Study of the possibility of carrying out reactions on the interaction of sulfur with metals. Iron - a general characteristic of the element, the chemical properties of iron and its compounds The level of reactions of iron with oxygen
DEFINITION
Iron- an element of the eighth group of the fourth period of the Periodic system of chemical elements of D. I. Mendeleev.
And the languid number is 26. The symbol is Fe (lat. “ferrum”). One of the most common metals in the earth's crust (second place after aluminum).
Physical properties of iron
Iron is a gray metal. In its pure form, it is quite soft, malleable and ductile. The electronic configuration of the external energy level is 3d 6 4s 2 . In its compounds, iron exhibits the oxidation states "+2" and "+3". The melting point of iron is 1539C. Iron forms two crystalline modifications: α- and γ-iron. The first of them has a cubic body-centered lattice, the second has a cubic face-centered one. α-Iron is thermodynamically stable in two temperature ranges: below 912 and from 1394C to the melting point. Between 912 and 1394C, γ-iron is stable.
The mechanical properties of iron depend on its purity - the content in it of even very small amounts of other elements. Solid iron has the ability to dissolve many elements in itself.
Chemical properties of iron
In moist air, iron quickly rusts, i.e. covered with a brown coating of hydrated iron oxide, which, due to its friability, does not protect iron from further oxidation. In water, iron corrodes intensively; with abundant access of oxygen, hydrated forms of iron oxide (III) are formed:
2Fe + 3/2O 2 + nH 2 O = Fe 2 O 3 × H 2 O.
With a lack of oxygen or with difficult access, a mixed oxide (II, III) Fe 3 O 4 is formed:
3Fe + 4H 2 O (v) ↔ Fe 3 O 4 + 4H 2.
Iron dissolves in hydrochloric acid of any concentration:
Fe + 2HCl \u003d FeCl 2 + H 2.
Similarly, dissolution occurs in dilute sulfuric acid:
Fe + H 2 SO 4 \u003d FeSO 4 + H 2.
In concentrated solutions of sulfuric acid, iron is oxidized to iron (III):
2Fe + 6H 2 SO 4 \u003d Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O.
However, in sulfuric acid, the concentration of which is close to 100%, iron becomes passive and there is practically no interaction. In dilute and moderately concentrated solutions of nitric acid, iron dissolves:
Fe + 4HNO 3 \u003d Fe (NO 3) 3 + NO + 2H 2 O.
At high concentrations of nitric acid, dissolution slows down and iron becomes passive.
Like other metals, iron reacts with simple substances. The reactions of the interaction of iron with halogens (regardless of the type of halogen) proceed when heated. The interaction of iron with bromine proceeds at an increased vapor pressure of the latter:
2Fe + 3Cl 2 \u003d 2FeCl 3;
3Fe + 4I 2 = Fe 3 I 8.
The interaction of iron with sulfur (powder), nitrogen and phosphorus also occurs when heated:
6Fe + N 2 = 2Fe 3 N;
2Fe + P = Fe 2 P;
3Fe + P = Fe 3 P.
Iron is able to react with non-metals such as carbon and silicon:
3Fe + C = Fe 3 C;
Among the reactions of the interaction of iron with complex substances, the following reactions play a special role - iron is able to reduce metals that are in the activity series to the right of it, from salt solutions (1), to reduce iron (III) compounds (2):
Fe + CuSO 4 \u003d FeSO 4 + Cu (1);
Fe + 2FeCl 3 = 3FeCl 2 (2).
Iron, at elevated pressure, reacts with a non-salt-forming oxide - CO to form substances of complex composition - carbonyls - Fe (CO) 5, Fe 2 (CO) 9 and Fe 3 (CO) 12.
Iron, in the absence of impurities, is stable in water and in dilute alkali solutions.
Getting iron
The main way to obtain iron is from iron ore (hematite, magnetite) or electrolysis of solutions of its salts (in this case, “pure” iron is obtained, i.e. iron without impurities).
