Aluminum is a transition metal. Regional company

Video lesson 1: Inorganic chemistry. Metals: alkali, alkaline earth, aluminum

Video lesson 2: transition metals

Lecture: Characteristic chemical properties and production of simple substances - metals: alkali, alkaline earth, aluminum; transition elements (copper, zinc, chromium, iron)

Chemical properties of metals

All metals in chemical reactions manifest themselves as reducing agents. They easily part with valence electrons, being oxidized at the same time. Recall that the further to the left a metal is located in the electrochemical series of tension, the stronger the reducing agent it is. Therefore, the strongest is lithium, the weakest is gold and vice versa, gold is the strongest oxidizing agent, and lithium is the weakest.

Li→Rb→K→Ba→Sr→Ca→Na→Mg→Al→Mn→Cr→Zn→Fe→Cd→Co→Ni→Sn→Pb→H→Sb→Bi→Cu→Hg→Ag→Pd→ Pt→Au

All metals displace other metals from the salt solution, i.e. restore them. All except alkaline and alkaline earth as they interact with water. Metals located before H displace it from solutions of dilute acids, and they themselves dissolve in them.

Consider some general chemical properties of metals:

The interaction of metals with oxygen forms basic (CaO, Na 2 O, 2Li 2 O, etc.) or amphoteric (ZnO, Cr 2 O 3, Fe 2 O 3, etc.) oxides.
The interaction of metals with halogens (the main subgroup of group VII) forms hydrohalic acids (HF - hydrogen fluoride, HCl - hydrogen chloride, etc.).
The interaction of metals with non-metals forms salts (chlorides, sulfides, nitrides, etc.).
The interaction of metals with metals forms intermetallic compounds (MgB 2 , NaSn, Fe 3 Ni, etc.).
The interaction of active metals with hydrogen forms hydrides (NaH, CaH 2, KH, etc.).
The interaction of alkali and alkaline earth metals with water forms alkalis (NaOH, Ca (OH) 2, Cu (OH) 2, etc.).
The interaction of metals (only those standing in the electrochemical series up to H) with acids forms salts (sulfates, nitrites, phosphates, etc.). It should be borne in mind that metals react with acids quite reluctantly, while they almost always interact with bases and salts. In order for the reaction of the metal with the acid to take place, the metal must be active and the acid strong.

Chemical properties of alkali metals

The group of alkali metals includes the following chemical elements: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr). As they move from top to bottom in group I of the Periodic Table, their atomic radii increase, which means that their metallic and reducing properties increase.

Consider the chemical properties of alkali metals:

They do not have signs of amphotericity, as they have negative values of electrode potentials.
The strongest reducing agents among all metals.
In compounds, they exhibit only the +1 oxidation state.
Giving a single valence electron, the atoms of these chemical elements are converted into cations.
They form numerous ionic compounds.
Almost all are soluble in water.

Interaction of alkali metals with other elements:

1. With oxygen, forming individual compounds, so the oxide forms only lithium (Li 2 O), sodium forms peroxide (Na 2 O 2), and potassium, rubidium and cesium form superoxides (KO 2, RbO 2 , CsO 2).

2. With water, forming alkalis and hydrogen. Remember, these reactions are explosive. Without an explosion, only lithium reacts with water:

2Li + 2H 2 O → 2LiO H + H 2.

3. With halogens, forming halides (NaCl - sodium chloride, NaBr - sodium bromide, NaI - sodium iodide, etc.).

4. With hydrogen when heated, forming hydrides (LiH, NaH, etc.)

5. With sulfur when heated, forming sulfides (Na 2 S, K 2 S, etc.). They are colorless and highly soluble in water.

6. With phosphorus when heated, forming phosphides (Na 3 P, Li 3 P, etc.), they are very sensitive to moisture and air.

7. With carbon, when heated, carbides form only lithium and sodium (Li 2 CO 3, Na 2 CO 3), while potassium, rubidium and cesium do not form carbides, they form binary compounds with graphite (C 8 Rb, C 8 Cs, etc.) .

8. Under normal conditions, only lithium reacts with nitrogen, forming Li 3 N nitride, with other alkali metals, the reaction is possible only when heated.

9. They react explosively with acids, so carrying out such reactions is very dangerous. These reactions are ambiguous, because the alkali metal actively reacts with water, forming an alkali, which is then neutralized by an acid. This creates competition between alkali and acid.

10. With ammonia, forming amides - analogues of hydroxides, but stronger bases (NaNH 2 - sodium amide, KNH 2 - potassium amide, etc.).

11. With alcohols, forming alcoholates.

Francium is a radioactive alkali metal, one of the rarest and least stable of all radioactive elements. Its chemical properties are not well understood.

Getting alkali metals:

To obtain alkali metals, they mainly use the electrolysis of melts of their halides, most often chlorides, which form natural minerals:

NaCl → 2Na + Cl 2 .

There are other ways to obtain alkali metals:
Sodium can also be obtained by calcining soda with coal in closed crucibles:

Na 2 CO 3 + 2C → 2Na + 3CO.

A known method for producing lithium from its oxide in a vacuum at 300°C:

2Li 2 O + Si + 2CaO → 4Li + Ca 2 SiO 4 .

Potassium is obtained by passing sodium vapor through a potassium chloride melt at 800 ° C, emitting potassium vapor condenses:

KCl + Na → K + NaCl.

Chemical properties of alkaline earth metals

Alkaline earth metals include elements of the main subgroup of group II: calcium (Ca), strontium (Sr), barium (Ba), radium (Ra). The chemical activity of these elements grows in the same way as that of alkali metals, i.e. increasing down the subgroup.

Chemical properties of alkaline earth metals:

The structure of the valence shells of atoms of these elements ns 2 .

Giving two valence electrons, the atoms of these chemical elements are converted into cations.
The compounds exhibit an oxidation state of +2.
The charges of atomic nuclei are greater by one than those of alkaline elements of the same periods, which leads to a decrease in the radius of atoms and an increase in ionization potentials.

Interaction of alkaline earth metals with other elements:

1. With oxygen, all alkaline earth metals, except for barium, form oxides, barium forms peroxide BaO 2. Of these metals, beryllium and magnesium, coated with a thin protective oxide film, interact with oxygen only at very high t. Basic oxides of alkaline earth metals react with water, with the exception of beryllium oxide BeO, which has amphoteric properties. The reaction of calcium oxide and water is called the lime slaking reaction. If the reagent is CaO, quicklime is formed, if Ca(OH) 2, slaked. Also, basic oxides react with acidic oxides and acids. For example:

3CaO + P 2 O 5 → Ca 3 (PO 4) 2 .

2. With water, alkaline earth metals and their oxides form hydroxides - white crystalline substances, which, in comparison with alkali metal hydroxides, are less soluble in water. Hydroxides of alkaline earth metals are alkalis, except for the amphoteric Be(OH ) 2 and weak base Mg(OH)2. Since beryllium does not react with water, Be (OH ) 2 can be obtained in other ways, for example, by hydrolysis of nitride:

Be 3 N 2+ 6H 2 O → 3 Be (OH)2+ 2N N 3.

