copper compounds. Copper and its compounds Copper oxide 2 precipitate color
§one. Chemical properties of a simple substance (st. ok. = 0).
a) Relation to oxygen.
Unlike its subgroup neighbors, silver and gold, copper reacts directly with oxygen. Copper exhibits little activity towards oxygen, but in humid air it gradually oxidizes and becomes covered with a greenish film, consisting of basic copper carbonates:
In dry air, oxidation is very slow, a thin layer of copper oxide forms on the copper surface:
Outwardly, copper does not change, since copper (I) oxide, like copper itself, is pink. In addition, the oxide layer is so thin that it transmits light, i.e. shines through. In a different way, copper oxidizes when heated, for example, at 600-800 0 C. In the first seconds, oxidation goes to copper (I) oxide, which from the surface turns into black copper (II) oxide. A two-layer oxide coating is formed.
Q formation (Cu 2 O) = 84935 kJ.
Figure 2. The structure of the copper oxide film.
b) Interaction with water.
The metals of the copper subgroup are at the end of the electrochemical series of voltages, after the hydrogen ion. Therefore, these metals cannot displace hydrogen from water. At the same time, hydrogen and other metals can displace copper subgroup metals from solutions of their salts, for example:
This reaction is redox, as there is a transfer of electrons:
Molecular hydrogen displaces the metals of the copper subgroup with great difficulty. This is explained by the fact that the bond between hydrogen atoms is strong and a lot of energy is spent on breaking it. The reaction takes place only with hydrogen atoms.
Copper in the absence of oxygen practically does not interact with water. In the presence of oxygen, copper slowly reacts with water and becomes covered with a green film of copper hydroxide and basic carbonate:
c) Interaction with acids.
Being in a series of voltages after hydrogen, copper does not displace it from acids. Therefore, hydrochloric and dilute sulfuric acid do not act on copper.
However, in the presence of oxygen, copper dissolves in these acids to form the corresponding salts:
The only exception is hydroiodic acid, which reacts with copper to release hydrogen and form a very stable copper (I) complex:
2 Cu + 3 HI → 2 H[ CuI 2 ] + H 2
Copper also reacts with acids - oxidizing agents, for example, with nitric acid:
Cu+4HNO 3( conc .) → Cu(NO 3 ) 2 +2NO 2 +2H 2 O
3Cu + 8HNO 3( having diluted .) → 3Cu(NO 3 ) 2 +2NO+4H 2 O
And also with concentrated cold sulfuric acid:
Cu + H 2 SO 4(conc.) → CuO + SO 2 + H 2 O
With hot concentrated sulfuric acid :
Cu+2H 2 SO 4( conc ., hot ) → CuSO 4 + SO 2 + 2H 2 O
With anhydrous sulfuric acid at a temperature of 200 0 C, copper (I) sulfate is formed:
2Cu+2H 2 SO 4( anhydrous .) 200°C → Cu 2 SO 4 ↓+SO 2 + 2H 2 O
d) Relation to halogens and some other non-metals.
Q formation (CuCl) = 134300 kJ
Q formation (CuCl 2) = 111700 kJ
Copper reacts well with halogens, gives two types of halides: CuX and CuX 2 .. Under the action of halogens at room temperature, no visible changes occur, but a layer of adsorbed molecules first forms on the surface, and then a very thin layer of halides. When heated, the reaction with copper is very violent. We heat the copper wire or foil and lower it hot into a jar of chlorine - brown vapors will appear near the copper, consisting of copper (II) chloride CuCl 2 mixed with copper (I) chloride CuCl. The reaction occurs spontaneously due to the release of heat. Monovalent copper halides are obtained by reacting metallic copper with a solution of divalent copper halide, for example:
In this case, the monochloride precipitates out of solution in the form of a white precipitate on the copper surface.
Copper also reacts quite easily with sulfur and selenium when heated (300-400 ° C):
2Cu+S→Cu 2 S
2Cu+Se→Cu 2 Se
But copper does not react with hydrogen, carbon and nitrogen even at high temperatures.
e) Interaction with oxides of non-metals
When heated, copper can displace simple substances from some non-metal oxides (for example, sulfur (IV) oxide and nitrogen (II, IV) oxides), while forming a thermodynamically more stable copper (II) oxide):
4Cu+SO 2 600-800°C →2CuO + Cu 2 S
4Cu+2NO 2 500-600°C →4CuO + N 2
2 Cu+2 NO 500-600° C →2 CuO + N 2
§2. Chemical properties of monovalent copper (st.c. = +1)
In aqueous solutions, the Cu + ion is very unstable and disproportionate:
Cu + ↔ Cu 0 + Cu 2+
However, copper in the oxidation state (+1) can be stabilized in compounds with very low solubility or through complexation.