Examples of problem solving
EXAMPLE 1
| Exercise | Iron scale Fe 3 O 4 weighing 10 g was first treated with 150 ml of hydrochloric acid solution (density 1.1 g/ml) with a mass fraction of hydrogen chloride 20%, and then an excess of iron was added to the resulting solution. Determine the composition of the solution (in % by weight). |
| Solution | We write the reaction equations according to the condition of the problem: 8HCl + Fe 3 O 4 \u003d FeCl 2 + 2FeCl 3 + 4H 2 O (1); 2FeCl 3 + Fe = 3FeCl 2 (2). Knowing the density and volume of a hydrochloric acid solution, you can find its mass: m sol (HCl) = V(HCl) × ρ (HCl); m sol (HCl) \u003d 150 × 1.1 \u003d 165 g. Calculate the mass of hydrogen chloride: m(HCl)=msol(HCl)×ω(HCl)/100%; m(HCl) = 165 x 20%/100% = 33 g. The molar mass (mass of one mol) of hydrochloric acid, calculated using the table of chemical elements of D.I. Mendeleev - 36.5 g / mol. Find the amount of hydrogen chloride substance: v(HCl) = m(HCl)/M(HCl); v (HCl) \u003d 33 / 36.5 \u003d 0.904 mol. Molar mass (mass of one mole) of scale, calculated using the table of chemical elements of D.I. Mendeleev - 232 g/mol. Find the amount of scale substance: v (Fe 3 O 4) \u003d 10/232 \u003d 0.043 mol. According to equation 1, v(HCl): v(Fe 3 O 4) \u003d 1: 8, therefore, v (HCl) \u003d 8 v (Fe 3 O 4) \u003d 0.344 mol. Then, the amount of hydrogen chloride substance calculated according to the equation (0.344 mol) will be less than that indicated in the condition of the problem (0.904 mol). Therefore, hydrochloric acid is in excess and another reaction will proceed: Fe + 2HCl = FeCl 2 + H 2 (3). Let's determine the amount of iron chloride substance formed as a result of the first reaction (indices denote a specific reaction): v 1 (FeCl 2): v (Fe 2 O 3) = 1:1 = 0.043 mol; v 1 (FeCl 3): v (Fe 2 O 3) = 2:1; v 1 (FeCl 3) = 2 × v (Fe 2 O 3) = 0.086 mol. Let's determine the amount of hydrogen chloride that did not react in reaction 1 and the amount of iron (II) chloride substance formed during reaction 3: v rem (HCl) \u003d v (HCl) - v 1 (HCl) \u003d 0.904 - 0.344 \u003d 0.56 mol; v 3 (FeCl 2): v rem (HCl) = 1:2; v 3 (FeCl 2) \u003d 1/2 × v rem (HCl) \u003d 0.28 mol. Let's determine the amount of FeCl 2 substance formed during reaction 2, the total amount of FeCl 2 substance and its mass: v 2 (FeCl 3) = v 1 (FeCl 3) = 0.086 mol; v 2 (FeCl 2): v 2 (FeCl 3) = 3:2; v 2 (FeCl 2) = 3/2× v 2 (FeCl 3) = 0.129 mol; v sum (FeCl 2) \u003d v 1 (FeCl 2) + v 2 (FeCl 2) + v 3 (FeCl 2) \u003d 0.043 + 0.129 + 0.28 \u003d 0.452 mol; m (FeCl 2) \u003d v sum (FeCl 2) × M (FeCl 2) \u003d 0.452 × 127 \u003d 57.404 g. Let us determine the amount of substance and the mass of iron that entered into reactions 2 and 3: v 2 (Fe): v 2 (FeCl 3) = 1:2; v 2 (Fe) \u003d 1/2 × v 2 (FeCl 3) \u003d 0.043 mol; v 3 (Fe): v rem (HCl) = 1:2; v 3 (Fe) = 1/2×v rem (HCl) = 0.28 mol; v sum (Fe) \u003d v 2 (Fe) + v 3 (Fe) \u003d 0.043 + 0.28 \u003d 0.323 mol; m(Fe) = v sum (Fe) ×M(Fe) = 0.323 ×56 = 18.088 g. Let us calculate the amount of substance and the mass of hydrogen released in reaction 3: v (H 2) \u003d 1/2 × v rem (HCl) \u003d 0.28 mol; m (H 2) \u003d v (H 2) × M (H 2) \u003d 0.28 × 2 \u003d 0.56 g. We determine the mass of the resulting solution m ' sol and the mass fraction of FeCl 2 in it: m’ sol \u003d m sol (HCl) + m (Fe 3 O 4) + m (Fe) - m (H 2); |
Introduction
The study of the chemical properties of individual elements is an integral part of the course of chemistry in the modern school, which allows, on the basis of the inductive approach, to make an assumption about the features of the chemical interaction of elements based on their physicochemical characteristics. However, the capabilities of the school chemical laboratory do not always fully allow to demonstrate the dependence of the chemical properties of an element on its position in the periodic system of chemical elements, structural features of simple substances.
