Aluminum was first produced only at the beginning of the 19th century. This was done by physicist Hans Oersted. He conducted his experiment with potassium amalgam, aluminum chloride and.

By the way, the name of this silvery material comes from the Latin word “alum”, because it is from them that this element is mined.

Alum is a natural metal-based mineral that combines sulfuric acid salts in its composition.

Previously, it was considered a precious metal and was worth an order of magnitude more expensive than gold. This was explained by the fact that the metal was quite difficult to separate from impurities. So only rich and influential people could afford aluminum jewelry.

But in 1886, Charles Hall came up with a method for extracting aluminum on an industrial scale, which dramatically reduced the cost of this metal and made it possible to use it in metallurgical production. The industrial method involved the electrolysis of molten cryolite in which aluminum oxide was dissolved.

Aluminum is a very popular metal, because many things that people use in everyday life are made from it.

Application of aluminum

Due to its malleability and lightness, as well as its resistance to corrosion, aluminum is a valuable metal in modern industry. Not only kitchenware is made from aluminum - it is widely used in automobile and aircraft construction.

Aluminum is also one of the most inexpensive and economical materials, as it can be used endlessly by melting down unwanted aluminum items, such as cans.

Aluminum metal is safe, but its compounds can be toxic to humans and animals (especially aluminum chloride, acetate and sulfate).

Physical properties of aluminum

Aluminum is a fairly light, silver-colored metal that can form alloys with most metals, especially copper and silicon. It is also very plastic; it can easily be turned into a thin plate or foil. The melting point of aluminum = 660 °C and the boiling point is 2470 °C.

Chemical properties of aluminum

At room temperature, the metal is coated with a durable film of aluminum oxide Al₂O₃, which protects it from corrosion.

Aluminum practically does not react with oxidizing agents due to the oxide film that protects it. However, it can be easily destroyed so that the metal exhibits active restorative properties. The aluminum oxide film can be destroyed with a solution or melt of alkalis, acids, or with the help of mercury chloride.

Due to its reducing properties, aluminum has found application in industry for the production of other metals. This process is called aluminothermy. This feature of aluminum is its interaction with oxides of other metals.

For example, consider the reaction with chromium oxide:

Cr₂O₃ + Al = Al₂O₃ + Cr.

Aluminum reacts well with simple substances. For example, with halogens (except fluorine), aluminum can form aluminum iodide, chloride, or bromide:

2Al + 3Cl₂ → 2AlCl₃

With other non-metals such as fluorine, sulfur, nitrogen, carbon, etc. aluminum can only react when heated.

Silver metal also reacts with complex chemicals. For example, with alkalis it forms aluminates, that is, complex compounds that are actively used in the paper and textile industries. Moreover, it reacts as aluminum hydroxide

Al(OH)₃ + NaOH = Na),

and metallic aluminum or aluminum oxide:

2Al + 2NaOH + 6Н₂О = 2Na + ЗН₂.

Al₂O₃ + 2NaOH + 3H₂O = 2Na

Aluminum reacts quite calmly with aggressive acids (for example, sulfuric and hydrochloric acids), without ignition.

If you dip a piece of metal into hydrochloric acid, the reaction will be slow - the oxide film will dissolve at first - but then it will speed up. Aluminum is dissolved in hydrochloric acid to release mercury for two minutes, and then rinse it well. The result is an amalgam, an alloy of mercury and aluminum:

3HgCI₂ + 2Al = 2AlCI₃ + 3Hg

Moreover, it does not stick to the metal surface. Now, by immersing the purified metal in water, you can observe a slow reaction, which is accompanied by the release of hydrogen and the formation of aluminum hydroxide:

2Al + 6H₂O = 2Al(OH)₃ + 3H₂.

When determining the concentration of impurities in drinking and natural water, attention is paid to the volume of nitrates, sulfates, nitrites, chlorides, forgetting about aluminum - the most common metal in nature. Under normal conditions, aluminum dissolves in water to form various compounds that actively react with other impurities. As a result, the substance is saturated with aluminum hydrochloride, salts and other compounds. And this leads to a change in water quality - a deterioration in the chemical composition, organoleptic properties, microbiological, bacterial indicators.
The official MPC of aluminum in water for drinking and natural reservoirs is calculated by WHO and environmental organizations. But this parameter does not take into account the numerous ways metal enters natural sources and the human body. Therefore, accurate determination of aluminum in water is important.

Aluminum in natural bodies of water

Natural saturation of water with metal occurs due to the ingress of aluminosilicates and some types of clay into it. After their dissolution, the interaction of aluminum with water begins, directly depending on its pH. Dissolution under natural conditions occurs slowly, but always with the release of hydroxide, bauxite, hydrochloride and other compounds. The substances and aluminum itself are contained in both sea water and river water. But this is under normal conditions.

Metal enters natural waters from:

• drains of industrial and domestic waters;
• wastewater from chemical production (any production increases the concentration of aluminum in wastewater by 2-5 times);
• construction waste and emissions.

Every year there are more and more such emissions into the environment, and control over the degree of their pollution is becoming less and less. In dirty wastewater with a high content of impurities and suspended matter, aluminum dissolves faster in water. It enters water bodies in the form of suspended forms, ions and colloids. It is the ions and oxides that have increased toxicity. They have a detrimental effect on most living organisms living in natural sources. According to standards, the concentration of aluminum in natural waters should not exceed 0.5 mg/dm3.

Aluminum in drinking water

The most common metal on the planet will certainly be contained in drinking water. According to the standards and requirements of GOST, aluminum in water must contain:

• no more than 0.5 mg/l in tap water;
• within 0.2-0.3 mg/l in bottled water;
• within 0.1-0.2 mg/l in filtered water.

