Cheat sheet on inorganic chemistry. Iron (III) compounds Iron hydroxide 2 when heated decomposes into

Plan.
Introduction.

Conclusion.
Bibliography.

Introduction.
Iron (II) hydroxide is an inorganic substance with the formula Fe(OH) 2, an iron compound. Occurs in nature in the form of the mineral amakinite. This mineral contains impurities of magnesium and manganese (empirical formula Fe 0.7 Mg 0.2 Mn 0.1 (OH) 2). The color of the mineral is yellow-green or light green, Mohs hardness 3.5-4, density 2.925-2.98 g/cm?. Amphoteric hydroxide with a predominance of basic properties. The crystalline substance is white (sometimes with a greenish tint) and darkens over time in air. It is one of the intermediate compounds in the rusting of iron. Iron (II) hydroxide is used in the manufacture of the active mass of iron-nickel batteries.
The purpose of this work is to obtain iron (II) hydroxide and study its properties.
During the work, the following tasks were set:

The bases are classified according to a number of characteristics.

The division into soluble and insoluble bases almost completely coincides with the division into strong and weak bases, or hydroxides of metals and transition elements

1. Exchange reaction between salt and alkali in solution.
This is the most common method of obtaining both soluble (alkalis) and insoluble bases, for which it is the only laboratory method of preparation.
Preparation of strong alkali:
Na 2 CO 3 + Ca(OH) 2 = CaCO 3 + 2NaOH
Preparation of an insoluble base:
CuSO 4 + 2KOH = Cu(OH) 2 + K 2 SO 4
2. Hydration of basic oxides.
This method can only produce strong alkalis, i.e. hydroxides of alkali and alkaline earth metals. For example:
BaO + H 2 O = Ba(OH) 2
3. Interaction of metals with water.
Under normal conditions, only alkali and alkaline earth metals react with water. In this case, the corresponding alkali and hydrogen are formed:
Ba + 2H 2 O = Ba(OH) 2 + H 2
4. Electrolysis of aqueous salt solutions.
In industry, NaOH and KOH are produced by electrolysis of aqueous solutions of potassium and sodium chlorides.
KCl + 2H 2 O = 2KOH + H 2 + Cl 2

Physical properties.
Alkalis (sodium, potassium, lithium hydroxides) form hard, white, very hygroscopic crystals. The melting point of NaOH is 322°C, KOH is 405°C, and LiOH is 473°C. The crystal lattices of sodium hydroxide are cubic, like NaCl, and those of potassium hydroxide are tetragonal.
Hydroxides of calcium, magnesium, beryllium, barium form white powders, also quite hygroscopic, but not as much as alkalis. They form a hexagonal crystal lattice; their melting temperatures are not high due to decomposition into oxide and water.
Hydroxides of other metals (aluminum, copper, zinc, etc.) form precipitates of different colors, most often white. Colored hydroxides are used as pigments in the production of enamels and glazes.
Only alkalis are well soluble in water, significantly less than the base of metals of the second group (main subgroup), and all the rest are practically insoluble in water.
Chemical properties.
Metal hydroxides exhibit different chemical properties depending on the activity of the metal that is included in the hydroxide.
Bases react with acids to form salt and water. This reaction is called a neutralization reaction, because after its completion the medium becomes close to neutral:
2KOH+H 2 SO 4 =K 2 SO 4 +2H 2 O
If the base is soluble in water, then it reacts with acidic and amphoteric oxides, forming salt and water:
2KOH+SO 3 =K 2 SO 4 +H 2 O
2RbOH+ZnO=Rb 2 ZnO 2 +H 2 O.
Also, water-soluble bases can react with salts to form a new salt and a new base, provided that the new base is insoluble:
2NaOH+CuSO 4 =Cu(OH) 2 +Na 2 SO 4
A special group of hydroxides consists of amphoteric hydroxides. During dissociation, they simultaneously form both H + cations and OH - hydroxide ions. These include, for example, Zn(OH) 2, Al(OH) 3, Be(OH) 2, Pb(OH) 2 and others.
Amphoteric hydroxides react with both acid and alkali solutions. When interacting with bases, they exhibit the properties of acids, and when interacting with acids, they exhibit the properties of bases:
Zn(OH) 2 +H 2 SO 4 =ZnSO 4 +2H 2 O
Cr(OH) 3 + 3NaOH = Na 3 (sodium hexahydroxochromate (III))
Al(OH) 3 + NaOH = Na (sodium tetrahydroxoaluminate (III))
From the point of view of the theory of electrolytic dissociation, the properties of base solutions (changes in the color of indicators, soapiness to the touch, interaction with acids, acid oxides and salts) are determined by the presence of OH - hydroxide ions. The bases are colored by the indicators phenolphthalein - crimson, litmus - blue.
Insoluble bases decompose when heated into metal oxide and water
2Fe(OH) 3 = Fe 2 O 3 + 3H 2 O
Cu(OH) 2 = CuO + H 2 O

