Copper (II) oxide, properties, production, chemical reactions. Copper and its compounds Copper oxide 2 chemical formula
Chemical properties copper(II) oxide
Brief description of copper oxide (II):
copper oxide(ii)- inorganic matter black color.
2. reaction of copper (II) oxide with carbon:
CuO + C → Cu + CO (t = 1200 o C).
carbon.
3.copper oxide reaction(II) with gray:
CuO + 2S → Cu + S 2 O (t = 150-200 o C).
The reaction takes place in a vacuum. As a result of the reaction, copper and oxide are formed sulfur.
4. copper oxide reaction(II) with aluminum:
3CuO + 2Al → 3Cu + Al 2 O 3 (t = 1000-1100 o C).
As a result of the reaction, copper and oxide are formed aluminum.
5.copper oxide reaction(II) with copper:
CuO + Cu → Cu 2 O (t = 1000-1200 o C).
As a result of the reaction, copper (I) oxide is formed.
6. copper oxide reaction(II) With lithium oxide:
CuO + Li 2 O → Li 2 CuO 2 (t = 800-1000 o C, O 2).
The reaction takes place in a flow of oxygen. As a result of the reaction, lithium cuprate is formed.
7. copper oxide reaction(II) with sodium oxide:
CuO + Na 2 O → Na 2 CuO 2 (t = 800-1000 o C, O 2).
The reaction takes place in a flow of oxygen. As a result of the reaction, sodium cuprate is formed.
8.copper oxide reaction(II) with carbon monoxide:
CuO + CO → Cu + CO 2.
As a result of the reaction, copper and carbon monoxide (carbon dioxide) are formed.
9. copper oxide reaction(II) with oxide gland:
CuO + Fe 2 O 3 → CuFe 2 O 4 (t o).
As a result of the reaction, a salt is formed - copper ferrite. The reaction proceeds when the reaction mixture is calcined.
10. copper oxide reaction(II) with hydrofluoric acid:
CuO + 2HF → CuF 2 + H 2 O.
As a result of a chemical reaction, a salt is obtained - copper fluoride and water.
11.copper oxide reaction(II) with nitric acid:
CuO + 2HNO 3 → 2Cu(NO 3) 2 + H 2 O.
As a result of a chemical reaction, a salt is obtained - copper nitrate and water .
Copper oxide reacts similarly(II) and with other acids.
12. copper oxide reaction(II) with hydrogen bromide (hydrogen bromide):
CuO + 2HBr → CuBr 2 + H 2 O.
As a result of a chemical reaction, a salt is obtained - copper bromide and water .
13. copper oxide reaction(II) with hydrogen iodine:
CuO + 2HI → CuI 2 + H 2 O.
As a result of a chemical reaction, a salt is obtained - copper iodide and water .
14. copper oxide reaction(II) With sodium hydroxide :
CuO + 2NaOH → Na 2 CuO 2 + H 2 O.
As a result of a chemical reaction, a salt is obtained - sodium cuprate and water .
15.copper oxide reaction(II) With potassium hydroxide :
CuO + 2KOH → K 2 CuO 2 + H 2 O.
As a result of a chemical reaction, a salt is obtained - potassium cuprate and water .
16.copper oxide reaction(II) with sodium hydroxide and water:
CuO + 2NaOH + H 2 O → Na 2 2 (t = 100 o C).
Sodium hydroxide is dissolved in water. A solution of sodium hydroxide in water 20-30%. The reaction proceeds at boiling. As a result of a chemical reaction, sodium tetrahydroxocuprate is obtained.
17.copper oxide reaction(II) with potassium superoxide:
2CuO + 2KO 2 → 2KCuO 2 + O 2 (t = 400-500 o C).
As a result of a chemical reaction, a salt is obtained - potassium cuprate (III) and
Application
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
This oxide is used to color green or Blue colour enamel and glass, and also produce a copper-ruby variety of the latter. 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. Necessary condition reactions - the alkaline environment in which it should 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 valency. There are other elements that have several oxides, but we'll talk about them another time.
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 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, 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 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 low 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 conducting a reaction to outdoors, bivalent 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 USE assignments, as a question on the topic to what degree of oxidation iron will be oxidized under the action of dilute and concentrated hydrochloric acid on it. 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
Please note that diluted 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..
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. Even though they are made up 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". His education requires long time. 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. Irritating when swallowed gastrointestinal tract 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 applied as food additive 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.