The chemical properties of sulfur are used both at the beginning of the study of a chemistry course to demonstrate the difference between chemical phenomena and physical ones, and in the study of the characteristics of individual chemical elements. The most frequently recommended demonstration in guidelines is the interaction of sulfur with iron, as an example of chemical phenomena and an example of the oxidizing properties of sulfur. But in most cases, this reaction either does not proceed at all, or the results of its course cannot be assessed with the naked eye. Various options for conducting this experiment are often characterized by low reproducibility of the results, which does not allow them to be used systematically in characterizing the above processes. Therefore, it is relevant to search for options that can constitute an alternative to demonstrating the process of interaction of iron with sulfur, adequate to the characteristics of a school chemical laboratory.
Target: Investigate the possibility of carrying out reactions on the interaction of sulfur with metals in a school laboratory.
Tasks:
Determine the main physical and chemical characteristics of sulfur;
Analyze the conditions for the conduct and flow of reactions of interaction of sulfur with metals;
To study known methods for the implementation of the interaction of sulfur with metals;
Select systems for carrying out reactions;
Assess the adequacy of the selected reactions to the conditions of the school chemical laboratory.
Object of study: reactions of interaction of sulfur with metals
Subject of study: the feasibility of interaction reactions between sulfur and metals in a school laboratory.
Hypothesis: An alternative to the interaction of iron with sulfur in the conditions of a school chemical laboratory will be a chemical reaction that meets the requirements of clarity, reproducibility, relative safety and availability of reactants.
We want to start our work with a brief description of sulfur:
Position in the periodic system: sulfur is in period 3, group VI, main (A) subgroup, belongs to s-elements.
The atomic number of sulfur is 16, therefore, the charge of the sulfur atom is + 16, the number of electrons is 16. Three electronic levels in the outer level are 6 electrons
Scheme of arrangement of electrons by levels:
16S )))
2 8 6
The nucleus of the 32 S sulfur atom contains 16 protons (equal to the nuclear charge) and 16 neutrons (atomic mass minus the number of protons: 32 - 16 = 16).
Electronic formula: 1s 2 2s 2 2p 6 3s 2 3p 4
Table 1
The values of the ionization potentials of the sulfur atom
Ionization potential
Energy (eV)
Sulfur in the cold rather inert (vigorously connects only with fluorine), but when heated it becomes very reactive - it reacts with halides(except iodine), oxygen, hydrogen and almost all metals. As a result reactions of the latter type, the corresponding sulfur compounds are formed.
The reactivity of sulfur, like any other element, when interacting with metals depends on:
activity of reacting substances. For example, sulfur will most actively interact with alkali metals
on the reaction temperature. This is explained by the thermodynamic features of the process.
The thermodynamic possibility of spontaneous chemical reactions under standard conditions is determined by the standard Gibbs energy of the reaction:
ΔG 0 T< 0 – прямая реакция протекает
ΔG 0 T > 0 - direct reaction is impossible
on the degree of grinding of the reacting substances, since both sulfur and metals react mainly in the solid state.
Thermodynamic characteristics of some reactions of interaction of sulfur with metals are given in slide 4
It can be seen from the table that it is thermodynamically possible for sulfur to interact both with metals of the beginning of a series of stresses and with low-activity metals.
Thus, sulfur is a rather active non-metal when heated, capable of reacting with metals of both high activity (alkaline) and low activity (silver, copper).
Study of the interaction of sulfur with metals
Selection of systems for research
To study the interaction of sulfur with metals, systems were selected, including metals located in different places of the Beketov series, having different activities.
The following criteria were determined as selection conditions: speed of carrying out, visibility, completeness of the reaction, relative safety, reproducibility of the result, substances should differ markedly in physical properties, the presence of substances in the school laboratory, there are successful attempts to conduct interactions of sulfur with specific metals.