3. Under normal conditions, everything reacts with halogens, except for beryllium. The latter reacts only at high t. Halides are formed (MgI 2 - magnesium iodide, CaI 2 - calcium iodide, CaBr 2 - calcium bromide, etc.).

4. All alkaline earth metals, except beryllium, react with hydrogen when heated. Hydrides are formed (BaH 2 , CaH 2 , etc.). For the reaction of magnesium with hydrogen, in addition to high t, an increased hydrogen pressure is also required.

5. Sulfur forms sulfides. For example:

Ca + S → CaS.

Sulfides are used to obtain sulfuric acid and the corresponding metals.

6. They form nitrides with nitrogen. For example:

3Be + N 2 → Be 3 N 2.

7. With acids, forming salts of the corresponding acid and hydrogen. For example:

Be + H 2 SO 4 (razb.) → BeSO 4 + H 2.

These reactions proceed in the same way as in the case of alkali metals.

Obtaining alkaline earth metals:

Beryllium is obtained by reduction of fluoride:

BeF 2 + Mg –t o → Be + MgF 2

Barium is obtained by oxide reduction:

3BaO + 2Al –t o → 3Ba + Al 2 O 3

The remaining metals are obtained by electrolysis of chloride melts:

CaCl 2 → Ca + Cl 2

Chemical properties of aluminum

Aluminum is an active, light metal, number 13 in the table. In nature, the most common of all metals. And of the chemical elements, it occupies the third position in terms of distribution. High heat and electrical conductor. Resistant to corrosion, as it is covered with an oxide film. The melting point is 660 0 С.

Consider the chemical properties and interaction of aluminum with other elements:

1. In all compounds, aluminum is in the +3 oxidation state.

2. It exhibits reducing properties in almost all reactions.

3. Amphoteric metal exhibits both acidic and basic properties.

4. Restores many metals from oxides. This method of obtaining metals is called aluminothermy. Example of getting chromium:

2Al + Cr 2 O 3 → Al 2 O 3 + 2Cr.

5. Reacts with all dilute acids to form salts and release hydrogen. For example:

2Al + 6HCl → 2AlCl 3 + 3H 2;

2Al + 3H2SO4 → Al 2 (SO 4) 3 + 3H 2.

In concentrated HNO 3 and H 2 SO 4 aluminum is passivated. Thanks to this, it is possible to store and transport these acids in containers made of aluminum.

6. Interacts with alkalis, as they dissolve the oxide film.

7. Reacts with all non-metals except hydrogen. To carry out the reaction with oxygen, finely divided aluminum is needed. The reaction is only possible at high t:

4Al + 3O 2 → 2Al 2 O 3 .

According to its thermal effect, this reaction is exothermic. Interaction with sulfur forms aluminum sulfide Al 2 S 3 , with phosphorus phosphide AlP, with nitrogen nitride AlN, with carbon carbide Al 4 C 3 .

8. It interacts with other metals, forming aluminides (FeAl 3 CuAl 2, CrAl 7, etc.).

Receiving aluminum:

Metallic aluminum is obtained by electrolysis of a solution of alumina Al 2 O 3 in molten cryolite Na 2 AlF 6 at 960–970°C.

2Al2O3 → 4Al + 3O 2 .

Chemical properties of transition elements

Transitional elements include elements of secondary subgroups of the Periodic Table. Consider the chemical properties of copper, zinc, chromium and iron.

Chemical properties of copper

1. In the electrochemical series, it is located to the right of H, so this metal is inactive.

2. Weak reducer.

3. In compounds, it exhibits oxidation states +1 and +2.

4. Reacts with oxygen when heated to form:

copper oxide (I) 2Cu + O 2 → 2CuO(at t 400 0 C)
or copper(II) oxide: 4Cu + O2 → 2Cu2O(at t 200 0 C).

Oxides have basic properties. When heated in an inert atmosphere, Cu 2 O disproportionates: Cu2O → CuO + Cu. Copper (II) oxide CuO forms cuprates in reactions with alkalis, for example: CuO + 2NaOH → Na 2 CuO 2 + H 2 O.

5. Copper hydroxide Cu (OH) 2 is amphoteric, the main properties prevail in it. It dissolves easily in acids:

Cu (OH) 2 + 2HNO 3 → Cu(NO 3) 2 + 2H 2 O,

and in concentrated solutions of alkalis with difficulty:

Сu(OH) 2 + 2NaOH → Na 2.

6. The interaction of copper with sulfur under various temperature conditions also forms two sulfides. When heated to 300-400 0 C in a vacuum, copper (I) sulfide is formed:

2Cu+S → Cu2S.

At room temperature, by dissolving sulfur in hydrogen sulfide, copper (II) sulfide can be obtained:

Cu+S → CuS.

7. Of the halogens, it interacts with fluorine, chlorine and bromine, forming halides (CuF 2 , CuCl 2 , CuBr 2), iodine, forming copper (I) iodide CuI; does not interact with hydrogen, nitrogen, carbon, silicon.

8. It does not react with acids - non-oxidizing agents, because they oxidize only metals located to hydrogen in the electrochemical series. This chemical element reacts with oxidizing acids: dilute and concentrated nitric and concentrated sulfuric:

3Cu + 8HNO 3 (diff) → 3Cu(NO 3) 2 + 2NO + 4H 2 O;

Cu + 4HNO 3 (conc) → Cu(NO 3) 2 + 2NO 2 + 2H 2 O;

Cu + 2H 2 SO 4 (conc) → CuSO 4 + SO 2 + 2H 2 O.

9. Interacting with salts, copper displaces from their composition the metals located to the right of it in the electrochemical series. For example,

2FeCl 3 + Cu → CuCl 2 + 2FeCl 2 .

Here we see that copper went into solution, and iron (III) was reduced to iron (II). This reaction is of great practical importance and is used to remove copper deposited on plastic.

Chemical properties of zinc

1. The most active after the alkaline earth metals.

2. It has pronounced reducing properties and amphoteric properties.

3. In compounds, it exhibits an oxidation state of +2.

4. In air, it is covered with an oxide film of ZnO.

5. Interaction with water is possible at a temperature of red heat. As a result, zinc oxide and hydrogen are formed:

Zn + H 2 O → ZnO + H 2.

6. Interacts with halogens, forming halides (ZnF 2 - zinc fluoride, ZnBr 2 - zinc bromide, ZnI 2 - zinc iodide, ZnCl 2 - zinc chloride).

7. With phosphorus it forms the phosphides Zn 3 P 2 and ZnP 2 .

8. With sulfur chalcogenide ZnS.

9. Does not directly react with hydrogen, nitrogen, carbon, silicon and boron.

10. It interacts with non-oxidizing acids, forming salts and displacing hydrogen. For example:

H 2 SO 4 + Zn → ZnSO 4 + H 2
Zn + 2HCl → ZnCl 2 + H 2 .

It also reacts with acids - oxidizing agents: with conc. sulfuric acid forms zinc sulfate and sulfur dioxide:

Zn + 2H 2 SO 4 → ZnSO 4 + SO 2 + 2H 2 O.

