Aluminum is an amphoteric metal. The electronic configuration of the aluminum atom is 1s 2 2s 2 2p 6 3s 2 3p 1 . Thus, it has three valence electrons on the outer electron layer: 2 - on the 3s- and 1 - on the 3p-sublevel. In connection with this structure, it is characterized by reactions, as a result of which the aluminum atom loses three electrons from the external level and acquires an oxidation state of +3. Aluminum is a highly active metal and exhibits very strong reducing properties.

Interaction of aluminum with simple substances

with oxygen

Upon contact of absolutely pure aluminum with air, the aluminum atoms located in the surface layer instantly interact with the oxygen of the air and form the thinnest, several tens of atomic layers thick, strong oxide film of Al 2 O 3 composition, which protects aluminum from further oxidation. It is also impossible to oxidize large samples of aluminum even at very high temperatures. However, fine aluminum powder burns quite easily in a burner flame:

4Al + 3O 2 \u003d 2Al 2 O 3

with halogens

Aluminum reacts very vigorously with all halogens. Thus, the reaction between mixed powders of aluminum and iodine proceeds already at room temperature after adding a drop of water as a catalyst. The equation for the interaction of iodine with aluminum:

2Al + 3I 2 \u003d 2AlI 3

With bromine, which is a dark brown liquid, aluminum also reacts without heating. It is enough to simply introduce a sample of aluminum into liquid bromine: a violent reaction immediately begins with the release of a large amount of heat and light:

2Al + 3Br 2 = 2AlBr 3

The reaction between aluminum and chlorine proceeds when heated aluminum foil or fine aluminum powder is introduced into a flask filled with chlorine. Aluminum burns effectively in chlorine according to the equation:

2Al + 3Cl 2 = 2AlCl 3

with sulfur

When heated to 150-200 ° C or after igniting a mixture of powdered aluminum and sulfur, an intense exothermic reaction begins between them with the release of light:

— sulfide aluminum

with nitrogen

When aluminum interacts with nitrogen at a temperature of about 800 o C, aluminum nitride is formed:

with carbon

At a temperature of about 2000 o C, aluminum interacts with carbon and forms aluminum carbide (methanide), containing carbon in the -4 oxidation state, as in methane.

Interaction of aluminum with complex substances

with water

As mentioned above, a stable and durable oxide film of Al 2 O 3 does not allow aluminum to oxidize in air. The same protective oxide film makes aluminum inert to water as well. When removing the protective oxide film from the surface by methods such as treatment with aqueous solutions of alkali, ammonium chloride or mercury salts (amalgation), aluminum begins to react vigorously with water to form aluminum hydroxide and hydrogen gas:

with metal oxides

After ignition of a mixture of aluminum with oxides of less active metals(to the right of aluminum in the activity series) an extremely violent, strongly exothermic reaction begins. So, in the case of the interaction of aluminum with iron oxide (III), a temperature of 2500-3000 ° C develops. As a result of this reaction, high-purity molten iron is formed:

2AI + Fe 2 O 3 \u003d 2Fe + Al 2 O 3

This method of obtaining metals from their oxides by reduction with aluminum is called aluminothermy or aluminothermy.

with non-oxidizing acids

The interaction of aluminum with non-oxidizing acids, i.e. practically all acids, except concentrated sulfuric and nitric acids, leads to the formation of an aluminum salt of the corresponding acid and hydrogen gas:

a) 2Al + 3H 2 SO 4 (razb.) \u003d Al 2 (SO 4) 3 + 3H 2

2Al 0 + 6H + = 2Al 3+ + 3H 2 0;

b) 2AI + 6HCl = 2AICl 3 + 3H 2

with oxidizing acids

- concentrated sulfuric acid

The interaction of aluminum with concentrated sulfuric acid under normal conditions, as well as at low temperatures, does not occur due to an effect called passivation. When heated, the reaction is possible and leads to the formation of aluminum sulfate, water and hydrogen sulfide, which is formed as a result of the reduction of sulfur, which is part of sulfuric acid:

Such a deep reduction of sulfur from the oxidation state +6 (in H 2 SO 4) to the oxidation state -2 (in H 2 S) occurs due to the very high reducing ability of aluminum.

