, in addition to zero, oxidation states +1 and +2 are possible. The oxidation state +1 is prone to disproportionation and is stable only in insoluble compounds such as CuI, CuCl, Cu 2 O, etc., as well as in complex compounds, for example, Cl and OH. Copper compounds in the +1 oxidation state do not have a specific color. So, copper (I) oxide, depending on the size of the crystals, can be dark red (large crystals) and yellow (small crystals), CuCl and CuI - white, and Cu 2 S - black and blue. More chemically stable is the oxidation state of copper, equal to +2. Salts containing copper in this oxidation state are blue and blue-green in color.
Copper is a very soft, ductile and ductile metal with high electrical and thermal conductivity. The color of metallic copper is red-pink. Copper is in the line of metal activity to the right of hydrogen, i.e. refers to low-activity metals.
with oxygen
Under normal conditions, copper does not interact with oxygen. For the reaction to take place between them, heating is required. Depending on the excess or lack of oxygen and temperature conditions, it can form copper (II) oxide and copper (I) oxide:
with gray
The reaction of sulfur with copper, depending on the operating conditions, can lead to the formation of both copper (I) sulfide and copper (II) sulfide. When a mixture of powdered Cu and S is heated to a temperature of 300-400 ° C, copper (I) sulfide is formed:
With a lack of sulfur and the reaction is carried out at a temperature of more than 400 ° C, copper (II) sulfide is formed. However, more in a simple way obtaining copper (II) sulfide from simple substances is the interaction of copper with sulfur dissolved in carbon disulfide:
This reaction takes place at room temperature.
with halogens
Copper reacts with fluorine, chlorine and bromine, forming halides with the general formula CuHal 2, where Hal is F, Cl or Br:
Cu + Br 2 = CuBr 2
In the case of iodine, the weakest oxidizing agent among halogens, copper (I) iodide is formed:
Copper does not interact with hydrogen, nitrogen, carbon and silicon.
with non-oxidizing acids
Almost all acids are non-oxidizing acids, except for concentrated sulfuric acid and nitric acid of any concentration. Since non-oxidizing acids are able to oxidize only metals that are in the range of activity to hydrogen; this means that copper does not react with such acids.
with oxidizing acids
- concentrated sulfuric acid
Copper reacts with concentrated sulfuric acid both when heated and at room temperature. When heated, the reaction proceeds in accordance with the equation:
Since copper is not a strong reducing agent, sulfur is reduced in this reaction only to the +4 oxidation state (in SO 2).
- with diluted nitric acid
The reaction of copper with dilute HNO 3 leads to the formation of copper (II) nitrate and nitrogen monoxide:
3Cu + 8HNO 3 (dil.) = 3Cu (NO 3) 2 + 2NO + 4H 2 O
- with concentrated nitric acid
Concentrated HNO 3 reacts readily with copper under normal conditions. The difference between the reaction of copper with concentrated nitric acid and the reaction with dilute nitric acid lies in the product of nitrogen reduction. In the case of concentrated HNO 3, nitrogen is reduced to a lesser extent: instead of nitric oxide (II), nitrogen oxide (IV) is formed, which is associated with greater competition between nitric acid molecules in concentrated acid for electrons of the reducing agent (Cu):
Cu + 4HNO 3 = Cu (NO 3) 2 + 2NO 2 + 2H 2 O
with oxides of non-metals
Copper reacts with some non-metal oxides. For example, with such oxides as NO 2, NO, N 2 O, copper is oxidized to copper (II) oxide, and nitrogen is reduced to oxidation state 0, i.e. a simple substance N 2 is formed:
In the case of sulfur dioxide, instead of a simple substance (sulfur), copper (I) sulfide is formed. This is due to the fact that copper with sulfur, unlike nitrogen, reacts:
with metal oxides
When sintering metallic copper with copper (II) oxide at a temperature of 1000-2000 ° C, copper (I) oxide can be obtained:
Also metallic copper can reduce iron (III) oxide to iron (II) oxide when calcined:
with metal salts
Copper displaces less active metals (to the right in the row of activity) from solutions of their salts:
Cu + 2AgNO 3 = Cu (NO 3) 2 + 2Ag ↓
An interesting reaction also takes place in which copper dissolves in the salt of a more active metal - iron in the +3 oxidation state. However, there are no contradictions, since copper does not displace iron from its salt, but only restores it from the +3 oxidation state to the +2 oxidation state:
Fe 2 (SO 4) 3 + Cu = CuSO 4 + 2FeSO 4
Cu + 2FeCl 3 = CuCl 2 + 2FeCl 2
The latter reaction is used in the manufacture of microcircuits at the stage of etching copper plates.