11. It actively reacts with alkalis, since zinc is an amphoteric metal. With alkali solutions, it forms tetrahydroxozincates and releases hydrogen:

Zn + 2NaOH + 2H 2 O → Na 2 + H 2 .

Gas bubbles appear on the zinc granules after the reaction. With anhydrous alkalis, when fused, it forms zincates and releases hydrogen:

Zn+ 2NaOH → Na 2 ZnO 2 + H 2.

Chemical properties of chromium

1. Under normal conditions, it is inert, but active when heated.

3. Forms colored compounds.

4. In compounds, it exhibits oxidation states +2 (basic oxide CrO black), +3 (amphoteric oxide Cr 2 O 3 and hydroxide Cr (OH) 3 green) and +6 (acid chromium oxide (VI) CrO 3 and acids: chromic H 2 CrO 4 and two-chrome H 2 Cr 2 O 7, etc.).

5. It interacts with fluorine at t 350-400 0 C, forming chromium (IV) fluoride:

Cr+2F 2 → CrF 4 .

6. With oxygen, nitrogen, boron, silicon, sulfur, phosphorus and halogens at t 600 0 C:

connection with oxygen forms chromium oxide (VI) CrO 3 (dark red crystals),
nitrogen compound - chromium nitride CrN (black crystals),
compound with boron - chromium boride CrB (yellow crystals),
compound with silicon - chromium silicide CrSi,
connection with carbon - chromium carbide Cr 3 C 2 .

7. It reacts with water vapor, being in a hot state, forming chromium (III) oxide and hydrogen:

2Cr + 3H 2 O → Cr 2 O 3 + 3H 2 .

8. It does not react with alkali solutions, but slowly reacts with their melts, forming chromates:

2Cr + 6KOH → 2KCrO 2 + 2K 2 O + 3H 2 .

9. It dissolves in dilute strong acids to form salts. If the reaction takes place in air, Cr 3+ salts are formed, for example:

2Cr + 6HCl + O 2 → 2CrCl 3 + 2H 2 O + H 2 .

Cr + 2HCl → CrCl 2 + H 2 .

10. With concentrated sulfuric and nitric acids, as well as with aqua regia, it reacts only when heated, because. at low temperatures, these acids passivate chromium. Reactions with acids when heated look like this:

2Cr + 6H 2 SO 4 (conc) → Cr 2 (SO 4) 3 + 3SO 2 + 6H 2 O

Cr + 6HNO 3 (conc) → Cr(NO 3) 3 + 3NO 2 + 3H 2 O

Chromium(II) oxide CrO- solid black or red, insoluble in water.

Chemical properties:

It has basic and restorative properties.
When heated to 100 0 C in air, it oxidizes to Cr 2 O 3 - chromium (III) oxide.
It is possible to restore chromium with hydrogen from this oxide: CrO + H 2 → Cr + H 2 O or coke: CrO + C → Cr + CO.
Reacts with hydrochloric acid, while releasing hydrogen: 2CrO + 6HCl → 2CrCl 3 + H 2 + 2H 2 O.
Does not react with alkalis, dilute sulfuric and nitric acids.

Chromium oxide (III) Cr 2 O 3- a refractory substance, dark green in color, insoluble in water.

Chemical properties:

It has amphoteric properties.
How basic oxide interacts with acids: Cr 2 O 3 + 6HCl → CrCl 3 + 3H 2 O.
How acidic oxide interacts with alkalis: Cr 2 O 3 + 2KOH → 2KCrO 3 + H 2 O.
Strong oxidizing agents oxidize Cr 2 O 3 to chromate H 2 CrO 4 .
Strong reducing agents restoreCr out Cr2O3.

Chromium(II) hydroxide Cr(OH) 2 - solid yellow or brown color, poorly soluble in water.

Chemical properties:

Weak base, exhibits basic properties.
In the presence of moisture in air, it oxidizes to Cr(OH) 3 - chromium (III) hydroxide.
Reacts with concentrated acids to form blue chromium (II) salts: Cr(OH) 2 + H 2 SO 4 → CrSO 4 + 2H 2 O.
Does not react with alkalis and dilute acids.

Chromium (III) hydroxide Cr(OH) 3 - a gray-green substance, insoluble in water.

Chemical properties:

It has amphoteric properties.
How basic hydroxide interacts with acids: Cr(OH) 3 + 3HCl → CrCl 3 + 3H 2 O.
How acid hydroxide interacts with alkalis: Cr(OH) 3 + 3NaOH → Na 3 [Cr(OH)6].

Chemical properties of iron

1. Active metal with high reactivity.

2. It has restorative properties, as well as pronounced magnetic properties.

3. In compounds, it exhibits the main oxidation states +2 (with weak oxidizing agents: S, I, HCl, salt solutions), +3 (with strong oxidizing agents: Br and Cl) and less characteristic +6 (with O and H 2 O). In weak oxidizing agents, iron takes the oxidation state +2, in stronger ones +3. +2 oxidation states correspond to black oxide FeO and green hydroxide Fe (OH) 2, which have basic properties. +3 oxidation states correspond to red-brown oxide Fe 2 O 3 and brown hydroxide Fe (OH) 3, which have weakly pronounced amphoteric properties. Fe (+2) is a weak reducing agent, and Fe (+3) is often a weak oxidizing agent. When the redox conditions change, the oxidation states of iron can change with each other.

4. In air at t 200 0 C, it is covered with an oxide film. Under normal atmospheric conditions, it is easily corroded. P When oxygen is passed through an iron melt, FeO oxide is formed. When iron is burned in air, oxide Fe 2 O 3 is formed. When burned in pure oxygen, an oxide is formed - iron scale:

3Fe + 2O 2 → Fe 3 O 4.

5. Reacts with halogens when heated:

connection with chlorine forms iron (III) chloride FeCl 3,
compound with bromine - iron (III) bromide FeBr 3,
compound with iodine - iron (II,III) iodide Fe 3 I 8,
compound with fluorine - iron (II) fluoride FeF 2, iron (III) fluoride FeF 3.

6. It also reacts with sulfur, nitrogen, phosphorus, silicon and carbon when heated:

connection with sulfur forms iron(II) sulfide FeS,
connection with nitrogen - iron nitride Fe 3 N,
compound with phosphorus - phosphides FeP, Fe 2 P and Fe 3 P,
compound with silicon - iron silicide FeSi,
compound with carbon - iron carbide Fe 3 C.

2Fe + 4H 2 SO 4 → Fe 2 (SO 4) 3 + SO 2 + 4H 2 O

9. It does not react with alkali solutions, but slowly reacts with alkali melts, which are strong oxidizing agents:

Fe + KClO 3 + 2KOH → K 2 FeO 4 + KCl + H 2 O.

10. Restores metals located in the electrochemical row to the right:

Fe + SnCl 2 → FeCl 2 + Sn.

Getting iron: In industry, iron is obtained from iron ore, mainly from hematite (Fe 2 O 3) and magnetite (FeO·Fe 2 O 3).

3Fe2O3 + CO → CO 2 + 2Fe 3 O 4,
Fe 3 O 4 + CO → CO 2 + 3FeO,
FeO + CO → CO 2 + Fe.