- concentrated nitric acid

Concentrated nitric acid also passivates aluminum under normal conditions, which makes it possible to store it in aluminum containers. Just as in the case of concentrated sulfuric, the interaction of aluminum with concentrated nitric acid becomes possible with strong heating, while the reaction proceeds predominantly:

- dilute nitric acid

The interaction of aluminum with dilute compared to concentrated nitric acid leads to products of a deeper reduction of nitrogen. Instead of NO, depending on the degree of dilution, N 2 O and NH 4 NO 3 can be formed:

8Al + 30HNO 3 (razb.) \u003d 8Al (NO 3) 3 + 3N 2 O + 15H 2 O

8Al + 30HNO 3 (highly diluted) = 8Al (NO 3) 3 + 3NH 4 NO 3 + 9H 2 O

with alkalis

Aluminum reacts both with aqueous solutions of alkalis:

2Al + 2NaOH + 6H 2 O = 2Na + 3H 2

and with pure alkalis during fusion:

In both cases, the reaction begins with the dissolution of the protective film of aluminum oxide:

Al 2 O 3 + 2NaOH + 3H 2 O \u003d 2Na

Al 2 O 3 + 2NaOH \u003d 2NaAlO 2 + H 2 O

In the case of an aqueous solution, aluminum, purified from the protective oxide film, begins to react with water according to the equation:

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

The resulting aluminum hydroxide, being amphoteric, reacts with an aqueous solution of sodium hydroxide to form soluble sodium tetrahydroxoaluminate:

Al(OH) 3 + NaOH = Na

Aluminum - destruction of metal under the influence of the environment.

For the reaction Al 3+ + 3e → Al, the standard electrode potential of aluminum is -1.66 V.

The melting point of aluminum is 660 °C.

The density of aluminum is 2.6989 g / cm 3 (under normal conditions).

Aluminum, although it is an active metal, has fairly good corrosion properties. This can be explained by the ability to be passivated in many aggressive environments.

The corrosion resistance of aluminum depends on many factors: the purity of the metal, the corrosive environment, the concentration of aggressive impurities in the environment, temperature, etc. The pH of solutions has a strong influence. Aluminum oxide on the metal surface is formed only in the pH range from 3 to 9!

Its purity greatly affects the corrosion resistance of Al. For the manufacture of chemical aggregates, equipment, only high-purity metal (without impurities) is used, for example, aluminum grades AB1 and AB2.

Corrosion of aluminum is not observed only in those environments where a protective oxide film is formed on the metal surface.

When heated, aluminum can react with some non-metals:

2Al + N 2 → 2AlN - interaction of aluminum and nitrogen with the formation of aluminum nitride;

4Al + 3С → Al 4 С 3 - reaction of interaction of aluminum with carbon with the formation of aluminum carbide;

2Al + 3S → Al 2 S 3 - the interaction of aluminum and sulfur with the formation of aluminum sulfide.

Corrosion of aluminum in air (atmospheric corrosion of aluminum)

Aluminum, when interacting with air, passes into a passive state. When pure metal comes into contact with air, a thin protective film of aluminum oxide instantly appears on the aluminum surface. Further, the growth of the film slows down. The formula of aluminum oxide is Al 2 O 3 or Al 2 O 3 H 2 O.

The reaction of interaction of aluminum with oxygen:

4Al + 3O 2 → 2Al 2 O 3 .

The thickness of this oxide film is between 5 and 100 nm (depending on operating conditions). Aluminum oxide has good adhesion to the surface, satisfies the condition of the continuity of oxide films. When stored in a warehouse, the thickness of aluminum oxide on the metal surface is about 0.01 - 0.02 microns. When interacting with dry oxygen - 0.02 - 0.04 microns. During heat treatment of aluminum, the thickness of the oxide film can reach 0.1 µm.