Iron has been known since ancient times. The oldest iron objects found by archaeologists date back to 4 thousand BC. e. It is believed that the material from which man made the first iron products was meteorite iron. It is no coincidence that in many languages iron was called “heavenly metal”, “dripping from the sky”, etc. The first scientific evidence that “iron stones fall from the sky” was provided in 1775 by the St. Petersburg academic geographer and traveler Peter Simon Pallas (1741–1811), who brought a block of iron meteorite weighing 600 kg to St. Petersburg. The largest iron meteorite found on Earth is the Gobe meteorite, weighing about 60 tons, which was discovered in 1920 in South-West Africa. The largest iron meteorite that was observed falling is located in Moscow in the Museum of the Russian Academy of Sciences. When it fell (October 18, 1816, Far East), the meteorite broke and two fragments weighing 256 kg were found. There was a time when iron on earth was valued much more than gold. Soviet historian G. Areshyan studied the influence of iron on the ancient culture of the Mediterranean countries. He gives the following proportion: 1: 160: 1280: 6400. This is the ratio of the values of copper, silver, gold and iron among the ancient Hittites. As Homer testifies in the Odyssey, the winner of the games arranged by Achilles was rewarded with a piece of gold and a piece of iron. Iron was equally necessary for both the warrior and the plowman, and practical need, as we know, is the best engine of production and technical progress.
The term “Iron Age” was introduced into science in the middle of the 19th century. Danish archaeologist K.Yu. Thomsen. “Official” boundaries of this period of human history: from IX...VII centuries. BC. when iron metallurgy began to develop among many peoples and tribes of Europe and Asia, and before the emergence of a class society and state among these tribes. But if eras are named by the main material of the tools, then, obviously, the Iron Age continues today. How did our distant ancestors obtain iron? First, the so-called cheese-blowing method. Cheese furnaces were built directly on the ground, usually on the slopes of ravines and ditches. They looked like a pipe. This pipe was filled with charcoal and iron ore. The coal was lit, and the wind blowing into the slope of the ravine kept the coal burning. Iron ore was reduced, and a soft crust was obtained - iron with slag inclusions. Such iron was called welding iron; it contained some carbon and impurities transferred from the ore. Kritsa was forged. Pieces of slag fell off, and iron, riddled with slag threads, remained under the hammer. Various tools were forged from it. The age of wrought iron was long, but people of antiquity and the early Middle Ages were also familiar with other iron. The famous Damascus steel (or damask steel) was made in the East back in the time of Aristotle (IV century BC). But the technology of its production, as well as the process of making damask blades, was kept secret. Iron ores began to be smelted in Africa in the 1st millennium BC. Here iron ores come to the surface of the earth. Perhaps they were found in river sediments. In the river basin Zambezi archaeologists discovered clay blast furnaces, abandoned iron ore mines, and piles of slag. Local tribes moved from the Stone Age directly to the Iron Age, bypassing the Bronze Age. Over time, iron replaced other metals everywhere and became the main material for the manufacture of tools, weapons, mechanisms and other products. The “Iron Age” that began in those distant times continues to this day. Iron and its alloys account for about 95% of all metal products produced in the world. Now the bulk of iron is smelted in the form of cast iron and steel.

Iron is quite widespread in the earth's crust - it accounts for about 4.1% of the mass of the earth's crust (4th place among all elements, 2nd among metals). In the mantle and crust, iron is concentrated mainly in silicates, while its content is significant in basic and ultrabasic rocks, and low in acidic and intermediate rocks.
A large number of ores and minerals containing iron are known. Of greatest practical importance are red iron ore (hematite, Fe 2 O 3; contains up to 70% Fe), magnetic iron ore (magnetite, FeFe 2 O 4, Fe 3 O 4; contains 72.4% Fe), brown iron ore or limonite (goethite and hydrogoethite, respectively FeOOH and FeOOH·nH 2 O). Goethite and hydrogoethite are most often found in weathering crusts, forming so-called “iron hats”, the thickness of which reaches several hundred meters. They can also be of sedimentary origin, falling out of colloidal solutions in lakes or coastal areas of the seas. In this case, oolitic, or legume, iron ores are formed. Vivianite Fe 3 (PO 4) 2 8H 2 O is often found in them, forming black elongated crystals and radial aggregates.
Iron sulfides are also widespread in nature - pyrite FeS 2 (sulfur or iron pyrite) and pyrrhotite. They are not iron ore - pyrite is used to make sulfuric acid, and pyrrhotite often contains nickel and cobalt.
Russia ranks first in the world in terms of iron ore reserves. The iron content in sea water is 1·10?5 -1·10?8%.