Iron(II) oxide FeO - a black crystalline substance (wustite) that does not dissolve in water.

Chemical properties:

Has basic properties.
Reacts with dilute hydrochloric acid: FeO + 2HCl → FeCl 2 + H 2 O.
Reacts with concentrated nitric acid:FeO + 4HNO 3 → Fe(NO 3) 3 + NO 2 + 2H 2 O.
Does not react with water and salts.
With hydrogen at t 350 0 C it is reduced to pure metal: FeO + H 2 → Fe + H 2 O.
It is also reduced to pure metal when combined with coke: FeO + C → Fe + CO.
This oxide can be obtained in various ways, one of them is heating Fe at low pressure O: 2Fe + O 2 → 2FeO.

Iron(III) oxideFe2O3- brown powder (hematite), a substance insoluble in water. Other names: iron oxide, iron minium, food coloring E172, etc.

Chemical properties:

Fe 2 O 3 + 6HCl → 2 FeCl 3 + 3H 2 O.
It does not react with alkali solutions, it reacts with their melts, forming ferrites: Fe 2 O 3 + 2NaOH → 2NaFeO 2 + H 2 O.
When heated with hydrogen, it exhibits oxidizing properties:Fe 2 O 3 + H 2 → 2FeO + H 2 O.
Fe 2 O 3 + 3KNO 3 + 4KOH → 2K 2 FeO 4 + 3KNO 2 + 2H 2 O.

Iron oxide (II, III) Fe 3 O 4 or FeO Fe 2 O 3 - a grayish-black solid (magnetite, magnetic iron ore), a substance insoluble in water.

Chemical properties:

Decomposes when heated above 1500 0 С: 2Fe 3 O 4 → 6FeO + O 2.
Reacts with dilute acids: Fe 3 O 4 + 8HCl → FeCl 2 + 2FeCl 3 + 4H 2 O.
Does not react with alkali solutions, reacts with their melts: Fe 3 O 4 + 14NaOH → Na 3 FeO 3 + 2Na 5 FeO 4 + 7H 2 O.
When reacting with oxygen, it oxidizes: 4Fe 3 O 4 + O 2 → 6Fe 2 O 3.
With hydrogen, when heated, it is restored:Fe 3 O 4 + 4H 2 → 3Fe + 4H 2 O.
It is also reduced when combined with carbon monoxide: Fe 3 O 4 + 4CO → 3Fe + 4CO 2.

Iron(II) hydroxide Fe(OH) 2 - white, rarely greenish crystalline substance, insoluble in water.

Chemical properties:

It has amphoteric properties with a predominance of basic ones.
It enters into the neutralization reaction of the non-oxidizing acid, showing the main properties: Fe(OH) 2 + 2HCl → FeCl 2 + 2H 2 O.
When interacting with nitric or concentrated sulfuric acids, it exhibits reducing properties, forming iron (III) salts: 2Fe(OH) 2 + 4H 2 SO 4 → Fe 2 (SO 4) 3 + SO 2 + 6H 2 O.
When heated, it reacts with concentrated alkali solutions: Fe (OH) 2 + 2NaOH → Na 2.

Iron hydroxide (I I I) Fe (OH) 3- brown crystalline or amorphous substance, insoluble in water.

Chemical properties:

It has mild amphoteric properties with a predominance of basic ones.
Easily interacts with acids: Fe(OH) 3 + 3HCl → FeCl 3 + 3H 2 O.
With concentrated alkali solutions it forms hexahydroxoferrates (III): Fe (OH) 3 + 3NaOH → Na 3.
It forms ferrates with alkali melts:2Fe(OH) 3 + Na 2 CO 3 → 2NaFeO 2 + CO 2 + 3H 2 O.
In an alkaline environment with strong oxidizing agents, it exhibits reducing properties: 2Fe(OH) 3 + 3Br 2 + 10KOH → 2K 2 FeO 4 + 6NaBr + 8H 2 O.

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Lesson Objectives: consider the distribution of aluminum in nature, its physical and chemical properties, as well as the properties of the compounds it forms.

Progress

2. Learning new material. Aluminum

The main subgroup of group III of the periodic system is boron (B), aluminum (Al), gallium (Ga), indium (In) and thallium (Tl).

As can be seen from the above data, all these elements were discovered in the 19th century.

Discovery of metals of the main subgroup III groups


1806	1825	1875	1863	1861
G. Lussac,	G.H. Oersted	L. de Boisbaudran	F. Reich,	W. Crooks
L. Tenard	(Denmark)	(France)	I. Richter	(England)
(France)			(Germany)

Boron is a nonmetal. Aluminum is a transition metal, while gallium, indium and thallium are full metals. Thus, with an increase in the atomic radii of the elements of each group of the periodic system, the metallic properties of simple substances increase.

In this lecture, we will take a closer look at the properties of aluminum.

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MUNICIPAL BUDGET EDUCATIONAL INSTITUTION

GENERAL EDUCATIONAL SCHOOL № 81

Aluminum. The position of aluminum in the periodic system and the structure of its atom. Finding in nature. Physical and chemical properties of aluminum.

chemistry teacher

MBOU secondary school №81

2013

Lesson topic: Aluminum. The position of aluminum in the periodic system and the structure of its atom. Finding in nature. Physical and chemical properties of aluminum.

Lesson Objectives: consider the distribution of aluminum in nature, its physical and chemical properties, as well as the properties of the compounds it forms.

Progress

1. Organizational moment of the lesson.

2. Learning new material. Aluminum

The main subgroup of group III of the periodic system is boron (B),aluminum (Al), gallium (Ga), indium (In) and thallium (Tl).

As can be seen from the above data, all these elements were discovered in the 19th century.

Discovery of metals of the main subgroup of group III


1806	1825	1875	1863	1861
G. Lussac,	G.H. Oersted	L. de Boisbaudran	F. Reich,	W. Crooks
L. Tenard	(Denmark)	(France)	I. Richter	(England)
(France)			(Germany)

In this lecture, we will take a closer look at the properties of aluminum.

1. The position of aluminum in the table of D. I. Mendeleev. The structure of the atom, the oxidation states shown.

The element aluminum is located in group III, main “A” subgroup, 3rd period of the periodic system, serial number No. 13, relative atomic mass Ar (Al) \u003d 27. Its neighbor on the left in the table is magnesium - a typical metal, and on the right - silicon - already a non-metal . Therefore, aluminum must exhibit properties of some intermediate nature and its compounds are amphoteric.

Al +13) 2 ) 8 ) 3 , p is an element,

Basic state

1s 2 2s 2 2p 6 3s 2 3p 1

excited state

1s 2 2s 2 2p 6 3s 1 3p 2

Aluminum exhibits an oxidation state of +3 in compounds:

Al 0 - 3 e - → Al +3

2. Physical properties

Free form aluminum is a silvery-white metal with high thermal and electrical conductivity. Melting point 650 about C. Aluminum has a low density (2.7 g/cm 3 ) - about three times less than that of iron or copper, and at the same time it is a durable metal.

3. Being in nature

In terms of prevalence in nature, it occupies1st among metals and 3rd among elementssecond only to oxygen and silicon. The percentage of aluminum content in the earth's crust, according to various researchers, ranges from 7.45 to 8.14% of the mass of the earth's crust.

