Being in nature
Copper(I) sulfide occurs naturally as the mineral chalcocite (obsolete names: chalcocite, redrutite, copper luster). Also, the composition of this compound is similar to the mineral jurleite, sometimes called jarleite (for more details, see Copper sulfides). The main properties are presented in the table.
| Chalcozine | Djurleit | |
| Compound | Cu2S | Cu1.94S |
| Color | dark grey | black |
| singonia | rhombic | monoclinic |
| Density, g/cm³ | 5,5—5,8 | 5,5—5,7 |
| Hardness | 2,5—3,0 | 2,2—2,4 |
Physical properties
Copper(I) sulfide is a dark gray solid, insoluble in water and ethanol.
Cu2S exists in three crystalline modifications:
- α-Cu2S, orthorhombic system, space group Ab
2m, cell parameters
a
= 1.1190 nm,
b
= 2.728 nm,
c
= 1.341 nm,
Z
= 96,
d
= 5.81 g/cm3; - β-Cu2S, hexagonal system, space group P
63/mmc, cell parameters
a
= 0.389 nm,
c
= 0.668 nm,
Z
= 2,
d
= 5.78 g/cm3; - γ-Cu2S, cubic system, space group F
m3m, cell parameters
a
= 0.5735 nm,
Z
= 4,
d
= 5.60 g/cm3.
Phase transition temperatures: α → β 103 °C, β → γ 437 °C.
Copper(I) sulfide tends to form non-stoichiometric compounds, the properties of which can differ greatly from stoichiometric ones.
Iron. Properties of iron and its compounds
Iron Fe: chemical properties, methods of obtaining iron, interaction with simple substances (oxygen, sulfur) and with complex substances (acids, water, strong oxidizing agents). Iron (II) oxide FeO, iron (III) oxide Fe2O3, iron scale (Fe3O4) - production methods and chemical properties. Iron (II) hydroxide Fe(OH)2, iron (III) hydroxide Fe(OH)3 - methods of preparation and chemical properties.
Position of iron in the periodic table of chemical elements Electronic structure of iron Physical properties Occurrence in nature Methods of preparation Qualitative reactions Chemical properties 1. Interaction with simple substances 1.1. Interaction with halogens 1.2. Interaction with sulfur 1.3. Interaction with phosphorus 1.4. Interaction with nitrogen 1.5. Interaction with carbon 1.6. Combustion 2. Interaction with complex substances 2.1. Interaction with water 2.2. Interaction with mineral acids 2.3. Interaction with sulfuric acid 2.4. Interaction with nitric acid 2.5. Interaction with strong oxidizing agents 2.6. Interaction with oxides and salts
Iron (II) oxide Preparation methods Chemical properties 1. Interaction with acid oxides 2. Interaction with acids 3. Interaction with water 4. Interaction with oxidizing agents 5. Interaction with acids 6. Interaction with reducing agents
Iron (III) oxide Preparation methods Chemical properties 1. Interaction with acidic oxides and acids 2. Interaction with alkalis and basic oxides 3. Interaction with water 4. Interaction with oxidizing agents 5. Oxidative properties of iron (III) oxide
6. Interaction with salts of more volatile acids
Iron (II, III) oxide Preparation methods Chemical properties 1. Interaction with acidic oxides and acids 2. Interaction with strong oxidizing acids 3. Interaction with water 4. Interaction with oxidizing agents 5. Oxidative properties of iron (II, III) oxide
Iron (II) hydroxide Methods of preparation Chemical properties 1. Interaction with acids 2. Interaction with acid oxides 3. Reducing properties 4. Decomposition when heated
Iron (III) hydroxide Methods of preparation Chemical properties 1. Interaction with acids 2. Interaction with acid oxides 3. Interaction with alkalis 4. Decomposition when heated
Iron salts
Iron
Position in the periodic table of chemical elements
The element iron is located in the secondary subgroup of group VIII (or in group 8 in the modern form of PSHE) and in the fourth period of the periodic table of chemical elements D.I. Mendeleev.
Electronic structure of the iron atom
Electronic configuration of iron in its ground state :
+26Fe 1s 2 2s22p63s23p64s23d6
Iron
exhibits pronounced magnetic properties.
Physical properties
Iron
– a metal of silver-white color, with high chemical activity and high malleability. Has high thermal and electrical conductivity.