Every day, the human body should receive no more than 90 mg of metal. But after the reaction of aluminum with water is completed, toxic impurities appear in it. Therefore, tap water, as well as well and borehole substances, should be checked for the concentration of unsafe impurities and components. Below is a table of maximum concentration limits for aluminum in drinking water and other substances important for human health.

Why should you drink water with a minimum concentration of aluminum?

Having found out where aluminum appears in water, it is worthwhile to look at the other ways it enters the body. This will help control the daily metal requirement. The bulk of the chemical element comes from food.
The metal is also contained in:

• cosmetic preparations;
• utensils made of the same metal;
• medicines;
• deodorants, etc.

With the standard aluminum content in water there will be no effect on the body. With excessive concentration, the nervous system suffers, memory decreases, depression and irritability appear. The consequences do not occur immediately. This is due to the fact that not the entire volume of metal is absorbed by the body. Scientists have also proven that a high aluminum content in water leads to neurological diseases and disruption of calcium-phosphorus metabolism, which inhibits the production of hemoglobin. Therefore, it is recommended to use a drinking substance with a metal volume of no more than 0.3 mg/l. With this content of dissolved aluminum in water, daily consumption will not exceed 50 mg/l. Household filter systems are used for cleaning.

Water purification using coagulation method

In order for a liquid suitable for drinking or technical needs to flow from the taps, it must first be cleaned. Both groundwater and surface water must undergo this procedure before any use. It is described above what happens when aluminum interacts with water - an unpleasant odor, unwanted impurities are formed, the substance becomes cloudy, and a sediment appears. By deteriorating the organoleptic qualities of the liquid, some metal compounds can act as excellent coagulants - elements that bind dangerous and unnecessary particles in the substance. They are effectively used to improve the quality of liquids in water treatment systems.

Aluminum sulfate is most often used to purify water for any need. The coagulum is most active in an environment with an acidity of 4.4-6.1 pH. But they also apply to substances with a pH from 7 to 8. The water treatment procedure is as follows:

• adding aluminum sulfate to the liquid;
• mixing media - complete mixing occurs within 1-3 minutes;
• coagulation, in which the medium passes from one reservoir to another (the process lasts from 30 minutes to 1 hour);
• settling of bound sediment;
• filtration of the purified medium.

At the moment, water purification with aluminum is an affordable and effective way to remove suspended particles from liquids. During coagulation, the removal of sodium and calcium bicarbonates and carbonates is also observed. Upon completion of the water treatment procedure, the consumer receives clean and pleasant-smelling water.

The main raw material for aluminum production is bauxite containing alumina. The most important aluminum ores also include alunite and nepheline.

The USSR has reserves of aluminum. In addition to bauxite, deposits of which we have in the Urals, in the Bashkir Autonomous Soviet Socialist Republic and in Kazakhstan, the richest source of aluminum is nepheline, which occurs together with apatite in the Khibiny Mountains. Significant deposits of aluminum raw materials are available in Siberia.

Aluminum was first obtained by Wöhler in 1827 by the action of potassium metal on aluminum chloride. However, despite its widespread occurrence in nature, aluminum was one of the rare metals until the end of the 19th century.

Currently, aluminum is produced in huge quantities from aluminum oxide by the electrolytic method. The aluminum oxide used for this must be sufficiently pure, since impurities are difficult to remove from smelted aluminum. Purified bauxite is obtained by processing natural bauxite.

Producing aluminum is a complex process fraught with great difficulties. The main starting material - aluminum oxide - does not conduct electricity and has a very high melting point (about 2050). Therefore, a molten mixture of cryolite and aluminum oxide is subjected to electrolysis.

A mixture containing about (mass.) melts at and has electrical conductivity, density and viscosity that are most favorable for the process. To further improve these characteristics, additives and are added to the mixture. Thanks to this, electrolysis is possible at .

The electrolyzer for smelting aluminum is an iron casing lined with refractory bricks on the inside. Its bottom (under), assembled from blocks of compressed coal, serves as a cathode. Anodes (one or more) are located on top: these are aluminum frames filled with coal briquettes. In modern factories, electrolyzers are installed in series; each series consists of 150 or more electrolysers.

During electrolysis, aluminum is released at the cathode and oxygen at the anode. Aluminum, which has a higher density than the original melt, is collected in an electrolyzer; from here he is periodically released. As the metal is released, new portions of aluminum oxide are added to the melt. The oxygen released during electrolysis interacts with the carbon of the anode, which burns out, forming CO and.

Aluminum was not produced in pre-revolutionary Russia. The first aluminum smelter in the USSR (Volkhovsky) came into operation in 1932, and already in 1935 our country took third place in the world in aluminum production.

The identical structure of the outer electronic layer of the boron and aluminum atom determines the similarity in the properties of these elements. Thus, aluminum, like boron, is characterized only by its oxidation state. However, when going from boron to aluminum, the radius of the atom increases greatly (from 0.091 to ) and, in addition, another intermediate eight-electron layer appears, shielding the nucleus. All this leads to a weakening of the connection between the outer electrons and the nucleus and to a decrease in the ionization energy of the atom (see Table 35). Therefore, aluminum has much more pronounced metallic properties than boron. However, the chemical bonds formed by aluminum with other elements are primarily covalent in nature.

Another feature of aluminum (as well as its analogues - gallium, indium and thallium) compared to boron is the existence of free -sublevels in the outer electronic layer of its atom. Due to this, the coordination number of aluminum in its compounds can be not only four, like boron, but also six.

Rice. 165. Scheme of the spatial structure of a molecule: black circles are aluminum atoms, light circles are chlorine atoms.