a) Copper oxide (I) Cu 2 O
amphoteric oxide. Brown-red crystalline substance. It occurs naturally as the mineral cuprite. It can be artificially obtained by heating a solution of a copper (II) salt with alkali and some strong reducing agent, for example, formalin or glucose. Copper(I) oxide does not react with water. Copper(I) oxide is transferred into a solution with concentrated hydrochloric acid to form a chloride complex:
Cu 2 O+4 HCl→2 H[ CuCl2]+ H 2 O
We also dissolve in a concentrated solution of ammonia and ammonium salts:
Cu 2 O+2NH 4 + →2 +
In dilute sulfuric acid, it disproportionates to divalent copper and metallic copper:
Cu 2 O+H 2 SO 4(dil.) →CuSO 4 + Cu 0 ↓+H 2 O
Also, copper(I) oxide enters into the following reactions in aqueous solutions:
1. Slowly oxidized by oxygen to copper (II) hydroxide:
2 Cu 2 O+4 H 2 O+ O 2 →4 Cu(Oh) 2 ↓
2. Reacts with dilute hydrohalic acids to form the corresponding copper(I) halides:
Cu 2 O+2 HG→2CuG↓ +H 2 O(G=Cl, Br, J)
3.Reduced to metallic copper with typical reducing agents, for example, sodium hydrosulfite in a concentrated solution:
2 Cu 2 O+2 NaSO 3 →4 Cu↓+ Na 2 SO 4 + H 2 SO 4
Copper(I) oxide is reduced to metallic copper in the following reactions:
1. When heated up to 1800 °C (decomposition):
2 Cu 2 O - 1800° C →2 Cu + O 2
2. When heated in a stream of hydrogen, carbon monoxide, aluminum and other typical reducing agents:
Cu 2 O+H 2 - >250°C →2Cu+H 2 O
Cu 2 O+CO - 250-300°C →2Cu+CO 2
3 Cu 2 O + 2 Al - 1000° C →6 Cu + Al 2 O 3
Also, at high temperatures, copper (I) oxide reacts:
1. With ammonia (copper(I) nitride is formed)
3 Cu 2 O + 2 NH 3 - 250° C →2 Cu 3 N + 3 H 2 O
2. With alkali metal oxides:
Cu 2 O+M 2 O- 600-800°C →2 MCuO (M= Li, Na, K)
In this case, cuprates of copper (I) are formed.
Copper(I) oxide reacts markedly with alkalis:
Cu 2 O+2 NaOH (conc.) + H 2 O↔2 Na[ Cu(Oh) 2 ]
b) Copper hydroxide (I) CuOH
Copper(I) hydroxide forms a yellow substance and is insoluble in water.
Easily decomposes when heated or boiled:
2 CuOH → Cu 2 O + H 2 O
c) HalidesCuF, CuFROMl, CuBrandCuJ
All these compounds are white crystalline substances, poorly soluble in water, but readily soluble in excess NH 3 , cyanide ions, thiosulfate ions, and other strong complexing agents. Iodine forms only the compound Cu +1 J. In the gaseous state, cycles of the (CuГ) 3 type are formed. Reversibly soluble in the corresponding hydrohalic acids:
CuG + HG ↔H[ CuG 2 ] (G=Cl, Br, J)
Copper (I) chloride and bromide are unstable in moist air and gradually turn into basic copper (II) salts:
4 CuD +2H 2 O + O 2 →4 Cu(Oh)G (G=Cl, Br)
d) Other copper compounds (I)
1. Copper (I) acetate (CH 3 COOCu) - a copper compound, has the form of colorless crystals. In water, it slowly hydrolyzes to Cu 2 O, in air it oxidizes to divalent copper acetate; Receive CH 3 COOSu by reduction (CH 3 COO) 2 Cu with hydrogen or copper, sublimation (CH 3 COO) 2 Cu in a vacuum or interaction (NH 3 OH)SO 4 with (CH 3 COO) 2 Cu in p-re in the presence of H 3 COOH 3 . The substance is toxic.
2. Copper(I) acetylenide - red-brown, sometimes black crystals. When dry, the crystals detonate on impact or heat. Wet resistant. Detonation in the absence of oxygen produces no gaseous substances. Decomposes under the action of acids. It is formed as a precipitate when acetylene is passed into ammonia solutions of copper(I) salts:
FROM 2 H 2 +2[ Cu(NH 3 ) 2 ](Oh) → Cu 2 C 2 ↓ +2 H 2 O+2 NH 3
This reaction is used for the qualitative detection of acetylene.
3. Copper nitride - an inorganic compound with the formula Cu 3 N, dark green crystals.
Decomposes on heating:
2 Cu 3 N - 300° C →6 Cu + N 2
Reacts violently with acids:
2 Cu 3 N +6 HCl - 300° C →3 Cu↓ +3 CuCl 2 +2 NH 3
§3. Chemical properties of bivalent copper (st.c. = +2)
The most stable oxidation state of copper and the most characteristic of it.