Corrosion of copper
Copper corrodes over time when it comes into contact with moisture, carbon dioxide and oxygen in the air:
2Cu + H 2 O + CO 2 + O 2 = (CuOH) 2 CO 3
As a result of this reaction, copper products are covered with a loose blue-green bloom of copper (II) hydroxycarbonate.
Zinc chemical properties
Zinc Zn is in the IIB group of the IV-th period. The electronic configuration of the valence orbitals of the atoms of a chemical element in the ground state is 3d 10 4s 2. For zinc, only one single oxidation state is possible, equal to +2. Zinc oxide ZnO and zinc hydroxide Zn (OH) 2 have pronounced amphoteric properties.
Zinc, when stored in air, tarnishes, covered with a thin layer of ZnO oxide. Oxidation proceeds especially easily at high humidity and in the presence of carbon dioxide due to the reaction:
2Zn + H 2 O + O 2 + CO 2 → Zn 2 (OH) 2 CO 3
Zinc vapor burns in air, and a thin strip of zinc, after heating in a burner flame, burns in it with a greenish flame:
When heated, zinc metal also interacts with halogens, sulfur, phosphorus:
Zinc does not directly react with hydrogen, nitrogen, carbon, silicon and boron.
Zinc reacts with non-oxidizing acids to release hydrogen:
Zn + H 2 SO 4 (20%) → ZnSO 4 + H 2
Zn + 2HCl → ZnCl 2 + H 2
Technical zinc is especially easily soluble in acids, since it contains impurities of other less active metals, in particular, cadmium and copper. High-purity zinc is resistant to acids for certain reasons. To speed up the reaction, a sample of high purity zinc is brought into contact with copper or a little copper salt is added to the acid solution.
At a temperature of 800-900 o C (red heat), metallic zinc, being in a molten state, interacts with superheated steam, releasing hydrogen from it:
Zn + H 2 O = ZnO + H 2
Zinc also reacts with oxidizing acids: concentrated sulfuric and nitric.
Zinc as an active metal can form sulfur dioxide, elemental sulfur and even hydrogen sulfide with concentrated sulfuric acid.
Zn + 2H 2 SO 4 = ZnSO 4 + SO 2 + 2H 2 O
The composition of the nitric acid reduction products is determined by the concentration of the solution:
Zn + 4HNO 3 (conc.) = Zn (NO 3) 2 + 2NO 2 + 2H 2 O
3Zn + 8HNO 3 (40%) = 3Zn (NO 3) 2 + 2NO + 4H 2 O
4Zn + 10HNO 3 (20%) = 4Zn (NO 3) 2 + N 2 O + 5H 2 O
5Zn + 12HNO 3 (6%) = 5Zn (NO 3) 2 + N 2 + 6H 2 O
4Zn + 10HNO 3 (0.5%) = 4Zn (NO 3) 2 + NH 4 NO 3 + 3H 2 O
The direction of the process is also influenced by the temperature, the amount of acid, the purity of the metal, and the reaction time.
Zinc reacts with alkali solutions to form tetrahydroxozincates and hydrogen:
Zn + 2NaOH + 2H 2 O = Na 2 + H 2
Zn + Ba (OH) 2 + 2H 2 O = Ba + H 2
When alloyed with anhydrous alkalis, zinc forms zincates and hydrogen:
In a highly alkaline environment, zinc is an extremely strong reducing agent capable of reducing nitrogen in nitrates and nitrites to ammonia:
4Zn + NaNO 3 + 7NaOH + 6H 2 O → 4Na 2 + NH 3
Due to complexation, zinc slowly dissolves in ammonia solution, reducing hydrogen:
Zn + 4NH 3 H 2 O → (OH) 2 + H 2 + 2H 2 O
Zinc also reduces less active metals (to the right of it in the row of activity) from aqueous solutions of their salts:
Zn + CuCl 2 = Cu + ZnCl 2
Zn + FeSO 4 = Fe + ZnSO 4
Chemical properties of chromium
Chromium is an element of the VIB group of the periodic table. The electronic configuration of the chromium atom is written as 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1, i.e. in the case of chromium, as well as in the case of the copper atom, the so-called "electron slip" is observed
The most common oxidation states of chromium are +2, +3 and +6. They should be remembered, and within the framework of the USE program in chemistry, it can be assumed that chromium has no other oxidation states.