Aluminum is connected like similar boron compounds; in individual molecules of such compounds there are only six electrons in the outer electron layer of the aluminum atom. Therefore, here the aluminum atom is capable of being an acceptor of electron pairs. In particular, aluminum halides are characterized by the formation of dimers, carried out according to the donor-acceptor method (in diagram D - halogen atom):

As can be seen, such dimeric molecules contain two “bridge” halogen atoms. The spatial structure is shown in Fig. 165. Aluminum halides exist in the form of dimeric molecules in melts and vapors. However, traditionally their composition is usually expressed in the form . Below we will also adhere to this method of writing formulas for aluminum halides.

Aluminum hydride is also an electron-deficient compound. However, the hydrogen atom, unlike halogen atoms in molecules, does not have a lone electron pair and cannot play the role of an electron donor. Therefore, here individual molecules are connected to each other through “bridging” hydrogen atoms by three-center bonds, similar to the bonds in borohydride molecules (see p. 612). As a result, a solid polymer is formed, the composition of which can be expressed by the formula.

Aluminum is a silvery-white light metal. It is easily drawn into wire and rolled into thin sheets.

At room temperature, aluminum does not change in air, but only because its surface is covered with a thin film of oxide, which has a very strong protective effect. The destruction of this film, for example, by amalgamating aluminum, causes rapid oxidation of the metal, accompanied by noticeable heating.

The standard electrode potential of aluminum is -1.663 V. Despite its negative value, aluminum, due to the formation of a protective oxide film on its surface, does not displace hydrogen from water. However, amalgamated aluminum, which does not form a dense oxide layer, reacts vigorously with water to release hydrogen.

Dilute hydrochloric and sulfuric acids easily dissolve aluminum, especially when heated. Highly diluted and cold concentrated nitric acid does not dissolve aluminum.

When aqueous solutions of alkalis act on aluminum, the oxide layer dissolves, and aluminates are formed - salts containing aluminum as part of the anion:

sodium tetrahydroxyaluminate

Aluminum, devoid of a protective film, interacts with water, displacing hydrogen from it:

The resulting aluminum hydroxide reacts with excess alkali, forming hydroxoaluminate:

By doubling the last equation and adding it to the previous one, we obtain the total equation for the dissolution of aluminum in an aqueous alkali solution:

Aluminum dissolves noticeably in solutions of salts that, due to their hydrolysis, have an acidic or alkaline reaction, for example, in a solution.

If aluminum powder (or thin aluminum foil) is heated strongly, it ignites and burns with a blinding white flame, forming aluminum oxide.

The main use of aluminum is the production of alloys based on it. Alloying additives (for example, copper, silicon, magnesium, zinc, manganese) are added to aluminum mainly to increase its strength. Dura homins containing copper and magnesium, silumins, in which the main additive is silicon, and magnalium (an alloy of aluminum and magnesium) are widespread. The main advantages of all aluminum alloys are their low density, high strength (per unit mass), satisfactory resistance to atmospheric corrosion, comparative cheapness and ease of production and processing. Aluminum alloys are used in rocketry, aircraft, auto, shipbuilding and instrument making, in the production of tableware and in many other industries. Aluminum alloys occupy second place in terms of the breadth of application after steel and cast iron.

Aluminum is one of the most common additives in alloys based on copper, magnesium, titanium, nickel, zinc, and iron.

In the form of pure metal, aluminum is used for the manufacture of chemical equipment, electrical wires, and capacitors. Although the electrical conductivity of aluminum is less than that of copper (about the electrical conductivity of copper), this is compensated by the lightness of aluminum, which allows the wires to be made thicker: with the same electrical conductivity, an aluminum wire weighs half as much as a copper wire.

It is important to use aluminum for aluminizing, which consists of saturating the surface of steel or cast iron products with aluminum in order to protect the base material from oxidation under high heat. In metallurgy, aluminum is used to produce calcium, barium, lithium and some other metals by aluminothermy (see § 192).

Aluminum oxide, also called alumina, occurs naturally in crystalline form, forming the mineral corundum. Corundum has very high hardness. Its transparent crystals, colored red or blue by impurities, are the precious stones ruby ​​and sapphire. Rubies are now produced artificially by melting alumina in an electric furnace. They are used not so much for jewelry as for technical purposes, for example, for the manufacture of parts for precision instruments, stones in watches, etc. Ruby crystals containing a small impurity are used as quantum generators - lasers that create a directed beam of monochromatic radiation.

Corundum and its fine-grained variety containing a large amount of impurities, emery, are used as abrasive materials.

Aluminum hydroxide precipitates in the form of a gelatinous precipitate under the action of alkalis on solutions of aluminum salts and easily forms colloidal solutions.

Aluminum hydroxide is a typical amphoteric hydroxide. With acids it forms salts containing aluminum cation, with alkalis - aluminates. When aluminum hydroxide reacts with aqueous solutions of alkalis or when metallic aluminum is dissolved in alkali solutions, hydroxoaluminates are formed, as mentioned above, for example. When aluminum oxide is fused with the corresponding oxides or hydroxides, meta-aluminum acid derivatives are obtained, for example:

Both aluminum salts and aluminates in solutions are highly hydrolyzed. Therefore, aluminum salts and weak acids in solutions are converted into basic salts or undergo complete hydrolysis. For example, when an aluminum salt reacts in a solution with aluminum, it is not aluminum carbonate that is formed, but its hydroxide, and carbon dioxide is released:

Aluminum chloride. Anhydrous aluminum chloride is obtained by direct interaction of chlorine with aluminum. It is widely used as a catalyst in various organic syntheses.

It dissolves in water, releasing a large amount of heat. When the solution is evaporated, hydrolysis occurs, hydrogen chloride is released and aluminum hydroxide is obtained. If evaporation is carried out in the presence of excess hydrochloric acid, then crystals of the composition can be obtained.

As already indicated on page 614, the chemical bonds formed by the aluminum atom are predominantly covalent in nature. This affects the properties of the compounds it forms. Thus, at normal atmospheric pressure, anhydrous aluminum chloride already sublimates at, and at high pressures it melts at, and in the molten state it does not conduct electric current. Therefore, the melt cannot be used for the electrolytic production of aluminum.

