Oxidation of iron hydroxide 3. Reducing properties
The human body contains about 5 g of iron, most of it (70%) is part of the hemoglobin in the blood.
Physical properties
In the free state, iron is a silvery-white metal with a grayish tinge. Pure iron is ductile and has ferromagnetic properties. In practice, iron alloys are commonly used - cast irons and steels.
Fe is the most important and most common element of the nine d-metals of the secondary subgroup of group VIII. Together with cobalt and nickel, it forms the "iron family".
When forming compounds with other elements, it often uses 2 or 3 electrons (B \u003d II, III).
Iron, like almost all d-elements of group VIII, does not show a higher valency equal to the group number. Its maximum valency reaches VI and is extremely rare.
The most typical compounds are those in which the Fe atoms are in the +2 and +3 oxidation states.
Methods for obtaining iron
1. Commercial iron (in an alloy with carbon and other impurities) is obtained by carbothermal reduction of its natural compounds according to the 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 \u003d Fe + CO 2
The cast iron resulting from this process contains more than 2% carbon. In the future, steels are obtained from cast iron - iron alloys containing less than 1.5% carbon.
2. Very pure iron is obtained in one of the following ways:
a) decomposition of pentacarbonyl Fe
Fe(CO) 5 = Fe + 5CO
b) hydrogen reduction of pure FeO
FeO + H 2 \u003d Fe + H 2 O
c) electrolysis of aqueous solutions of Fe +2 salts
FeC 2 O 4 \u003d Fe + 2СO 2
iron(II) oxalate
Chemical properties
Fe - a metal of medium activity, 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 noticeably react only at T > 150°C; when calcined, “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) \u003d 4H 2
The rusting process in its mechanism 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 the iron from further destruction.
Types of corrosion
Corrosion protection of iron
1. Interaction with halogens and sulfur at high temperature.
2Fe + 3Cl 2 = 2FeCl 3
2Fe + 3F 2 = 2FeF 3
Fe + I 2 \u003d 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 N 2 and H 2, but dissolves them).
Fe + P = Fe x P y
Fe + C = Fe x C y
Fe + Si = FexSiy
Substances of variable composition are formed, since berthollides (the covalent nature of the bond prevails 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+ \u003d -0.44V), it is able to displace H 2 from ordinary acids.
Fe + 2HCl \u003d FeCl 2 + H 2
Fe + H 2 SO 4 \u003d 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 release of H 2).
In razb. 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 \u003d Fe (NO 3) 3 + NO + 2H 2 O
It dissolves very well in a mixture of HCl and HNO 3
5. Attitude 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 \u003d 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) \u003d Fe 0 (CO) 5 iron pentacarbonyl
Fe(III) compounds
Fe 2 O 3 - iron oxide (III).
Red-brown powder, n. R. in H 2 O. In nature - "red iron ore".
Ways to get:
1) decomposition of iron hydroxide (III)
2Fe(OH) 3 = Fe 2 O 3 + 3H 2 O
2) pyrite roasting
4FeS 2 + 11O 2 \u003d 8SO 2 + 2Fe 2 O 3
3) decomposition of nitrate
Chemical properties
Fe 2 O 3 is a basic oxide with signs of amphoterism.
I. The main properties are manifested in the ability to react with acids:
Fe 2 O 3 + 6H + = 2Fe 3+ + ZN 2 O
Fe 2 O 3 + 6HCI \u003d 2FeCI 3 + 3H 2 O
Fe 2 O 3 + 6HNO 3 \u003d 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 are formed:
Fe 2 O 3 + CaO \u003d Ca (FeO 2) 2
Fe 2 O 3 + 2NaOH \u003d 2NaFeO 2 + H 2 O
Fe 2 O 3 + MgCO 3 \u003d Mg (FeO 2) 2 + CO 2
III. Fe 2 O 3 - feedstock for iron production in metallurgy:
Fe 2 O 3 + ZS \u003d 2Fe + ZSO or Fe 2 O 3 + ZSO \u003d 2Fe + ZSO 2
Fe (OH) 3 - iron (III) hydroxide
Ways to get:
Obtained by the action of alkalis on soluble salts Fe 3+:
FeCl 3 + 3NaOH \u003d Fe (OH) 3 + 3NaCl
At the time of receipt of Fe(OH) 3 - red-brown mucosamorphous precipitate.
Fe (III) hydroxide is also formed during the oxidation of Fe and Fe (OH) 2 in humid air:
4Fe + 6H 2 O + 3O 2 \u003d 4Fe (OH) 3
4Fe(OH) 2 + 2Н 2 O + O 2 = 4Fe(OH) 3
Fe(III) hydroxide is the end product of 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 proceed easily:
2) a fresh precipitate of Fe(OH) 3 is dissolved in hot conc. solutions of KOH or NaOH with the formation of hydroxo complexes:
Fe (OH) 3 + 3KOH \u003d K 3
In an alkaline solution, Fe (OH) 3 can be oxidized to ferrates (salts of iron acid H 2 FeO 4 not isolated 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 \u003d Fe 4 3 Prussian blue (dark blue precipitate)
b) Fe 3+ + 3SCN - \u003d Fe (SCN) 3 Fe (III) thiocyanate (blood red solution)
The inorganic compound iron hydroxide 3 has the chemical formula Fe(OH)2. It belongs to a number of amphoteric in which the properties characteristic of bases predominate. In appearance, this substance is a white crystals, which gradually darken with prolonged exposure to the open air. There are options for crystals of a greenish tint. In everyday life, the substance can be observed by everyone 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 one or another 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. So, for example, having the properties of a base, it comes into contact with various acids. In particular, sulfuric iron 3 in the course of the reaction leads to the production of (III). Since this reaction can take place by conventional open-air roasting, this inexpensive sulfate is used in both laboratory and industrial settings.
During the reaction with its result is the formation of iron (II) chloride.
In some cases, iron hydroxide 3 can also exhibit acidic properties. So, for example, when interacting with a highly concentrated (the concentration should be at least 50%) sodium hydroxide solution, sodium tetrahydroxoferrate (II) is obtained, which precipitates. True, for such a reaction to proceed, it is necessary to provide rather complex conditions: the reaction must occur under conditions of solution boiling in a nitrogen atmosphere.
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 proceeding with the experiment, it is imperative to recall the safety rules for conducting such experiments. These rules apply to all handling of acid-base solutions. The main thing here is to provide reliable protection and avoid drops of solutions on the mucous membranes and skin.
So, you can get hydroxide during the reaction, in which iron (III) chloride and KOH - potassium hydroxide interact. This method is the most common for the formation of insoluble bases. When these substances interact, the usual exchange reaction occurs, as a result of which a brown precipitate is obtained. This precipitate is the desired substance.
The use of iron hydroxide in industrial production is quite wide. The most common is its use as an active substance in iron-nickel type batteries. In addition, the compound is used in metallurgy for the production of various metal alloys, as well as in electroplating and automotive manufacturing.
Iron(III) oxide
Iron(II) hydroxide
Ferrous compounds
Chemical properties
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
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:
Fe + S – t ° ® FeS
4) Iron dissolves easily in hydrochloric and dilute sulfuric acids:
Fe + 2HCl ® FeCl 2 + H 2
Fe + H 2 SO 4 (razb.) ® FeSO 4 + H 2
In concentrated oxidizing acids, iron dissolves only when heated.
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 iron).
5) Iron displaces metals to the right of it in a series of stresses from solutions of their salts.
Fe + CuSO 4 ® FeSO 4 + Cu¯
It is formed by the action of alkali solutions on iron (II) salts without air access:
FeCl + 2KOH ® 2KCl + Fe(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
When Fe (OH) 2 is calcined without air access, iron oxide (II) FeO is formed:
Fe(OH) 2 - t ° ® FeO + H 2 O
In the presence of atmospheric oxygen, a 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
Iron (II) compounds have reducing properties, they are easily converted into iron (III) compounds under the action 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 (coordination number = 6):
FeCl 2 + 6NH 3 ® Cl 2
Fe(CN) 2 + 4KCN ® K 4 (yellow blood salt)
Qualitative reaction for Fe 2+
Under the action of potassium hexacyanoferrate (III) K 3 (red blood salt) on solutions of ferrous salts, a blue precipitate (turnbull blue) is formed:
3FeSO 4 + 2K 3 ® Fe 3 2 ¯ + 3K 2 SO 4
3Fe 2+ + 3SO 4 2- +6K + + 2 3- ® Fe 3 2 ¯ + 6K + + 3SO 4 2-
3Fe 2+ + 2 3- ® Fe 3 2 ¯
Ferric compounds
It is formed during the combustion of iron sulfides, for example, during the firing of pyrite:
4FeS 2 + 11O 2 ® 2Fe 2 O 3 + 8SO 2
or when calcining iron salts:
2FeSO 4 - t ° ® Fe 2 O 3 + SO 2 + SO 3
Fe 2 O 3 - basic oxide, showing amphoteric properties to a small extent
Fe 2 O 3 + 6HCl - t ° ® 2FeCl 3 + 3H 2 O
Fe 2 O 3 + 6H + - t ° ® 2Fe 3+ + 3H 2 O
Fe 2 O 3 + 2NaOH + 3H 2 O - t ° ® 2Na
Fe 2 O 3 + 2OH - + 3H 2 O ® 2 -
It is formed by the action of alkali solutions on ferric iron salts: it precipitates as a red-brown precipitate
Fe(NO 3) 3 + 3KOH ® Fe(OH) 3 ¯ + 3KNO 3
Fe 3+ + 3OH - ® Fe(OH) 3 ¯
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+ holds 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 very strongly hydrolyzed. For a better understanding of the materials in this section, it is recommended to watch the video clip (only available on CDROM). Hydrolysis also explains the color of solutions of Fe (III) salts: despite the fact that the Fe 3+ ion is almost colorless, the 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 a weakly pronounced amphoterism: it dissolves in dilute acids and in 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 - ® -
Iron (III) compounds are weak oxidizing agents, they react with strong reducing agents:
2Fe +3 Cl 3 + H 2 S -2 ® S 0 + 2Fe +2 Cl 2 + 2HCl
Qualitative reactions for Fe 3+
1) Under the action of potassium hexacyanoferrate (II) K 4 (yellow blood salt) on solutions of ferric salts, a blue precipitate (Prussian blue) is formed:
4FeCl 3 +3K 4 ® Fe 4 3 ¯ + 12KCl
4Fe 3+ + 12C l - + 12K + + 3 4- ® Fe 4 3 ¯ + 12K + + 12C l -
4Fe 3+ + 3 4- ® Fe 4 3 ¯
2) When potassium or ammonium thiocyanate is added to a solution containing Fe 3+ ions, an intense blood-red color of iron(III) thiocyanate appears:
FeCl 3 + 3NH 4 CNS « 3NH 4 Cl + Fe(CNS) 3
(when interacting with Fe 2+ ions with thiocyanates, the solution remains almost colorless).
Since Fe2+ is easily oxidized to Fe+3:
Fe+2 – 1e = Fe+3
So, 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 environment to manganese (II) sulfate:
10Fe+2SO4 + 2KMn+7O4 + 8H2SO4 = 5Fe+32(SO4)3 + 2Mn+2SO4 + K2SO4 + 8H2O.
Qualitative reaction to the iron (II) cation.
The reagent for determining the iron cation Fe2+ is hexacyano(III) potassium ferrate (red blood salt) K3:
3FeSO4 + 2K3 = Fe32¯ + 3K2SO4.
When 3- ions interact with Fe2+ iron cations, a dark blue precipitate is formed - turnbull blue:
3Fe2+ +23- = Fe32¯
Iron(III) compounds
Iron oxide (III) Fe2O3- brown powder, insoluble in water. Iron oxide (III) 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 oxide (III) exhibits amphoteric properties:
A) interacts with solid alkalis NaOH and KOH and with sodium and potassium carbonates at high temperature:
Fe2O3 + 2NaOH = 2NaFeO2 + H2O,
Fe2O3 + 2OH- = 2FeO2- + H2O,
Fe2O3 + Na2CO3 = 2NaFeO2 + CO2.
sodium ferrite
Iron(III) hydroxide obtained from iron (III) salts when they interact with alkalis:
FeCl3 + 3NaOH = Fe(OH)3¯ + 3NaCl,
Fe3+ + 3OH- = Fe(OH)3¯.
Iron hydroxide (III) 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 alkali solutions proceed 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 an iron oxidation state of +3 exhibit oxidizing properties, since under the action of reducing agents Fe + 3 turns into Fe + 2:
Fe+3 + 1e = Fe+2.
So, for example, iron (III) chloride oxidizes potassium iodide to free iodine:
2Fe+3Cl3 + 2KI = 2Fe+2Cl2 + 2KCl + I20
Qualitative reactions to the iron (III) cation
A) The reagent for detecting the Fe3+ cation is hexacyano(II) potassium 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-dissociating blood-red iron (III) thiocyanate 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 irons and steels - are the main structural materials in almost all branches of modern production.
Iron (III) chloride FeCl3 is used for water treatment. 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 trace elements in the human and animal body (in the body of an adult it 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 the hematopoietic organs.
List of used literature:
1. “Chemistry. Allowance tutor. Rostov-on-Don. "Phoenix". 1997
2. "Handbook for applicants to universities." Moscow. "High School", 1995.
3. E.T. Oganesyan. "A guide to chemistry entering universities." Moscow. 1994