(image from the portal vchemraznica.ru)
Melting point 1538oC, boiling point 2861oC.
Being in nature
Iron
quite common in the earth's crust (about 4% of the mass of the earth's crust). In terms of abundance on Earth, iron ranks 4th among all elements and 2nd among metals. The content in the earth's crust is about 8%.
In nature, iron is mainly found in the form of compounds:
Red iron ore Fe2O3 (hematite).
(image from the portal karatto.ru)
Magnetic iron ore Fe3O4 or FeO·Fe2O3 (magnetite).
(image from emchi-med.ru portal)
Iron sulfides are also widespread in nature, for example, pyrite FeS2 .
(image from the portal livemaster.ru)
There are also other minerals containing iron.
Methods of obtaining
Iron
industrially obtained from iron ore, hematite
Fe2O3
or magnetite (
Fe3O4 or FeO·Fe2O3 ).
1. One of the main methods of iron production is the blast furnace process . The blast furnace process is based on the reduction of iron from oxide with carbon in a blast furnace .
Ore, coke and fluxes are loaded into the furnace.
Charge –
a mixture of starting materials, and in some cases fuel, in a certain proportion, which is processed in a furnace.
Coal coke is a solid, gray, porous product obtained by coking coal at temperatures of 950–1100 °C without air access. Contains 96-98% carbon.
Fluxes are inorganic substances that are added to ore during metal smelting to lower the melting point and more easily separate the metal from the waste rock.
Slag –
melt (and after solidification
- a glassy mass) covering the surface of the liquid metal.
Slag consists of floating waste rock products with fluxes and protects the metal from the harmful effects of the gaseous environment of the furnace and removes impurities. In the furnace, coke is oxidized to carbon monoxide (II):
2C + O2 → 2CO
The heated carbon monoxide then reduces iron(III) oxide:
3CO + Fe2O3 → 3CO2 + 2Fe
The process of obtaining iron is multi-stage and depends on temperature.
At the top, where the temperature is usually between 200 °C and 700 °C, the following reaction takes place:
3Fe2O3 + CO → 2Fe3O4 + CO2
Below in the furnace, at temperatures of approximately 850 °C, the reduction of mixed iron (II, III) oxide to iron (II) oxide occurs:
Fe3O4 + CO → 3FeO + CO2
Counter flows of gases heat the charge, and limestone decomposes:
CaCO3 → CaO + CO2
Iron (II) oxide falls into a region of higher temperatures (up to 1200oC), where the following reaction occurs:
FeO + CO → Fe + CO2
Carbon dioxide rises and reacts with coke to form carbon monoxide:
CO2 + C → 2CO
(image from the portal 900igr.net)
2. Iron is also obtained by direct reduction from the oxide with hydrogen:
Fe2O3 + 3H2 → 2Fe + 3H2O
This produces purer iron, because the resulting iron is not contaminated with sulfur and phosphorus, which are impurities in coal.
3. Another way to obtain iron in industry is electrolysis of solutions of iron salts.
Qualitative reactions
Qualitative reactions to iron ions +2.
– interaction of iron (II) salts with alkalis . This results in the formation of gray-green gelatinous precipitate of iron (II) hydroxide .
For example , iron(II) chloride reacts with sodium hydroxide:
2NaOH + FeCl2 → Fe(OH)2 + 2NaCl
A video experience of the interaction of a solution of iron (II) sulfate with a solution of sodium hydroxide (qualitative reaction to iron (II) ions) can be viewed here.
Iron (II) hydroxide turns brown in air as it oxidizes to iron (III) hydroxide:
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3
– iron +2 color the solution a light yellow-green color .
– interaction with red blood salt K3[Fe(CN)6] – also a qualitative reaction to iron ions +2. This produces a blue precipitate “Turnboole blue.”
A video experience of the interaction of a solution of iron (II) chloride with a solution of potassium hexacyanoferrate (III) (qualitative reaction to iron (II) ions) can be viewed here.
Qualitative reactions to iron ions +3
– interaction of iron (III) salts with alkalis . In this case, a brown precipitate of iron (III) hydroxide .
For example , iron(III) chloride reacts with sodium hydroxide:
3NaOH + FeCl3 → Fe(OH)3 + 3NaCl
A video experience of the interaction of a solution of iron (III) chloride with a solution of sodium hydroxide (qualitative reaction to iron (III) ions) can be viewed here.
– iron +3 color the solution light yellow-orange .