Main deposits.

According to the US Geological Survey, the world's proven reserves of iron ore amount to about 178 billion tons. The main iron deposits are located in Brazil, Australia, the USA, Canada, Sweden, Venezuela, Liberia, Ukraine, France, and India. In Russia, iron is mined in the Kursk Magnetic Anomaly (KMA), the Kola Peninsula, Karelia and Siberia. Bottom ocean deposits, in which iron, together with manganese and other valuable metals, are found in nodules, have recently acquired a significant role.

Receipt.

In industry, iron is obtained from iron ore, mainly from hematite (Fe 2 O 3) and magnetite (FeO Fe 2 O 3).
There are various ways to extract iron from ores. The most common is the domain process.
The first stage of production is the reduction of iron with carbon in a blast furnace at a temperature of 2000°C. In a blast furnace, carbon in the form of coke, iron ore in the form of agglomerate or pellets, and flux (such as limestone) are fed from above, and are met by a stream of forced hot air from below.
In the furnace, carbon in the form of coke is oxidized to carbon monoxide. This oxide is formed during combustion in a lack of oxygen:
2C + O = 2CO
In turn, carbon monoxide reduces iron from the ore. To make this reaction go faster, heated carbon monoxide is passed through iron (III) oxide:
3CO + Fe 2 O 3 = 2Fe + 3CO 2
Flux is added to get rid of unwanted impurities (primarily silicates; for example, quartz) in the mined ore. A typical flux contains limestone (calcium carbonate) and dolomite (magnesium carbonate). To remove other impurities, other fluxes are used.
The effect of flux (in this case calcium carbonate) is that when it is heated, it decomposes to its oxide:
CaCO 3 = CaO + CO 2
Calcium oxide combines with silicon dioxide, forming slag - calcium metasilicate:
CaO + SiO 2 = CaSiO 3
Slag, unlike silicon dioxide, is melted in a furnace. Slag, lighter than iron, floats on the surface - this property allows the slag to be separated from the metal. The slag can then be used in construction and agriculture. The molten iron produced in a blast furnace contains quite a lot of carbon (cast iron). Except in cases where cast iron is used directly, it requires further processing.
Excess carbon and other impurities (sulfur, phosphorus) are removed from cast iron by oxidation in open-hearth furnaces or converters. Electric furnaces are also used for smelting alloy steels.
In addition to the blast furnace process, the process of direct iron production is common. In this case, pre-crushed ore is mixed with special clay, forming pellets. The pellets are fired and treated in a shaft furnace with hot methane conversion products, which contain hydrogen. Hydrogen easily reduces iron:
Fe 2 O 3 + 3H 2 = 2Fe + 3 H 2 O
in this case, the iron does not become contaminated with such impurities as sulfur and phosphorus, which are common impurities in coal. Iron is obtained in solid form and is subsequently melted in electric furnaces.
Chemically pure iron is obtained by electrolysis of solutions of its salts.
Use of iron.
Iron is the most important metal of modern technology. In its pure form, iron is practically not used due to its low strength, although in everyday life steel or cast iron products are often called “iron”. The bulk of iron is used in the form of alloys with very different compositions and properties. Iron alloys account for approximately 95% of all metal products. Carbon-rich alloys (over 2% by weight) - cast irons - are smelted in blast furnaces from iron-enriched ores. Various grades of steel (carbon content less than 2% by weight) are smelted from cast iron in open-hearth and electric furnaces and converters by oxidizing (burning out) excess carbon, removing harmful impurities (mainly S, P, O) and adding alloying elements. High-alloy steels (with a high content of nickel, chromium, tungsten and other elements) are smelted in electric arc and induction furnaces. For the production of steels and iron alloys for special purposes, new processes are used - vacuum, electroslag remelting, plasma and electron beam melting, and others. Methods are being developed for steel smelting in continuously operating units that ensure high quality metal and automation of the process.
Iron-based materials are created that can withstand high and low temperatures, vacuum and high pressures, aggressive environments, high alternating voltages, nuclear radiation, etc. The production of iron and its alloys is constantly growing.
Iron as an artistic material has been used since ancient times in Egypt, Mesopotamia, and India. Since the Middle Ages, numerous highly artistic iron products have been preserved in European countries (England, France, Italy, Russia and others) - forged fences, door hinges, wall brackets, weather vanes, chest frames, and lights. Forged through products made from rods and products made from expanded sheet iron (often with a mica lining) are distinguished by their flat shapes, a clear linear graphic silhouette and are effectively visible against a light-air background. In the 20th century, iron was used to make grilles, fences, openwork interior partitions, candlesticks, and monuments.

