The inorganic compound iron hydroxide 3 has the chemical formula Fe(OH)2. It belongs to a number of amphoteric compounds in which the properties characteristic of bases predominate. In appearance, this substance is white crystals, which gradually darken when left in the open air for a long time. There are options for crystals with a greenish tint. In everyday life, everyone can observe the substance in the form of a greenish coating on metal surfaces, which indicates the beginning of the rusting process - iron hydroxide 3 acts as one of the intermediate stages of this process.
In nature, the compound is found in the form of amakinite. This crystalline mineral, in addition to iron itself, also contains impurities of magnesium and manganese; all these substances give amakinite different shades - from yellow-green to pale green, depending on the percentage of a particular element. The hardness of the mineral is 3.5-4 units on the Mohs scale, and the density is approximately 3 g/cm³.
The physical properties of the substance should also include its extremely low solubility. When iron hydroxide 3 is heated, it decomposes.
This substance is very active and interacts with many other substances and compounds. For example, having the properties of a base, it interacts with various acids. In particular, iron sulfur 3 during the reaction leads to the production of (III). Since this reaction can occur by conventional calcination in open air, this inexpensive sulfate is used in both laboratory and industrial settings.
During the reaction, the result is the formation of iron (II) chloride.
In some cases, iron hydroxide 3 may also exhibit acidic properties. For example, when interacting with a highly concentrated (concentration must be at least 50%) solution of sodium hydroxide, sodium tetrahydroxoferrate (II) is obtained, which precipitates. True, for such a reaction to occur, it is necessary to provide rather complex conditions: the reaction must occur under conditions of boiling solution in a nitrogen atmospheric environment.
As already mentioned, when heated, the substance decomposes. The result of this decomposition is (II), and, in addition, metallic iron and its derivatives are obtained in the form of impurities: diiron oxide (III), the chemical formula of which is Fe3O4.
How to produce iron hydroxide 3, the production of which is associated with its ability to react with acids? Before you begin the experiment, you should be sure to remember the safety rules when conducting such experiments. These rules apply to all cases of handling acid-base solutions. The main thing here is to provide reliable protection and avoid contact of drops of solutions with mucous membranes and skin.
So, hydroxide can be obtained through a reaction in which iron (III) chloride and KOH - potassium hydroxide react. This method is the most common for the formation of insoluble bases. When these substances interact, a normal exchange reaction occurs, resulting in a brown precipitate. This precipitate is the desired substance.
The use of iron hydroxide in industrial production is quite widespread. The most common is its use as an active substance in iron-nickel batteries. In addition, the compound is used in metallurgy to produce various metal alloys, as well as in electroplating production and the automotive industry.
Ferrous compounds
I . Iron(II) hydroxide
Formed by the action of alkali solutions on iron (II) salts without air access:
FeCl 2 + 2 KOH = 2 KCl + F e (OH) 2 ↓
Fe(OH) 2 is a weak base, soluble in strong acids:
Fe(OH) 2 + H 2 SO 4 = FeSO 4 + 2H 2 O
Fe(OH) 2 + 2H + = Fe 2+ + 2H 2 O
Additional material:
Fe(OH) 2 – also exhibits weak amphoteric properties, reacts with concentrated alkalis:
Fe( OH) 2 + 2 NaOH = Na 2 [ Fe( OH) 4 ]. tetrahydroxoferrate salt is formed ( II) sodium
When Fe(OH) 2 is calcined without air access, iron (II) oxide FeO is formed -black connection:
Fe(OH) 2 t˚C → FeO + H 2 O
In the presence of atmospheric oxygen, the white precipitate Fe(OH) 2, oxidizing, turns brown - forming iron (III) hydroxide Fe(OH) 3:
4Fe(OH) 2 + O 2 + 2H 2 O = 4Fe(OH) 3 ↓
Additional material:
Iron (II) compounds have reducing properties; they are easily converted into iron (III) compounds under the influence of oxidizing agents:
10FeSO 4 + 2KMnO 4 + 8H 2 SO 4 = 5Fe 2 (SO 4) 3 + K 2 SO 4 + 2MnSO 4 + 8H 2 O
6FeSO 4 + 2HNO 3 + 3H 2 SO 4 = 3Fe 2 (SO 4) 3 + 2NO + 4H 2 O
Iron compounds are prone to complex formation:
FeCl 2 + 6NH 3 = Cl 2
Fe(CN) 2 + 4KCN = K 4 (yellow blood salt)
Qualitative reaction to Fe 2+
When in action potassium hexacyanoferrate (III) K 3 (red blood salt) on solutions of divalent iron salts is formed blue precipitate (Turnboole blue):
3 Fe 2+ Cl 2 + 3 K 3 [ Fe 3+ ( CN) 6 ] → 6 KCl + 3 KFe 2+ [ Fe 3+ ( CN) 6 ]↓
(Turnbull blue - hexacyanoferrate ( III ) iron ( II )-potassium)
Turnbull blue its properties are very similar to Prussian blue and also served as a dye. Named after one of the founders of the Scottish dyeing company Arthur and Turnbull.