a) Copper oxide (II) CuO
CuO is the basic oxide of divalent copper. Black crystals, under normal conditions quite stable, practically insoluble in water. In nature, it occurs in the form of the mineral tenorite (melaconite) of black color. Copper(II) oxide reacts with acids to form the corresponding salts of copper(II) and water:
CuO + 2 HNO 3 → Cu(NO 3 ) 2 + H 2 O
When CuO is fused with alkalis, cuprates of copper (II) are formed:
CuO+2 KOH- t ° → K 2 CuO 2 + H 2 O
When heated to 1100 °C, it decomposes:
4CuO- t ° →2 Cu 2 O + O 2
b) Copper (II) hydroxideCu(Oh) 2
Copper(II) hydroxide is a blue amorphous or crystalline substance, practically insoluble in water. When heated to 70-90 ° C, Cu (OH) 2 powder or its aqueous suspensions decompose to CuO and H 2 O:
Cu(Oh) 2 → CuO + H 2 O
It is an amphoteric hydroxide. Reacts with acids to form water and the corresponding copper salt:
It does not react with dilute alkali solutions, but dissolves in concentrated ones, forming bright blue tetrahydroxocuprates (II):
Copper (II) hydroxide with weak acids forms basic salts. It dissolves very easily in excess ammonia to form copper ammonia:
Cu(OH) 2 +4NH 4 OH→(OH) 2 +4H 2 O
Copper ammonia has an intense blue-violet color, so it is used in analytical chemistry to determine small amounts of Cu 2+ ions in solution.
c) Copper salts (II)
Simple salts of copper (II) are known for most anions, except for cyanide and iodide, which, when interacting with the Cu 2+ cation, form covalent copper (I) compounds that are insoluble in water.
Copper salts (+2) are mostly water soluble. The blue color of their solutions is associated with the formation of the 2+ ion. They often crystallize as hydrates. Thus, tetrahydrate crystallizes from an aqueous solution of copper (II) chloride below 15 0 C, trihydrate at 15-26 0 C, and dihydrate above 26 0 C. In aqueous solutions, copper(II) salts are subject to hydrolysis to a small extent, and basic salts often precipitate out of them.
1. Copper (II) sulfate pentahydrate (copper sulfate)
CuSO 4 * 5H 2 O, called copper sulphate, is of the greatest practical importance. Dry salt has a blue color, however, when slightly heated (200 0 C), it loses water of crystallization. Anhydrous white salt. Upon further heating to 700 0 C, it turns into copper oxide, losing sulfur trioxide:
CuSO 4 -- t ° → CuO+ SO 3
Copper sulphate is prepared by dissolving copper in concentrated sulfuric acid. This reaction is described in the section "Chemical Properties of a Simple Substance". Copper sulfate is used in the electrolytic production of copper, in agriculture to control pests and plant diseases, and to obtain other copper compounds.
2. Copper (II) chloride dihydrate.
These are dark green crystals, easily soluble in water. Concentrated solutions of copper chloride are green, and dilute solutions are blue. This is due to the formation of a green chloride complex:
Cu 2+ +4 Cl - →[ CuCl 4 ] 2-
And its further destruction and the formation of a blue aquacomplex.
3. Copper (II) nitrate trihydrate.
Blue crystalline solid. Obtained by dissolving copper in nitric acid. When heated, the crystals first lose water, then decompose with the release of oxygen and nitrogen dioxide, turning into copper (II) oxide:
2Cu(NO 3 ) 2 -- t° →2CuO+4NO 2 +O 2
4. Hydroxomedi(II) carbonate.
Copper carbonates are unstable and almost never used in practice. Of some importance for the production of copper is only the basic copper carbonate Cu 2 (OH) 2 CO 3, which occurs in nature in the form of the mineral malachite. When heated, it easily decomposes with the release of water, carbon monoxide (IV) and copper oxide (II):
Cu 2 (OH) 2 CO 3 -- t° →2CuO+H 2 O+CO 2
§4. Chemical properties of trivalent copper (st.c. = +3)
This oxidation state is the least stable for copper, and therefore copper(III) compounds are the exception rather than the "rule". However, some trivalent copper compounds exist.
a) Copper oxide (III) Cu 2 O 3
It is a crystalline substance, dark garnet color. Does not dissolve in water.
Obtained by oxidation of copper (II) hydroxide with potassium peroxodisulfate in an alkaline medium at low temperatures:
2Cu(OH) 2 +K 2 S 2 O 8 +2KOH -- -20°C →Cu 2 O 3 ↓+2K 2 SO 4 +3H 2 O
This substance decomposes at a temperature of 400 0 C:
Cu 2 O 3 -- t ° →2 CuO+ O 2
Copper(III) oxide is a strong oxidizing agent. When interacting with hydrogen chloride, chlorine is reduced to free chlorine:
Cu 2 O 3 +6 HCl-- t ° →2 CuCl 2 + Cl 2 +3 H 2 O
b) Copper cuprates (W)
These are black or blue substances, they are not stable in water, they are diamagnetic, the anion is a ribbon of squares (dsp 2). Formed by the interaction of copper (II) hydroxide and alkali metal hypochlorite in an alkaline environment:
2 Cu(Oh) 2 + MClO + 2 NaOH→2MCuO 3 + NaCl +3 H 2 O (M= Na- Cs)
c) Potassium hexafluorocuprate(III)
Green substance, paramagnetic. Octahedral structure sp 3 d 2 . Copper fluoride complex CuF 3, which decomposes in the free state at -60 0 C. It is formed by heating a mixture of potassium and copper chlorides in a fluorine atmosphere:
3KCl + CuCl + 3F 2 → K 3 + 2Cl 2
Decomposes water with the formation of free fluorine.