Physical properties.
Iron is a typical metal; in its free state it is silvery-white in color with a grayish tint. Pure metal is ductile; various impurities increase its hardness and brittleness. It has pronounced magnetic properties. The so-called “iron triad” is often distinguished - a group of three metals that have similar physical properties, atomic radii and electronegativity values.
Iron is characterized by polymorphism; it has four crystalline modifications:

Metallurgy does not distinguish?-Fe as a separate phase, and considers it as a variety of?-Fe. When iron or steel is heated above the Curie point, the thermal movement of ions upsets the orientation of the spin magnetic moments of the electrons, the ferromagnet becomes paramagnetic - a second-order phase transition occurs, but a first-order phase transition with a change in the basic physical parameters of the crystals does not occur.
For pure iron at normal pressure, from the point of view of metallurgy, there are the following stable modifications:

The presence of carbon and alloying elements in steel significantly changes the temperatures of phase transitions. A solid solution of carbon in α- and β-iron is called ferrite. Sometimes a distinction is made between high-temperature?-ferrite and low-temperature?-ferrite, although their atomic structures are the same. A solid solution of carbon in α-iron is called austenite.

The phenomenon of polymorphism is extremely important for steel metallurgy. Precisely thanks to?-? Heat treatment of steel occurs at crystal lattice transitions. Without this phenomenon, iron as the basis of steel would not have received such widespread use.
Iron is refractory and belongs to the metals of medium activity. The melting point of iron is 1539 °C, the boiling point is 2862 °C.
Chemical properties.
Iron exhibits moderate chemical activity. It burns in an oxygen atmosphere, forming the oxide Fe 2 O 3. In a finely crushed state, the metal is pyrophoric, i.e. capable of spontaneous combustion in air. Fine iron powder can be obtained by thermal decomposition of iron oxalate in a hydrogen atmosphere.
When stored in air at temperatures up to 200°C, iron is gradually covered with a dense film of oxide, which prevents further oxidation of the metal. In humid air, iron becomes covered with a loose layer of rust, which does not prevent the access of oxygen and moisture to the metal and its destruction. Rust does not have a constant chemical composition; approximately its chemical formula can be written as Fe 2 O 3.
Iron reacts with molten sulfur, forming sulfide, and actively interacts with chlorine, bromine and iodine to form trichloride, tribromide and diiodide. Iron reacts weakly with fluorine due to the formation of a dense, low-volatile trifluoride film on the surface. At temperatures above 500° C, the metal reacts reversibly with carbon:
3Fe+C<=>Fe3C
Iron carbide of this composition is called cementite. It is found in cast iron and steel.
Iron reacts with oxygen when heated. When iron burns in air, Fe 2 O 3 oxide is formed, when burned in pure oxygen, Fe 3 O 4 oxide is formed. If oxygen or air is passed through molten iron, FeO oxide is formed.
When heated, iron reacts with nitrogen, forming iron nitride Fe3N, with phosphorus, forming phosphides FeP, Fe 2 P and Fe 3 P, with carbon, forming carbide Fe 3 C, with silicon, forming several silicides, for example, FeSi. At elevated pressure, metallic iron reacts with carbon monoxide CO, and liquid, under normal conditions, highly volatile iron pentacarbonyl Fe(CO) 5 is formed. Iron carbonyls of the compositions Fe 2 (CO) 9 and Fe 3 (CO) 12 are also known. Iron carbonyls serve as starting materials in the synthesis of organoiron compounds, including the composition ferrocene.
Pure metallic iron is stable in water and dilute alkali solutions. Iron does not dissolve in concentrated sulfuric and nitric acids, since a strong oxide film passivates its surface. With hydrochloric and dilute (approximately 20%) sulfuric acids, iron reacts to form iron(II) salts:
Fe + 2HCl = FeCl 2 + H 2
Fe + H 2 SO 4 = FeSO 4 + H 2
Iron dissolves in dilute and moderately concentrated solutions of nitric acid:
Fe + 4HNO 3 = Fe(NO 3) 3 + NO ^ + 2H 2 O
When iron reacts with approximately 70% sulfuric acid, the reaction proceeds to form iron (III) sulfate:
2Fe + 4H 2 SO 4 = Fe 2 (SO 4) 3 + SO 2 + 4H 2 O
Under the influence of atmospheric moisture and air, iron corrodes (rusts):
4Fe + 2H 2 O + 3O 2 = 4FeO(OH)
Up to 10% of all iron produced is lost annually due to corrosion.
Very pure iron, containing less than 0.01% impurities of sulfur, carbon and phosphorus, is resistant to corrosion. Near the city of Delhi in India there is an iron column erected in the 9th century. BC, which shows no signs of rust. It is made of very pure metal with an iron content of 99.72%. The climatic features of this area can play an important role in the corrosion resistance of the material of the famous column.
Metallic iron reacts when heated with concentrated (more than 30%) solutions of alkalis, forming hydroxo complexes. Under the influence of strong oxidizing agents when heated, iron can form compounds in the oxidation state (+VI) - ferrates:
Fe + 2KNO 3 = K 2 FeO 4 + 2NO
For iron, oxides and hydroxides are known in oxidation states (II) and (III).
Iron forms simple salts with almost all anions. Nitrates, sulfates, halides (except fluorides), acetates, etc. are soluble in water. The iron (II) cation can be oxidized by many oxidizing agents to the iron (III) cation. Solutions of iron (II) salts and its solid salts gradually oxidize even simply when stored in air:
4FeCO 3 + 2H 2 O + O 2 = 4FeO(OH) + 2CO 2
4FeS + 6H 2 O + O 2 = 4FeO(OH) + 4H 2 S
When heated, iron sulfates, nitrates, carbonates and oxalates decompose. In this case, iron (II) is usually oxidized to iron (III), for example:
2FeSO 4 = Fe 2 O 3 + SO 3 + SO 2
Iron (III) salts undergo severe hydrolysis.