Aluminum sulfate is obtained by the action of hot sulfuric acid on aluminum oxide or kaolin. Used for water purification (see page 598), as well as in the preparation of certain types of paper.

Potassium alum is used in large quantities for tanning, as well as in dyeing as a mordant for cotton fabrics. In the latter case, the effect of alum is based on the fact that aluminum hydroxide formed as a result of its hydrolysis is deposited in the fabric fibers in a finely dispersed state and, adsorbing the dye, firmly holds it on the fiber.

Aluminum is an element with serial number 13, relative atomic mass - 26.98154. Located in period III, group III, main subgroup. Electronic configuration: 1s 2 2s 2 2p 6 3s 2 3p 1 3d 0 . The stable oxidation state of aluminum is “+3”. The resulting cation has a noble gas shell, which contributes to its stability, but the ratio of charge to radius, that is, the charge concentration, is quite high, which increases the energy of the cation. This feature leads to the fact that aluminum, along with ionic compounds, forms a number of covalent compounds, and its cation undergoes significant hydrolysis in solution.

Aluminum can exhibit valency I only at temperatures above 1500 o C. Al 2 O and AlCl are known.

In terms of physical properties, aluminum is a typical metal, with high thermal and electrical conductivity, second only to silver and copper. The ionization potential of aluminum is not very high, so one would expect high chemical activity from it, but it is significantly reduced due to the fact that the metal is passivated in air due to the formation of a strong oxide film on its surface. If the metal is activated: a) mechanically remove the film, b) amalgamate (react with mercury), c) use powder, then such a metal becomes so reactive that it even interacts with moisture and oxygen in the air, collapsing in accordance with the process:

4(Al,Hg) +3O 2 + 6H 2 O = 4Al(OH) 3 + (Hg)

Interaction with simple substances.

1. Powdered aluminum reacts when heated strongly with oxygen. These conditions are necessary due to passivation, and the reaction of formation of aluminum oxide itself is highly exothermic - 1676 kJ/mol of heat is released.

2. With chlorine and bromine reacts under standard conditions and can even ignite in their environment. Only does not respond with fluorine, because Aluminum fluoride, like oxide, forms a protective salt film on the metal surface. With iodine reacts when heated and in the presence of water as a catalyst.

3. With sulfur reacts upon fusion, giving aluminum sulfide of the composition Al 2 S 3.

4. It also reacts with phosphorus when heated to form phosphide: AlP.

5. Directly with hydrogen aluminum does not react.

6. With nitrogen reacts at 800 o C, giving aluminum nitride (AlN). It should be said that the combustion of aluminum in air occurs at approximately the same temperatures, so the combustion products (taking into account the composition of the air) are both oxide and nitride.

7. With carbon aluminum interacts at an even higher temperature: 2000 o C. Aluminum carbide of the composition Al 4 C 3 belongs to the methanides, it does not contain C-C bonds, and during hydrolysis methane is released: Al 4 C 3 + 12H 2 O = 4Al(OH ) 3 + 3CH 4

Interaction with complex substances

1. With water activated (devoid of a protective film) aluminum actively interacts with the release of hydrogen: 2Al (act.) + 6H 2 O = 2Al(OH) 3 + 3H 2 Aluminum hydroxide is obtained in the form of a white loose powder; the absence of a film does not interfere with the completion of the reaction.

2. Interaction with acids: a) Aluminum actively interacts with non-oxidizing acids in accordance with the equation: 2Al + 6H 3 O + + 6H 2 O = 2 3+ + 3H 2,

b) Interaction with oxidizing acids occurs with the following features. Concentrated nitric and sulfuric acids, as well as very dilute nitric acid, passivate aluminum (rapid oxidation of the surface leads to the formation of an oxide film) in the cold. When heated, the film is disrupted and the reaction takes place, but only the products of their minimal reduction are released from concentrated acids when heated: 2Al + 6H 2 SO 4 (conc) = Al 2 (SO 4) 3 + 3SO 2 6H 2 O Al + 6HNO 3 ( conc) = Al(NO 3) 3 + 3NO 2 + 3H 2 O With moderately dilute nitric acid, depending on the reaction conditions, you can get NO, N 2 O, N 2, NH 4 +.

3. Interaction with alkalis. Aluminum is an amphoteric element (in terms of chemical properties), because has a fairly high electronegativity for metals - 1.61. Therefore, it dissolves quite easily in alkali solutions with the formation of hydroxo complexes and hydrogen. The composition of the hydroxo complex depends on the ratio of the reagents: 2Al + 2NaOH + 6H 2 O = 2Na + 3H 2 2Al + 6NaOH + 6H 2 O = 2Na 3 + 3H 2 The ratio of aluminum and hydrogen is determined by the electronic balance of the redox reaction occurring between them and on the ratio of the reagents does not depend.

4. Low ionization potential and high affinity for oxygen (high oxide stability) lead to the fact that aluminum actively interacts with oxides of many metals, restoring them. The reactions take place during initial heating with further release of heat, so that the temperature rises to 1200 o - 3000 o C. A mixture of 75% aluminum powder and 25% (by weight) Fe 3 O 4 is called “thermite”. Previously, the combustion reaction of this mixture was used to weld rails. The reduction of metals from oxides using aluminum is called aluminothermy and is used in industry as a method for producing metals such as manganese, chromium, vanadium, tungsten, and ferroalloys.