Ferric compounds
I . Iron(III) oxide
Formed by burning iron sulfides, for example, by roasting pyrite:
4 FeS 2 + 11 O 2 t ˚ C → 2 Fe 2 O 3 + 8 SO 2
or when calcining iron salts:
2FeSO 4 t˚C → Fe 2 O 3 + SO 2 + SO 3
Fe 2 O 3 - oxide red-brown color, exhibiting amphoteric properties to a small extent
Fe 2 O 3 + 6HCl t˚C → 2FeCl 3 + 3H 2 O
Fe 2 O 3 + 6H + t˚C → 2Fe 3+ + 3H 2 O
Fe 2 O 3 + 2 NaOH + 3 H 2 O t ˚ C → 2 Na [ Fe (OH ) 4 ],a salt is formed - tetrahydroxoferrate ( III) sodium
Fe 2 O 3 + 2OH - + 3H 2 O t˚C → 2 -
When fused with basic oxides or carbonates of alkali metals, ferrites are formed:
Fe 2 O 3 + Na 2 O t˚C → 2NaFeO 2
Fe 2 O 3 + Na 2 CO 3 = 2NaFeO 2 + CO 2
II. Iron hydroxide ( III )
Formed by the action of alkali solutions on ferric iron salts: it precipitates in the form of a red-brown precipitate
Fe(NO 3) 3 + 3KOH = Fe(OH) 3 ↓ + 3KNO 3
Fe 3+ + 3OH - = Fe(OH) 3 ↓
Additionally:
Fe(OH) 3 is a weaker base than iron (II) hydroxide.
This is explained by the fact that Fe 2+ has a smaller ion charge and a larger radius than Fe 3+, and therefore Fe 2+ retains hydroxide ions weaker, i.e. Fe(OH) 2 dissociates more easily.
In this regard, iron (II) salts are hydrolyzed slightly, and iron (III) salts are hydrolyzed very strongly.
Hydrolysis also explains the color of solutions of Fe(III) salts: despite the fact that the Fe 3+ ion is almost colorless, solutions containing it are colored yellow-brown, which is explained by the presence of iron hydroxoions or Fe(OH) 3 molecules, which are formed due to hydrolysis :
Fe 3+ + H 2 O ↔ 2+ + H +
2+ + H 2 O ↔ + + H +
+ + H 2 O ↔ Fe(OH) 3 + H +
When heated, the color darkens, and when acids are added it becomes lighter due to the suppression of hydrolysis.