To assess the reproducibility of the reactions carried out, each experiment was carried out three times.
Based on these criteria, the following reaction systems were selected for the experiment:
SULFUR AND COPPER Cu + S = CuS + 79 kJ/mol
Methodology and expected effect
Let's take 4 g of sulfur in a powder state and pour it into a test tube. Heat sulfur in a test tube to a boil. Then take a copper wire and heat it over a flame. When the sulfur melts and boils, put copper wire into it
Expected Result:The test tube is filled with brown vapors, the wire heats up and "burns out" with the formation of brittle sulfide.
2. Interaction of sulfur with copper.
The reaction turned out to be not very clear, spontaneous heating of copper also did not occur. When hydrochloric acid was added, no special gas evolution was observed.
SULFUR AND IRON Fe + S = FeS + 100.4 kJ/mol
Methodology and expected effect
Take 4 g of powdered sulfur and 7 g of powdered iron and mix. Pour the resulting mixture into a test tube. We heat the substances in the test tube
Expected Result:There is a strong spontaneous heating of the mixture. The resulting iron sulfide is sintered. The substance is not separated by water and does not react to a magnet.
1. Interaction of sulfur with iron.
It is practically impossible to carry out a reaction to obtain iron sulfide without a residue in laboratory conditions, it is very difficult to determine when the substances have completely reacted, spontaneous heating of the reaction mixture is not observed. The resulting substance was checked to see if it was iron sulfide. For this we used HCl. When we dropped hydrochloric acid on the substance, it began to foam, hydrogen sulfide was released.
SULFUR AND SODIUM 2Na + S \u003d Na 2 S + 370.3 kJ / mol
Methodology and expected effect
Take 4 g of powdered sulfur and pour it into a mortar, grind it well
Let's cut off a piece of sodium weighing about 2 g. Cut off the oxide film, grind them together.
Expected Result:The reaction proceeds violently, self-ignition of the reagents is possible.
3. Interaction of sulfur with sodium.
The interaction of sulfur with sodium is itself a dangerous and memorable experiment. After a few seconds of rubbing, the first sparks flew, sodium and sulfur flared up in the mortar and began to burn. When the product interacts with hydrochloric acid, hydrogen sulfide is actively released.
SULFUR AND ZINC Zn + S = ZnS + 209 kJ/mol
Methodology and expected effect
Take powdered sulfur and zinc, 4 g each, mix the substances. Pour the finished mixture onto an asbestos mesh. We bring a hot torch to the substances
Expected Result:The reaction does not proceed immediately, but violently, a greenish-blue flame is formed.
4. Interaction of sulfur with zinc.
The reaction is very difficult to start, it requires the use of strong oxidizing agents or high temperature to initiate it. Substances flash with a greenish-blue flame. When the flame goes out, a residue remains in this place; when interacting with hydrochloric acid, hydrogen sulfide is slightly released.
SULFUR AND ALUMINUM 2Al + 3S \u003d Al 2 S 3 + 509.0 kJ / mol
Methodology and expected effect
Take powdered sulfur weighing 4 g and aluminum weighing 2.5 g and mix. We place the resulting mixture on an asbestos mesh. Ignite the mixture with burning magnesium
Expected Result:The reaction is a flash.
5. Interaction of sulfur with aluminum.
The reaction requires the addition of a strong oxidizing agent as an initiator. After ignition with burning magnesium, there was a powerful flash of yellowish-white color, hydrogen sulfide is released quite actively.
SULFUR AND MAGNESIUM Mg + S = MgS + 346.0 kJ/mol
Methodology and expected effect
Take magnesium shavings 2.5 g and powdered sulfur 4 g and mix
The resulting mixture will be placed on an asbestos mesh. We bring the splinter to the resulting mixture.
Expected Result:During the reaction, a powerful flash occurs.
4. Interaction of sulfur with magnesium.
The reaction requires the addition of pure magnesium as an initiator. There is a powerful flash of a whitish color, hydrogen sulfide is actively released.
Conclusion
The reaction to obtain iron sulfide was not completed, since a residue remained in the form of a mixture of plastic sulfur and iron.
The most active release of hydrogen sulfide was manifested in sodium sulfide and magnesium and aluminum sulfides.