5. With salt solutions aluminum reacts in two different ways. 1. If, as a result of hydrolysis, the salt solution has an acidic or alkaline environment, hydrogen is released (with acidic solutions, the reaction occurs only with significant heating, since the protective oxide film dissolves better in alkalis than in acids). 2Al + 6KHSO 4 + (H 2 O) = Al 2 (SO 4) 3 + 3K 2 SO 4 + 3H 2 2Al + 2K 2 CO 3 + 8H 2 O = 2K + 2KHCO 3 + 3H 2. 2. Aluminum can displace from the salt composition metals that are in the voltage series to the right of it, i.e. will actually be oxidized by cations of these metals. Due to the oxide film, this reaction does not always take place. For example, chloride anions can disrupt the film, and the reaction 2Al + 3FeCl 2 = 2AlCl 3 + 3Fe takes place, but a similar reaction with sulfates at room temperature will not work. With activated aluminum, any interaction that does not contradict the general rule will work.

Aluminum connections.

1. Oxide (Al 2 O 3). Known in the form of several modifications, most of which are very durable and chemically inert. The modification α-Al 2 O 3 occurs in nature in the form of the mineral corundum. In the crystal lattice of this compound, aluminum cations are sometimes partially replaced by cations of other metals, which gives the mineral its color. The admixture of Cr(III) gives a red color, such corundum is already a ruby ​​gemstone. The admixture of Ti(III) and Fe(III) produces blue sapphire. The amorphous modification is chemically active. Aluminum oxide is a typical amphoteric oxide, reacting both with acids and acidic oxides, and with alkalis and basic oxides, with alkalis being preferable. The reaction products in solution and in the solid phase during fusion are different: Na 2 O + Al 2 O 3 = 2NaAlO 2 (fusion) - sodium metaaluminate, 6NaOH + Al 2 O 3 = 2Na 3 AlO 3 + 3H 2 O (fusion) - orthoaluminate sodium, Al 2 O 3 + 3CrO 3 = Al 2 (CrO 4) 3 (fusion) - aluminum chromate. In addition to oxides and solid alkalis, aluminum during fusion reacts with salts formed by volatile acid oxides, displacing them from the salt composition: K 2 CO 3 + Al 2 O 3 = 2KAlO 2 + CO 2 Reactions in solution: Al 2 O 3 + 6HCl = 2 3+ + 6Cl 1- + 3H 2 O Al 2 O 3 +2 NaOH + 3H 2 O =2 Na – sodium tetrahydroxyaluminate. The tetrahydroxoaluminate anion is actually the 1- tetrahydroxodiaquaanion, because coordination number 6 is preferable for aluminum. With an excess of alkali, hexahydroxoaluminate is formed: Al 2 O 3 + 6NaOH + 3H 2 O = 2Na 3. In addition to acids and alkalis, reactions with acidic salts can be expected: 6KHSO 4 + Al 2 O 3 = 3K 2 SO 4 + Al 2 (SO 4) 3 + 3H 2 O.

3. Aluminum hydroxides. There are two known aluminum hydroxides - metahydroxide -AlO(OH) and orthohydroxide - Al(OH) 3. Both of them are insoluble in water, but are also amphoteric, therefore they dissolve in solutions of acids and alkalis, as well as salts that have an acidic or alkaline environment as a result of hydrolysis. When fused, hydroxides react similarly to oxides. Like all insoluble bases, aluminum hydroxides decompose when heated: 2Al(OH) 3 = Al 2 O 3 + 3H 2 O. Dissolving in alkaline solutions, aluminum hydroxides do not dissolve in aqueous ammonia, so they can be precipitated with ammonia from a soluble salt: Al(NO 3) 3 + 3NH 3 + 2H 2 O = AlO(OH)↓ + 3NH 4 NO 3, this reaction produces metahydroxide. It is difficult to precipitate hydroxide by the action of alkalis, because the resulting precipitate easily dissolves, and the total reaction has the form: AlCl 3 + 4 NaOH = Na + 3NaCl

4. Aluminum salts. Almost all aluminum salts are highly soluble in water. AlPO 4 phosphate and AlF 3 fluoride are insoluble. Because the aluminum cation has a high charge concentration, its aqua complex acquires the properties of a cationic acid: 3+ + H 2 O = H 3 O + + 2+, i.e. aluminum salts undergo strong cation hydrolysis. In the case of salts of weak acids, due to the mutual enhancement of hydrolysis at the cation and anion, hydrolysis becomes irreversible. In solution, aluminum carbonate, sulfite, sulfide and silicate are completely decomposed by water or cannot be obtained by exchange reaction: Al 2 S 3 + 6H 2 O = 2Al(OH) 3 ↓ + 3H 2 S 2Al(NO 3) 3 + 3K 2 CO 3 + 3H 2 O = 2Al(OH) 3 ↓ + 3CO 2 + 6KNO 3. For some salts, hydrolysis becomes irreversible when heated. When heated, wet aluminum acetate decomposes in accordance with the equation: 2Al(OOCCH 3) 3 + 3H 2 O = Al 2 O 3 + 6CH 3 COOH In the case of aluminum halides, the decomposition of the salt is facilitated by a decrease in the solubility of gaseous hydrogen halides when heated: AlCl 3 + 3H 2 O = Al(OH) 3 ↓ + 3HCl. Of the aluminum halides, only fluoride is an ionic compound, the remaining halides are covalent compounds, their melting points are significantly lower than those of fluoride, aluminum chloride is capable of sublimation. At very high temperatures, the vapor contains single molecules of aluminum halides, which have a flat triangular structure due to sp 2 hybridization of the atomic orbitals of the central atom. The ground state of these compounds in vapors and in some organic solvents is dimers, for example, Al 2 Cl 6 . Aluminum halides are strong Lewis acids because have a vacant atomic orbital. Dissolution in water therefore occurs with the release of a large amount of heat. An interesting class of aluminum compounds (as well as other trivalent metals) are alum - 12-aqueous double sulfates M I M III (SO 4) 2, which, when dissolved like all double salts, give a mixture of the corresponding cations and anions.