Fe(OH) 3 has weak amphoteric properties: it dissolves in dilute acids and concentrated alkali solutions:
Fe(OH) 3 + 3HCl = FeCl 3 + 3H 2 O
Fe(OH) 3 + 3H + = Fe 3+ + 3H 2 O
Fe(OH)3 + NaOH = Na
Fe(OH) 3 + OH - = -
Additional material:
Iron (III) compounds are weak oxidizing agents, react with strong reducing agents:
2Fe +3 Cl 3 + H 2 S -2 = S 0 ↓ + 2Fe +2 Cl 2 + 2HCl
FeCl 3 + KI = I 2 ↓ + FeCl 2 + KCl
Qualitative reactions to Fe 3+
Experience
1) During action potassium hexacyanoferrate (II) K 4 (yellow blood salt) on solutions of ferric iron salts is formed blue precipitate (Prussian blue):
4 Fe 3+ Cl 3 + 4 K 4 [ Fe 2+ ( CN) 6 ] → 12 KCl + 4 KFe 3+ [ Fe 2+ ( CN) 6 ]↓
(Prussian blue - hexacyanoferrate ( II ) iron ( III )-potassium)
Prussian blue was obtained by chance at the beginning of the 18th century in Berlin by the dyer Diesbach. Disbach bought an unusual potash (potassium carbonate) from a merchant: a solution of this potash when added with iron salts turned out blue. When checking the potash, it turned out that it was calcined with ox blood. The paint turned out to be suitable for fabrics: bright, durable and inexpensive. Soon the recipe for making paint became known: potash was fused with dried animal blood and iron filings. By leaching such an alloy, yellow blood salt was obtained. Nowadays Prussian blue is used to produce printing ink and tint polymers.
It has been established that Prussian blue and Turnboole blue are the same substance, since the complexes formed in the reactions are in equilibrium with each other:
KFe III[ FeII( CN) 6 ] ↔ KFe II[ Fe III( CN) 6 ]
2) When potassium or ammonium thiocyanate is added to a solution containing Fe 3+ ions, an intense blood-red color appears solution iron(III) thiocyanate:
2FeCl 3 + 6KCNS = 6KCl + Fe III[ Fe III( CNS) 6 ]
(when interacting with thiocyanates, Fe 2+ ions, the solution remains almost colorless).
Exercise equipment
Trainer No. 1 - Recognition of compounds containing Fe (2+) ion
Trainer No. 2 - Recognition of compounds containing Fe (3+) ion
Tasks for consolidation
№1.
Carry out the transformations:
FeCl 2 -> Fe(OH) 2 -> FeO -> FeSO 4
Fe -> Fe(NO 3) 3 -> Fe(OH) 3 -> Fe 2 O 3 -> NaFeO 2
No. 2. Write down reaction equations that can be used to obtain:
a) iron (II) salts and iron (III) salts;
b) iron (II) hydroxide and iron (III) hydroxide;
c) iron oxides.
4Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3.
Iron(III) oxide Fe2O3 - brown powder, insoluble in water.
Iron (III) oxide is obtained by decomposition of iron (III) hydroxide:
2Fe(OH)3 = Fe2O3 + 3H2O
Iron (III) oxide exhibits amphoteric properties:
Reacts with acids and solid alkalis NaOH and KOH, as well as with sodium and potassium carbonates at high temperatures:
Fe2O3 + 2NaOH = 2NaFeO2 + H2O,
Fe2O3 + 2OH - = 2FeO2- + H2O,
Fe2O3 + Na2CO3 = 2NaFeO2 + CO2.
Sodium ferrite
Iron(III) hydroxide obtained from iron (III) salts by reacting them with alkalis:
FeCl3 + 3NaOH = Fe(OH)3 + 3NaCl,
Iron(III) hydroxide is a weaker base than Fe(OH)2 and exhibits amphoteric properties (with a predominance of the main ones). When interacting with dilute acids, Fe(OH)3 easily forms the corresponding salts:
Fe(OH)3 + 3HCl = FeCl3 + H2O
2Fe(OH)3 + 3H2SO4 = Fe2(SO4)3 + 6H2O
Reactions with concentrated solutions of alkalis occur only with prolonged heating:
Fe(OH)3 + KOH = K
Compounds with iron oxidation state +3 exhibit oxidizing properties , since under the influence of reducing agents Fe+3 turns into Fe+2: Fe+3 + 1e = Fe+2.
For example, iron (III) chloride oxidizes potassium iodide to free iodine:
2FeCl3 + 2KI = 2FeCl2 + 2KCl + I20
Chromium.
Chromium is in the secondary subgroup of group VI of the Periodic table. Structure of the electronic shell of chromium: Cr 3d54s1. Oxidation states range from +1 to +6, but the most stable are +2, +3, +6.
The mass fraction of chromium in the earth's crust is 0.02%. The most important minerals that make up chromium ores are chromite, or chromium iron ore, and its varieties in which iron is partially replaced by magnesium and chromium by aluminum.