– interaction with yellow blood salt K4[Fe(CN)6] iron ions +3. This produces a blue precipitate called Prussian blue.
A video experience of the interaction of a solution of iron (III) chloride with a solution of potassium hexacyanoferrate (II) (qualitative reaction to iron (III) ions) can be viewed here.
Recently, data have been obtained that indicate that the molecules of Prussian blue are identical in structure to the molecules of Turnboule blue. The composition of the molecules of both of these substances can be expressed by the formula Fe4[Fe2(CN)6]3.
– when iron (III) salts interact with thiocyanates , the solution turns blood red.
For example , iron(III) chloride reacts with sodium thiocyanate :
FeCl3 + 3NaCN S → Fe(CN S )3 + 3NaCl
A video experience of the interaction of a solution of iron (III) chloride with a solution of potassium thiocyanate (qualitative reaction to iron (III) ions) can be viewed here.
Chemical properties
1. Under normal conditions, iron is inactive , but when heated, especially in a finely crushed state, it becomes active and reacts with almost all non-metals .
1.1. Iron reacts with halogens to form halides. In this case, active non-metals (fluorine, chlorine and bromine) oxidize iron to the oxidation state +3:
2Fe + 3Cl2 → 2FeCl3
Less active iodine oxidizes iron to oxidation state +2:
Fe + I2 → FeI2
1.2. Iron reacts with sulfur to form iron(II) sulfide :
Fe + S → FeS
1.3. Iron reacts with phosphorus . In this case, a binary compound is formed - iron phosphide :
Fe + P → FeP
1.4. With nitrogen iron
reacts under specific conditions.
1.5. Iron reacts with carbon and silicon to form carbide and silicide.
1.6. When interacting with oxygen , iron forms scale - double iron oxide (II, III):
3Fe + 2O2 → Fe3O4
When oxygen is passed through molten iron, iron (II) oxide can be formed:
2Fe + O2 → 2FeO
2. Iron interacts with complex substances.
2.1. Under normal conditions, iron practically does not react with water Hot iron can react at a temperature of 700-900°C with water vapor:
3Fe0 + 4H2+O → Fe+33O4 + 4H20
In water in the presence of oxygen or in humid air, iron slowly oxidizes (corrodes):
4Fe + 3O2 + 6H2O → 4Fe(OH)3
2.2. Iron reacts with mineral acids (hydrochloric, phosphoric and dilute sulfuric acid). In this case, an iron salt with an oxidation state of +2 and hydrogen are formed.
For example , iron reacts violently with hydrochloric acid :
Fe + 2HCl → FeCl2 + H2↑
2.3. Under normal conditions, iron does not react with concentrated sulfuric acid due to passivation - the formation of a dense oxide film. When heated, the reaction proceeds, sulfur (IV) oxide , iron (III) sulfate and water :
2Fe + 6H2SO4(conc.) → Fe2(SO4)3 + 3SO2 + 6H2O
2.4. Iron does not react under normal conditions with concentrated nitric acid also due to passivation. When heated, the reaction occurs with the formation of iron (III) nitrate, nitrogen (IV) oxide and water:
Fe + 6HNO3(conc.) → Fe(NO3)3 + 3NO2↑ + 3H2O
with dilute nitric acid to form nitric oxide (II):
Fe + 4HNO3(dil.hor.) → Fe(NO3)3 + NO + 2H2O
When iron reacts with very dilute nitric acid, ammonium nitrate is formed :
8Fe + 30HNO3(ultra dil.) → 8Fe(NO3)3 + 3NH4NO3 + 9H2O
2.5. Iron can react with alkaline solutions or melts of strong oxidizing agents . In this case, iron oxidizes to the oxidation state +6, forming a salt (ferrate) .
For example , when iron reacts with molten potassium nitrate in the presence of potassium hydroxide, iron is oxidized to potassium ferrate, and nitrogen is reduced to either potassium nitrite or ammonia:
Fe + 2KOH + 3KNO3 → 3KNO2 + K2FeO4 + H2O
2.6. Iron reduces less active metals from oxides and salts .
For example , iron replaces copper
from
copper(II) sulfate. Exothermic reaction:
Fe + CuSO4 → FeSO4 + Cu
Another example : the simple substance iron reduces iron to the oxidation state +2 when interacting with iron compounds +3:
2Fe(NO3)3 + Fe → 3Fe(NO3)2
2FeCl3 + Fe → 3FeCl2
Fe2(SO4)3 + Fe → 3FeSO4
Iron(II) oxide
Iron (II) oxide is a solid, black, water-insoluble substance.