5. Complex connections. Let's consider hydroxo complexes of aluminum. These are salts in which the complex particle is an anion. All salts are soluble. They are destroyed when interacting with acids. In this case, strong acids dissolve the resulting orthohydroxide, and weak or corresponding acidic oxides (H 2 S, CO 2, SO 2) precipitate it: K + 4HCl = KCl + AlCl 3 + 4H 2 O K + CO 2 = Al(OH) 3 ↓ + KHCO 3

When calcined, hydroxoaluminates transform into ortho- or meta-aluminates, losing water.

Iron

An element with atomic number 26, with a relative atomic mass of 55.847. Belongs to the 3d family of elements, has an electronic configuration: 3d 6 4s 2 and is in the IV period, VIII group, secondary subgroup in the periodic table. In compounds, iron predominantly exhibits oxidation states +2 and +3. The Fe 3+ ion has a half-filled d-electron shell, 3d 5, which gives it additional stability. It is much more difficult to achieve oxidation states +4, +6, +8.

According to its physical properties, iron is a silvery-white, shiny, relatively soft, malleable, easily magnetized and demagnetized metal. Melting point 1539 o C. It has several allotropic modifications, differing in the type of crystal lattice.

Properties of a simple substance.

1. When burned in air, it forms a mixed oxide Fe 3 O 4, and when interacting with pure oxygen - Fe 2 O 3. Powdered iron is pyrophoric - spontaneously ignites in air.

2. Fluorine, chlorine and bromine easily react with iron, oxidizing it to Fe 3+. FeJ 2 is formed with iodine, since the trivalent iron cation oxidizes the iodide anion, and therefore the FeJ 3 compound does not exist.

3. For a similar reason, the Fe 2 S 3 compound does not exist, and the interaction of iron and sulfur at the melting point of sulfur leads to the FeS compound. With an excess of sulfur, pyrite is obtained - iron (II) disulfide - FeS 2. Non-stoichiometric compounds are also formed.

4. Iron reacts with other non-metals under strong heating, forming solid solutions or metal-like compounds. You can give a reaction that occurs at 500 o C: 3Fe + C = Fe 3 C. This compound of iron and carbon is called cementite.

5. Iron forms alloys with many metals.

6. In air at room temperature, iron is covered with an oxide film, so it does not interact with water. Interaction with superheated steam gives the following products: 3Fe + 4H 2 O (steam) = Fe 3 O 4 + 4H 2. In the presence of oxygen, iron even interacts with air moisture: 4Fe + 3O 2 + 6H 2 O = 4Fe(OH) 3. The above equation reflects the rusting process, which up to 10% of metal products undergo per year.

7. Since iron is in the voltage series before hydrogen, it easily reacts with non-oxidizing acids, but is oxidized only to Fe 2+.

8. Concentrated nitric and sulfuric acids passivate iron, but the reaction occurs when heated. Dilute nitric acid also reacts at room temperature. With all oxidizing acids, iron produces iron (III) salts (according to some reports, the formation of iron (II) nitrate is possible with dilute nitric acid), and reduces HNO 3 (diluted) to NO, N 2 O, N 2, NH 4 + depending on conditions, and HNO 3 (conc.) - to NO 2 due to the heating that is necessary for the reaction to occur.

9. Iron is capable of reacting with concentrated (50%) alkalis when heated: Fe + 2KOH + 2H 2 O = K 2 + H 2

10. Reacting with solutions of salts of less active metals, iron removes these metals from the composition of the salt, turning into a divalent cation: CuCl 2 + Fe = FeCl 2 + Cu.

Properties of iron compounds.

Fe 2+ The charge to radius ratio of this cation is close to that of Mg 2+, so the chemical behavior of the oxide, hydroxide and salts of ferrous iron is similar to the behavior of the corresponding magnesium compounds. In an aqueous solution, the divalent iron cation forms an aqua complex 2+ of a pale green color. This cation is easily oxidized even directly in solution by atmospheric oxygen. The FeCl 2 solution contains complex particles 0. The charge concentration of such a cation is low, so the hydrolysis of salts is moderate.

1. FeO - the main oxide, black in color, does not dissolve in water. Easily dissolves in acids. When heated above 500 0 C, it disproportionates: 4FeO = Fe + Fe 3 O 4. It can be obtained by careful calcination of the corresponding hydroxide, carbonate and oxalate, while thermal decomposition of other Fe 2+ salts leads to the formation of ferric oxide: FeC 2 O 4 = FeO + CO + CO 2, but 2 FeSO 4 = Fe 2 O 3 + SO 2 + SO 3 4Fe(NO 3) 2 = 2Fe 2 O 3 + 8NO 2 + O 2 Iron (II) oxide itself can act as an oxidizing agent, for example, when heated, the reaction occurs: 3FeO + 2NH 3 = 3Fe + N 2 +3H 2 O

2. Fe(OH) 2 – iron (II) hydroxide – insoluble base. Reacts with acids. With oxidizing acids, an acid-base interaction and oxidation to ferric iron occur simultaneously: 2Fe(OH) 2 + 4H 2 SO 4 (conc) = Fe 2 (SO 4) 3 + SO 2 + 4H 2 O. Can be obtained by exchange reactions from soluble salt. This is a white compound that first turns green in air due to interaction with air moisture, and then turns brown due to oxidation by air oxygen: 4Fe(OH) 2 + 2H 2 O + O 2 = 4Fe(OH) 3.

3. Salts. As already mentioned, most Fe(II) salts oxidize slowly in air or in solution. The most resistant to oxidation is Mohr's salt - double iron (II) and ammonium sulfate: (NH 4) 2 Fe(SO 4) 2. 6H 2 O. The Fe 2+ cation is easily oxidized to Fe 3+, therefore most oxidizing agents, in particular oxidizing acids, oxidize ferrous iron salts. When iron sulfide and disulfide are fired, iron (III) oxide and sulfur (IV) oxide are obtained: 4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2 Iron (II) sulfide also dissolves in strong acids: FeS + 2HCl = FeCl 2 + 2H 2 S Iron (II) carbonate is insoluble, whereas bicarbonate is soluble in water.