Chrome is a silvery gray metal. Pure chromium is quite ductile, and technical chromium is the hardest of all metals.
Chromium is chemically inactive . Under normal conditions, it reacts only with fluorine (from non-metals), forming a mixture of fluorides. At high temperatures (above 600°C) it interacts with oxygen, halogens, nitrogen, silicon, boron, sulfur, phosphorus:
4Cr + 3O2 = 2Cr2O3
2Cr + 3Cl2 = 2CrCl3
2Cr + N2 = 2CrN
2Cr + 3S = Cr2S3
It passivates in nitric and concentrated sulfuric acids, covered with a protective oxide film. It dissolves in hydrochloric and dilute sulfuric acids, and if the acid is completely freed from dissolved oxygen, chromium(II) salts are obtained, and if the reaction occurs in air, chromium(III) salts are obtained: Cr + 2HCl = CrCl2 + H2; 2 Cr + 6 HCl + O 2 = 2 CrCl 3 + 2 H 2 O + H 2
MANGANESE
Mn, chemical element with atomic number 25, atomic mass 54.9. Chemical symbol for the element Mn
pronounced the same way as the name of the element itself. Natural manganese consists only of the nuclide 55Mn. The configuration of the two outer electronic layers of the manganese atom is 3s2p6d54s2. In the periodic table, manganese is included in group VIIB and is located in the 4th period. Forms compounds in oxidation states from +2 to +7, the most stable oxidation states are +2 and +7. Manganese, like many other transition metals, also has compounds containing manganese atoms in oxidation state 0.
Manganese in its compact form is a hard, silvery-white, brittle metal.
Chemical properties
Manganese is an active metal.
1. Interaction with non-metals
When metallic manganese reacts with various non-metals, manganese (II) compounds are formed:
Mn + C2 = MnCl2 (manganese (II) chloride);
Mn + S = MnS (manganese (II) sulfide);
3Mn + 2 P = Mn3P2 (manganese (II) phosphide);
3Mn + N2 = Mn3N2 (manganese (II) nitride);
2Mn + N2 = Mn2Si (manganese (II) silicide).
2. Interaction with water
At room temperature it reacts very slowly with water, when heated at a moderate speed:
Mn + 2H2O = MnO2 + 2H2
3. Interaction with acids
In the electrochemical voltage series of metals, manganese is located before hydrogen, it displaces hydrogen from solutions of non-oxidizing acids, and manganese (II) salts are formed:
Mn + 2HCl = MnCl2 + H2;
Mn + H2SO4 = MnSO4 + H2;
with dilute nitric acid forms manganese (II) nitrate and nitric oxide (II):
3Mn + 8HNO3 = 3Mn(NO3)2 + 2NO + 4H2O.
Concentrated nitric and sulfuric acids passivate manganese. Manganese dissolves in them only when heated, manganese (II) salts and acid reduction products are formed:
Mn + 2H2SO4 = MnSO4 + SO2 + 2H2O;
Mn + 4HNO3 = Mn(NO3)2 + 2NO2 + 2H2O
4. Recovery of metals from oxides
Manganese is an active metal, capable of displacing metals from their oxides:
5Mn + Nb2O5 = 5MnO + 2Nb.
font-size:14.0pt;color:#262626">If concentrated sulfuric acid is added to potassium permanganate KMnO4, the acidic oxide Mn2O7 is formed, which has strong oxidizing properties:
2KMnO4 + 2H2SO4 = 2KHSO4 + Mn2O7 + H2O.
Several acids correspond to manganese, of which the most important are the strong unstable permanganic acid H2MnO4 and permanganic acid HMnO4, the salts of which are manganates (for example, sodium manganate Na2MnO4) and permanganates (for example, potassium permanganate KMnO4), respectively.
Manganates (only alkali metal and barium manganates are known) can exhibit properties as oxidizing agents (more often) 2 NaI + Na 2 MnO 4 + 2 H 2 O = MnO 2 + I 2 + 4 NaOH , and reducing agents 2K2MnO4 + Cl2 = 2KMnO4 + 2KCl.