Iron (II) oxide – FeO. A black crystalline substance, the molecule has an ionic structure. Exhibits basic properties (although it interacts with alkali melts, exhibiting weak amphotericity). It does not react with water under normal conditions, but in the presence of atmospheric oxygen and low heating it slowly reacts with water vapor. Shows properties of a weak reducing agent. When heated, it decomposes, but upon further heating it forms again. Interacts with acids. Oxidized by oxygen to mixed iron oxide. Reduced by hydrogen, carbon, carbon monoxide:
FeO + 2HCl = FeCl 2 + H 2 O,
FeO + 4NaOH = Na 4 FeO 3 + 2H 2 O
4FeO + 6H 2 O+ O 2 = 4Fe(OH) 3
FeO Fe 3 O 4 +Fe FeO
6FeO + O 2 2Fe 3 O 4,
FeO + H 2 Fe + H 2 O,
FeO + C Fe + CO,
FeO + CO Fe + CO 2 .
FeO is obtained by reduction of mixed iron oxide with carbon monoxide or decomposition of divalent iron compounds in an inert atmosphere:
Fe 3 O 4 + CO 3FeO + CO 2,
Fe(OH) 2 FeO + H 2 O,
FeCO 3 FeO + CO 2 .

Iron(II) hydroxide occurs naturally as the mineral amakinite. This mineral contains impurities of magnesium and manganese (empirical formula Fe 0.7 Mg 0.2 Mn 0.1 (OH) 2). The color of the mineral is yellow-green or light green, Mohs hardness 3.5-4, density 2.925-2.98 g/cm?.
Pure iron (II) hydroxide is a white crystalline substance. Sometimes it has a greenish tint due to impurities of iron salts. Over time, it darkens in air due to oxidation. Insoluble in water (solubility 5.8·10?6 mol/l). Decomposes when heated. It has a trigonal crystal lattice system.
Iron (II) hydroxide exhibits the properties of a base - it easily enters into neutralization reactions with dilute acids, for example hydrochloric acid (a solution of iron (II) chloride is formed):
Fe(OH) 2 + 2HCl 2 = 2 H 2 O + FeCl 2
Under more severe conditions, it exhibits acidic properties, for example, with concentrated (more than 50%) sodium hydroxide when boiling in a nitrogen atmosphere, it forms a precipitate of sodium tetrahydroxoferrate (II):
Fe(OH) 2 + 2NaOH = Na 2
Does not react with ammonia hydrate. When heated, it reacts with concentrated solutions of ammonium salts, for example, ammonium chloride:
Fe(OH) 2 + 2NH 4 Cl = FeCl 2 + 2NH 3 + 2H 2 O
When heated, it decomposes to form iron (II) oxide: Fe(OH) 2 = FeO + H 2 O
In this reaction, metallic iron and diiron(III)-iron(II) oxide (Fe 3 O 4) are formed as impurities.
In the form of a suspension, when boiled in the presence of atmospheric oxygen, it is oxidized to iron metahydroxide. When heated with the latter, it forms diiron(III)-iron(II) oxide:
4Fe(OH) 2 + O 2 = 4FeO(OH) + 2H 2 O
Fe(OH) 2 + 2FeO(OH) = (FeFe 2)O + 2H 2 O
These reactions also occur (slowly) during the rusting process of iron.
Iron (II) hydroxide can be obtained in the form of a precipitate in exchange reactions of solutions of iron (II) salts with alkali, for example:
FeSO 4 + 2KOH = Fe(OH) 2 + K 2 SO 4
The formation of iron (II) hydroxide is one of the stages of iron rusting:
2Fe + 2H 2 O + O 2 = 2 Fe(OH) 2
Iron (II) hydroxide is used in the manufacture of the active mass of iron-nickel batteries.