Fe 3+ Charge to radius ratio this cation corresponds to the aluminum cation , therefore, the properties of iron(III) cation compounds are similar to the corresponding aluminum compounds.

Fe 2 O 3 is hematite, an amphoteric oxide in which basic properties predominate. Amphotericity is manifested in the possibility of fusion with solid alkalis and alkali metal carbonates: Fe 2 O 3 + 2NaOH = H 2 O + 2NaFeO 2 - yellow or red, Fe 2 O 3 + Na 2 CO 3 = 2NaFeO 2 + CO 2. Ferrates (II) decompose with water, releasing Fe 2 O 3. nH2O.

Fe3O4- magnetite, a black substance that can be considered either as a mixed oxide - FeO. Fe 2 O 3, or as iron (II) oxometaferrate (III): Fe(FeO 2) 2. When interacting with acids, it gives a mixture of salts: Fe 3 O 4 + 8HCl = FeCl 2 + 2FeCl 3 + 4H 2 O.

Fe(OH) 3 or FeO(OH) is a red-brown gelatinous precipitate, amphoteric hydroxide. In addition to interactions with acids, it reacts with a hot concentrated alkali solution and fuses with solid alkalis and carbonates: Fe(OH) 3 + 3KOH = K 3 .

Salt. Most ferric salts are soluble. Just like aluminum salts, they undergo strong hydrolysis at the cation, which in the presence of anions of weak and unstable or insoluble acids can become irreversible: 2FeCl 3 + 3Na 2 CO 3 + 3H 2 O = 2Fe(OH) 3 + 3CO 2 + 6NaCl. By boiling a solution of iron (III) chloride, hydrolysis can also be made irreversible, because the solubility of hydrogen chloride, like any gas, decreases when heated and it leaves the reaction sphere: FeCl 3 + 3H 2 O = Fe(OH) 3 + 3HCl (when heated).

The oxidizing capacity of this cation is very high, especially in relation to the conversion into the Fe 2+ cation: Fe 3+ + ē = Fe 2+ φ o = 0.77v. Resulting in:

a) solutions of ferric iron salts oxidize all metals up to copper: 2Fe(NO 3) 3 + Cu = 2Fe(NO 3) 2 + Cu(NO 3) 2,

b) exchange reactions with salts containing easily oxidized anions take place simultaneously with their oxidation: 2FeCl 3 + 2KJ = FeCl 2 + J 2 + 2KCl 2FeCl 3 + 3Na 2 S = 2FeS + S + 6NaCl

Like other trivalent cations, iron (III) is capable of forming alum - double sulfates with alkali metal or ammonium cations, for example: NH 4 Fe (SO 4) 2. 12H2O.

Complex connections. Both iron cations tend to form anionic complexes, especially iron(III). FeCl 3 + KCl = K, FeCl 3 + Cl 2 = Cl + -. The latter reaction reflects the action of iron (III) chloride as a catalyst for electrophilic chlorination. Cyanide complexes are of interest: 6KCN + FeSO 4 = K 4 – potassium hexacyanoferrate (II), yellow blood salt. 2K 4 + Cl 2 = 2K 3 + 2KCl – potassium hexacyanoferrate (III), red blood salt. The ferrous iron complex gives a blue precipitate or solution with the ferric salt, depending on the ratio of the reagents. The same reaction occurs between red blood salt and any ferrous salt. In the first case, the precipitate was called Prussian blue, in the second - Turnbull blue. Later it turned out that at least the solutions have the same composition: K – potassium iron (II,III) hexacyanoferrate. The described reactions are qualitative for the presence of the corresponding iron cations in the solution. A qualitative reaction to the presence of ferric cation is the appearance of a blood-red color when interacting with potassium thiocyanate (rhodanide): 2FeCl 3 + 6KCNS = 6KCl + Fe.

Fe +6. The oxidation state +6 for iron is unstable. It is possible to obtain only the FeO 4 2- anion, which exists only at pH>7-9, but is a strong oxidizing agent.

Fe 2 O 3 + 4KOH + 3KNO 3 = 2K 2 FeO 4 + 3KNO 2 + 2H 2 O

Fe (sawdust) + H 2 O + KOH + KNO 3 = K 2 FeO 4 + KNO 2 + H 2

2Fe(OH) 3 + 3Cl 2 + 10KOH = 2K 2 FeO 4 + 6KCl + 6H 2 O

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

4K 2 FeO 4 + 6H 2 O = 4FeO(OH)↓ + 8KOH + 3O 2

4BaFeO 4 (heating) = 4BaO + 2Fe 2 O 3 + 3O 2

2K 2 FeO 4 + 2CrCl 3 + 2HCl = FeCl 3 + K 2 Cr 2 O 7 + 2KCl + H 2 O

Obtaining iron in industry:

A) domain process: Fe 2 O 3 + C = 2FeO + CO

FeO + C = Fe + CO

FeO + CO = Fe + CO 2

B) aluminothermy: Fe 2 O 3 + Al = Al 2 O 3 + Fe

CHROMIUM – element with atomic number 24, with a relative atomic mass of 51.996. It belongs to the 3d family of elements, has an electronic configuration of 3d 5 4s 1 and is in period IV, group VI, a secondary subgroup in the periodic table. Possible oxidation states: +1, +2, +3, +4, +5, +6. Of these, the most stable are +2, +3, +6, and +3 has the minimum energy.