Permanganates are strong oxidizing agents. For example, potassium permanganate KMnO4 in an acidic environment oxidizes sulfur dioxide SO2 to sulfate:
2KMnO4 + 5SO2 +2H2O = K2SO4 + 2MnSO4 + 2H2SO4.
Application:more than 90% of the manganese produced goes to the ferrous metallurgy. Manganese is used as an additive to steels for their deoxidation, desulfurization (this removes unwanted impurities from the steel - oxygen, sulfur and others), as well as for alloying steels, i.e. improving their mechanical and corrosion properties. Manganese is also used in copper, aluminum and magnesium alloys. Manganese coatings on metal surfaces provide anti-corrosion protection. To apply thin manganese coatings, the highly volatile and thermally unstable binuclear decacarbonyl Mn2(CO)10 is used.
The concept of alloys.
A characteristic feature of metals is their ability to form alloys with each other or with non-metals. To create an alloy, a mixture of metals is usually melted and then cooled at different rates, which are determined by the nature of the components and how they interact with temperature. Sometimes alloys are produced by sintering fine metal powders without resorting to melting (powder metallurgy). So alloys are products of the chemical interaction of metals.
The crystal structure of alloys is in many ways similar to pure metals, which, interacting with each other during melting and subsequent crystallization, form: a) chemical compounds called intermetallic compounds; b) solid solutions; c) a mechanical mixture of component crystals.
Modern technology uses a huge number of alloys, and in the vast majority of cases they consist not of two, but of three, four or more metals. It is interesting that the properties of alloys often differ sharply from the properties of the individual metals that form them. Thus, an alloy containing 50% bismuth, 25% lead, 12.5% tin and 12.5% cadmium melts at only 60.5 degrees Celsius, while the alloy components have melting points of 271, 327, 232, and 321 degrees Celsius. The hardness of tin bronze (90% copper and 10% tin) is three times that of pure copper, and the coefficient of linear expansion of iron-nickel alloys is 10 times less than that of pure components.
However, some impurities deteriorate the quality of metals and alloys. It is known, for example, that cast iron (an alloy of iron and carbon) does not have the strength and hardness that are characteristic of steel. In addition to carbon, the properties of steel are affected by the addition of sulfur and phosphorus, which increase its brittleness.
Among the properties of alloys, the most important for practical use are heat resistance, corrosion resistance, mechanical strength, etc. For aviation, light alloys based on magnesium, titanium or aluminum are of great importance, for the metalworking industry - special alloys containing tungsten, cobalt, and nickel. In electronic technology, alloys are used, the main component of which is copper. Super-powerful magnets were obtained using the products of the interaction of cobalt, samarium and other rare earth elements, and alloys that superconduct at low temperatures were based on intermetallic compounds formed by niobium with tin, etc.
Tasks to consolidate and test knowledge
Control questions:
1. How to determine the oxidation states of metals of secondary subgroups?
2. What oxidation states are most characteristic of iron?
3. Give the formulas of iron oxides and their corresponding hydroxides.
4. Describe the acid-base properties of iron (II) and iron hydroxides
(III)?
5. What oxidation states are characteristic of chromium? Which ones are the most stable?
6. Name the formulas of chromium oxides and hydroxides and characterize their acid-base properties.
7. How do the redox properties of chromium compounds change with
An increase in its oxidation state?
8. Write the formulas of chromic and dichromic acids.
9. What oxidation states does manganese exhibit in compounds? Which ones are the most stable?
10. Write the formulas of chromium oxides and hydroxides and characterize their acid-base properties and redox properties.
11. How do the redox properties of manganese compounds change with increasing degree of oxidation?
Since Fe2+ is easily oxidized to Fe+3:
Fe+2 – 1e = Fe+3
Thus, a freshly obtained greenish precipitate of Fe(OH)2 in air very quickly changes color - turns brown. The color change is explained by the oxidation of Fe(OH)2 to Fe(OH)3 by atmospheric oxygen:
4Fe+2(OH)2 + O2 + 2H2O = 4Fe+3(OH)3.