Methods of obtaining
Iron (II) oxide can be obtained by various methods:
1. Partial reduction of iron (III) oxide .
For example, by partial reduction of iron (III) oxide with hydrogen:
Fe2O3 + H2 → 2FeO + H2O
Or by partial reduction of iron (III) oxide with carbon monoxide:
Fe2O3 + CO → 2FeO + CO2
Another example : reduction of iron (III) oxide with iron:
Fe2O3 + Fe → 3FeO
2. Decomposition of iron (II) hydroxide when heated:
Fe(OH)2 → FeO + H2O
Chemical properties
Iron(II) oxide is a typical basic oxide .
1. When iron (II) oxide reacts with acidic oxides, salts are formed.
For example , iron (II) oxide reacts with sulfur (VI) oxide :
FeO + SO3 → FeSO4
2. Iron (II) oxide reacts with soluble acids. In this case, the corresponding salts are also formed.
For example , iron(II) oxide reacts with hydrochloric acid
:
FeO + 2HCl → FeCl2 + H2O
3. Iron (II) oxide does not react with water.
4. Iron (II) oxide is unstable and easily oxidizes to iron (III) compounds.
For example , when reacting with concentrated nitric acid, iron (III) nitrate, nitric oxide (IV) and water are formed :
FeO + 4HNO3(conc.) → NO2 + Fe(NO3)3 + 2H2O
When reacting with dilute nitric acid, nitric oxide (II) is formed. The reaction occurs when heated:
3FeO + 10HNO3(dil.) → 3Fe(NO3)3 + NO + 5H2O
5. Iron (II) oxide exhibits weak oxidizing properties .
For example , iron(II) oxide reacts with carbon monoxide when heated:
FeO + CO → Fe + CO2
Iron(III) oxide
Iron (III) oxide is a solid, water-insoluble, red-brown substance.
Methods of obtaining
Iron(III) oxide can be obtained by various methods:
1. Oxidation of iron (II) oxide with oxygen .
4FeO + O2 → 2Fe2O3
2. Decomposition of iron (III) hydroxide when heated:
2Fe(OH)3 → Fe2O3 + 3H2O
Chemical properties
Iron (III) oxide is amphoteric .
1. When iron (III) oxide reacts with acidic oxides and acids, salts are formed.
For example , iron(III) oxide reacts with nitric acid :
Fe2O3 + 6HNO3 → 2Fe(NO3)3 + 3H2O
2. Iron (III) oxide reacts with alkalis and basic oxides. The reaction takes place in the melt, and the corresponding salt (ferrite) is formed.
For example , iron(III) oxide reacts with sodium hydroxide
:
Fe2O3 + 2NaOH → 2NaFeO2 + H2O
3. Iron (III) oxide does not react with water.
4. Iron (III) oxide is oxidized by strong oxidizing agents to form iron (VI) compounds.
For example , potassium chlorate in an alkaline medium oxidizes iron (III) oxide to ferrate :
Fe2O3 + KClO3 + 4KOH → 2K2FeO4 + KCl + 2H2O
Nitrates and nitrites in an alkaline environment also oxidize iron (III) oxide:
Fe2O3 + 3KNO3 + 4KOH → 2K2FeO4 + 3KNO2 + 2H2O
5. Iron (III) oxide exhibits oxidizing properties .
For example , iron(III) oxide reacts with carbon monoxide when heated. In this case, it is possible to reduce both to pure iron and to iron (II) oxide or iron scale:
Fe2O3 + 3СO → 2Fe + 3CO2
Also iron(III) oxide is reduced by hydrogen :
Fe2O3 + 3H2 → 2Fe + 3H2O
Iron can only reduce iron oxide to iron (II) oxide:
Fe2O3 + Fe → 3FeO
Iron(III) oxide reacts with more reactive metals .
For example , with aluminum ( aluminothermy ):
Fe2O3 + 2Al → 2Fe + Al2O3
Iron(III) oxide also reacts with some other strong reducing agents .
For example , with sodium hydride:
Fe2O3 + 3NaH → 3NaOH + 2Fe
6. Iron (III) oxide is solid, non-volatile and amphoteric. Consequently, it displaces more volatile oxides (usually carbon dioxide) from the salts during fusion.