§five. Copper compounds in oxidation state (+4)
So far, only one substance is known to science, where copper is in the +4 oxidation state, this is cesium hexafluorocuprate (IV) - Cs 2 Cu +4 F 6 - an orange crystalline substance, stable in glass ampoules at 0 0 C. It reacts violently with water. Obtained by fluorination at high pressure and temperature of a mixture of cesium and copper chlorides:
CuCl 2 +2CsCl +3F 2 -- t ° p → Cs 2 CuF 6 +2Cl 2
Copper (Cu) belongs to the d-elements and is located in the IB group of the periodic table of D.I. Mendeleev. The electronic configuration of the copper atom in the ground state is written as 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 instead of the expected 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 jump” from the 4s sublevel to the 3d sublevel is observed. For copper, in addition to zero, oxidation states +1 and +2 are possible. The +1 oxidation state 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 are white, and Cu 2 S is black-blue. More chemically stable is the oxidation state of copper, equal to +2. Salts containing copper in a given oxidation state are blue and blue-green in color.
Copper is a very soft, malleable and ductile metal with high electrical and thermal conductivity. The color of metallic copper is red-pink. Copper is in the activity series of metals to the right of hydrogen, i.e. refers to low-active metals.
with oxygen
Under normal conditions, copper does not interact with oxygen. Heat is required for the reaction between them to proceed. Depending on the excess or lack of oxygen and temperature conditions, it can form copper (II) oxide and copper (I) oxide:
with sulfur
The reaction of sulfur with copper, depending on the conditions of carrying out, 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, a simpler way to obtain copper (II) sulfide from simple substances is the interaction of copper with sulfur dissolved in carbon disulfide:
This reaction proceeds 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 activity series up 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 dilute nitric acid
The reaction of copper with dilute HNO 3 leads to the formation of copper (II) nitrate and nitrogen monoxide:
3Cu + 8HNO 3 (diff.) = 3Cu(NO 3) 2 + 2NO + 4H 2 O
- with concentrated nitric acid
Concentrated HNO 3 readily reacts with copper under normal conditions. The difference between the reaction of copper with concentrated nitric acid and the interaction 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), nitric oxide (IV) is formed, which is associated with greater competition between nitric acid molecules in concentrated acid for the electrons of the reducing agent (Cu):
Cu + 4HNO 3 \u003d Cu (NO 3) 2 + 2NO 2 + 2H 2 O
with non-metal oxides
Copper reacts with some non-metal oxides. For example, with oxides such 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 oxide (II) at a temperature of 1000-2000 ° C, copper oxide (I) can be obtained:
Also, metallic copper can reduce iron (III) oxide upon calcination to iron (II) oxide:
with metal salts
Copper displaces less active metals (to the right of it in the activity series) from solutions of their salts:
Cu + 2AgNO 3 \u003d Cu (NO 3) 2 + 2Ag ↓
An interesting reaction also takes place, in which copper is dissolved in a salt of a more active metal - iron in the +3 oxidation state. However, there are no contradictions, because 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 \u003d CuSO 4 + 2FeSO 4
Cu + 2FeCl 3 = CuCl 2 + 2FeCl 2
The latter reaction is used in the production of microcircuits at the stage of etching of copper boards.
Corrosion of copper
Copper corrodes over time when exposed to moisture, carbon dioxide and atmospheric oxygen:
2Cu + H 2 O + CO 2 + O 2 \u003d (CuOH) 2 CO 3
As a result of this reaction, copper products are covered with a loose blue-green coating of copper (II) hydroxocarbonate.
Chemical properties of zinc
Zinc Zn is in the IIB group of the IVth period. Electronic configuration of valence orbitals of atoms of a chemical element in the ground state 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 tarnishes when stored in air, becoming 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 glowing in a burner flame, burns in it with a greenish flame:
When heated, metallic zinc 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
Industrial 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 small amount of 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 water vapor, releasing hydrogen from it:
Zn + H 2 O \u003d 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 \u003d ZnSO 4 + SO 2 + 2H 2 O
The composition of the products of nitric acid reduction 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 affected 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 \u003d Na 2 + H 2
Zn + Ba (OH) 2 + 2H 2 O \u003d Ba + H 2
With anhydrous alkalis, zinc, when fused, 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 an ammonia solution, reducing hydrogen:
Zn + 4NH 3 H 2 O → (OH) 2 + H 2 + 2H 2 O
Zinc also restores less active metals (to the right of it in the activity series) from aqueous solutions of their salts:
Zn + CuCl 2 \u003d Cu + ZnCl 2
Zn + FeSO 4 \u003d 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 frequently exhibited oxidation states of chromium are +2, +3 and +6. They should be remembered, and within the framework of the USE program in chemistry, we can assume 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
Heated to a temperature of more than 600 o C, powdered metallic chromium 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 a red heat temperature (850-900 o C):
2Cr + 3Br 2 = o t=> 2CrBr 3
with nitrogen
Metallic chromium interacts with nitrogen at temperatures above 1000 o C:
2Cr + N 2 = ot=> 2CrN
with sulfur
With sulfur, chromium can form both chromium (II) sulfide and chromium (III) sulfide, depending 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 belongs to the metals of medium activity (located in the activity series of metals between aluminum and hydrogen). This means that the reaction proceeds between red-hot chromium and superheated water vapor:
2Cr + 3H 2 O = o t=> Cr 2 O 3 + 3H 2
Interaction with acids
Chromium is passivated under normal conditions with concentrated sulfuric and nitric acids, however, it dissolves in them during boiling, while being oxidized to an oxidation state of +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 a simple substance N 2:
10Cr + 36HNO 3 (razb) \u003d 10Cr (NO 3) 3 + 3N 2 + 18H 2 O
Chromium is located in the activity series 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 access to atmospheric oxygen, chromium (II) salts are formed:
Cr + 2HCl \u003d CrCl 2 + H 2
Cr + H 2 SO 4 (razb.) \u003d CrSO 4 + H 2
When carrying out the reaction in the open air, divalent chromium is instantly oxidized by oxygen contained in the air to an oxidation state of +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 chromium metal is fused with strong oxidizing agents in the presence of alkalis, chromium is oxidized to an oxidation state of +6, forming chromates:
Chemical properties of iron
Iron Fe, a chemical element in group VIIIB and having serial number 26 in the periodic table. The distribution of electrons in an 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 characteristic of two oxidation states +2 and +3. FeO oxide and Fe(OH) 2 hydroxide are dominated by basic properties, Fe 2 O 3 oxide and Fe(OH) 3 hydroxide are markedly amphoteric. So the oxide and hydroxide of iron (lll) dissolve to some extent when boiled in concentrated solutions of alkalis, and also react with anhydrous alkalis during fusion. It should be noted that the oxidation state of iron +2 is very unstable, and easily passes into the oxidation state +3. Iron compounds are also known in a rare oxidation state of +6 - ferrates, salts of the non-existent "iron acid" H 2 FeO 4. These compounds are relatively stable only in the solid state or in strongly alkaline solutions. With insufficient alkalinity of the medium, ferrates 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 actually representing a mixed oxide, the composition of which can be conditionally 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 sulfur
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=> FeS2
With halogens
With all halogens except iodine, metallic iron is oxidized to an oxidation state of +3, forming iron halides (lll):
2Fe + 3F 2 = t o=> 2FeF 3 - iron fluoride (lll)
2Fe + 3Cl 2 = t o=> 2FeCl 3 - iron chloride (lll)
Iodine, as the weakest oxidizing agent among halogens, oxidizes iron only to the +2 oxidation state:
Fe + I 2 = t o=> FeI 2 - iron iodide (ll)
It should be noted that ferric iron compounds easily oxidize iodide ions in an aqueous solution to free iodine I 2 while recovering to the +2 oxidation state. 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 \u003d 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
Interaction with acids
With non-oxidizing acids
Since iron is located in the activity series 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 (diff.) \u003d FeSO 4 + H 2
Fe + 2HCl \u003d 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 be oxidized 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.o. +3) in the case of its interaction with concentrated hydrochloric acid.
Interaction with oxidizing acids
Under normal conditions, iron does not react with concentrated sulfuric and nitric acids 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
Note that dilute sulfuric acid oxidizes iron to an oxidation state of +2, and concentrated to +3.
Corrosion (rusting) of iron
In moist air, iron rusts very quickly:
4Fe + 6H 2 O + 3O 2 \u003d 4Fe (OH) 3
Iron does not react with water in the absence of oxygen either under normal conditions or when boiled. The reaction with water proceeds only at a temperature above the red heat temperature (> 800 ° C). those..
There are a lot of representatives of each of them, but oxides undoubtedly occupy the leading position. One chemical element can have several different binary compounds with oxygen at once. Copper also has this property. She has three oxides. Let's look at them in more detail.
Copper(I) oxide
Its formula is Cu 2 O. In some sources, this compound may be called copper hemioxide, dicopper oxide, or cuprous oxide.
Properties
It is a crystalline substance having a brown-red color. This oxide is insoluble in water and ethanol. It can melt without decomposing at a temperature of just over 1240 ° C. This substance does not interact with water, but can be transferred into solution if the participants in the reaction with it are concentrated hydrochloric acid, alkali, nitric acid, ammonia hydrate, ammonium salts, sulfuric acid .
Obtaining copper oxide (I)
It can be obtained by heating metallic copper, or in an environment where oxygen has a low concentration, as well as in a stream of certain nitrogen oxides and together with copper (II) oxide. In addition, it can become a reaction product of the thermal decomposition of the latter. Copper (I) oxide will also be obtained if copper (I) sulfide is heated in a stream of oxygen. There are other, more complex ways to obtain it (for example, the reduction of one of the copper hydroxides, the ion exchange of any monovalent copper salt with alkali, etc.), but they are practiced only in laboratories.
Application
Needed as a pigment when painting ceramics, glass; component of paints that protect the underwater part of the vessel from fouling. Also used as a fungicide. Copper oxide valves cannot do without it.
Copper(II) oxide
Its formula is CuO. In many sources it can be found under the name of copper oxide.
Properties
It is the highest copper oxide. The substance has the appearance of black crystals, which are almost insoluble in water. It reacts with acid and during this reaction forms the corresponding salt of divalent copper, as well as water. When it is fused with alkali, the reaction products are represented by cuprates. The decomposition of copper oxide (II) occurs at a temperature of about 1100 o C. Ammonia, carbon monoxide, hydrogen and coal are able to extract metallic copper from this compound.
