Transitional aluminum. Aluminum - general characteristics of the element, chemical properties

General characteristics.

The term 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 usually called the main transition elements. Their atoms are characterized by internal building up of d-subshells. The point is that the s-orbital of their outer shell is usually filled 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), falls not on the outer shell, but on the inner subshell preceding 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 transitional rows - in the 4th, 5th and 6th periods, respectively. The first transition series includes 10 elements, from scandium to zinc. It is characterized by internal building-ups - orbitals (Table 15.1). The orbital fills up earlier than the orbital, because it has less energy (see Klechkovsky's rule, § 2).

However, the existence of two anomalies should be noted. Chromium and copper have only one electron on their -orbitals. This is because semi-filled or filled subshells are more stable than partially filled subshells.

In the chromium atom, there is one electron on each of the five α-orbitals that form the β-subshell. This subshell is half full. In a copper atom, there is a pair of electrons on each of the five -orbitals. A similar anomaly is observed for silver.

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

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 data presented, all these elements were discovered in the 19th century.

Discovery of metals of the main subgroup III group

Boron is a non-metal. Aluminum is a transition metal, while gallium, indium and thallium are high grade metals. Thus, with an increase in the radii of the atoms of the elements of each group of the periodic table, the metallic properties of simple substances increase.

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

Download:

Preview:

MUNICIPAL BUDGETARY EDUCATIONAL INSTITUTION

GENERAL EDUCATIONAL SCHOOL number 81

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

chemistry teacher

MBOU OSH №81

2013

Lesson topic: Aluminum. The position of aluminum in the periodic table and the structure of its atom. Being 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 formed by it.

Progress

1. Organizational moment of the lesson.

2. Learning new material. Aluminum

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

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

Discovery of metals of the main subgroup of group III

Boron is a non-metal. Aluminum is a transition metal, while gallium, indium and thallium are high-grade metals. Thus, with an increase in the radii of the atoms of the elements of each group of the periodic table, the metallic properties of simple substances increase.

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. Atomic structure, exhibited oxidation states.

The element aluminum is located in the III group, the main "A" subgroup, the 3rd period of the periodic system, serial number 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 - already a non-metal ... Consequently, aluminum must exhibit properties of some intermediate character and its compounds are amphoteric.

Al +13) 2) 8) 3, p - element,

Aluminum exhibits an oxidation state of +3 in compounds:

Al 0 - 3 e - → Al +3

2. Physical properties

Free aluminum is a silvery white metal with high thermal and electrical conductivity. Melting point 650 O 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 strong metal.

3. Being in nature

In terms of prevalence in nature, it occupies1st among metals and 3rd among elements, second only to oxygen and silicon. The percentage of aluminum 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 is found only in compounds(minerals).

Some of them:

Bauxite - Al 2 O 3 H 2 O (with admixtures of SiO 2, Fe 2 O 3, CaCO 3)

Nepheline - KNa 3 4

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

Alumina (mixtures of kaolin 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 - 3ВеО Al 2 О 3 6SiO 2

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 look).

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, dipping into an alkali solution to remove the oxide film, and then into a solution of mercury salts to form an alloy of aluminum with mercury - amalgam).

I. Interaction with simple substances

Already at room temperature, aluminum 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:

2Аl + 3S = Аl 2 S 3 (aluminum sulfide),

2Аl + N 2 = 2АlN (aluminum nitride),

Al + P = AlP (aluminum phosphide),

4Аl + 3С = Аl 4 С 3 (aluminum carbide).

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

All these compounds are completely hydrolyzed to form aluminum hydroxide and, accordingly, hydrogen sulfide, ammonia, phosphine and methane:

Al 2 S 3 + 6H 2 O = 2Al (OH) 3 + 3H 2 S

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

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

4Аl + 3O 2 = 2Аl 2 О 3 + 1676 kJ.

II. Interaction with complex substances

Interaction with water:

2 Al + 6 H 2 O = 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. Stands in the line of activity immediately after alkaline earth metals. That's whyrestores metals from their oxides... Such a reaction - alumothermy - is used to obtain pure rare metals such as tungsten, vanadium, etc.