Aluminum is quite resistant both in clean rural air and in an industrial atmosphere (containing sulfur vapor, hydrogen sulfide, gaseous ammonia, dry hydrogen chloride, etc.). Because aluminum corrosion in gaseous media is not affected by sulfur compounds - it is used for the manufacture of sour oil processing plants, rubber vulcanization devices.

Corrosion of aluminum in water

Corrosion of aluminum is almost not observed when interacting with clean fresh, distilled water. Increasing the temperature to 180 °C has no particular effect. Hot water vapor also has no effect on aluminum corrosion. If a little alkali is added to water, even at room temperature, the rate of aluminum corrosion in such an environment will slightly increase.

The interaction of pure aluminum (not coated with an oxide film) with water can be described using the reaction equation:

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

When interacting with sea water, pure aluminum begins to corrode, because. sensitive to dissolved salts. To exploit aluminum in sea water, a small amount of magnesium and silicon is introduced into its composition. Corrosion resistance of aluminum and its alloys, when exposed to sea ​​water, is significantly reduced if copper is included in the composition of the metal.

Corrosion of aluminum in acids

As the purity of aluminum increases, its resistance to acids increases.

Corrosion of aluminum in sulfuric acid

For aluminum and its alloys is very dangerous sulfuric acid(possesses oxidizing properties) medium concentrations. The reaction with dilute sulfuric acid is described by the equation:

2Al + 3H 2 SO 4 (razb) → Al 2 (SO 4) 3 + 3H 2.

Concentrated cold sulfuric acid has no effect. And when heated, aluminum corrodes:

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

This forms a soluble salt - aluminum sulfate.

Al is stable in oleum (fuming sulfuric acid) at temperatures up to 200 °C. Due to this, it is used for the production of chlorosulfonic acid (HSO 3 Cl) and oleum.

Corrosion of aluminum in hydrochloric acid

In hydrochloric acid, aluminum or its alloys quickly dissolve (especially with increasing temperature). Corrosion equation:

2Al + 6HCl → 2AlCl 3 + 3H 2 .

Solutions of hydrobromic (HBr), hydrofluoric (HF) acids act similarly.

Corrosion of aluminum in nitric acid

A concentrated solution of nitric acid has high oxidizing properties. Aluminum in nitric acid at normal temperature is exceptionally stable (higher resistance than stainless steel 12X18H9). It is even used to produce concentrated nitric acid by direct synthesis.

When heated, the corrosion of aluminum in nitric acid proceeds according to the reaction:

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

Corrosion of aluminum in acetic acid

Aluminum has a sufficiently high resistance to acetic acid of any concentration, but only if the temperature does not exceed 65 ° C. It is used for the production of formaldehyde and acetic acid. At higher temperatures, aluminum dissolves (with the exception of acid concentrations of 98 - 99.8%).

In bromine, weak solutions of chromic (up to 10%), phosphoric (up to 1%) acids at room temperature, aluminum is stable.

Citric, butyric, malic, tartaric, propionic acids, wine, fruit juices have a weak effect on aluminum and its alloys.

Oxalic, formic, organochlorine acids destroy the metal.

The corrosion resistance of aluminum is greatly affected by vaporous and droplet-liquid mercury. After a short contact, the metal and its alloys corrode intensively, forming amalgams.

Corrosion of aluminum in alkalis

Alkalis easily dissolve the protective oxide film on the surface of aluminum, it begins to react with water, as a result of which the metal dissolves with the release of hydrogen (corrosion of aluminum with hydrogen depolarization).

2Al + 2NaOH + 6H 2 O → 2Na + 3H 2;

2(NaOH H 2 O) + 2Al → 2NaAlO 2 + 3H 2.

aluminates are formed.

Also, the oxide film is destroyed by salts of mercury, copper and chloride ions.

1) Silicon was burned in an atmosphere of chlorine. The resulting chloride was treated with water. The precipitate thus formed was calcined. Then it was fused with calcium phosphate and coal. Write the equations for the four described reactions.