For example , from sodium carbonate :
Fe2O3 + Na2CO3 → 2NaFeO2 + CO2
Iron (II, III) oxide
Iron (II, III) oxide (iron scale, magnetite) is a solid, black, water-insoluble substance.
Photo from wikipedia.ru
Methods of obtaining
Iron (II, III) oxide can be obtained by various methods:
1. combustion in air:
3Fe + 2O2 → Fe3O4
2. Partial reduction of iron (III) oxide with hydrogen or carbon monoxide:
3Fe2O3 + H2 → 2Fe3O4 + H2O
3. At high temperatures, hot iron reacts with water , forming iron (II, III) double oxide:
3Fe + 4H2O(steam) → Fe3O4 + 4H2
Chemical properties
The properties of iron (II, III) oxide are determined by the properties of the two oxides of which it consists: basic iron (II) oxide and amphoteric iron (III) oxide.
1. When iron (II, III) oxide reacts with acidic oxides and acids, iron (II) and iron (III) salts are formed.
For example , iron (II, III) oxide reacts with hydrochloric acid. In this case, two salts are formed - iron (II) chloride and iron (III) chloride:
Fe3O4 + 8HCl → FeCl2 + 2FeCl3 + 4H2O
Another example : iron (II, III) oxide reacts with dilute sulfuric acid.
Fe3O4 + 4H2SO4(dil.) → Fe2(SO4)3 + FeSO4 + 4H2O
2. Iron oxide (II, III) interacts with strong oxidizing acids (concentrated sulfuric and nitric).
For example , iron scale is oxidized with concentrated nitric acid :
Fe3O4 + 10HNO3(conc.) → NO2↑ + 3Fe(NO3)3 + 5H2O
With dilute nitric acid, scale is oxidized when heated:
3Fe3O4 + 28HNO3(dil.) → 9Fe(NO3)3 + NO + 14H2O
Also, iron (II, III) oxide is oxidized with concentrated sulfuric acid :
2Fe3O4 + 10H2SO4(conc.) → 3Fe2(SO4)3 + SO2 + 10H2O
Scale is also oxidized oxygen :
4Fe3O4 + O2(air) → 6Fe2O3
3. Iron (II, III) oxide does not interact with water.
4. Iron (II, III) oxide is oxidized by strong oxidizing agents to iron (VI) compounds, like other iron oxides (see above).
5. Iron scale exhibits oxidizing properties .
For example , iron(II,III) oxide reacts with carbon monoxide when heated. In this case, reduction is possible both to pure iron and to iron (II) oxide:
Fe3O4 + 4CO → 3Fe + 4CO2
Also, iron scale is reduced with hydrogen :
Fe3O4 + 4H2 → 3Fe + 4H2O
Iron(II,III) oxide reacts with more reactive metals .
For example , with aluminum ( aluminothermy ):
3Fe3O4 + 8Al → 9Fe + 4Al2O3
Iron (II, III) oxide also reacts with some other strong reducing agents (iodides and sulfides).
For example , with hydrogen iodide:
Fe3O4 + 8HI → 3FeI2 + I2 + 4H2O
Iron(II) hydroxide
Methods of obtaining
1. Iron (II) hydroxide can be obtained by reacting an ammonia with iron (II) salts .
For example , iron(II) chloride reacts with an aqueous solution of ammonia to form iron(II) hydroxide and ammonium chloride :
FeCl2 + 2NH3 + 2H2O → Fe(OH)2 + 2NH4Cl
2. Iron (II) hydroxide can be obtained by the action of alkali on iron (II) salts .
For example , iron(II) chloride reacts with potassium hydroxide to form iron(II) hydroxide and potassium chloride :
FeCl2 + 2KOH → Fe(OH)2↓ + 2KCl
Chemical properties
1. Iron (II) hydroxide exhibits basic properties , namely, it reacts with acids . In this case, the corresponding salts .
For example , iron(II) hydroxide reacts with hydrochloric acid to form iron(II) chloride
:
Fe(OH)2 + 2HCl → FeCl2 + 2H2O
Fe(OH)2 + H2SO4 → FeSO4 + 2H2O
Fe(OH)2 + 2HBr → FeBr2 + 2H2O
2. Iron (II) hydroxide reacts with acidic oxides of strong acids .
For example , iron(II) hydroxide reacts with sulfur(VI) oxide to form iron(II) sulfate
:
Fe(OH)2 + SO3 → FeSO4 + 2H2O
3. Iron (II) hydroxide exhibits strong reducing properties and reacts with oxidizing agents. In this case, iron (III) compounds
.