Divalent iron salts also exhibit reducing properties, especially when exposed to oxidizing agents in an acidic environment. For example, iron (II) sulfate reduces potassium permanganate in a sulfuric acid medium to manganese (II) sulfate:
10Fe+2SO4 + 2KMn+7O4 + 8H2SO4 = 5Fe+32(SO4)3 + 2Mn+2SO4 + K2SO4 + 8H2O.
Qualitative reaction to iron (II) cation.
The reagent for determining the iron cation Fe2+ is potassium hexacyano(III) ferrate (red blood salt) K3:
3FeSO4 + 2K3 = Fe32¯ + 3K2SO4.
When 3- ions interact with iron cations Fe2+, a dark blue precipitate is formed - Turnbull blue:
3Fe2+ +23- = Fe32¯
Iron(III) compounds
Iron(III) oxide Fe2O3– brown powder, insoluble in water. Iron (III) oxide is obtained:
A) decomposition of iron (III) hydroxide:
2Fe(OH)3 = Fe2O3 + 3H2O
B) oxidation of pyrite (FeS2):
4Fe+2S2-1 + 11O20 = 2Fe2+3O3 + 8S+4O2-2.
Fe+2 – 1e ® Fe+3
2S-1 – 10e ® 2S+4
O20 + 4e ® 2O-2 11e
Iron (III) oxide exhibits amphoteric properties:
A) interacts with solid alkalis NaOH and KOH and with sodium and potassium carbonates at high temperatures:
Fe2O3 + 2NaOH = 2NaFeO2 + H2O,
Fe2O3 + 2OH- = 2FeO2- + H2O,
Fe2O3 + Na2CO3 = 2NaFeO2 + CO2.
Sodium ferrite
Iron(III) hydroxide obtained from iron (III) salts by reacting them with alkalis:
FeCl3 + 3NaOH = Fe(OH)3¯ + 3NaCl,
Fe3+ + 3OH- = Fe(OH)3¯.
Iron (III) hydroxide is a weaker base than Fe(OH)2 and exhibits amphoteric properties (with a predominance of basic ones). When interacting with dilute acids, Fe(OH)3 easily forms the corresponding salts:
Fe(OH)3 + 3HCl « FeCl3 + H2O
2Fe(OH)3 + 3H2SO4 « Fe2(SO4)3 + 6H2O
Fe(OH)3 + 3H+ « Fe3+ + 3H2O
Reactions with concentrated solutions of alkalis occur only with prolonged heating. In this case, stable hydrocomplexes with a coordination number of 4 or 6 are obtained:
Fe(OH)3 + NaOH = Na,
Fe(OH)3 + OH- = -,
Fe(OH)3 + 3NaOH = Na3,
Fe(OH)3 + 3OH- = 3-.
Compounds with the oxidation state of iron +3 exhibit oxidizing properties, since under the influence of reducing agents Fe+3 is converted into Fe+2:
Fe+3 + 1e = Fe+2.
For example, iron (III) chloride oxidizes potassium iodide to free iodine:
2Fe+3Cl3 + 2KI = 2Fe+2Cl2 + 2KCl + I20
Qualitative reactions to iron (III) cation
A) The reagent for detecting the Fe3+ cation is potassium hexacyano(II) ferrate (yellow blood salt) K2.
When 4- ions interact with Fe3+ ions, a dark blue precipitate is formed - Prussian blue:
4FeCl3 + 3K4 « Fe43¯ +12KCl,
4Fe3+ + 34- = Fe43¯.
B) Fe3+ cations are easily detected using ammonium thiocyanate (NH4CNS). As a result of the interaction of CNS-1 ions with iron (III) cations Fe3+, low-dissociation iron (III) thiocyanate of blood-red color is formed:
FeCl3 + 3NH4CNS « Fe(CNS)3 + 3NH4Cl,
Fe3+ + 3CNS1- « Fe(CNS)3.
Application and biological role of iron and its compounds.
The most important iron alloys - cast iron and steel - are the main structural materials in almost all branches of modern production.
Iron (III) chloride FeCl3 is used for water purification. In organic synthesis, FeCl3 is used as a catalyst. Iron nitrate Fe(NO3)3 9H2O is used for dyeing fabrics.