In nature, aluminum occurs only in compounds(minerals).

Some of them:

Bauxites - Al 2 O 3 H 2 O (with impurities SiO 2, Fe 2 O 3, CaCO 3)

Nephelines - KNa 3 4

Alunites - KAl(SO 4 ) 2 2Al(OH) 3

Alumina (mixtures of kaolins with sand SiO 2 , limestone CaCO 3 , magnesite MgCO 3 )

Corundum - Al 2 O 3

Feldspar (orthoclase) - K 2 O × Al 2 O 3 × 6 SiO 2

Kaolinite - Al 2 O 3 ×2SiO 2 × 2H 2 O

Alunite - (Na,K) 2 SO 4 × Al 2 (SO 4 ) 3 × 4Al (OH) 3

Beryl - 3BeO Al 2 O 3 6SiO 2

Bauxite
Al2O3	Corundum
	Ruby
	Sapphire

4. Chemical properties of aluminum and its compounds

Aluminum easily interacts with oxygen under normal conditions and is covered with an oxide film (it gives a matte appearance).

Its thickness is 0.00001 mm, but thanks to it, aluminum does not corrode. To study the chemical properties of aluminum, the oxide film is removed. (Using sandpaper, or chemically: first lowering into an alkali solution to remove the oxide film, and then into a solution of mercury salts to form an aluminum alloy with mercury - amalgam).

I. Interaction with simple substances

Aluminum already at room temperature actively reacts with all halogens, forming halides. When heated, it interacts with sulfur (200 °C), nitrogen (800 °C), phosphorus (500 °C) and carbon (2000 °C), with iodine in the presence of a catalyst - water:

2Al + 3S \u003d Al 2 S 3 (aluminum sulfide),

2Al + N 2 = 2AlN (aluminum nitride),

Al + P = AlP (aluminum phosphide),

4Al + 3C \u003d Al 4 C 3 (aluminum carbide).

2 Al + 3 I 2 = 2 AlI 3 (aluminum iodide)

All these compounds are completely hydrolyzed with the formation of aluminum hydroxide and, accordingly, hydrogen sulfide, ammonia, phosphine and methane:

Al 2 S 3 + 6H 2 O \u003d 2Al (OH) 3 + 3H 2 S

Al 4 C 3 + 12H 2 O \u003d 4Al (OH) 3 + 3CH 4

In the form of shavings or powder, it burns brightly in air, releasing a large amount of heat:

4Al + 3O 2 = 2Al 2 O 3 + 1676 kJ.

II. Interaction with complex substances

Interaction with water:

2 Al + 6 H 2 O \u003d 2 Al (OH) 3 + 3 H 2

without oxide film

Interaction with metal oxides:

Aluminum is a good reducing agent, as it is one of the active metals. It is in the activity series right after the alkaline earth metals. That's whyrestores metals from their oxides. Such a reaction - aluminothermy - is used to obtain pure rare metals, such as tungsten, vanadium, etc.

3 Fe 3 O 4 + 8 Al \u003d 4 Al 2 O 3 + 9 Fe + Q

Thermite mixture Fe 3 O 4 and Al (powder) - also used in thermite welding.

Cr 2 O 3 + 2Al \u003d 2Cr + Al 2 O 3

Interaction with acids:

With sulfuric acid solution: 2 Al + 3 H 2 SO 4 \u003d Al 2 (SO 4) 3 + 3 H 2

It does not react with cold concentrated sulfuric and nitrogenous (passivates). Therefore, nitric acid is transported in aluminum tanks. When heated, aluminum is able to reduce these acids without releasing hydrogen:

2Al + 6H 2 SO 4 (conc) \u003d Al 2 (SO 4) 3 + 3SO 2 + 6H 2 O,

Al + 6HNO 3 (conc) \u003d Al (NO 3) 3 + 3NO 2 + 3H 2 O.

Interaction with alkalis.

2 Al + 2 NaOH + 6 H 2 O \u003d 2 NaAl (OH) 4 + 3 H 2

Na [Al (OH) 4] - sodium tetrahydroxoaluminate

At the suggestion of the chemist Gorbov, during the Russo-Japanese War, this reaction was used to produce hydrogen for balloons.

With salt solutions:

2Al + 3CuSO 4 \u003d Al 2 (SO 4) 3 + 3Cu

If the surface of aluminum is rubbed with mercury salt, then the following reaction occurs:

2Al + 3HgCl 2 = 2AlCl 3 + 3Hg

The released mercury dissolves the aluminum, forming an amalgam.

5. Application of aluminum and its compounds

The physical and chemical properties of aluminum have led to its widespread use in technology.The aviation industry is a major consumer of aluminum.: 2/3 aircraft is made of aluminum and its alloys. An aircraft made of steel would be too heavy and could carry far fewer passengers.Therefore, aluminum is called the winged metal.Cables and wires are made from aluminum: with the same electrical conductivity, their mass is 2 times less than the corresponding copper products.

Considering the corrosion resistance of aluminum, itmanufacture parts of apparatuses and containers for nitric acid. Aluminum powder is the basis for the manufacture of silver paint to protect iron products from corrosion, as well as to reflect heat rays, such paint is used to cover oil storage facilities and firefighters' suits.

Aluminum oxide is used to produce aluminum and also as a refractory material.

Aluminum hydroxide is the main component of the well-known drugs Maalox, Almagel, which lower the acidity of gastric juice.

Aluminum salts are highly hydrolyzed. This property is used in the process of water purification. Aluminum sulfate and a small amount of slaked lime are added to the water to be purified to neutralize the resulting acid. As a result, a volumetric precipitate of aluminum hydroxide is released, which, settling, takes with it suspended particles of turbidity and bacteria.

Thus, aluminum sulfate is a coagulant.

6. Obtaining aluminum

1) The modern cost-effective method for producing aluminum was invented by the American Hall and the Frenchman Héroux in 1886. It consists in the electrolysis of a solution of aluminum oxide in molten cryolite. Molten cryolite Na 3 AlF 6 dissolves Al 2 O 3, how water dissolves sugar. The electrolysis of a "solution" of aluminum oxide in molten cryolite proceeds as if cryolite were only a solvent, and aluminum oxide was an electrolyte.

2Al 2 O 3 electric current → 4Al + 3O 2

In the English Encyclopedia for Boys and Girls, an article about aluminum begins with the following words: “On February 23, 1886, a new metal age began in the history of civilization - the age of aluminum. On this day, Charles Hall, a 22-year-old chemist, showed up in his first teacher's laboratory with a dozen small balls of silvery-white aluminum in his hand, and with the news that he had found a way to manufacture this metal cheaply and in large quantities. So Hall became the founder of the American aluminum industry and an Anglo-Saxon national hero, as a man who made a great business out of science.