3) Some amount of iron(II) sulfide was divided into two parts. One of them was treated with hydrochloric acid, and the other was fired in air. During the interaction of the evolved gases, a simple yellow substance was formed. The resulting substance was heated with concentrated nitric acid, and a brown gas was released. Write the equations for the four described reactions.

4) When aluminum oxide reacted with nitric acid, a salt was formed. The salt was dried and calcined. The solid residue formed during calcination was subjected to electrolysis in molten cryolite. The metal obtained by electrolysis was heated with a concentrated solution containing potassium nitrate and potassium hydroxide, and a gas with a pungent odor was released. Write the equations for the four described reactions.

5) Chromium(VI) oxide reacted with potassium hydroxide. The resulting substance was treated with sulfuric acid, an orange salt was isolated from the resulting solution. This salt was treated with hydrobromic acid. The resulting simple substance reacted with hydrogen sulfide. Write the equations for the four described reactions.

6) The magnesium powder was heated under a nitrogen atmosphere. When the resulting substance interacts with water, a gas is released. The gas was passed through an aqueous solution of chromium(III) sulfate, resulting in a gray precipitate. The precipitate was separated and treated with heating with a solution containing hydrogen peroxide and potassium hydroxide. Write the equations for the four described reactions.

7) Ammonia was passed through hydrobromic acid. Silver nitrate solution was added to the resulting solution. The precipitate formed was separated and heated with zinc powder. The metal formed during the reaction was treated with a concentrated solution of sulfuric acid, and a gas with a pungent odor was released. Write the equations for the four described reactions.

8) Potassium chlorate was heated in the presence of a catalyst, and a colorless gas was released. By burning iron in an atmosphere of this gas, iron scale was obtained. It was dissolved in an excess of hydrochloric acid. To the solution thus obtained was added a solution containing sodium dichromate and hydrochloric acid. Write the equations for the four described reactions.

9) Sodium was heated in a hydrogen atmosphere. When water was added to the resulting substance, gas evolution and the formation of a clear solution were observed. A brown gas was passed through this solution, which was obtained as a result of the interaction of copper with a concentrated solution of nitric acid. Write the equations for the four described reactions.

10) Aluminum reacted with sodium hydroxide solution. The evolved gas was passed over heated copper(II) oxide powder. The resulting simple substance was dissolved by heating in concentrated sulfuric acid. The resulting salt was isolated and added to a solution of potassium iodide. Write the equations for the four described reactions.

11) Spent the electrolysis of a solution of sodium chloride. Iron(III) chloride was added to the resulting solution. The precipitate that formed was filtered off and calcined. The solid residue was dissolved in hydroiodic acid. Write the equations for the four described reactions.

12) Aluminum powder was added to the sodium hydroxide solution. An excess of carbon dioxide was passed through the solution of the obtained substance. The precipitate that formed was separated and calcined. The resulting product was fused with sodium carbonate. Write the equations for the four described reactions.

The chemical properties of aluminum are determined by its position in the periodic table of chemical elements.

Below are the main chemical reactions of aluminum with other chemical elements. These reactions determine the main Chemical properties aluminum.

What does aluminum react with?

Simple substances:

• halogens (fluorine, chlorine, bromine and iodine)
• phosphorus
• carbon
• oxygen (combustion)

Compound substances:

• mineral acids (hydrochloric, phosphoric)
• sulfuric acid
• Nitric acid
• alkalis
• oxidizers
• oxides of less active metals (aluminothermy)

What does aluminum not react with?

Aluminum does not react:

• with hydrogen
• under normal conditions - with concentrated sulfuric acid (due to passivation - the formation of a dense oxide film)
• under normal conditions - with concentrated nitric acid (also due to passivation)

aluminum and air

Usually the surface of aluminum is always covered with a thin layer of aluminum oxide, which protects it from exposure to air, more precisely, oxygen. Therefore, it is believed that aluminum does not react with air. If this oxide layer is damaged or removed, then the fresh aluminum surface reacts with atmospheric oxygen. Aluminum can burn in oxygen with a dazzling white flame to form aluminum oxide Al2O3.

The reaction of aluminum with oxygen:

• 4Al + 3O 2 -> 2Al 2 O 3

aluminum and water

Aluminum reacts with water in the following reactions:

• 2Al + 6H 2 O = 2Al(OH) 3 + 3H 2 (1)
• 2Al + 4H 2 O \u003d 2AlO (OH) + 3H 2 (2)
• 2Al + 3H 2 O \u003d Al 2 O 3 + 3H 2 (3)

As a result of these reactions, the following are formed, respectively:

• aluminum hydroxide modification bayerite and hydrogen (1)
• modification of aluminum hydroxide bohemite and hydrogen (2)
• aluminum oxide and hydrogen (3)

These reactions, by the way, are of great interest in the development of compact plants for the production of hydrogen for vehicles that run on hydrogen.

All these reactions are thermodynamically possible at temperatures from room temperature to the melting point of aluminum 660 ºС. All of them are also exothermic, that is, they occur with the release of heat:

• At temperatures from room temperature to 280 ºС, the most stable reaction product is Al(OH) 3 .
• At temperatures from 280 to 480 ºС, AlO(OH) is the most stable reaction product.
• At temperatures above 480 ºС, the most stable reaction product is Al 2 O 3 .

Thus, aluminum oxide Al 2 O 3 becomes thermodynamically more stable than Al(OH) 3 at elevated temperature. The reaction product of aluminum with water at room temperature will be aluminum hydroxide Al(OH) 3 .

Reaction (1) shows that aluminum should spontaneously react with water at room temperature. However, in practice, a piece of aluminum immersed in water does not react with water at room temperature and even in boiling water. The fact is that aluminum has a thin coherent layer of aluminum oxide Al 2 O 3 on the surface. This oxide film adheres strongly to the aluminum surface and prevents it from reacting with water. Therefore, in order to start and maintain the reaction of aluminum with water at room temperature, it is necessary to constantly remove or destroy this oxide layer.

Aluminum and halogens

Aluminum reacts violently with all halogens - these are:

• fluorine F
• chlorine Cl
• bromine Br and
• iodine (iodine) I,

with education respectively:

• fluoride AlF 3
• chloride AlCl 3
• bromide Al 2 Br 6 and
• iodide Al 2 Br 6.

Reactions of hydrogen with fluorine, chlorine, bromine and iodine:

• 2Al + 3F 2 → 2AlF 3
• 2Al + 3Cl 2 → 2AlCl 3
• 2Al + 3Br 2 → Al 2 Br 6
• 2Al + 3l 2 → Al 2 I 6

aluminum and acids

Aluminum actively reacts with dilute acids: sulfuric, hydrochloric and nitric, with the formation of the corresponding salts: aluminum sulfate Al 2 SO 4 , aluminum chloride AlCl 3 and aluminum nitrate Al(NO 3) 3 .

Reactions of aluminum with dilute acids:

• 2Al + 3H 2 SO 4 -> Al 2 (SO 4) 3 + 3H 2
• 2Al + 6HCl -> 2AlCl 3 + 3H 2
• 2Al + 6HNO 3 -> 2Al(NO 3) 3 + 3H 2

It does not interact with concentrated sulfuric and hydrochloric acids at room temperature; when heated, it reacts with the formation of salt, oxides and water.

Aluminum and alkalis

Aluminum in an aqueous solution of alkali - sodium hydroxide - reacts to form sodium aluminate.

The reaction of aluminum with sodium hydroxide is:

• 2Al + 2NaOH + 10H 2 O -> 2Na + 3H 2

Sources:

1. Chemical Elements. The first 118 elements, ordered alphabetically / ed. Wikipedians-2018

2. Reaction of Aluminum with Water to Produce Hydrogen / John Petrovic and George Thomas, U.S. Department of Energy, 2008