For example , iron(II) hydroxide reacts with oxygen in the presence of water:
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3↓
Iron (II) hydroxide reacts with hydrogen peroxide
:
2Fe(OH)2 + H2O2 → 2Fe(OH)3
When Fe(OH)2 is dissolved in nitric or concentrated sulfuric acids, iron (III) salts are formed:
2Fe(OH)2 + 4H2SO4(conc.) → Fe2(SO4)3 + SO2 + 6H2O
4. Iron (II) hydroxide decomposes when heated:
Fe(OH)2 → FeO + H2O
Iron(III) hydroxide
Methods of obtaining
1. Iron (III) hydroxide can be obtained by reacting an ammonia with iron (III) salts .
For example , iron(III) chloride reacts with an aqueous solution of ammonia to form iron(III) hydroxide and ammonium chloride :
FeCl3 + 3NH3 + 3H2O = Fe(OH)3 + 3NH4Cl
2. Oxidation of iron (II) hydroxide with oxygen or hydrogen peroxide
:
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3↓
2Fe(OH)2 + H2O2 → 2Fe(OH)3
3. Iron (III) hydroxide can be obtained by the action of alkali on a solution of iron (III) salt .
For example , iron(III) chloride reacts with a solution of potassium hydroxide to form iron(III) hydroxide and potassium chloride :
FeCl3 + 3KOH → Fe(OH)3↓ + 3KCl
A video of the production of iron (III) hydroxide by the reaction of iron (III) chloride and potassium hydroxide can be viewed here.
4. Iron (III) hydroxide is also formed by the interaction of soluble iron (III) salts with solutions of carbonates and sulfites . Iron(III) carbonates and sulfites are irreversibly hydrolyzed in aqueous solution.
For example: iron(III) bromide reacts with sodium carbonate . In this case, a precipitate of iron (III) hydroxide precipitates, carbon dioxide is released and sodium bromide is formed:
2FeBr3 + 3Na2CO3 + 3H2O = 2Fe(OH)3↓ + CO2↑ + 6NaBr
But there is an exception! The interaction of iron (III) salts with sulfites in the Unified State Examination in Chemistry is an oxidation-reduction reaction. Iron (III) compounds oxidize sulfites, as well as sulfides and iodides.
The interaction of iron (III) chloride with sulfite, for example, potassium, is a very interesting reaction. Firstly, some sources indicate that irreversible hydrolysis may still occur in it. But for the Unified State Exam, it is better to assume that OVR occurs in this case. Secondly, OVR can be written in different forms:
2FeCl3 + Na2SO3 + H2O = 2FeCl2 + Na2SO4 + 2HCl
The following entry is also acceptable:
2FeCl3 + Na2SO3 + H2O = FeSO4 + 2NaCl + FeCl2 + 2HCl
Chemical properties
1. Iron (III) hydroxide exhibits weak amphoteric properties , with a predominance of basic ones. As a base, iron(III) hydroxide reacts with soluble acids .
For example , iron(III) hydroxide reacts with nitric acid to form iron(III) nitrate
:
Fe(OH)3 + 3HNO3 → Fe(NO3)3 + 3H2O
Fe(OH)3 + 3HCl → FeCl3 + 3H2O
2Fe(OH)3 + 3H2SO4 → Fe2(SO4)3 + 6H2O
Fe(OH)3 + 3HBr → FeBr3 + 3H2O
2. Iron (III) hydroxide reacts with acidic oxides of strong acids .
For example , iron(III) hydroxide reacts with sulfur(VI) oxide to form iron(III) sulfate
:
2Fe(OH)3 + 3SO3 → Fe2(SO4)3 + 3H2O
3. Iron (III) hydroxide reacts with soluble bases (alkalis). In this case, salts — are formed in the melt , but the reaction practically does not occur in the solution. In this case, iron (III) hydroxide exhibits acidic properties .
For example , iron(III) hydroxide reacts with potassium hydroxide
in the melt with the formation of
potassium ferrite and water :
KOH + Fe(OH)3 → KFeO2 + 2H2O
4. Iron (III) hydroxide decomposes when heated:
2Fe(OH)3 → Fe2O3 + 3H2O
A video experiment of the interaction of iron (III) hydroxide with hydrochloric acid can be viewed here.