Receipt
It can be obtained by heating metallic copper in air under one condition - the heating temperature must be below 1100 ° C. Copper (II) oxide can also be obtained by heating carbonate, nitrate, divalent copper hydroxide.
Application
With the help of this oxide, enamel and glass are colored green or blue, and a copper-ruby variety of the latter is also produced. In the laboratory, this oxide is used to discover the reducing properties of substances.
Copper(III) oxide
Its formula is Cu 2 O 3. It has a traditional name, which probably sounds a little unusual - copper oxide.
Properties
It has the appearance of red crystals that do not dissolve in water. The decomposition of this substance occurs at a temperature of 400 ° C, the products of this reaction are copper (II) oxide and oxygen.
Receipt
It can be obtained by oxidizing divalent copper hydroxide with potassium peroxydisulphate. A necessary condition for the reaction is an alkaline environment in which it must occur.
Application
This substance is not used by itself. In science and industry, the products of its decomposition - copper (II) oxide and oxygen - are more widely used.
Conclusion
That's all copper oxides. There are several of them due to the fact that copper has a variable valence. There are other elements that have several oxides, but we'll talk about them another time.
Cuprum (Cu) is one of the low-active metals. It is characterized by the formation of chemical compounds with oxidation states +1 and +2. So, for example, two oxides, which are a compound of two elements Cu and oxygen O: with an oxidation state of +1 - copper oxide Cu2O and an oxidation state of +2 - copper oxide CuO. Despite the fact that they consist of the same chemical elements, but each of them has its own special characteristics. In the cold, the metal interacts very weakly with atmospheric oxygen, becoming covered with a film, which is copper oxide, which prevents further oxidation of cuprum. When heated, this simple substance with serial number 29 in the periodic table is completely oxidized. In this case, copper (II) oxide is also formed: 2Cu + O2 → 2CuO.
The nitrous oxide is a brownish red solid with a molar mass of 143.1 g/mol. The compound has a melting point of 1235°C, a boiling point of 1800°C. It is insoluble in water, but soluble in acids. Copper (I) oxide is diluted in (concentrated), and a colorless complex + is formed, which is easily oxidized in air to a blue-violet ammonium complex 2+, which dissolves in hydrochloric acid to form CuCl2. In the history of semiconductor physics, Cu2O is one of the most studied materials.
Copper(I) oxide, also known as hemioxide, has basic properties. It can be obtained by metal oxidation: 4Cu + O2 → 2 Cu2O. Impurities such as water and acids affect the rate of this process as well as further oxidation to the divalent oxide. Copper oxide can dissolve in this form pure metal and salt: H2SO4 + Cu2O → Cu + CuSO4 + H2O. According to a similar scheme, an oxide with a degree of +1 interacts with other oxygen-containing acids. In the interaction of hemioxide with halogen-containing acids, monovalent metal salts are formed: 2HCl + Cu2O → 2CuCl + H2O.
Oxide of copper (I) occurs in nature in the form of red ore (this is an outdated name, along with such as ruby Cu), called the mineral "Cuprite". It takes a long time to educate. It can be produced artificially at high temperatures or under high oxygen pressure. Hemioxide is commonly used as a fungicide, as a pigment, as an antifouling agent in underwater or marine paint, and as a catalyst.
However, the effect of this substance with the chemical formula Cu2O on the body can be dangerous. If inhaled, it causes dyspnoea, coughing, and ulceration and perforation of the respiratory tract. If ingested, it irritates the gastrointestinal tract, which is accompanied by vomiting, pain and diarrhea.
H2 + CuO → Cu + H2O;
CO + CuO → Cu + CO2.
Copper(II) oxide is used in ceramics (as a pigment) to produce glazes (blue, green, and red, and sometimes pink, gray, or black). It is also used as a dietary supplement in animals to reduce cuprum deficiency in the body. It is an abrasive material that is necessary for polishing optical equipment. It is used for the production of dry cells, for the production of other Cu salts. The CuO compound is also used in the welding of copper alloys.
Exposure to the chemical compound CuO can also be dangerous to the human body. Causes lung irritation if inhaled. Copper(II) oxide can cause metal vapor fever (MFF). Cu oxide provokes a change in skin color, vision problems may appear. When ingested, like hemioxide, it leads to poisoning, which is accompanied by symptoms in the form of vomiting and pain.
COPPER AND ITS COMPOUNDS
LESSON IN THE 11th NATURAL SCIENCE CLASS
To increase the cognitive activity and independence of students, we use the lessons of the collective study of the material. At such lessons, each student (or a pair of students) receives a task, the completion of which he must report on in the same lesson, and his report is recorded by the rest of the class in notebooks and is an element of the content of the lesson's educational material. Each student contributes to the study of the topic by the class.
During the lesson, the mode of work of students changes from intraactive (a mode in which information flows are closed within the students, typical for independent work) to interactive (a mode in which information flows are two-way, i.e. information goes both from the student and to the student, information is exchanged). At the same time, the teacher acts as the organizer of the process, corrects and supplements the information provided by the students.
The lessons of collective study of the material consist of the following stages:
1st stage - installation, in which the teacher explains the goals and program of work in the lesson (up to 7 minutes);
Stage 2 - independent work of students according to the instructions (up to 15 minutes);
Stage 3 - exchange of information and summing up the lesson (takes all the remaining time).