2) 2Al 2 O 3 + 3 C \u003d 4 Al + 3 CO 2

IT IS INTERESTING:

Metallic aluminum was first isolated in 1825 by the Danish physicist Hans Christian Oersted. By passing gaseous chlorine through a layer of hot alumina mixed with coal, Oersted isolated aluminum chloride without the slightest trace of moisture. To restore metallic aluminum, Oersted needed to treat aluminum chloride with potassium amalgam. After 2 years, the German chemist Friedrich Wöller. He improved the method by replacing potassium amalgam with pure potassium.
In the 18th and 19th centuries, aluminum was the main jewelry metal. In 1889, in London, D.I. Mendeleev was awarded a valuable gift for his services to the development of chemistry - scales made of gold and aluminum.
By 1855, the French scientist Saint-Clair Deville had developed a process for producing aluminum metal on an industrial scale. But the method was very expensive. Deville enjoyed the special patronage of Napoleon III, Emperor of France. As a sign of his devotion and gratitude, Deville made for Napoleon's son, the newborn prince, an elegantly engraved rattle - the first "consumer product" made of aluminum. Napoleon even intended to equip his guardsmen with aluminum cuirasses, but the price was prohibitive. At that time, 1 kg of aluminum cost 1000 marks, i.e. 5 times more expensive than silver. It wasn't until the invention of the electrolytic process that aluminum became as valuable as conventional metals.
Did you know that aluminum, entering the human body, causes a disorder of the nervous system. With its excess, metabolism is disturbed. And protective agents are vitamin C, calcium, zinc compounds.
When aluminum burns in oxygen and fluorine, a lot of heat is released. Therefore, it is used as an additive to rocket fuel. The Saturn rocket burns 36 tons of aluminum powder during its flight. The idea of using metals as a component of rocket fuel was first proposed by F.A. Zander.

3. Consolidation of the studied material

No. 1. To obtain aluminum from aluminum chloride, calcium metal can be used as a reducing agent. Make an equation for this chemical reaction, characterize this process using electronic balance.
Think! Why can't this reaction be carried out in an aqueous solution?

No. 2. Complete the equations of chemical reactions:
Al+H 2 SO 4 (solution) ->
Al + CuCl 2 ->
Al + HNO 3 (conc) - t ->
Al + NaOH + H 2 O ->

Number 3. Solve the problem:
An aluminum-copper alloy was exposed to an excess of concentrated sodium hydroxide solution while being heated. 2.24 liters of gas (n.o.s.) were released. Calculate the percentage composition of the alloy if its total mass was 10 g?

4. Homework slide 2

AL Element III (A) of the table group D.I. Mendeleev Element with serial number 13, its Element of the 3rd period The third most common in the earth's crust, the name is derived from lat. "Aluminis" - alum

Danish physicist Hans Oersted (1777-1851) For the first time, aluminum was obtained by him in 1825 by the action of potassium amalgam on aluminum chloride, followed by distillation of mercury.

Modern production of aluminum The modern production method was developed independently by the American Charles Hall and the Frenchman Paul Héroux in 1886. It consists in dissolving alumina in a cryolite melt followed by electrolysis using consumable coke or graphite electrodes.

As a student at Oberlin College, he learned that you can get rich and receive the gratitude of mankind if you invent a method for producing aluminum on an industrial scale. Like a man possessed, Charles conducted experiments on the production of aluminum by electrolysis of a cryolite-alumina melt. On February 23, 1886, a year after graduating from college, Charles produced the first aluminum by electrolysis. Hall Charles (1863 - 1914) American chemical engineer

Paul Héroux (1863-1914) - French chemical engineer In 1889 he opened an aluminum plant in Fron (France), becoming its director, he designed an electric arc furnace for steel smelting, named after him; he also developed an electrolytic method for producing aluminum alloys

8 Aluminum 1. From the history of the discovery Main Next During the discovery of aluminum, the metal was more expensive than gold. The British wanted to honor the great Russian chemist D.I. Mendeleev with a rich gift, they gave him a chemical balance, in which one cup was made of gold, the other - of aluminum. A cup made of aluminum has become more expensive than gold. The resulting "silver from clay" interested not only scientists, but also industrialists and even the emperor of France. Further

9 Aluminum 7. Content in the earth's crust main Next

Finding in nature The most important aluminum mineral today is bauxite. The main chemical component of bauxite is alumina (Al 2 O 3) (28 - 80%).

11 Aluminum 4. Physical properties Color - silver-white t pl. = 660 °C. t b.p. ≈ 2450 °C. Electrically conductive, thermally conductive Lightweight, density ρ = 2.6989 g/cm 3 Soft, ductile. home next

12 Aluminum 7. Found in nature Bauxite – Al 2 O 3 Alumina – Al 2 O 3 main Next

13 Aluminum main Insert the missing words Aluminum is an element of group III, the main subgroup. The charge of the nucleus of an aluminum atom is +13. There are 13 protons in the nucleus of an aluminum atom. There are 14 neutrons in the nucleus of an aluminum atom. There are 13 electrons in an aluminum atom. The aluminum atom has 3 energy levels. The electron shell has a structure of 2e, 8e, 3e. At the outer level, there are 3 electrons in an atom. The oxidation state of an atom in compounds is +3. The simple substance aluminum is a metal. Aluminum oxide and hydroxide are amphoteric in nature. Further

14 Aluminum 3 . The structure of a simple substance Metal Bond - metallic Crystal lattice - metallic, cubic face-centered main More

15 Aluminum 2. Electronic structure 27 A l +13 0 2e 8e 3e P + = 13 n 0 = 14 e - = 13 1 s 2 2 s 2 2p 6 3s 2 3p 1 Short electronic record 1 s 2 2 s 2 2p 6 3s 2 3p 1 Filling order main Next

Aluminum \u003d 2AlCl 3 4Al + 3C \u003d Al 4 C 3 C non-metals (with halogens, with carbon) (Remove the oxide film) 2 Al + 6 H 2 O \u003d 2Al (OH) 2 + H 2 C with water 2 Al + 6 HCl \u003d 2AlCl 3 + H 2 2Al + 3H 2 SO 4 \u003d Al 2 (SO 4) 3 + H 2 C acids and 2 Al + 6NaOH + 6H 2 O \u003d 2Na 3 [Al (OH ) 6] + 3H 2 2Al + 2NaOH + 2H 2 O \u003d 2NaAlO 2 + 3H 2 C with alkalis and 8Al + 3Fe 3 O 4 \u003d 4Al 2 O 3 + 9Fe 2Al + WO 3 \u003d Al 2 O 3 + W C oxi d a m e t a l l

17 Aluminum 8. Obtaining 1825 H. Oersted: AlCl 3 + 3K = 3KCl + Al: Electrolysis (t pl. = 2050 ° C): 2Al 2 O 3 = 4 Al + 3O 2 Electrolysis (in melting cryolite Na 3 AlF 6, t pl ≈ 1000 ° С): 2Al 2 O 3 \u003d 4 Al + 3O 2 main Next

At the end of the 90s, the Electrical Installation Rules (PUE) of the 7th edition were put into effect in Russia, according to which the electrical installation of internal networks of buildings from aluminum cables and wires with a cross section of less than 16 mm2 is prohibited, and it is prescribed to be made from copper wire. The reason for the change in regulatory requirements was some of the properties of aluminum.

aluminum as an electrical conductor

Aluminum cables and wires have long been widely used both for wiring internal power networks in buildings for various purposes, and for laying outdoor power lines. This is due to the following properties of aluminum:

low specific gravity, which is three times lighter than that of copper;
ease of processing;
low material cost;
good electrical conductivity, per unit mass;
high corrosion resistance.

However, other features of aluminum are: high fluidity, which does not provide sufficient quality of contacts for a long time; low strength under mechanical impact on fracture; low heat resistance, leading to an increase in brittleness during overheating - led to the introduction of a ban on the electrical installation of aluminum wires of small cross section for internal power supply networks.

One of the main reasons that influenced the change in the requirements of the PUE is that during operation a thin oxide film is formed on the surface of aluminum wires, which has a much worse electrical conductivity than the base metal. As a result, a higher contact resistance is formed at the junction of the wires, which significantly increases the possibility of heating the contacts, the risk of their destruction and fire.

Copper used as a material for electrical cables and wires, despite the higher cost, is devoid of the listed disadvantages of aluminum and has a number of advantages: higher conductivity; does not form an oxide film on the surface; higher flexibility, this allows the production of wires with a very small cross section of up to 0.3 mm2, which cannot be made from aluminum.

Connection of aluminum and copper wires

Since many buildings of the old construction retain electric networks made of aluminum wires, during repairs it often becomes necessary to connect wiring from different materials - copper and aluminum. According to the same Electrical Installation Rules, the connection of aluminum and copper wires can be done in several ways:

with the help of “nuts” type connections, consisting of three plates, between which wires are clamped with bolts;
by means of WAGO clamps. The ends of the wires to be connected are stripped by 10-15 mm, inserted into different holes of the terminal block, then clamped with lowering blocks;
using terminal blocks, which are a bar with two holes. The ends of the connected wires are inserted into the holes from different ends and clamped with a screw
using a simple bolted connection, when the wires are clamped with a nut with a metal washer laid between them. This method is considered temporary, as it is not suitable for rooms with high humidity and is not used for outdoor connections.

The article was prepared based on materials from the site http://energy-systems.ru/

General characteristics.

The concept of a transition element is usually used to refer to any element with valence d- or f-electrons. These elements occupy a transitional position in the periodic table between electropositive s-elements and electronegative p-elements (see § 2, 3).

d-Elements are called the main transition elements. Their atoms are characterized by internal building up of d-subshells. The fact is that the s-orbital of their outer shell is usually filled already before the filling of the d-orbitals in the previous electron shell begins. This means that each new electron added to the electron shell of the next d-element, in accordance with the filling principle (see § 2), does not fall on the outer shell, but on the inner subshell that precedes it. The chemical properties of these elements are determined by the participation of electrons in the reactions of both of these shells.

d-Elements form three transition series - in the 4th, 5th and 6th periods, respectively. The first transitional series includes 10 elements, from scandium to zinc. It is characterized by internal development -orbitals (Table 15.1). The orbital fills up earlier than the orbital because it has less energy (see Klechkovsky's rule, § 2).

However, two anomalies should be noted. Chromium and copper have only one electron each in their -orbitals. This is because half-filled or filled subshells are more stable than partially filled subshells.

In the chromium atom, each of the five -orbitals that form the -subshell has one electron. Such a subshell is half-filled. In a copper atom, each of the five -orbitals has a pair of electrons. A similar anomaly is observed in silver.

(A l ), gallium (Ga ), indium (In ) and thallium (T l ).

As can be seen from the given data, all these elements were opened in XIX century.

Discovery of metals of the main subgroup III groups

AT	Al	Ga	In	Tl
1806	1825	1875	1863	1861
G. Lussac,	G.H. Oersted	L. de Boisbaudran	F. Reich,	W. Crooks
L. Tenard	(Denmark)	(France)	I. Richter	(England)
(France)			(Germany)

In this lecture, we will take a closer look at the properties of aluminum.

1. The position of aluminum in the table of D. I. Mendeleev. The structure of the atom, the oxidation states shown.

The aluminum element is located in III group, main "A" subgroup, 3rd period of the periodic system, serial number No. 13, relative atomic mass Ar (Al ) = 27. Its neighbor on the left in the table is magnesium, a typical metal, and on the right, silicon, which is no longer a metal. Therefore, aluminum must exhibit properties of some intermediate nature and its compounds are amphoteric.

Al +13) 2) 8) 3 , p is an element,

Basic state

1s 2 2s 2 2p 6 3s 2 3p 1

excited state

1s 2 2s 2 2p 6 3s 1 3p 2

Aluminum exhibits an oxidation state of +3 in compounds:

Al 0 - 3 e - → Al +3

2. Physical properties

Free form aluminum is a silvery-white metal with high thermal and electrical conductivity.The melting temperature is 650 ° C. Aluminum has a low density (2.7 g / cm 3) - about three times less than that of iron or copper, and at the same time it is a durable metal.

3. Being in nature

In terms of prevalence in nature, it occupies 1st among metals and 3rd among elements second only to oxygen and silicon. The percentage of aluminum content in the earth's crust, according to various researchers, ranges from 7.45 to 8.14% of the mass of the earth's crust.

In nature, aluminum occurs only in compounds (minerals).

Some of them:

· Bauxites - Al 2 O 3 H 2 O (with impurities SiO 2, Fe 2 O 3, CaCO 3)

· Nephelines - KNa 3 4

· Alunites - KAl(SO 4) 2 2Al(OH) 3

· Alumina (mixtures of kaolins with sand SiO 2, limestone CaCO 3, magnesite MgCO 3)

· Corundum - Al 2 O 3

· Feldspar (orthoclase) - K 2 O × Al 2 O 3 × 6SiO 2

· Kaolinite - Al 2 O 3 ×2SiO 2 × 2H 2 O

· Alunite - (Na,K) 2 SO 4 × Al 2 (SO 4) 3 × 4Al (OH) 3

· Beryl - 3BeO Al 2 O 3 6SiO 2

Bauxite
Al2O3	Corundum
	Ruby
	Sapphire

4. Chemical properties of aluminum and its compounds

Aluminum easily interacts with oxygen under normal conditions and is covered with an oxide film (it gives a matte appearance).

DEMONSTRATION OF OXIDE FILM

Its thickness is 0.00001 mm, but thanks to it, aluminum does not corrode. To study the chemical properties of aluminum, the oxide film is removed. (Using sandpaper, or chemically: first by lowering into an alkali solution to remove the oxide film, and then into a solution of mercury salts to form an aluminum-mercury alloy - an amalgam).

I. Interaction with simple substances

2A l + 3 S \u003d A l 2 S 3 (aluminum sulfide),

2A l + N 2 \u003d 2A lN (aluminum nitride),

A l + P = A l P (aluminum phosphide),

4A l + 3C \u003d A l 4 C 3 (aluminum carbide).

2 Al +3 I 2 \u003d 2 A l I 3 (aluminum iodide) AN EXPERIENCE

All these compounds are completely hydrolyzed with the formation of aluminum hydroxide and, accordingly, hydrogen sulfide, ammonia, phosphine and methane:

Al 2 S 3 + 6H 2 O \u003d 2Al (OH) 3 + 3H 2 S

Al 4 C 3 + 12H 2 O \u003d 4Al (OH) 3 + 3CH 4

In the form of shavings or powder, it burns brightly in air, releasing a large amount of heat:

4A l + 3 O 2 \u003d 2A l 2 O 3 + 1676 kJ.