Less active release of hydrogen sulfide was in copper sulfide.
Conducting experiments to obtain sodium sulfide is dangerous and not recommended in a school laboratory.
Reactions for the production of aluminum, magnesium and zinc sulfides are most suitable for conducting in school conditions.
The expected and actual results coincided with the interaction of sulfur with sodium, magnesium and aluminum.
Conclusion
Despite the existing recommendations for demonstrating the interaction of iron with sulfur as an example illustrating the chemical phenomena and oxidizing properties of sulfur in a general school chemistry course, the actual implementation of such an experiment is often not accompanied by a visible effect.
When determining an alternative to this demonstration, systems were selected that met the requirements for visibility, safety, and the availability of reactants in the school laboratory. As possible options, the reaction systems of sulfur with copper, iron, zinc, magnesium, aluminum, sodium were chosen, allowing to evaluate the effectiveness of using the reaction of interaction of sulfur with various metals as demonstration experiments in chemistry lessons.
According to the results of the experiments, it was determined that it is most optimal for these purposes to use the reaction systems of sulfur with metals of medium-high activity (magnesium, aluminum).
Based on the experiments, a video was created that demonstrates the oxidizing properties of sulfur using the example of its interaction with metals, which makes it possible to describe these properties without conducting a full-scale experiment. A website has been created as an additional aid ( ), which presents, among other things, the results of the study in a visual form.
The results of the study can become the basis for a deeper study of the features of the chemical properties of non-metals, chemical kinetics and thermodynamics.
Chemical properties of iron let's consider the example of its interaction with typical non-metals - sulfur and oxygen.
Mix iron and sulfur crushed to a powdery state in a Petri dish. Let's heat a steel needle in a flame and touch it with a mixture of reagents. A violent reaction between iron and sulfur is accompanied by the release of heat and light energy. The solid product of the interaction of these substances - iron (II) sulfide - is black. Unlike iron, it is not attracted by a magnet.
Iron reacts with sulfur to form iron(II) sulfide. Let's write the reaction equation:
The reaction of iron with oxygen also requires pre-heating. Pour quartz sand into a thick-walled vessel. Let us heat a bundle of very thin iron wire, the so-called iron wool, in the flame of a burner. Let's bring the red-hot wire into the vessel with oxygen. Iron burns with a dazzling flame, scattering sparks - red-hot particles of iron scale Fe 3 O 4.
The same reaction occurs in air, when the steel is strongly heated by friction during machining.
When iron is burned in oxygen or in air, iron scale is formed:
3Fe + 2O 2 \u003d Fe 3 O 4, material from the site
or 3Fe + 2O 2 \u003d FeO. Fe2O3.
Iron oxide is a compound in which iron has different valence values.
The passage of both reactions of the connection is accompanied by the release of thermal and light energy.
On this page, material on the topics:
What type of reaction is iron sulfide with oxygen
Write an equation between iron and sulfur
Equation of the reactions of iron with oxygen
An example of a chemical reaction of the combination of iron with sulfur
The equation for the interaction of oxygen with iron
Questions about this item:
Iron is an element of a side subgroup of the eighth group of the fourth period of the periodic system of chemical elements of D. I. Mendeleev with atomic number 26. It is designated by the symbol Fe (lat. Ferrum). One of the most common metals in the earth's crust (second place after aluminum). Medium activity metal, reducing agent.
Main oxidation states - +2, +3
A simple substance iron is a malleable silver-white metal with a high chemical reactivity: iron quickly corrodes at high temperatures or high humidity in the air. In pure oxygen, iron burns, and in a finely dispersed state, it ignites spontaneously in air.