Iron (II) hydroxide is a yellow-green or light green mineral, Mohs hardness 3.5-4, density 2.925-2.98 g/cm?. Amphoteric hydroxide with a predominance of basic properties.

In iron (II) salts, due to its partial oxidation in air, iron (III) cations are always present. Therefore, to study the properties of Fe 2+ cations, instead of iron (II) sulfate, you should take the most stable double crystalline Mohr's salt (NH 4) 2 SO 4 · FeSO 4 · 6H 2 O or use a freshly prepared solution of iron (II) sulfate. Since the stability of iron (II) in the crystalline state is higher than in solution, for research it is necessary to take a freshly prepared salt solution.

Equipment and reagents: pipette, test tubes, beaker, filter paper, scissors; Mohr's salt, sodium hydroxide, sulfuric acid.

An aqueous solution of sodium hydroxide is added to the Mohr's salt solution until a green precipitate forms. The separated precipitate is filtered and divided into three test tubes. One test tube is left to stand in air, stirring the sediment with a glass rod. After 2–3 minutes, the color of the precipitate will begin to change due to the oxidation of iron (II) hydroxide to iron (III) hydroxide. Add a few drops of a dilute solution of hydrochloric acid to the second test tube, and excess alkali to the third.

The drug is obtained by the interaction of alkali and iron salt +2 (Mohr's salt):

Studying properties:
Fe(OH) 2 + NaOH = reaction does not occur, because Fe(OH) 2 exhibits basic properties
Fe(OH) 2 + H 2 SO 4 = FeSO 4 + 2H 2 O color changes to dirty green
4Fe(OH) 2 + O 2 = 2Fe 2 O 3 + 4H 2 O the precipitate in air oxidizes (rusts) and turns into iron (III) hydroxide
In order to get 6 gr. Fe(OH) 2 let’s calculate each substance that reacted.
Calculations:
(NH 4) 2 SO 4 FeSO 4 6H 2 O + 2NaOH = Fe(OH) 2 v + Na 2 SO 4 + NH 4 O 2
M(Fe(OH) 2) = 53 g/mol
n(Fe(OH) 2) = 0.067 mol
M(NaOH) = 40 g/mol
m(NaOH) = 0.067 mol? 40 g/mol? 2=5.36g
M((NH 4) 2 SO 4 FeSO 4 6H 2 O) = 392 g/mol
m((NH 4) 2 SO 4 FeSO 4 6H 2 O) = 26g
? = (me/mtheor)?100% = (5.63/6)?100% =93.8%

Conclusion.
In the course of this course work, the physical and chemical properties of hydroxides as a class of inorganic compounds, iron and its compounds in the oxidation state +2 were studied; their history of discovery, distribution in nature, production is considered; the optimal method for obtaining iron (II) hydroxide was selected; Iron (II) hydroxide was obtained and its properties were studied.

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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

Propertiesd-elements (part 2).

Theoretical part

Iron, cobalt, and nickel make up the iron “family.”

They exhibit oxidation states mainly +2 and +3. The oxidation state +3 is more typical for iron, +2 for cobalt and nickel.

Related features of these metals are manifested in their inherent ferromagneticity, catalytic activity, ability to form colored ions, and complex formation. However, despite the similarity of properties, iron stands out clearly in the triad in its magnetic properties. The reduction activity of iron is much greater than that of cobalt and nickel (see electrode potentials). All these metals do not interact with alkalis. When dissolved in non-oxidizing acids, Fe 2+, Co 2+, Ni 2+ ions are formed

In dilute nitric acid (a strong oxidizing agent), Fe 3+, Co 3+, Ni 3+ ions are formed

A strong oxidizing environment: H 2 SO 4 (conc.), HNO 3 (conc.) passivates iron and it begins to react only when heated:

In nitric acid containing salts NaNO 2 and NaNO 3, iron is passivated to form a film of oxide of the highest oxidation state FeO 3

In the series of hydroxides: Fe(OH) 2 - Co(OH) 2 - Ni(OH) 2

Regenerative ability decreases

Iron (II) hydroxide is easily oxidized by atmospheric oxygen:

Oxidation of Co 2+ ions is more difficult and occurs slowly:

The process occurs more intensively when hydrogen peroxide is added to the solution:

Spontaneous oxidation of Ni(OH) 2 with atmospheric oxygen does not occur; H 2 O 2 is also not a strong enough oxidizing agent and the oxidation process of Ni(OH) 2 becomes possible only when using a stronger oxidizing agent, for example bromine water:

Hydroxides of Fe (III), Co (II), Co (III), Ni (II), Ni (III) are basic in nature under normal conditions. When dissolved in acids, Co(OH) 3 and Ni(OH) 3 exhibit strong oxidizing properties and are reduced to Co 2+ and Ni 2+ cations.