The lesson "Copper and its compounds" is designed for classes with an in-depth study of chemistry (4 hours of chemistry per week), is held for two academic hours, the lesson updates students' knowledge on the following topics: "General properties of metals", "Attitude towards metals with concentrated sulfuric acid, nitric acid", "Qualitative reactions to aldehydes and polyhydric alcohols", "Oxidation of saturated monohydric alcohols with copper (II) oxide", "Complex compounds".
Before the lesson, students receive homework: to review the topics listed. The preliminary preparation of the teacher for the lesson consists in compiling instructional cards for students and preparing sets for laboratory experiments.
DURING THE CLASSES
Installation stage
The teacher puts in front of the students the purpose of the lesson: based on existing knowledge about the properties of substances, predict, confirm in practice, generalize information about copper and its compounds.
Students make up the electronic formula of the copper atom, find out what oxidation states copper can exhibit in compounds, what properties (redox, acid-base) copper compounds will have.
A table appears in the students' notebooks.
Properties of copper and its compounds
| Metal | Cu 2 O - basic oxide | CuO - basic oxide |
| Reducing agent | CuOH is an unstable base | Cu (OH) 2 - insoluble base |
| CuCl - insoluble salt | CuSO 4 - soluble salt | |
| Possess redox duality | Oxidizers |
Stage of independent work
To confirm and supplement the assumptions, students perform laboratory experiments according to the instructions and write down the equations of the reactions performed.
Instructions for independent work in pairs
1. Ignite the copper wire in a flame. Note how its color has changed. Place the hot calcined copper wire in ethyl alcohol. Note the change in its color. Repeat these manipulations 2-3 times. Check if the smell of ethanol has changed.
Write down two reaction equations corresponding to the transformations carried out. What properties of copper and its oxide are confirmed by these reactions?2. Add hydrochloric acid to copper(I) oxide.
What are you watching? Write down the reaction equations, given that copper (I) chloride is an insoluble compound. What properties of copper(I) are confirmed by these reactions?3. a) Place a zinc granule into the copper(II) sulfate solution. If no reaction occurs, heat the solution. b) Add 1 ml of sulfuric acid to copper (II) oxide and heat.
What are you watching? Write down the reaction equations. What properties of copper compounds are confirmed by these reactions?4. Place a universal indicator strip into the copper(II) sulfate solution.
Explain the result. Write down the ionic equation of hydrolysis for the first stage.
Add a solution of honey(II) sulfate to a solution of sodium carbonate.
What are you watching? Write the equation for the reaction of joint hydrolysis in molecular and ionic forms.5.
What are you watching?
Add ammonia solution to the resulting precipitate.
What changes have taken place? Write down the reaction equations. What properties of copper compounds are proved by the reactions carried out?6. Add a solution of potassium iodide to copper(II) sulfate.
What are you watching? Write an equation for the reaction. What property of copper(II) does this reaction prove?7. Place a small piece of copper wire into a test tube with 1 ml of concentrated nitric acid. Close the tube with a stopper.
What are you watching? (Take the test tube under draft.) Write down the reaction equation.
Pour hydrochloric acid into another test tube, place a small piece of copper wire in it.
What are you watching? Explain your observations. What properties of copper are confirmed by these reactions?8. Add an excess of sodium hydroxide to copper(II) sulfate.
What are you watching? Heat up the precipitate. What happened? Write down the reaction equations. What properties of copper compounds are confirmed by these reactions?9. Add an excess of sodium hydroxide to copper(II) sulfate.
What are you watching?
Add a solution of glycerin to the resulting precipitate.
What changes have taken place? Write down the reaction equations. What properties of copper compounds prove these reactions?10. Add an excess of sodium hydroxide to copper(II) sulfate.
What are you watching?
Pour the glucose solution to the resulting precipitate and heat.
What happened? Write the reaction equation using the general formula for aldehydes to denote glucoseWhat property of the copper compound does this reaction prove?
11. Add to copper(II) sulfate: a) ammonia solution; b) sodium phosphate solution.
What are you watching? Write down the reaction equations. What properties of copper compounds are proved by the reactions carried out?
Phase of communication and debriefing
The teacher asks a question concerning the properties of a particular substance. The students who performed the corresponding experiments report on the experiment and write down the reaction equations on the blackboard. Then the teacher and students complete the information about the chemical properties of the substance, which could not be confirmed by reactions in the conditions of the school laboratory.
The order of discussion of the chemical properties of copper compounds
1. How does copper react with acids, what other substances can copper react with?
The reactions of copper are written with:
Concentrated and dilute nitric acid:
Cu + 4HNO 3 (conc.) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O,
3Cu + 8HNO 3 (diff.) = 3Cu(NO 3) 2 + 2NO + 4H 2 O;
Concentrated sulfuric acid:
Cu + 2H 2 SO 4 (conc.) = CuSO 4 + SO 2 + 2H 2 O;
Oxygen:
2Cu + O 2 \u003d 2CuO;
Cu + Cl 2 \u003d CuCl 2;
Hydrochloric acid in the presence of oxygen:
2Cu + 4HCl + O 2 = 2CuCl 2 + 2H 2 O;
Iron(III) chloride:
2FeCl 3 + Cu \u003d CuCl 2 + 2FeCl 2.
2. What are the properties of copper(I) oxide and chloride?
Attention is drawn to the main properties, the ability to complex formation, redox duality. The equations of reactions of copper (I) oxide with:
Hydrochloric acid to form CuCl:
Cu 2 O + 2HCl = 2CuCl + H 2 O;
Excess HCl:
CuCl + HCl = H;
Reactions of reduction and oxidation of Cu 2 O:
Cu 2 O + H 2 \u003d 2Cu + H 2 O,
2Cu 2 O + O 2 \u003d 4CuO;
Disproportionation when heated:
Cu 2 O \u003d Cu + CuO,
2CuCl \u003d Cu + CuCl 2.
3. What are the properties of copper(II) oxide?
Attention is drawn to the basic and oxidizing properties. Equations for the reactions of copper(II) oxide with:
Acid:
CuO + 2H + = Cu 2+ + H 2 O;
Ethanol:
C 2 H 5 OH + CuO = CH 3 CHO + Cu + H 2 O;
Hydrogen:
CuO + H 2 \u003d Cu + H 2 O;
Aluminum:
3CuO + 2Al \u003d 3Cu + Al 2 O 3.
4. What are the properties of copper(II) hydroxide?
Attention is drawn to the oxidizing, basic properties, the ability to complex with organic and inorganic compounds. The reaction equations are written with:
Aldehyde:
RCHO + 2Cu(OH) 2 = RCOOH + Cu 2 O + 2H 2 O;
Acid:
Cu(OH) 2 + 2H + = Cu 2+ + 2H 2 O;
Ammonia:
Cu (OH) 2 + 4NH 3 \u003d (OH) 2;
Glycerin:
Decomposition reaction equation:
Cu (OH) 2 \u003d CuO + H 2 O.
5. What are the properties of copper(II) salts?
Attention is drawn to the reactions of ion exchange, hydrolysis, oxidizing properties, complexation. The equations for the reactions of copper sulfate are written with:
Sodium hydroxide:
Cu 2+ + 2OH - \u003d Cu (OH) 2;
Sodium Phosphate:
3Cu 2+ + 2= Cu 3 (PO 4) 2;
Cu 2+ + Zn \u003d Cu + Zn 2+;
Potassium iodide:
2CuSO 4 + 4KI = 2CuI + I 2 + 2K 2 SO 4 ;
Ammonia:
Cu 2+ + 4NH 3 \u003d 2+;
and reaction equations:
Hydrolysis:
Cu 2+ + HOH = CuOH + + H + ;
Co-hydrolysis with sodium carbonate to form malachite:
2Cu 2+ + 2 + H 2 O \u003d (CuOH) 2 CO 3 + CO 2.
In addition, you can tell students about the interaction of copper(II) oxide and hydroxide with alkalis, which proves their amphotericity:
Cu (OH) 2 + 2NaOH (conc.) \u003d Na 2,
Cu + Cl 2 \u003d CuCl 2,
Cu + HgCl 2 \u003d CuCl 2 + Hg,
2Cu + 4HCl + O 2 = 2CuCl 2 + 2H 2 O,
CuO + 2HCl \u003d CuCl 2 + H 2 O,
Cu(OH) 2 + 2HCl = CuCl 2 + 2H 2 O,
CuBr 2 + Cl 2 \u003d CuCl 2 + Br 2,
(CuOH) 2 CO 3 + 4HCl \u003d 2CuCl 2 + 3H 2 O + CO 2,
2CuCl + Cl 2 \u003d 2CuCl 2,
2CuCl \u003d CuCl 2 + Cu,
CuSO 4 + BaCl 2 \u003d CuCl 2 + BaSO 4.)
Exercise 3 Make chains of transformations corresponding to the following schemes and carry them out:
Task 1.
An alloy of copper and aluminum was treated first with an excess of alkali and then with an excess of dilute nitric acid. Calculate the mass fractions of metals in the alloy, if it is known that the volumes of gases released in both reactions (under the same conditions) are equal to each other
.
(Answer . Mass fraction of copper - 84%.)
Task 2. On calcination of 6.05 g of hydrated copper(II) nitrate, 2 g of residue was obtained. Determine the formula of the original salt.
(Answer. Cu(NO 3) 2 3H 2 O.)
Task 3. A copper plate weighing 13.2 g was lowered into 300 g of an iron (III) nitrate solution with a mass fraction of salt of 0.112. When it was taken out, it turned out that the mass fraction of iron(III) nitrate became equal to the mass fraction of the formed copper(II) salt. Determine the mass of the plate after it has been removed from the solution.
(Answer. 10 y.)
Homework. Learn the material written in the notebook. Compose a chain of transformations for copper compounds, containing at least ten reactions, and carry it out.
LITERATURE
1. Puzakov S.A., Popkov V.A. A manual on chemistry for university students. Programs. Questions, exercises, tasks. Samples of exam papers. M.: Higher school, 1999, 575 p.
2. Kuzmenko N.E., Eremin V.V. 2000 tasks and exercises in chemistry. For schoolchildren and entrants. M.: 1st Federal Book Trade Company, 1998, 512 p.