Chemical properties of a simple substance - iron:
Rusting and burning in oxygen
1) In air, iron is easily oxidized in the presence of moisture (rusting):
4Fe + 3O 2 + 6H 2 O → 4Fe(OH) 3
A heated iron wire burns in oxygen, forming scale - iron oxide (II, III):
3Fe + 2O 2 → Fe 3 O 4
3Fe + 2O 2 → (Fe II Fe 2 III) O 4 (160 ° С)
2) At high temperatures (700–900°C), iron reacts with water vapor:
3Fe + 4H 2 O - t ° → Fe 3 O 4 + 4H 2
3) Iron reacts with non-metals when heated:
2Fe+3Cl 2 →2FeCl 3 (200 °C)
Fe + S – t° → FeS (600 °С)
Fe + 2S → Fe +2 (S 2 -1) (700 ° С)
4) In a series of voltages, it is to the left of hydrogen, reacts with dilute acids Hcl and H 2 SO 4, while iron (II) salts are formed and hydrogen is released:
Fe + 2HCl → FeCl 2 + H 2 (reactions are carried out without air access, otherwise Fe +2 is gradually converted by oxygen into Fe +3)
Fe + H 2 SO 4 (diff.) → FeSO 4 + H 2
In concentrated oxidizing acids, iron dissolves only when heated, it immediately passes into the Fe 3+ cation:
2Fe + 6H 2 SO 4 (conc.) – t° → Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O
Fe + 6HNO 3 (conc.) – t° → Fe(NO 3) 3 + 3NO 2 + 3H 2 O
(in the cold, concentrated nitric and sulfuric acids passivate
An iron nail immersed in a bluish solution of copper sulphate is gradually covered with a coating of red metallic copper.
5) Iron displaces metals to the right of it in solutions of their salts.
Fe + CuSO 4 → FeSO 4 + Cu
Amphotericity of iron is manifested only in concentrated alkalis during boiling:
Fe + 2NaOH (50%) + 2H 2 O \u003d Na 2 ↓ + H 2
and a precipitate of sodium tetrahydroxoferrate(II) is formed.
Technical iron- alloys of iron with carbon: cast iron contains 2.06-6.67% C, steel 0.02-2.06% C, other natural impurities (S, P, Si) and artificially introduced special additives (Mn, Ni, Cr) are often present, which gives iron alloys technically useful properties - hardness, thermal and corrosion resistance, malleability, etc. .
Blast furnace iron production process
The blast-furnace process of iron production consists of the following stages:
a) preparation (roasting) of sulfide and carbonate ores - conversion to oxide ore:
FeS 2 → Fe 2 O 3 (O 2, 800 ° С, -SO 2) FeCO 3 → Fe 2 O 3 (O 2, 500-600 ° С, -CO 2)
b) burning coke with hot blast:
C (coke) + O 2 (air) → CO 2 (600-700 ° C) CO 2 + C (coke) ⇌ 2CO (700-1000 ° C)
c) reduction of oxide ore with carbon monoxide CO in succession:
Fe2O3 →(CO)(Fe II Fe 2 III) O 4 →(CO) FeO →(CO) Fe
d) carburization of iron (up to 6.67% C) and melting of cast iron:
Fe (t ) →(C(coke)900-1200°С) Fe (g) (cast iron, t pl 1145°C)
In cast iron, cementite Fe 2 C and graphite are always present in the form of grains.
Steel production
The redistribution of cast iron into steel is carried out in special furnaces (converter, open-hearth, electric), which differ in the method of heating; process temperature 1700-2000 °C. Blowing oxygen-enriched air burns out excess carbon from cast iron, as well as sulfur, phosphorus and silicon in the form of oxides. In this case, oxides are either captured in the form of exhaust gases (CO 2, SO 2), or are bound into an easily separated slag - a mixture of Ca 3 (PO 4) 2 and CaSiO 3. To obtain special steels, alloying additives of other metals are introduced into the furnace.
Receipt pure iron in industry - electrolysis of a solution of iron salts, for example:
FeCl 2 → Fe↓ + Cl 2 (90°C) (electrolysis)
(there are other special methods, including the reduction of iron oxides with hydrogen).
Pure iron is used in the production of special alloys, in the manufacture of cores of electromagnets and transformers, cast iron is used in the production of castings and steel, steel is used as structural and tool materials, including wear-, heat- and corrosion-resistant materials.