Iron (III) hydroxide, when boiled with a concentrated alkali solution, forms ferrites - salts of ferrous acid.

Thus, Fe(OH) 3 hydroxide exhibits amphoteric properties.

Hydroxides of Fe(II), Fe(III), Co(II), Co(III), Ni(II) are insoluble

Hydroxides of Fe (II), Fe (III), Co (II), Co (III), Ni (II) are weak bases, so their salts hydrolyze in aqueous solutions.

These processes take place without heating.

However, the hydrolysis of salts does not proceed to completion due to the accumulation of H + in the solution. With strong dilution and heating, hydrolysis proceeds further:

When boiling a FeCl 3 solution, the hydrolysis process can be carried out irreversibly:

When salt solutions are acidified with appropriate acids, the degree of their hydrolysis decreases, since the equilibrium shifts towards the starting substances.

With strong dilution, the degree of hydrolysis increases. When soluble carbonates are added, irreversible hydrolysis occurs, since the reaction products leave the sphere of interaction.

Fe(II) salts in air gradually transform into Fe(III) salts.

In complex compounds, Fe, Co, Ni are the central complexing ions with coordination numbers of 4 or 6.

PRACTICAL PART

Experiment No. 3 Preparation and properties of iron (II) hydroxide.

Add NaOH alkali solution to the iron (II) salt solution until a precipitate forms, then divide the precipitate into three parts:

a) add excess alkali to the first test tube;

b) in the second - acid;

c) filter the precipitate from the third test tube and let it stand in air.

In the presence of moisture and atmospheric oxygen, iron (II) hydroxide transforms into iron (III) hydroxide.

Experience No. 4 Characteristic reaction to Fe 2+ ion.

Pour some iron (II) salt into the test tube and add a small amount of red blood salt K 3 solution. This reaction is used to discover iron(II) ions.

Conclusion: Red blood salt K 3 is a reagent for iron (II) ion.

Experience No. 5 Preparation and hydrolysis of iron (II) carbonate.

Add sodium carbonate solution to the iron (II) salt solution. The resulting white carbon dioxide salt of iron (II) instantly undergoes hydrolysis to form iron (II) hydroxide.

Conclusion: the salt formed by a weak acid is hydrolyzed to the end, because H 2 CO 3 breaks down into H 2 O and CO 2 and  H 2 CO 3 is completely removed from the reaction sphere.

Experiment No. 6 Preparation of iron (III) hydroxide

Pour some iron (II) salt into the test tube. Add NaOH alkali solution until a precipitate appears.

Add lye. There is no reaction.

Add acid -

Experiment No. 8 Characteristic reactions to the Fe 3+ ion

A) Add a few drops of yellow blood salt to the solution of iron (III) salt. The result is a blue precipitate of Prussian blue.

B) Add a few drops of ammonium thiocyanate NH 4 CNS to the solution of iron (III) salt.

Conclusion: A blue precipitate forms.

Conclusion: A red solution is formed

Experiment No. 12 Study of the strength of cyanide and thiocyanate complexes

Add a concentrated solution of ammonium thiocyanate to the red blood salt solution.

No reaction

Conclusion: no red color of Fe 3+ is observed.

Experiment No. 13 Preparation of cobalt (II) hydroxide and study of its properties

Add a little sodium hydroxide solution to the cobalt(II) salt solution. A poorly soluble basic cobalt salt is formed. The basic salt transforms into pink cobalt (II) hydroxide. We divide the resulting precipitate into three parts: a) add acid to the first test tube; b) second - excess alkali; c) in the third test tube we observe a gradual browning of the precipitate due to the oxidation of Co(OH) 2 to Co(OH) 3 by atmospheric oxygen. The process of browning the sediment in air takes quite a long time (within 10 minutes). Let's separate part of the Co(OH) 2 precipitate and treat it with a solution of hydrogen peroxide.

Conclusion: the precipitate turns into pink Co(II) hydroxide

b)
R.N.I.