3 Fe 3 O 4 + 8 Al = 4 Al 2 O 3 + 9 Fe + Q

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

Cr 2 O 3 + 2Al = 2Cr + Al 2 O 3

5interactions with acids:

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

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 the evolution of hydrogen:

2Аl + 6Н 2 SO 4 (conc) = Аl 2 (SO 4) 3 + 3SO 2 + 6Н 2 О,

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

Interaction with alkalis.

2 Al + 2 NaOH + 6 H 2 O = 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 = Al 2 (SO 4) 3 + 3Cu

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

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

The released mercury dissolves the aluminum to form 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.: the plane is 2/3 composed of aluminum and its alloys. An airplane made of steel would be too heavy and could carry far fewer passengers.Therefore, aluminum is called a winged metal.Aluminum is used to make cables and wires: with the same electrical conductivity, their mass is 2 times less than the corresponding copper products.

Given the corrosion resistance of aluminum,manufacture parts for devices and containers for nitric acid... Aluminum powder is the basis for the manufacture of silvery paint to protect iron products from corrosion, as well as to reflect heat rays with this paint they cover oil storage tanks 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 treated to neutralize the resulting acid. As a result, a bulk precipitate of aluminum hydroxide is released, which, when settling, carries away suspended particles of turbidity and bacteria.

Thus, aluminum sulfate is a coagulant.

6. Obtaining aluminum

1) The modern cost-effective method of producing aluminum was invented by the American Hall and the Frenchman Eroux in 1886. It consists in the electrolysis of a solution of aluminum oxide in molten cryolite. Molten Na cryolite 3 AlF 6 dissolves Al 2 O 3, how water dissolves sugar. The electrolysis of the “solution” of alumina in molten cryolite occurs as if cryolite was only a solvent, and alumina 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 that day, Charles Hall, a 22-year-old chemist, came to 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 make this metal cheaply and in large quantities. ” Thus Hall became the founder of the American aluminum industry and the Anglo-Saxon national hero, as a man who made a great business out of science.

2) 2Al 2 O 3 + 3 C = 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 incandescent aluminum oxide 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 the potassium amalgam with pure potassium.
• In the 18th and 19th centuries, aluminum was the main jewelry metal. In 1889, D.I. Mendeleev in London for his merits in the development of chemistry was awarded a valuable gift - a balance made of gold and aluminum.
• By 1855, the French scientist Saint-Clair Deville had developed a method for producing metallic aluminum on a technical 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 exquisitely engraved rattle - the first "consumer goods" made of aluminum. Napoleon even intended to equip his guardsmen with aluminum cuirass, but the price turned out to be prohibitive. At that time, 1 kg of aluminum cost 1000 marks, i.e. 5 times more expensive than silver. Only after the invention of the electrolytic process did aluminum become equal in cost to 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 the protective agents are vitamin C, calcium compounds, zinc.
• When aluminum burns in oxygen and fluorine, a lot of heat is generated. Therefore, it is used as an additive to rocket fuel. The Saturn rocket burns 36 tons of aluminum powder during the flight. The idea of ​​using metals as a component of rocket fuel was first expressed by F. A. Tsander.

3. Consolidation of the studied material

# 1. To obtain aluminum from aluminum chloride, metallic calcium can be used as a reducing agent. Make an equation for a given chemical reaction, characterize this process using electronic balance.
Think! Why can't this reaction be carried out in aqueous solution?

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

No. 3. Solve the problem:
The aluminum-copper alloy was exposed to an excess of concentrated sodium hydroxide solution when heated. Allocated 2.24 liters of gas (n.o.). Calculate the percentage of the alloy if its total weight was 10 g?

4. Homework Slide 2

AL Element III (A) of the table group D.I. Mendeleev Element with ordinal number 13, its Element of the 3rd period The third most common in the earth's crust 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 of each other: the American Charles Hall and the Frenchman Paul Héroux in 1886. It consists in dissolving aluminum oxide 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 way to produce aluminum on an industrial scale. Like a man possessed, Charles experimented with the production of aluminum by electrolysis of a cryolite-alumina melt. On February 23, 1886, a year after graduating from college, Charles obtained 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 smelter in Frona (France), becoming its director, he designed an electric arc furnace for smelting steel, named after him; he also developed an electrolytic method for producing aluminum alloys

8 Aluminum 1. From the history of discovery Home Next In the period of discovery of aluminum - metal was more expensive than gold. The British wanted to honor the great Russian chemist D.I. Mendeleev with a rich gift, gave him a chemical balance, in which one cup was made of gold, the other - of aluminum. An aluminum cup has become more expensive than a gold one. 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 home Next

Being 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 - silvery-white t pl. = 660 ° C. t bale. ≈ 2450 ° C. Electrically conductive, thermally conductive Light, density ρ = 2.6989 g / cm 3 Soft, plastic. home Next

12 Aluminum 7. Finding in nature Bauxite - Al 2 O 3 Alumina - Al 2 O 3 home Next

13 Aluminum main Insert the missing words Aluminum is an element of the III group, 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. The aluminum atom has 13 electrons. The aluminum atom has 3 energy levels. The electron shell has a structure of 2 e, 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 - metal Crystal lattice - metal, cubic face-centered main More

15 Aluminum 2. Electronic structure 27 А 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 Order of filling home Next

16 Aluminum 6. Chemical properties 4А l + 3O 2 = 2Al 2 O 3 t 2Al + 3S = Al 2 S 3 C nonmetallam and (with oxygen, with sulfur) 2 А l + 3Cl 2 = 2AlCl 3 4Al + 3C = Al 4 C 3 C with non-metals (with halogens, with carbon) (Remove oxide film) 2 Al + 6 H 2 O = 2Al (OH) 2 + H 2 C water 2 Al + 6 HCl = 2AlCl 3 + H 2 2Al + 3H 2 SO 4 = Al 2 (SO 4) 3 + H 2 C to and with lot am and 2 Al + 6NaOH + 6H 2 O = 2Na 3 [Al (OH ) 6] + 3H 2 2Al + 2NaOH + 2H 2 O = 2NaAlO 2 + 3H 2 C about 8Al + 3Fe 3 O 4 = 4Al 2 O 3 + 9Fe 2Al + WO 3 = Al 2 O 3 + WC o c i d a m i m et al lo v home Next

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

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

Video tutorial 2: Transition metals

Lecture: Typical 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, oxidizing in the process. Let us recall that the more to the left the metal is located in the electrochemical series of tension, the more powerful a 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. Everything 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.

Let's take a look at some of the 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 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 rather reluctantly, while they almost always interact with bases and salts. In order for the reaction of a metal with an acid to take place, it is necessary for the metal to be active and the acid to be 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). Moving 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 have no signs of amphotericity, since they have negative values ​​of electrode potentials.
• The strongest reducing agent of all metals.
• The compounds exhibit only an oxidation state of +1.
• By donating a single valence electron, the atoms of these chemical elements are converted into cations.
• Form numerous ionic compounds.
• Almost everyone dissolves 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 - superoxides (KO 2, RbO 2, CsO 2).

2. With water, forming alkalis and hydrogen. Remember, these reactions are explosive. Only lithium reacts with water without explosion:

2Li + 2Н 2 О → 2LiO Н + Н 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 nitride Li 3 N, with the rest of the alkali metals, the reaction is possible only when heated.

9. They react with acids explosively, therefore 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 with an acid. This creates competition between alkali and acid.

10. With ammonia, forming amides - analogs 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 among all radioactive elements. Its chemical properties are not well understood.

To obtain alkali metals, electrolysis of melts of their halides is mainly used, most often chlorides, which form natural minerals:

• NaCl → 2Na + Cl 2.

• Na 2 CO 3 + 2C → 2Na + 3CO.

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

• 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 increases in the same way as that of alkali metals, i.e. with an increase down the subgroup.

Chemical properties of alkaline earth metals:

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

• By donating 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 one unit higher than that 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, covered 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 slaking reaction. If the reagent is CaO, quicklime is formed, if Ca (OH) 2, slaked lime. 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 that, in comparison with alkali metal hydroxides, are less soluble in water. Alkaline earth metal hydroxides are alkalis, except for amphoteric Be (OH ) 2 and weak base Mg (OH) 2. Since beryllium does not react with water, Be (OH ) 2 can be obtained by other methods, for example, by hydrolysis of nitride:

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

3. Under normal conditions, I react 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 pressure of hydrogen is also required.

5. Form sulfides with sulfur. For example:

• Ca + S → СaS.

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

6. Nitrides are formed 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 (dil.) → BeSO 4 + H 2.

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

Obtaining alkaline earth metals:

• BeF 2 + Mg –t o → Be + MgF 2

• 3BaO + 2Al –t о → 3Ba + Al 2 O 3

• CaCl 2 → Ca + Cl 2

Chemical properties of aluminum

Aluminum is an active, light metal, at number 13 in the table. The most abundant of all metals in nature. And of the chemical elements, it takes 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. Recovers many metals from oxides. This method of obtaining metals is called alumothermy. An example of getting chrome:

2Al + Cr 2 О 3 → Al 2 О 3 + 2Cr.

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

2Al + 6HCl → 2AlCl 3 + 3H 2;

2Al + 3H 2 SO 4 → 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. Interacts with all non-metals except hydrogen. To carry out the reaction with oxygen, finely crushed aluminum is needed. The reaction is possible only at high t:

• 4Al + 3O 2 → 2Al 2 O 3 .

In terms of 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. Interacts with other metals to form aluminides (FeAl 3 CuAl 2, CrAl 7, etc.).

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.

• 2Al 2 O 3 → 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 row, it is located to the right of H, therefore this metal is inactive.

2. Weak reducing agent.

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

4. Reacts with oxygen when heated, forming:

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

Oxides have basic properties. When heated in an inert atmosphere, Cu 2 O disproportionates: Cu 2 O → CuO + Cu... Copper (II) oxide CuO in reactions with alkalis forms cuprates, 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 different temperature conditions also forms two sulfides. When heated to 300-400 0 С in vacuum, copper (I) sulfide is formed:

• 2Cu + S → Cu 2 S.

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

• Cu + S → CuS.

7. Of 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-oxidants, because they oxidize only metals located before hydrogen in the electrochemical series. This chemical element reacts with acids - oxidizing agents: diluted and concentrated nitric and concentrated sulfuric:

3Cu + 8HNO 3 (decomp) → 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 sprayed on plastic.

Zinc chemical properties

2. Possesses pronounced restorative properties and amphoteric properties.

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

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

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

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

8. With gray ZnS chalcogenide.

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

10. Reacts 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. Reacts actively with alkalis, since zinc is an amphoteric metal. Forms tetrahydroxozincates with alkali solutions and releases hydrogen:

• Zn + 2NaOH + 2H 2 O → Na 2 + H 2 .

On granules of zinc, after reaction, gas bubbles appear. With anhydrous alkalis, when fusion forms zincates and releases hydrogen:

• Zn + 2NaOH → Na 2 ZnO 2 + H 2.

Chemical properties of chromium

2.

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 (acidic chromium (VI) oxide CrO 3 and acids: chromic H 2 CrO 4 and two-chromic 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:

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

7. Reacts with water vapor, being in a red-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, however, it slowly reacts with their melts, forming chromates:

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

9. It dissolves in dilute strong acids, forming 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 t these acids passivate chromium. Reactions with acids when heated look like this:

2Сr + 6Н 2 SO 4 (conc) → Сr 2 (SO 4) 3 + 3SO 2 + 6Н 2 О

Cr + 6НNО 3 (conc) → Сr (NO 3) 3 + 3NO 2 + 3Н 2 О

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

Chemical properties:

• Possesses basic and regenerating properties.
• When heated to 100 0 С in air, it is oxidized to Cr 2 O 3 - chromium (III) oxide.
• It is possible to reduce 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, diluted sulfuric and nitric acids.

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

Chemical properties:

• It has amphoteric properties.
• How does the basic oxide react with acids: Cr 2 O 3 + 6HCl → CrCl 3 + 3H 2 O.
• How acidic oxide interacts with alkalis: Cr 2 O 3 + 2KON → 2KCrO 3 + H 2 O.
• Strong oxidants oxidize Cr 2 O 3 to chromate H 2 CrO 4.
• Strong reducing agents restoreCr out Cr 2 O 3.

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

Chemical properties:

• Weak base, showing basic properties.
• In the presence of moisture in the air, it is oxidized 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 that does not dissolve in water.

Chemical properties:

• It has amphoteric properties.
• How does the basic hydroxide react 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].

Iron chemical properties

1. A highly reactive active metal.

2. Possesses reducing properties, as well as pronounced magnetic properties.

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

• 3Fe + 2O 2 → Fe 3 O 4.

5. Reacts with halogens when heated:

• compound 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.

• compound 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.

9. It does not react with alkali solutions, but reacts slowly 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.

• 3Fe 2 O 3 + 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), which does not dissolve in water.

Chemical properties:

• Possesses 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) oxideFe 2 O 3- powder of a brown color (hematite), a substance insoluble in water. Other names: iron oxide, red lead, food coloring E172, etc.

Chemical properties:

• Fe 2 O 3 + 6HCl → 2 FeCl 3 + 3H 2 O.
• Does not react with alkali solutions, 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 that does not dissolve in water.

Chemical properties:

• Decomposes on heating more than 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 is oxidized: 4Fe 3 O 4 + O 2 → 6Fe 2 O 3.
• With hydrogen, when heated, it is reduced: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 reaction of neutralization 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 weakly expressed amphoteric properties with a predominance of the main ones.
• Reacts easily with acids: Fe (OH) 3 + 3HCl → FeCl 3 + 3H 2 O.
• Forms hexahydroxoferrates (III) with concentrated alkali solutions: Fe (OH) 3 + 3NaOH → Na 3.
• Forms ferrates with alkali melts:2Fe (OH) 3 + Na 2 CO 3 → 2NaFeO 2 + CO 2 + 3H 2 O.
• In an alkaline medium with strong oxidants, it exhibits reducing properties: 2Fe (OH) 3 + 3Br 2 + 10KOH → 2K 2 FeO 4 + 6NaBr + 8H 2 O.

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

As you can see from the data provided, all these items were opened in XIX century.

Discovery of metals of the main subgroup III group

Boron is a non-metal. Aluminum is a transition metal, while gallium, indium and thallium are high grade metals. Thus, with an increase in the radii of the atoms of the elements of each group of the periodic table, the metallic properties of simple substances increase.

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. Atomic structure, exhibited oxidation states.

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 - already a non-metal. Consequently, aluminum must exhibit properties of some intermediate character and its compounds are amphoteric.

Al +13) 2) 8) 3, p - element,

Aluminum exhibits an oxidation state of +3 in compounds:

Al 0 - 3 e - → Al +3

2. Physical properties

Free aluminum is a silvery-white metal with high thermal and electrical conductivity.The melting point is 650 o 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 strong 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 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 is found only in compounds (minerals).

Some of them:

· Bauxite - Al 2 O 3 H 2 O (with admixtures of SiO 2, Fe 2 O 3, CaCO 3)

· Nepheline - KNa 3 4

· Alunites - KAl (SO 4) 2 2Al (OH) 3

· Alumina (mixtures of kaolin 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 - 3ВеО Al 2 О 3 6SiO 2

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 look).

OXIDE FILM DEMONSTRATION

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, dipping into an alkali solution to remove the oxide film, and then into a solution of mercury salts to form an alloy of aluminum with mercury - amalgam).

I... Interaction with simple substances

Already at room temperature, aluminum 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:

2А l + 3 S = А l 2 S 3 (aluminum sulfide),

2А l + N 2 = 2А lN (aluminum nitride),

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

4А l + 3С = А l 4 C 3 (aluminum carbide).

2 Аl +3 I 2 = 2 A l I 3 (aluminum iodide) AN EXPERIENCE

All these compounds are completely hydrolyzed to form aluminum hydroxide and, accordingly, hydrogen sulfide, ammonia, phosphine and methane:

Al 2 S 3 + 6H 2 O = 2Al (OH) 3 + 3H 2 S

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

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

4А l + 3 O 2 = 2А l 2 О 3 + 1676 kJ.

COMBUSTION OF ALUMINUM IN AIR

AN EXPERIENCE

II... Interaction with complex substances

Interaction with water :

2 Al + 6 H 2 O = 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. Stands in the line of activity immediately after alkaline earth metals. That's why restores metals from their oxides ... Such a reaction - alumothermy - is used to obtain pure rare metals such as tungsten, vanadium, etc.

3 Fe 3 O 4 +8 Al = 4 Al 2 O 3 +9 Fe + Q

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

С r 2 О 3 + 2А l = 2С r + А l 2 О 3

5interactions with acids :

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

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 the evolution of hydrogen:

2А l + 6Н 2 S О 4 (conc) = А l 2 (S О 4) 3 + 3 S О 2 + 6Н 2 О,

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

Interaction with alkalis .

2 Al + 2 NaOH + 6 H 2 O = 2 Na [ Al (OH) 4 ] +3 H 2

AN EXPERIENCE

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:

2 Al + 3 CuSO 4 = Al 2 (SO 4) 3 + 3 Cu

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

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

Released mercury dissolves aluminum to form 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.: the plane is 2/3 composed of aluminum and its alloys. An airplane made of steel would be too heavy and could carry far fewer passengers. Therefore, aluminum is called a winged metal. Aluminum is used to make cables and wires: with the same electrical conductivity, their mass is 2 times less than the corresponding copper products.

Given the corrosion resistance of aluminum, manufacture parts for devices 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 with this paint they cover oil storage tanks, 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 treated to neutralize the resulting acid. As a result, a bulk precipitate of aluminum hydroxide is released, which, when settling, carries away suspended particles of turbidity and bacteria.

Thus, aluminum sulfate is a coagulant.

6. Obtaining aluminum

1) The modern cost-effective method of producing aluminum was invented by the American Hall and the Frenchman Eroux 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 like water dissolves sugar. The electrolysis of the “solution” of alumina in molten cryolite occurs as if cryolite was only a solvent, and alumina 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 that day, Charles Hall, a 22-year-old chemist, came to 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 make this metal cheaply and in large quantities. ” Thus Hall became the founder of the American aluminum industry and the Anglo-Saxon national hero, as a man who made a great business out of science.

2) 2Al 2 O 3 +3 C = 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 incandescent aluminum oxide 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 the potassium amalgam with pure potassium.
• In the 18th and 19th centuries, aluminum was the main jewelry metal. In 1889, D.I. Mendeleev in London for his merits in the development of chemistry was awarded a valuable gift - a balance made of gold and aluminum.
• By 1855, the French scientist Saint-Clair Deville had developed a method for producing metallic aluminum on a technical scale. But the method was very expensive. Deville enjoyed the special patronage of Napoleon III, Emperor of France. As a token of his devotion and gratitude, Deville made for Napoleon's son, the newborn prince, an exquisitely engraved rattle - the first "consumer goods" made of aluminum. Napoleon even intended to equip his guardsmen with aluminum cuirass, but the price turned out to be prohibitive. At that time, 1 kg of aluminum cost 1000 marks, i.e. 5 times more expensive than silver. Only after the invention of the electrolytic process did aluminum become equal in cost to conventional metals.
• Did you know that when aluminum enters the human body, it causes a disorder of the nervous system. With its excess, metabolism is disturbed. And the protective agents are vitamin C, calcium compounds, zinc.
• When aluminum burns in oxygen and fluorine, a lot of heat is generated. Therefore, it is used as an additive to rocket fuel. The Saturn rocket burns 36 tons of aluminum powder during the flight. The idea of ​​using metals as a component of rocket fuel was first expressed by F. A. Tsander.

SIMULATORS

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

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

Simulator No. 3 - Chemical properties of aluminum

TASKS FOR ANCHORING

# 1. To obtain aluminum from aluminum chloride, metallic calcium can be used as a reducing agent. Make an equation for a given chemical reaction, characterize this process using electronic balance.
Think! Why can't this reaction be carried out in aqueous solution?

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

No. 3. Make transformations:
Al -> AlCl 3 -> Al -> Al 2 S 3 -> Al (OH) 3 - t -> Al 2 O 3 -> Al

No. 4. Solve the problem:
The aluminum-copper alloy was exposed to an excess of concentrated sodium hydroxide solution when heated. Allocated 2.24 liters of gas (n.o.). Calculate the percentage of the alloy if its total weight was 10 g?