Iron(II) oxide F EO . Amphoteric oxide with a large predominance of basic properties. Black, has an ionic structure of Fe 2+ O 2-. When heated, it first decomposes, then re-forms. It is not formed during the combustion of iron in air. Does not react with water. Decomposed by acids, fused with alkalis. Slowly oxidizes in moist air. Recovered by hydrogen, coke. Participates in the blast-furnace process of iron smelting. It is used as a component of ceramics and mineral paints. Equations of the most important reactions:
4FeO ⇌ (Fe II Fe 2 III) + Fe (560-700 ° С, 900-1000 ° С)
FeO + 2HC1 (razb.) \u003d FeC1 2 + H 2 O
FeO + 4HNO 3 (conc.) \u003d Fe (NO 3) 3 + NO 2 + 2H 2 O
FeO + 4NaOH \u003d 2H 2 O + Na 4FeO3(red.) trioxoferrate(II)(400-500 °С)
FeO + H 2 \u003d H 2 O + Fe (high purity) (350 ° C)
FeO + C (coke) \u003d Fe + CO (above 1000 ° C)
FeO + CO \u003d Fe + CO 2 (900 ° C)
4FeO + 2H 2 O (moisture) + O 2 (air) → 4FeO (OH) (t)
6FeO + O 2 \u003d 2 (Fe II Fe 2 III) O 4 (300-500 ° С)
Receipt in laboratories: thermal decomposition of iron (II) compounds without air access:
Fe (OH) 2 \u003d FeO + H 2 O (150-200 ° C)
FeSOz \u003d FeO + CO 2 (490-550 ° С)
Diiron oxide (III) - iron ( II ) ( Fe II Fe 2 III) O 4 . Double oxide. Black, has the ionic structure of Fe 2+ (Fe 3+) 2 (O 2-) 4. Thermally stable up to high temperatures. Does not react with water. Decomposed by acids. It is reduced by hydrogen, red-hot iron. Participates in the blast-furnace process of iron production. It is used as a component of mineral paints ( minium iron), ceramics, colored cement. The product of special oxidation of the surface of steel products ( blackening, bluing). The composition corresponds to brown rust and dark scale on iron. The use of the Fe 3 O 4 formula is not recommended. Equations of the most important reactions:
2 (Fe II Fe 2 III) O 4 \u003d 6FeO + O 2 (above 1538 ° С)
(Fe II Fe 2 III) O 4 + 8HC1 (razb.) \u003d FeC1 2 + 2FeC1 3 + 4H 2 O
(Fe II Fe 2 III) O 4 + 10HNO 3 (conc.) \u003d 3 Fe (NO 3) 3 + NO 2 + 5H 2 O
(Fe II Fe 2 III) O 4 + O 2 (air) \u003d 6Fe 2 O 3 (450-600 ° С)
(Fe II Fe 2 III) O 4 + 4H 2 \u003d 4H 2 O + 3Fe (high purity, 1000 ° C)
(Fe II Fe 2 III) O 4 + CO \u003d 3 FeO + CO 2 (500-800 ° C)
(Fe II Fe 2 III) O4 + Fe ⇌4 FeO (900-1000 ° С, 560-700 ° С)
Receipt: combustion of iron (see) in air.
magnetite.
Iron(III) oxide F e 2 O 3 . Amphoteric oxide with a predominance of basic properties. Red-brown, has an ionic structure (Fe 3+) 2 (O 2-) 3. Thermally stable up to high temperatures. It is not formed during the combustion of iron in air. Does not react with water, a brown amorphous hydrate Fe 2 O 3 nH 2 O precipitates from the solution. Slowly reacts with acids and alkalis. It is reduced by carbon monoxide, molten iron. Alloys with oxides of other metals and forms double oxides - spinels(technical products are called ferrites). It is used as a raw material in iron smelting in the blast furnace process, as a catalyst in the production of ammonia, as a component of ceramics, colored cements and mineral paints, in thermite welding of steel structures, as a sound and image carrier on magnetic tapes, as a polishing agent for steel and glass.