Conclusion: When hydrogen peroxide is added, the rate of transition to Co(OH) 3 increases.

Experiment No. 14 Preparation of cobalt(II) ammonia

To the solution of cobalt (II) salt, first add a little ammonia solution, and then an excess of it.

Conclusion: A light blue colloidal solution is observed.

Experience No. 18 Production of nickel ammonia

To the solution of nickel (II) salt, add a few drops of ammonia solution, and then an excess of it.

light green solution

Experience No. 19 Characteristic reaction to an ion

To the solution of nickel complex salt obtained in experiment 18, we add an alcohol solution of dimethylglyoxime, a pink-red precipitate of nickel dimethylglyoximate is formed according to the following reaction equation:

pink-red precipitate.

Conclusion: reaction is used for quantitative determination of Ni 2+ in a solution obtained by dissolving its alloys.

Control questions

Fe(III) hydroxide is produced because Fe2(CO3)3 is further hydrolyzed.

An irreversible process is taking place.

cannot exist together.

can exist together, because H 2 O 2 is not a strong enough oxidizing agent for Ni(OH) 2.

They cannot exist together.

68. Iron compounds

Iron (II) oxide FeO– a black crystalline substance, insoluble in water and alkalis. FeO matches the base Fe(OH)2.

Receipt. Iron (II) oxide can be obtained by incomplete reduction of magnetic iron ore with carbon (II) oxide:

Chemical properties. It is the main oxide. Reacting with acids, it forms salts:

Iron (II) hydroxide Fe(OH)2- white crystalline substance.

Receipt. Iron (II) hydroxide is obtained from divalent iron salts under the action of alkali solutions:

Chemical properties. Basic hydroxide. Reacts with acids:

In air, Fe(OH)2 is oxidized to Fe(OH)3:

Iron(III) oxide Fe2O3– a brown substance, found in nature in the form of red iron ore, insoluble in water.

Receipt. When firing pyrite:

Chemical properties. Exhibits weak amphoteric properties. When interacting with alkalis, it forms salts:

Iron (III) hydroxide Fe(OH)3– a red-brown substance, insoluble in water and excess alkali.

Receipt. Obtained by the oxidation of iron (III) oxide and iron (II) hydroxide.

Chemical properties. It is an amphoteric compound (with a predominance of basic properties). Precipitates under the action of alkalis on ferric iron salts:

Ferrous salts obtained by reacting metallic iron with appropriate acids. They are highly hydrolyzed, which is why their aqueous solutions are energetic reducing agents:

When heated above 480 °C, it decomposes, forming oxides:

When alkalis act on iron (II) sulfate, iron (II) hydroxide is formed:

Forms crystalline hydrate - FeSO4?7Н2О (iron sulfate). Iron (III) chloride FeCl3 – dark brown crystalline substance.

Chemical properties. Soluble in water. FeCl3 exhibits oxidizing properties.

Reducing agents - magnesium, zinc, hydrogen sulfide, oxidize without heating.

1. Oxygen oxidizes iron, resulting in the formation of iron scale - mixed oxide

Chlorine is a strong oxidizing agent, so it oxidizes iron to a higher oxidation state (+3), resulting in the formation of iron (III) chloride. 2. Oxygen and chlorine are oxidizing agents, iron is a reducing agent.

Interaction of iron with concentrated acids 1. Nitric and concentrated sulfuric acids are oxidizing acids, i.e. they exhibit strong oxidizing properties due to the acid residue. Nitrogen oxide (II) released during the reduction of nitric acid is easily oxidized by oxygen in the air to nitrogen oxide (IV).

Note: Iron does not react with concentrated nitric acid and concentrated sulfuric acid in the cold (passivates).

Preparation of iron (II) hydroxide and its interaction with acids

A) Actions: Add sodium hydroxide solution to a freshly prepared solution of iron (II) sulfate. Observations: A greenish precipitate forms. Reaction equations:

Conclusions: Iron (II) and (III) hydroxides can be obtained as a result of an exchange reaction between soluble salts of iron (II) and (III) with an alkali solution, because in this case, ion binding occurs:

b) Actions: Add a solution of hydrochloric acid to the precipitate. Observations: The precipitate dissolves. Reaction equations:

Conclusions: Because

is basic in nature, so it reacts with acids.

Preparation of salts of iron (III) hydroxide and its interaction with acids to form the corresponding salts

A) Actions: Add alkali solution to the solution of iron (III) chloride. Observations: A brown precipitate forms. Reaction equations:

Conclusions: Ions

can be determined using the reaction between their salts and alkali, because in this case, precipitation forms:

- green;

- brown. b) Actions: Add sulfuric acid to the precipitate. Observations: The precipitate dissolves. Reaction equations: