Properties of aluminum briefly. What is aluminum


There is a lot of aluminum in the earth's crust: 8.6% by weight. It ranks first among all metals and third among other elements (after oxygen and silicon). There is twice as much aluminum as iron, and 350 times more than copper, zinc, chromium, tin and lead combined! As he wrote more than 100 years ago in his classic textbook Basics of Chemistry D.I. Mendeleev, of all metals, “aluminum is the most common in nature; It is enough to point out that it is part of clay to make clear the universal distribution of aluminum in the earth’s crust. Aluminum, or alum metal (alumen), is also called clay because it is found in clay.”

The most important aluminum mineral is bauxite, a mixture of the basic oxide AlO(OH) and hydroxide Al(OH) 3. The largest bauxite deposits are located in Australia, Brazil, Guinea and Jamaica; industrial production is also carried out in other countries. Alunite (alum stone) (Na,K) 2 SO 4 Al 2 (SO 4) 3 4Al(OH) 3 and nepheline (Na, K) 2 O Al 2 O 3 2SiO 2 are also rich in aluminum. In total, more than 250 minerals are known that contain aluminum; most of them are aluminosilicates, from which the earth’s crust is mainly formed. When they weather, clay is formed, the basis of which is the mineral kaolinite Al 2 O 3 2SiO 2 2H 2 O. Iron impurities usually color the clay brown, but there is also white clay - kaolin, which is used to make porcelain and earthenware products.

Occasionally, an exceptionally hard (second only to diamond) mineral corundum is found - crystalline oxide Al 2 O 3, often colored by impurities in different colors. Its blue variety (an admixture of titanium and iron) is called sapphire, the red one (an admixture of chromium) is called ruby. Various impurities can also color the so-called noble corundum green, yellow, orange, purple and other colors and shades.

Until recently, it was believed that aluminum, as a highly active metal, could not occur in nature in a free state, but in 1978, native aluminum was discovered in the rocks of the Siberian Platform - in the form of thread-like crystals only 0.5 mm long (with a thread thickness of several micrometers). Native aluminum was also discovered in lunar soil brought to Earth from the regions of the Seas of Crisis and Abundance. It is believed that aluminum metal can be formed by condensation from gas. It is known that when aluminum halides - chloride, bromide, fluoride - are heated, they can evaporate with greater or less ease (for example, AlCl 3 sublimes already at 180 ° C). With a strong increase in temperature, aluminum halides decompose, transforming into a state with a lower metal valence, for example, AlCl. When such a compound condenses with a decrease in temperature and the absence of oxygen, a disproportionation reaction occurs in the solid phase: some of the aluminum atoms are oxidized and pass into the usual trivalent state, and some are reduced. Monivalent aluminum can only be reduced to metal: 3AlCl ® 2Al + AlCl 3 . This assumption is also supported by the thread-like shape of native aluminum crystals. Typically, crystals of this structure are formed due to rapid growth from the gas phase. It is likely that microscopic aluminum nuggets in the lunar soil were formed in a similar way.

The name aluminum comes from the Latin alumen (genus aluminis). This was the name of alum, double potassium-aluminum sulfate KAl(SO 4) 2 · 12H 2 O), which was used as a mordant when dyeing fabrics. The Latin name probably goes back to the Greek “halme” - brine, salt solution. It is curious that in England aluminum is aluminum, and in the USA it is aluminum.

Many popular books on chemistry contain a legend that a certain inventor, whose name has not been preserved by history, brought to the Emperor Tiberius, who ruled Rome in 14–27 AD, a bowl made of a metal resembling the color of silver, but lighter. This gift cost the master his life: Tiberius ordered his execution and the destruction of the workshop, because he was afraid that the new metal could depreciate the value of silver in the imperial treasury.

This legend is based on a story by Pliny the Elder, a Roman writer and scholar, author Natural history– encyclopedia of natural science knowledge of ancient times. According to Pliny, the new metal was obtained from "clayey earth." But clay does contain aluminum.

Modern authors almost always make a reservation that this whole story is nothing more than a beautiful fairy tale. And this is not surprising: aluminum in rocks is extremely tightly bound to oxygen, and a lot of energy must be spent to release it. However, recently new data have appeared on the fundamental possibility of obtaining metallic aluminum in ancient times. As spectral analysis showed, the decorations on the tomb of the Chinese commander Zhou-Zhu, who died at the beginning of the 3rd century. AD, are made of an alloy consisting of 85% aluminum. Could the ancients have obtained free aluminum? All known methods (electrolysis, reduction with metallic sodium or potassium) are automatically eliminated. Could native aluminum be found in ancient times, like, for example, nuggets of gold, silver, and copper? This is also excluded: native aluminum is a rare mineral that is found in negligible quantities, so the ancient craftsmen could not find and collect such nuggets in the required quantities.

However, another explanation for Pliny's story is possible. Aluminum can be recovered from ores not only with the help of electricity and alkali metals. There is a reducing agent available and widely used since ancient times - coal, with the help of which the oxides of many metals are reduced to free metals when heated. In the late 1970s, German chemists decided to test whether aluminum could have been produced in ancient times by reduction with coal. They heated a mixture of clay with coal powder and table salt or potash (potassium carbonate) in a clay crucible to red heat. Salt was obtained from sea water, and potash from plant ash, in order to use only those substances and methods that were available in ancient times. After some time, slag with aluminum balls floated to the surface of the crucible! The metal yield was small, but it is possible that it was in this way that the ancient metallurgists could obtain the “metal of the 20th century.”

Properties of aluminum.

The color of pure aluminum resembles silver; it is a very light metal: its density is only 2.7 g/cm 3 . The only metals lighter than aluminum are alkali and alkaline earth metals (except barium), beryllium and magnesium. Aluminum also melts easily - at 600 ° C (thin aluminum wire can be melted on a regular kitchen burner), but it boils only at 2452 ° C. In terms of electrical conductivity, aluminum is in 4th place, second only to silver (it is in first place), copper and gold, which, given the cheapness of aluminum, is of great practical importance. The thermal conductivity of metals changes in the same order. It is easy to verify the high thermal conductivity of aluminum by dipping an aluminum spoon into hot tea. And one more remarkable property of this metal: its smooth, shiny surface perfectly reflects light: from 80 to 93% in the visible region of the spectrum, depending on the wavelength. In the ultraviolet region, aluminum has no equal in this regard, and only in the red region is it slightly inferior to silver (in the ultraviolet, silver has a very low reflectivity).

Pure aluminum is a fairly soft metal - almost three times softer than copper, so even relatively thick aluminum plates and rods are easy to bend, but when aluminum forms alloys (there are a huge number of them), its hardness can increase tenfold.

The characteristic oxidation state of aluminum is +3, but due to the presence of unfilled 3 R- and 3 d-orbitals, aluminum atoms can form additional donor-acceptor bonds. Therefore, the Al 3+ ion with a small radius is very prone to complex formation, forming a variety of cationic and anionic complexes: AlCl 4 –, AlF 6 3–, 3+, Al(OH) 4 –, Al(OH) 6 3–, AlH 4 – and many others. Complexes with organic compounds are also known.

The chemical activity of aluminum is very high; in the series of electrode potentials it stands immediately behind magnesium. At first glance, such a statement may seem strange: after all, an aluminum pan or spoon is quite stable in the air and does not collapse in boiling water. Aluminum, unlike iron, does not rust. It turns out that when exposed to air, the metal is covered with a colorless, thin but durable “armor” of oxide, which protects the metal from oxidation. So, if you introduce a thick aluminum wire or plate 0.5–1 mm thick into the burner flame, the metal melts, but the aluminum does not flow, since it remains in a bag of its oxide. If you deprive aluminum of its protective film or make it loose (for example, by immersing it in a solution of mercury salts), aluminum will immediately reveal its true essence: already at room temperature it will begin to react vigorously with water, releasing hydrogen: 2Al + 6H 2 O ® 2Al(OH) 3 + 3H 2 . In air, aluminum, stripped of its protective film, turns into loose oxide powder right before our eyes: 2Al + 3O 2 ® 2Al 2 O 3 . Aluminum is especially active in a finely crushed state; When blown into a flame, aluminum dust burns instantly. If you mix aluminum dust with sodium peroxide on a ceramic plate and drop water on the mixture, the aluminum also flares up and burns with a white flame.

The very high affinity of aluminum for oxygen allows it to “take away” oxygen from the oxides of a number of other metals, reducing them (aluminothermy method). The most famous example is the thermite mixture, which, when burned, releases so much heat that the resulting iron melts: 8Al + 3Fe 3 O 4 ® 4Al 2 O 3 + 9Fe. This reaction was discovered in 1856 by N.N. Beketov. In this way, Fe 2 O 3, CoO, NiO, MoO 3, V 2 O 5, SnO 2, CuO, and a number of other oxides can be reduced to metals. When reducing Cr 2 O 3, Nb 2 O 5, Ta 2 O 5, SiO 2, TiO 2, ZrO 2, B 2 O 3 with aluminum, the heat of reaction is not enough to heat the reaction products above their melting point.

Aluminum easily dissolves in dilute mineral acids to form salts. Concentrated nitric acid, oxidizing the surface of aluminum, promotes thickening and strengthening of the oxide film (the so-called passivation of the metal). Aluminum treated in this way does not react even with hydrochloric acid. Using electrochemical anodic oxidation (anodizing), a thick film can be created on the surface of aluminum, which can be easily painted in different colors.

The displacement of less active metals by aluminum from solutions of salts is often hindered by a protective film on the surface of aluminum. This film is quickly destroyed by copper chloride, so the reaction 3CuCl 2 + 2Al ® 2AlCl 3 + 3Cu occurs easily, which is accompanied by strong heating. In strong alkali solutions, aluminum easily dissolves with the release of hydrogen: 2Al + 6NaOH + 6H 2 O ® 2Na 3 + 3H 2 (other anionic hydroxo complexes are also formed). The amphoteric nature of aluminum compounds is also manifested in the easy dissolution of its freshly precipitated oxide and hydroxide in alkalis. Crystalline oxide (corundum) is very resistant to acids and alkalis. When fused with alkalis, anhydrous aluminates are formed: Al 2 O 3 + 2NaOH ® 2NaAlO 2 + H 2 O. Magnesium aluminate Mg(AlO 2) 2 is a semi-precious spinel stone, usually colored by impurities in a wide variety of colors.

The reaction of aluminum with halogens occurs rapidly. If a thin aluminum wire is placed in a test tube with 1 ml of bromine, then after a short time the aluminum ignites and burns with a bright flame. The reaction of a mixture of aluminum and iodine powders is initiated by a drop of water (water with iodine forms an acid that destroys the oxide film), after which a bright flame appears with clouds of violet iodine vapor. Aluminum halides in aqueous solutions have an acidic reaction due to hydrolysis: AlCl 3 + H 2 O Al(OH)Cl 2 + HCl.

The reaction of aluminum with nitrogen occurs only above 800 ° C with the formation of nitride AlN, with sulfur - at 200 ° C (sulfide Al 2 S 3 is formed), with phosphorus - at 500 ° C (phosphide AlP is formed). When boron is added to molten aluminum, borides of the composition AlB 2 and AlB 12 are formed - refractory compounds resistant to acids. Hydride (AlH) x (x = 1.2) is formed only in vacuum at low temperatures in the reaction of atomic hydrogen with aluminum vapor. AlH 3 hydride, stable in the absence of moisture at room temperature, is obtained in a solution of anhydrous ether: AlCl 3 + LiH ® AlH 3 + 3LiCl. With an excess of LiH, salt-like lithium aluminum hydride LiAlH 4 is formed - a very strong reducing agent used in organic syntheses. It instantly decomposes with water: LiAlH 4 + 4H 2 O ® LiOH + Al(OH) 3 + 4H 2.

Production of aluminum.

The documented discovery of aluminum occurred in 1825. This metal was first obtained by the Danish physicist Hans Christian Oersted, when he isolated it by the action of potassium amalgam on anhydrous aluminum chloride (obtained by passing chlorine through a hot mixture of aluminum oxide and coal). Having distilled off the mercury, Oersted obtained aluminum, although it was contaminated with impurities. In 1827, the German chemist Friedrich Wöhler obtained aluminum in powder form by reducing hexafluoroaluminate with potassium:

Na 3 AlF 6 + 3K ® Al + 3NaF + 3KF. Later he managed to obtain aluminum in the form of shiny metal balls. In 1854, the French chemist Henri Etienne Saint-Clair Deville developed the first industrial method for producing aluminum - by reducing the melt of tetrachloroaluminate with sodium: NaAlCl 4 + 3Na ® Al + 4NaCl. However, aluminum continued to be an extremely rare and expensive metal; it was not much cheaper than gold and 1500 times more expensive than iron (now only three times). A rattle was made from gold, aluminum and precious stones in the 1850s for the son of French Emperor Napoleon III. When a large ingot of aluminum produced by a new method was exhibited at the World Exhibition in Paris in 1855, it was looked upon as if it were a jewel. The upper part (in the form of a pyramid) of the Washington Monument in the US capital was made from precious aluminum. At that time, aluminum was not much cheaper than silver: in the USA, for example, in 1856 it was sold at a price of 12 dollars per pound (454 g), and silver for 15 dollars. In the 1st volume of the famous Brockhaus Encyclopedic Dictionary published in 1890, Efron said that “aluminum is still used primarily for the manufacture of... luxury goods.” By that time, only 2.5 tons of metal were mined annually throughout the world. Only towards the end of the 19th century, when an electrolytic method for producing aluminum was developed, its annual production began to amount to thousands of tons, and in the 20th century. – million tons. This transformed aluminum from a semi-precious metal to a widely available metal.

The modern method of producing aluminum was discovered in 1886 by a young American researcher, Charles Martin Hall. He became interested in chemistry as a child. Having found his father's old chemistry textbook, he began to diligently study it and carry out experiments, once even receiving a scolding from his mother for damaging the dinner tablecloth. And 10 years later he made an outstanding discovery that made him famous throughout the world.

As a student at age 16, Hall heard from his teacher, F. F. Jewett, that if someone could develop a cheap way to produce aluminum, that person would not only do a great service to humanity, but also make a huge fortune. Jewett knew what he was saying: he had previously trained in Germany, worked with Wöhler, and discussed with him the problems of producing aluminum. Jewett also brought a sample of the rare metal with him to America, which he showed to his students. Suddenly Hall declared publicly: “I will get this metal!”

Six years of hard work continued. Hall tried to obtain aluminum using different methods, but without success. Finally, he tried to extract this metal by electrolysis. At that time there were no power plants; current had to be generated using large homemade batteries from coal, zinc, nitric and sulfuric acids. Hall worked in a barn where he set up a small laboratory. He was helped by his sister Julia, who was very interested in her brother’s experiments. She preserved all his letters and work journals, which make it possible to literally trace the history of the discovery day by day. Here is an excerpt from her memoirs:

“Charles was always in a good mood, and even on the worst days he was able to laugh at the fate of unlucky inventors. In times of failure, he found solace at our old piano. In his home laboratory he worked for long hours without a break; and when he could leave the set up for a while, he would rush across our long house to play a little... I knew that, playing with such charm and feeling, he was constantly thinking about his work. And music helped him with this.”

The most difficult thing was to select an electrolyte and protect the aluminum from oxidation. After six months of exhausting labor, several small silver balls finally appeared in the crucible. Hall immediately ran to his former teacher to tell him about his success. “Professor, I got it!” he exclaimed, holding out his hand: in his palm lay a dozen small aluminum balls. This happened on February 23, 1886. And exactly two months later, on April 23 of the same year, the Frenchman Paul Héroux took out a patent for a similar invention, which he made independently and almost simultaneously (two other coincidences are also striking: both Hall and Héroux were born in 1863 and died in 1914).

Now the first balls of aluminum produced by Hall are kept at the American Aluminum Company in Pittsburgh as a national relic, and at his college there is a monument to Hall, cast from aluminum. Jewett subsequently wrote: “My most important discovery was the discovery of man. It was Charles M. Hall who, at the age of 21, discovered a method of reducing aluminum from ore, and thus made aluminum that wonderful metal which is now widely used throughout the world.” Jewett's prophecy came true: Hall received wide recognition and became an honorary member of many scientific societies. But his personal life was unsuccessful: the bride did not want to come to terms with the fact that her fiancé spends all his time in the laboratory, and broke off the engagement. Hall found solace in his native college, where he worked for the rest of his life. As Charles’s brother wrote, “College was his wife, his children, and everything else—his whole life.” Hall bequeathed the majority of his inheritance to the college - $5 million. Hall died of leukemia at the age of 51.