Equations of the most important reactions:
6Fe 2 O 3 \u003d 4 (Fe II Fe 2 III) O 4 + O 2 (1200-1300 ° С)
Fe 2 O 3 + 6HC1 (razb.) → 2FeC1 3 + ZH 2 O (t) (600 ° C, p)
Fe 2 O 3 + 2NaOH (conc.) → H 2 O+ 2 NaFeO 2 (red)dioxoferrate(III)
Fe 2 O 3 + MO \u003d (M II Fe 2 II I) O 4 (M \u003d Cu, Mn, Fe, Ni, Zn)
Fe 2 O 3 + ZN 2 \u003d ZN 2 O + 2Fe (highly pure, 1050-1100 ° С)
Fe 2 O 3 + Fe \u003d ZFeO (900 ° C)
3Fe 2 O 3 + CO \u003d 2 (Fe II Fe 2 III) O 4 + CO 2 (400-600 ° С)
Receipt in the laboratory - thermal decomposition of iron (III) salts in air:
Fe 2 (SO 4) 3 \u003d Fe 2 O 3 + 3SO 3 (500-700 ° С)
4 (Fe (NO 3) 3 9 H 2 O) \u003d 2 Fe a O 3 + 12NO 2 + 3O 2 + 36H 2 O (600-700 ° С)
In nature - iron oxide ores hematite Fe 2 O 3 and limonite Fe 2 O 3 nH 2 O
Iron(II) hydroxide F e(OH) 2 . Amphoteric hydroxide with a predominance of basic properties. White (sometimes with a greenish tinge), Fe-OH bonds are predominantly covalent. Thermally unstable. Easily oxidizes in air, especially when wet (darkens). Insoluble in water. Reacts with dilute acids, concentrated alkalis. Typical restorer. An intermediate product in the rusting of iron. It is used in the manufacture of the active mass of iron-nickel batteries.
Equations of the most important reactions:
Fe (OH) 2 \u003d FeO + H 2 O (150-200 ° C, in atm.N 2)
Fe (OH) 2 + 2HC1 (razb.) \u003d FeC1 2 + 2H 2 O
Fe (OH) 2 + 2NaOH (> 50%) \u003d Na 2 ↓ (blue-green) (boiling)
4Fe(OH) 2 (suspension) + O 2 (air) → 4FeO(OH)↓ + 2H 2 O (t)
2Fe (OH) 2 (suspension) + H 2 O 2 (razb.) \u003d 2FeO (OH) ↓ + 2H 2 O
Fe (OH) 2 + KNO 3 (conc.) \u003d FeO (OH) ↓ + NO + KOH (60 ° С)
Receipt: precipitation from solution with alkalis or ammonia hydrate in an inert atmosphere:
Fe 2+ + 2OH (razb.) = Fe(OH) 2 ↓
Fe 2+ + 2 (NH 3 H 2 O) = Fe(OH) 2 ↓+ 2NH4
Iron metahydroxide F eO(OH). Amphoteric hydroxide with a predominance of basic properties. Light brown, Fe-O and Fe-OH bonds are predominantly covalent. When heated, it decomposes without melting. Insoluble in water. It precipitates from solution in the form of a brown amorphous polyhydrate Fe 2 O 3 nH 2 O, which, when kept under a dilute alkaline solution or when dried, turns into FeO (OH). Reacts with acids, solid alkalis. Weak oxidizing and reducing agent. Sintered with Fe(OH) 2 . An intermediate product in the rusting of iron. It is used as a base for yellow mineral paints and enamels, as an exhaust gas absorber, as a catalyst in organic synthesis.
Connection composition Fe(OH) 3 is not known (not obtained).
Equations of the most important reactions:
Fe 2 O 3 . nH 2 O→( 200-250 °С, —H 2 O) FeO(OH)→( 560-700°C in air, -H2O)→Fe 2 O 3
FeO (OH) + ZNS1 (razb.) \u003d FeC1 3 + 2H 2 O
FeO(OH)→ Fe 2 O 3 . nH 2 O-colloid(NaOH (conc.))
FeO(OH) → Na 3 [Fe(OH) 6 ]white, Na 5 and K 4, respectively; in both cases, a blue product of the same composition and structure, KFe III, precipitates. In the laboratory, this precipitate is called Prussian blue, or turnbull blue:
Fe 2+ + K + + 3- = KFe III ↓
Fe 3+ + K + + 4- = KFe III ↓
Chemical names of initial reagents and reaction product:
K 3 Fe III - potassium hexacyanoferrate (III)
K 4 Fe III - potassium hexacyanoferrate (II)
KFe III - hexacyanoferrate (II) iron (III) potassium
In addition, the thiocyanate ion NCS - is a good reagent for Fe 3+ ions, iron (III) combines with it, and a bright red (“bloody”) color appears:
Fe 3+ + 6NCS - = 3-
With this reagent (for example, in the form of KNCS salt), even traces of iron (III) can be detected in tap water if it passes through iron pipes covered with rust from the inside.