According to its physical properties, chromium is a grayish-white, shiny, hard metal with a melting point of 1890 o C. The strength of its crystal lattice is due to the presence of five unpaired d-electrons, capable of partial covalent bonding.

Chemical properties of a simple substance.

At low temperatures, chromium is inert due to the presence of an oxide film and does not interact with water and air.

1. It interacts with oxygen at temperatures above 600 o C. In this case, chromium (III) oxide – Cr 2 O 3 – is formed.

2. Interaction with halogens occurs in different ways: Cr + 2F 2 = CrF 4 (at room temperature), 2Cr + 3Cl 2 (Br 2) = 2CrCl 3 (Br 3), Cr + J 2 = CrJ 2 (with significant heating ). It should be said that chromium (III) iodide can exist and is obtained by an exchange reaction in the form of crystalline hydrate CrJ 3. 9H 2 O, but its thermal stability is low, and when heated it decomposes into CrJ 2 and J 2.

3. At temperatures above 120 o C, chromium reacts with molten sulfur, giving chromium (II) sulfide - CrS (black).

4. At temperatures above 1000 o C, chromium reacts with nitrogen and carbon, giving non-stoichiometric, chemically inert compounds. Among them, we can note carbide with an approximate composition of CrC, which is close to diamond in hardness.

5. Chromium does not react with hydrogen.

6. The reaction with water vapor is as follows: 2Cr + 3H 2 O = Cr 2 O 3 + 3H 2

7. The reaction with non-oxidizing acids occurs quite easily, resulting in the formation of an aqua complex 2+ of a sky blue color, which is stable only in the absence of air or in a hydrogen atmosphere. In the presence of oxygen, the reaction proceeds differently: 4Cr + 12HCl + 3O 2 = 4CrCl 3 + 6H 2 O. Dilute acids saturated with oxygen even passivate chromium due to the formation of a strong oxide film on the surface.

8. Oxidizing acids: nitric acid of any concentration, concentrated sulfuric acid and perchloric acid passivate chromium so that after treating the surface with these acids, it no longer reacts with other acids. Passivation is removed when heated. This produces chromium (III) salts and sulfur or nitrogen dioxides (chloride from perchloric acid). Passivation due to the formation of a salt film occurs when chromium reacts with phosphoric acid.

9. Chromium does not react directly with alkali, but reacts with alkaline melts with the addition of oxidizing agents: 2Cr + 2Na 2 CO 3 (l) + 3O 2 = 2Na 2 CrO 4 + 2CO 2

10. Chromium is capable of reacting with salt solutions, displacing less active metals (those to the right of it in the voltage series) from the salt composition. Chromium itself is converted into the Cr 2+ cation.

Aluminum – destruction of the 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 an active metal, has fairly good corrosion properties. This can be explained by the ability to passivate 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 forms on the metal surface only in the pH range from 3 to 9!

The corrosion resistance of Al is greatly influenced by its purity. For the manufacture of chemical units and equipment, only high-purity metal (without impurities), for example, AB1 and AB2 aluminum, is used.

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

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 + 3C → Al 4 C 3 – the reaction of aluminum with carbon to form aluminum carbide;

2Al + 3S → Al 2 S 3 – 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, becomes passive. When pure metal comes into contact with air, a thin protective film of aluminum oxide instantly appears on the aluminum surface. Further, film growth slows down. The formula of aluminum oxide is Al 2 O 3 or Al 2 O 3 H 2 O.

The reaction of aluminum with oxygen:

4Al + 3O 2 → 2Al 2 O 3.

The thickness of this oxide film ranges from 5 to 100 nm (depending on operating conditions). Aluminum oxide has good adhesion to the surface and satisfies the condition of 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. When heat treating aluminum, the thickness of the oxide film can reach 0.1 microns.

Aluminum is quite resistant both in clean rural air and in an industrial atmosphere (containing sulfur vapor, hydrogen sulfide, ammonia gas, dry hydrogen chloride, etc.). Because sulfur compounds do not have any effect on the corrosion of aluminum in gas environments - it is used for the manufacture of sour crude oil processing plants and rubber vulcanization devices.

Corrosion of aluminum in water

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

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

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

When interacting with sea water, pure aluminum begins to corrode, because... sensitive to dissolved salts. To use aluminum in seawater, a small amount of magnesium and silicon is added to its composition. The corrosion resistance of aluminum and its alloys when exposed to sea water is significantly reduced if the metal contains copper.

Corrosion of aluminum in acids

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

Corrosion of aluminum in sulfuric acid

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

2Al + 3H 2 SO 4 (dil) → 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.

In this case, a soluble salt is formed - 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

Aluminum or its alloys quickly dissolve in hydrochloric acid (especially when the temperature rises). Corrosion equation:

2Al + 6HCl → 2AlCl 3 + 3H 2.

Solutions of hydrobromic (HBr) and 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 temperatures is extremely resistant (resistance is higher than that of stainless steel 12Х18Н9). It is even used to produce concentrated nitric acid by direct synthesis.

When heated, 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 is quite resistant to acetic acid of any concentration, but only if the temperature does not exceed 65 °C. It is used to produce formaldehyde and acetic acid. At higher temperatures, aluminum dissolves (with the exception of acid concentrations of 98 - 99.8%).

Aluminum is stable in bromic and weak solutions of chromic (up to 10%), phosphoric (up to 1%) acids at room temperature.

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

Oxalic, formic, and organochlorine acids destroy metal.

The corrosion resistance of aluminum is greatly influenced by vapor and liquid mercury. After a short contact, the metal and its alloys intensively corrode, 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 (aluminum corrosion with hydrogen depolarization).

2Al + 2NaOH + 6H 2 O → 2Na + 3H 2;

2(NaOHH 2 O) + 2Al → 2NaAlO 2 + 3H 2.

Aluminates are formed.

Also, the oxide film is destroyed by mercury, copper and chlorine ions.