Under normal conditions, chromium is resistant to corrosion both in air and in water.
Interaction with non-metals
with oxygen
Powdered metallic chromium heated to a temperature of more than 600 o C burns in pure oxygen to form chromium (III) oxide:
4Cr + 3O 2 = o t=> 2Cr 2 O 3
with halogens
Chromium reacts with chlorine and fluorine at lower temperatures than with oxygen (250 and 300 o C, respectively):
2Cr + 3F 2 = o t=> 2CrF 3
2Cr + 3Cl 2 = o t=> 2CrCl 3
Chromium reacts with bromine at the temperature of red heat (850-900 o C):
2Cr + 3Br 2 = o t=> 2CrBr 3
with nitrogen
Metallic chromium interacts with nitrogen at temperatures above 1000 o С:
2Cr + N 2 = ot=> 2CrN
with gray
With sulfur, chromium can form both chromium (II) sulfide and chromium (III) sulfide, which depends on the proportions of sulfur and chromium:
Cr + S = o t=> CrS
2Cr + 3S = o t=> Cr 2 S 3
Chromium does not react with hydrogen.
Interaction with complex substances
Interaction with water
Chromium refers to metals of average activity (located in the row of metal activity between aluminum and hydrogen). This means that the reaction takes place between red-hot chromium and superheated steam:
2Cr + 3H 2 O = o t=> Cr 2 O 3 + 3H 2
5interaction with acids
Chromium under normal conditions is passivated with concentrated sulfuric and nitric acids, however, it dissolves in them during boiling, while oxidizing to the oxidation state +3:
Cr + 6HNO 3 (conc.) = t o=> Cr (NO 3) 3 + 3NO 2 + 3H 2 O
2Cr + 6H 2 SO 4 (conc) = t o=> Cr 2 (SO 4) 3 + 3SO 2 + 6H 2 O
In the case of dilute nitric acid, the main product of nitrogen reduction is the simple substance N 2:
10Cr + 36HNO 3 (diluted) = 10Cr (NO 3) 3 + 3N 2 + 18H 2 O
Chromium is located in the row of activity to the left of hydrogen, which means that it is able to release H 2 from solutions of non-oxidizing acids. In the course of such reactions in the absence of air oxygen access, chromium (II) salts are formed:
Cr + 2HCl = CrCl 2 + H 2
Cr + H 2 SO 4 (dil.) = CrSO 4 + H 2
When the reaction is carried out in the open air, bivalent chromium is instantly oxidized by the oxygen contained in the air to the oxidation state +3. In this case, for example, the equation with hydrochloric acid will take the form:
4Cr + 12HCl + 3O 2 = 4CrCl 3 + 6H 2 O
When alloying metallic chromium with strong oxidants in the presence of alkalis, chromium is oxidized to the oxidation state +6, forming chromates:
Iron chemical properties
Iron Fe, a chemical element in the VIIIB group and having serial number 26 in the periodic table. The distribution of electrons in the iron atom is as follows 26 Fe1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2, that is, iron belongs to d-elements, since the d-sublevel is filled in its case. It is most characterized by two oxidation states +2 and +3. In oxide FeO and hydroxide Fe (OH) 2, the basic properties prevail, in oxide Fe 2 O 3 and hydroxide Fe (OH) 3 amphoteric properties are noticeably expressed. Thus, iron oxide and hydroxide (lll) dissolve to some extent during boiling in concentrated alkali solutions, and also react with anhydrous alkalis during fusion. It should be noted that the oxidation state of iron +2 is very unstable, and easily transforms into the oxidation state +3. Also known are iron compounds in the rare oxidation state +6 - ferrates, salts of non-existing "iron acid" H 2 FeO 4. These compounds are relatively stable only in the solid state, or in strongly alkaline solutions. With an insufficient alkalinity of the medium, ferrates quite quickly oxidize even water, releasing oxygen from it.