Iron salts
Iron nitrates
iron(II) nitrate decomposes into iron(III) oxide , nitrogen(IV) oxide and oxygen :
4Fe(NO3)2 → 2Fe2O3 + 8NO2 + O2
When heated, iron (III) nitrate iron (III) oxide , nitrogen (IV) oxide and oxygen :
4Fe(NO3)3 → 2Fe2O3 + 12NO2 + 3O2
Hydrolysis of iron salts
Soluble iron salts formed by acidic residues of strong acids are hydrolyzed at the cation . Hydrolysis proceeds stepwise and reversibly , i.e. partially:
Stage I: Fe3+ + H2O ↔ FeOH2+ + H+
Stage II: FeOH2+ + H2O ↔ Fe(OH)2+ + H+
Stage III: Fe(OH)2+ + H2O ↔ Fe(OH)3 + H+
However, iron(III) sulfites and carbonates
and their
acid salts are hydrolyzed irreversibly , completely , i.e. do not exist in an aqueous solution, but decompose with water :
Fe2(SO4)3 + 6NaHSO3 → 2Fe(OH)3 + 6SO2 + 3Na2SO4
2FeBr3 + 3Na2CO3 + 3H2O → 2Fe(OH)3↓ + CO2↑ + 6NaBr
2Fe(NO3)3 + 3Na2CO3 + 3H2O → 2Fe(OH)3↓ + 6NaNO3 + 3CO2↑
2FeCl3 + 3Na2CO3 + 3H2O → 2Fe(OH)3↓ + 6NaCl + 3CO2↑
Fe2(SO4)3 + 3K2CO3 + 3H2O → 2Fe(OH)3↓ + 3CO2↑ + 3K2SO4
When iron (III) compounds interact with sulfides, ORR occurs:
2FeCl3 + 3Na2S → 2FeS + S + 6NaCl
You can read more about hydrolysis in the corresponding article.
Oxidative properties of iron (III)
Iron (III) salts exhibit fairly strong oxidizing properties. Thus, when iron (III) compounds interact with sulfides, an oxidation-reduction reaction occurs.
For example : iron (III) chloride reacts with sodium sulfide . This produces sulfur, sodium chloride and either a black precipitate of iron (II) sulfide (in excess sodium sulfide) or iron (II) chloride (in excess iron (III) chloride):
2FeCl3 + 3Na2S → 2FeS + S + 6NaCl
2FeCl3 + Na2S → 2FeCl2 + S + 2NaCl
By the same principle, iron (III) salts react with hydrogen sulfide:
2FeCl3 + H2S → 2FeCl2 + S + 2HCl
Iron(III) salts also undergo redox reactions with iodides .
For example , iron(III) chloride reacts with potassium iodide . This produces iron (II) chloride, molecular iodine and potassium chloride:
2FeCl3 + 2KI → 2FeCl2 + I2 + 2KCl
The reactions of iron(III) salts with metals are also of interest. We know that more active metals displace less active metals from salts . In other words, metals that are to the left in the electrochemical series can interact with metal salts that are located to the right in this series . Based on this rule, iron salts can only interact with metals that are located before iron. And they interact.
However, iron salts with an oxidation state of +3 are a slight exception in this series. After all, iron is characterized by two oxidation states: +2 and +3. And iron with an oxidation state of +3 is a stronger oxidizing agent. Thus, relatively speaking, iron with an oxidation state of +3 is located in the activity series after copper . And iron (III) salts can also react with metals that are located to the right of iron! But up to copper, inclusive. This is such a paradox.
And one moment. Iron (III) compounds will react with these metals, but iron (II) will not react with them. Thus, metals located in the activity series between iron and copper (including copper) when interacting with iron (III) salts reduce iron to the oxidation state +2. But metals located before iron in the activity series can reduce iron to a simple substance.
For example , iron(III) chloride reacts with copper . In this case, iron (II) chloride and copper (II) chloride :
2FeCl3 + Cu → 2FeCl2 + CuCl2
But the reaction of iron (III) nitrate with zinc proceeds according to the usual mechanism. And iron is reduced to a simple substance:
2Fe(NO3)3 + 3Zn → 2Fe + 3Zn(NO3)2