Iron is one of the most important microelements in the human and animal body (the adult human body contains about 4 g of Fe in the form of compounds). It is part of hemoglobin, myoglobin, various enzymes and other complex iron-protein complexes that are found in the liver and spleen. Iron stimulates the function of hematopoietic organs.
List of used literature:
1. “Chemistry. Tutor's allowance." Rostov-on-Don. "Phoenix". 1997
2. “Handbook for applicants to universities.” Moscow. "Higher School", 1995.
3. E.T. Oganesyan. “Guide to chemistry for university applicants.” Moscow. 1994
The human body contains about 5 g of iron, most of it (70%) is part of blood hemoglobin.
Physical properties
In its free state, iron is a silvery-white metal with a grayish tint. Pure iron is ductile and has ferromagnetic properties. In practice, iron alloys - cast iron and steel - are usually used.
Fe is the most important and most abundant element of the nine d-metals of the Group VIII subgroup. Together with cobalt and nickel it forms the “iron family”.
When forming compounds with other elements, it often uses 2 or 3 electrons (B = II, III).
Iron, like almost all d-elements of group VIII, does not exhibit a higher valency equal to the group number. Its maximum valency reaches VI and appears extremely rarely.
The most typical compounds are those in which the Fe atoms are in oxidation states +2 and +3.
Methods for obtaining iron
1. Technical iron (alloyed with carbon and other impurities) is obtained by carbothermic reduction of its natural compounds according to the following scheme:
Recovery occurs gradually, in 3 stages:
1) 3Fe 2 O 3 + CO = 2Fe 3 O 4 + CO 2
2) Fe 3 O 4 + CO = 3FeO + CO 2
3) FeO + CO = Fe + CO 2
The cast iron resulting from this process contains more than 2% carbon. Subsequently, cast iron is used to produce steel - iron alloys containing less than 1.5% carbon.
2. Very pure iron is obtained in one of the following ways:
a) decomposition of Fe pentacarbonyl
Fe(CO) 5 = Fe + 5СО
b) reduction of pure FeO with hydrogen
FeO + H 2 = Fe + H 2 O
c) electrolysis of aqueous solutions of Fe +2 salts
FeC 2 O 4 = Fe + 2CO 2
iron(II) oxalate
Chemical properties
Fe is a metal of medium activity and exhibits general properties characteristic of metals.
A unique feature is the ability to “rust” in humid air:
In the absence of moisture with dry air, iron begins to react noticeably only at T > 150°C; upon calcination, “iron scale” Fe 3 O 4 is formed:
3Fe + 2O 2 = Fe 3 O 4
Iron does not dissolve in water in the absence of oxygen. At very high temperatures, Fe reacts with water vapor, displacing hydrogen from water molecules:
3 Fe + 4H 2 O(g) = 4H 2
The mechanism of rusting is electrochemical corrosion. The rust product is presented in a simplified form. In fact, a loose layer of a mixture of oxides and hydroxides of variable composition is formed. Unlike the Al 2 O 3 film, this layer does not protect iron from further destruction.
Types of corrosion
Protecting iron from corrosion
1. Interaction with halogens and sulfur at high temperatures.
2Fe + 3Cl 2 = 2FeCl 3
2Fe + 3F 2 = 2FeF 3
Fe + I 2 = FeI 2
Compounds are formed in which the ionic type of bond predominates.
2. Interaction with phosphorus, carbon, silicon (iron does not directly combine with N2 and H2, but dissolves them).
Fe + P = Fe x P y
Fe + C = Fe x C y
Fe + Si = Fe x Si y
Substances of variable composition are formed, such as berthollides (the covalent nature of the bond predominates in the compounds)
3. Interaction with “non-oxidizing” acids (HCl, H 2 SO 4 dil.)
Fe 0 + 2H + → Fe 2+ + H 2
Since Fe is located in the activity series to the left of hydrogen (E° Fe/Fe 2+ = -0.44 V), it is capable of displacing H 2 from ordinary acids.
Fe + 2HCl = FeCl 2 + H 2
Fe + H 2 SO 4 = FeSO 4 + H 2
4. Interaction with “oxidizing” acids (HNO 3, H 2 SO 4 conc.)
Fe 0 - 3e - → Fe 3+
Concentrated HNO 3 and H 2 SO 4 “passivate” iron, so at ordinary temperatures the metal does not dissolve in them. With strong heating, slow dissolution occurs (without releasing H 2).
In section HNO 3 iron dissolves, goes into solution in the form of Fe 3+ cations and the acid anion is reduced to NO*:
Fe + 4HNO 3 = Fe(NO 3) 3 + NO + 2H 2 O
Very soluble in a mixture of HCl and HNO 3
5. Relation to alkalis
Fe does not dissolve in aqueous solutions of alkalis. It reacts with molten alkalis only at very high temperatures.
6. Interaction with salts of less active metals
Fe + CuSO 4 = FeSO 4 + Cu
Fe 0 + Cu 2+ = Fe 2+ + Cu 0
7. Interaction with gaseous carbon monoxide (t = 200°C, P)
Fe (powder) + 5CO (g) = Fe 0 (CO) 5 iron pentacarbonyl
Fe(III) compounds
Fe 2 O 3 - iron (III) oxide.
Red-brown powder, n. R. in H 2 O. In nature - “red iron ore”.
Methods of obtaining:
1) decomposition of iron (III) hydroxide
2Fe(OH) 3 = Fe 2 O 3 + 3H 2 O
2) pyrite firing
4FeS 2 + 11O 2 = 8SO 2 + 2Fe 2 O 3
3) nitrate decomposition
Chemical properties
Fe 2 O 3 is a basic oxide with signs of amphotericity.
I. The main properties are manifested in the ability to react with acids:
Fe 2 O 3 + 6H + = 2Fe 3+ + ZH 2 O
Fe 2 O 3 + 6HCI = 2FeCI 3 + 3H 2 O
Fe 2 O 3 + 6HNO 3 = 2Fe(NO 3) 3 + 3H 2 O
II. Weak acid properties. Fe 2 O 3 does not dissolve in aqueous solutions of alkalis, but when fused with solid oxides, alkalis and carbonates, ferrites form:
Fe 2 O 3 + CaO = Ca(FeO 2) 2
Fe 2 O 3 + 2NaOH = 2NaFeO 2 + H 2 O
Fe 2 O 3 + MgCO 3 = Mg(FeO 2) 2 + CO 2
III. Fe 2 O 3 - feedstock for the production of iron in metallurgy:
Fe 2 O 3 + ZS = 2Fe + ZSO or Fe 2 O 3 + ZSO = 2Fe + ZSO 2
Fe(OH) 3 - iron (III) hydroxide
Methods of obtaining:
Obtained by the action of alkalis on soluble Fe 3+ salts:
FeCl 3 + 3NaOH = Fe(OH) 3 + 3NaCl
At the time of preparation, Fe(OH) 3 is a red-brown mucous-amorphous sediment.
Fe(III) hydroxide is also formed during the oxidation of Fe and Fe(OH) 2 in moist air:
4Fe + 6H 2 O + 3O 2 = 4Fe(OH) 3
4Fe(OH) 2 + 2H 2 O + O 2 = 4Fe(OH) 3
Fe(III) hydroxide is the end product of the hydrolysis of Fe 3+ salts.
Chemical properties
Fe(OH) 3 is a very weak base (much weaker than Fe(OH) 2). Shows noticeable acidic properties. Thus, Fe(OH) 3 has an amphoteric character:
1) reactions with acids occur easily:
2) fresh precipitate of Fe(OH) 3 dissolves in hot conc. solutions of KOH or NaOH with the formation of hydroxo complexes:
Fe(OH) 3 + 3KOH = K 3
In an alkaline solution, Fe(OH) 3 can be oxidized to ferrates (salts of iron acid H 2 FeO 4 not released in the free state):
2Fe(OH) 3 + 10KOH + 3Br 2 = 2K 2 FeO 4 + 6KBr + 8H 2 O
Fe 3+ salts
The most practically important are: Fe 2 (SO 4) 3, FeCl 3, Fe(NO 3) 3, Fe(SCN) 3, K 3 4 - yellow blood salt = Fe 4 3 Prussian blue (dark blue precipitate)
b) Fe 3+ + 3SCN - = Fe(SCN) 3 thiocyanate Fe(III) (blood red solution)