Interaction with simple substances
With oxygen
When burned in pure oxygen, iron forms the so-called iron scale, having the formula Fe 3 O 4 and is actually a mixed oxide, the composition of which can be conventionally represented by the formula FeO ∙ Fe 2 O 3. The combustion reaction of iron has the form:
3Fe + 2O 2 = t o=> Fe 3 O 4
With gray
When heated, iron reacts with sulfur to form ferrous sulfide:
Fe + S = t o=> FeS
Or with an excess of sulfur iron disulfide:
Fe + 2S = t o=> FeS 2
With halogens
With all halogens, except for iodine, metallic iron is oxidized to the oxidation state +3, forming iron halides (lll):
2Fe + 3F 2 = t o=> 2FeF 3 - iron fluoride (lll)
2Fe + 3Cl 2 = t o=> 2FeCl 3 - ferric chloride (lll)
Iodine, as the weakest oxidizing agent among halogens, oxidizes iron only to the oxidation state +2:
Fe + I 2 = t o=> FeI 2 - iron iodide (ll)
It should be noted that ferric iron compounds easily oxidize iodide ions in aqueous solution to free iodine I 2 while reducing to the oxidation state +2. Examples of similar reactions from the FIPI bank:
2FeCl 3 + 2KI = 2FeCl 2 + I 2 + 2KCl
2Fe (OH) 3 + 6HI = 2FeI 2 + I 2 + 6H 2 O
Fe 2 O 3 + 6HI = 2FeI 2 + I 2 + 3H 2 O
With hydrogen
Iron does not react with hydrogen (only alkali metals and alkaline earth metals react with hydrogen from metals):
Interaction with complex substances
5interaction with acids
With non-oxidizing acids
Since iron is located in the row of activity to the left of hydrogen, this means that it is able to displace hydrogen from non-oxidizing acids (almost all acids except H 2 SO 4 (conc.) And HNO 3 of any concentration):
Fe + H 2 SO 4 (dil.) = FeSO 4 + H 2
Fe + 2HCl = FeCl 2 + H 2
It is necessary to pay attention to such a trick in the tasks of the exam, as a question on the topic to what degree of oxidation iron will oxidize when it is exposed to dilute and concentrated hydrochloric acid. The correct answer is up to +2 in both cases.
The trap here lies in the intuitive expectation of a deeper oxidation of iron (up to s.d. +3) in the case of its interaction with concentrated hydrochloric acid.
Interaction with oxidizing acids
Iron does not react with concentrated sulfuric and nitric acids under normal conditions due to passivation. However, it reacts with them when boiled:
2Fe + 6H 2 SO 4 = o t=> Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O
Fe + 6HNO 3 = o t=> Fe (NO 3) 3 + 3NO 2 + 3H 2 O
Please note that dilute sulfuric acid oxidizes iron to the +2 oxidation state, and concentrated iron to +3.
Corrosion (rusting) of iron
Iron will rust very quickly in humid air:
4Fe + 6H 2 O + 3O 2 = 4Fe (OH) 3
Iron does not react with water in the absence of oxygen either under normal conditions or during boiling. The reaction with water takes place only at temperatures above the red heat temperature (> 800 o C). those..
1.2H 2SO 4 (conc.) + Cu = CuSO 4 + SO 2 + 2H 2O
copper sulfate
H 2SO 4 (dil.) + Zn = ZnSO 4 + H 2
zinc sulfate
2. FeO + H 2 = Fe + H 2O
CuSO 4 + Fe = Cu ↓ + FeSO 4
3. Let's compose the salts of nitric acid:
formula of nitric acid HNO3 acid residue NO3- - nitrate
Let's compose the salt formulas:
Na + NO3- According to the table of solubility, we determine the charges of the ions. Since the sodium ion and nitrate ion have charges "+" and "-", respectively, the subscripts in this formula are unnecessary. You get the following formula:
Na + NO3- - sodium nitrate
Ca2 + NO3- - According to the solubility table, we determine the charges of the ions. Let us arrange the indices according to the rule of the cross, but since the nitrate ion is a complex ion with a charge "-", it must be enclosed in brackets:
Ca2 + (NO3) -2 - calcium nitrate
Al3 + NO3- - According to the solubility table, we determine the charges of the ions. Let us arrange the indices according to the rule of the cross, but since the nitrate ion is a complex ion with a charge "-", it must be enclosed in brackets:
Al3 + (NO3) -3 - aluminum nitrate
further metals
zinc chloride ZnCl2
aluminum nitrate Al (NO3) 3