COMBUSTION OF ALUMINUM IN AIR

AN EXPERIENCE

II. Interaction with complex substances

Interaction with water :

2 Al + 6 H 2 O \u003d 2 Al (OH) 3 +3 H 2

without oxide film

AN EXPERIENCE

Interaction with metal oxides:

Aluminum is a good reducing agent, as it is one of the active metals. It is in the activity series right after the alkaline earth metals. That's why restores metals from their oxides . Such a reaction - aluminothermy - is used to obtain pure rare metals, such as tungsten, vanadium, etc.

3 Fe 3 O 4 +8 Al \u003d 4 Al 2 O 3 +9 Fe + Q

Thermite mixture of Fe 3 O 4 and Al (powder) is also used in thermite welding.

C r 2 O 3 + 2A l \u003d 2C r + A l 2 O 3

Interaction with acids :

With a solution of sulfuric acid: 2 Al + 3 H 2 SO 4 \u003d Al 2 (SO 4) 3 +3 H 2

2A l + 6H 2 S O 4 (conc) \u003d A l 2 (S O 4) 3 + 3 S O 2 + 6H 2 O,

A l + 6H NO 3 (conc) \u003d A l (NO 3) 3 + 3 NO 2 + 3H 2 O.

Interaction with alkalis .

2 Al + 2 NaOH + 6 H 2 O \u003d 2 Na [ Al(OH)4 ] +3H2

AN EXPERIENCE

Na[BUTl(OH) 4] – sodium tetrahydroxoaluminate

At the suggestion of the chemist Gorbov, during the Russo-Japanese War, this reaction was used to produce hydrogen for balloons.

With salt solutions:

2 Al + 3 CuSO 4 \u003d Al 2 (SO 4) 3 + 3 Cu

If the surface of aluminum is rubbed with mercury salt, then the following reaction occurs:

2 Al + 3 HgCl 2 = 2 AlCl 3 + 3 hg

Released mercury dissolves aluminum, forming an amalgam .

Detection of aluminum ions in solutions : AN EXPERIENCE

5. Application of aluminum and its compounds

The physical and chemical properties of aluminum have led to its widespread use in technology. The aviation industry is a major consumer of aluminum.: 2/3 aircraft is made of aluminum and its alloys. An aircraft made of steel would be too heavy and could carry far fewer passengers. Therefore, aluminum is called the winged metal. Cables and wires are made from aluminum: with the same electrical conductivity, their mass is 2 times less than the corresponding copper products.

Considering the corrosion resistance of aluminum, it manufacture parts of apparatuses and containers for nitric acid. Aluminum powder is the basis for the manufacture of silver paint to protect iron products from corrosion, as well as to reflect thermal rays, such paint is used to cover oil storage facilities and firefighters' suits.

Aluminum oxide is used to produce aluminum and also as a refractory material.

Aluminum hydroxide is the main component of the well-known drugs Maalox, Almagel, which lower the acidity of gastric juice.

Aluminum salts are strongly hydrolyzed. This property is used in the process of water purification. Aluminum sulfate and a small amount of slaked lime are added to the water to be purified to neutralize the resulting acid. As a result, a volumetric precipitate of aluminum hydroxide is released, which, settling, takes with it suspended particles of turbidity and bacteria.

Thus, aluminum sulfate is a coagulant.

6. Obtaining aluminum

1) The modern cost-effective method for producing aluminum was invented by the American Hall and the Frenchman Héroux in 1886. It consists in the electrolysis of a solution of aluminum oxide in molten cryolite. Molten cryolite Na 3 AlF 6 dissolves Al 2 O 3 as water dissolves sugar. The electrolysis of a "solution" of aluminum oxide in molten cryolite proceeds as if cryolite were only a solvent, and aluminum oxide was an electrolyte.

2Al 2 O 3 electric current → 4Al + 3O 2

2) 2Al 2 O 3 +3 C \u003d 4 Al + 3 CO 2

IT IS INTERESTING:

Metallic aluminum was first isolated in 1825 by the Danish physicist Hans Christian Oersted. By passing gaseous chlorine through a layer of hot alumina mixed with coal, Oersted isolated aluminum chloride without the slightest trace of moisture. To restore metallic aluminum, Oersted needed to treat aluminum chloride with potassium amalgam. After 2 years, the German chemist Friedrich Wöller. He improved the method by replacing potassium amalgam with pure potassium.
In the 18th and 19th centuries, aluminum was the main jewelry metal. In 1889, in London, D.I. Mendeleev was awarded a valuable gift for his services to the development of chemistry - scales made of gold and aluminum.
By 1855, the French scientist Saint-Clair Deville had developed a process for producing aluminum metal on an industrial scale. But the method was very expensive. Deville enjoyed the special patronage of Napoleon III, Emperor of France. As a sign of his devotion and gratitude, Deville made for Napoleon's son, the newborn prince, an elegantly engraved rattle - the first "consumer product" made of aluminum. Napoleon even intended to equip his guardsmen with aluminum cuirasses, but the price was prohibitive. At that time, 1 kg of aluminum cost 1000 marks, i.e. 5 times more expensive than silver. It wasn't until the invention of the electrolytic process that aluminum became as valuable as conventional metals.
Did you know that aluminum, entering the human body, causes a disorder of the nervous system. When it is in excess, the metabolism is disturbed. And protective agents are vitamin C, calcium, zinc compounds.
When aluminum burns in oxygen and fluorine, a lot of heat is released. Therefore, it is used as an additive to rocket fuel. The Saturn rocket burns 36 tons of aluminum powder during its flight. The idea of using metals as a component of rocket fuel was first proposed by F.A. Zander.

SIMULATORS

Simulator No. 1 - Characteristics of aluminum by position in the Periodic system of elements of D. I. Mendeleev

Simulator No. 2 - Equations for the reactions of aluminum with simple and complex substances

Simulator No. 3 - Chemical properties of aluminum

TASKS FOR REINFORCEMENT

No. 1. To obtain aluminum from aluminum chloride, calcium metal can be used as a reducing agent. Make an equation for this chemical reaction, characterize this process using electronic balance.
Think! Why can't this reaction be carried out in an aqueous solution?

No. 2. Finish the equations of chemical reactions:
Al + H 2 SO 4 (solution ) ->
Al + CuCl 2 ->
Al + HNO 3 ( conc )-t ->
Al + NaOH + H 2 O ->

Number 3. Perform transformations:
Al -> AlCl 3 -> Al -> Al 2 S 3 -> Al(OH) 3 - t -> Al 2 O 3 -> Al

No. 4. Solve the problem:
An aluminum-copper alloy was exposed to an excess of concentrated sodium hydroxide solution while being heated. 2.24 liters of gas (n.o.s.) were released. Calculate the percentage composition of the alloy if its total mass was 10 g?

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aluminum as an electrical conductor

Connection of aluminum and copper wires

General characteristics.

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