Hall's method made it possible to produce relatively inexpensive aluminum on a large scale using electricity. If from 1855 to 1890 only 200 tons of aluminum were obtained, then over the next decade, using Hall’s method, 28,000 tons of this metal were already obtained worldwide! By 1930, global annual aluminum production reached 300 thousand tons. Now more than 15 million tons of aluminum are produced annually. In special baths at a temperature of 960–970 ° C, a solution of alumina (technical Al 2 O 3) in molten cryolite Na 3 AlF 6, which is partially mined in the form of a mineral, and partially specially synthesized, is subjected to electrolysis. Liquid aluminum accumulates at the bottom of the bath (cathode), oxygen is released at the carbon anodes, which gradually burn. At low voltage (about 4.5 V), electrolysers consume huge currents - up to 250,000 A! One electrolyzer produces about a ton of aluminum per day. Production requires a lot of electricity: it takes 15,000 kilowatt-hours of electricity to produce 1 ton of metal. This amount of electricity is consumed by a large 150-apartment building for a whole month. Aluminum production is environmentally hazardous, since the atmospheric air is polluted with volatile fluorine compounds.

Application of aluminum.

Even D.I. Mendeleev wrote that “metallic aluminum, having great lightness and strength and low variability in air, is very suitable for some products.” Aluminum is one of the most common and cheapest metals. It is difficult to imagine modern life without it. No wonder aluminum is called the metal of the 20th century. It lends itself well to processing: forging, stamping, rolling, drawing, pressing. Pure aluminum is a fairly soft metal; It is used to make electrical wires, structural parts, food foil, kitchen utensils and “silver” paint. This beautiful and lightweight metal is widely used in construction and aviation technology. Aluminum reflects light very well. Therefore, it is used to make mirrors using the method of metal deposition in a vacuum.

In aircraft and mechanical engineering, in the manufacture of building structures, much harder aluminum alloys are used. One of the most famous is an alloy of aluminum with copper and magnesium (duralumin, or simply “duralumin”; the name comes from the German city of Duren). After hardening, this alloy acquires special hardness and becomes approximately 7 times stronger than pure aluminum. At the same time, it is almost three times lighter than iron. It is obtained by alloying aluminum with small additions of copper, magnesium, manganese, silicon and iron. Silumins are widely used - casting alloys of aluminum and silicon. High-strength, cryogenic (frost-resistant) and heat-resistant alloys are also produced. Protective and decorative coatings are easily applied to products made of aluminum alloys. The lightness and strength of aluminum alloys are especially useful in aviation technology. For example, helicopter rotors are made from an alloy of aluminum, magnesium and silicon. Relatively cheap aluminum bronze (up to 11% Al) has high mechanical properties, it is stable in sea water and even in dilute hydrochloric acid. From 1926 to 1957, coins in denominations of 1, 2, 3 and 5 kopecks were minted from aluminum bronze in the USSR.

Currently, a quarter of all aluminum is used for construction needs, the same amount is consumed by transport engineering, approximately 17% is spent on packaging materials and cans, and 10% in electrical engineering.

Many flammable and explosive mixtures also contain aluminum. Alumotol, a cast mixture of trinitrotoluene and aluminum powder, is one of the most powerful industrial explosives. Ammonal is an explosive substance consisting of ammonium nitrate, trinitrotoluene and aluminum powder. Incendiary compositions contain aluminum and an oxidizing agent - nitrate, perchlorate. Zvezdochka pyrotechnic compositions also contain powdered aluminum.

A mixture of aluminum powder with metal oxides (thermite) is used to produce certain metals and alloys, for welding rails, and in incendiary ammunition.

Aluminum has also found practical use as rocket fuel. To completely burn 1 kg of aluminum, almost four times less oxygen is required than for 1 kg of kerosene. In addition, aluminum can be oxidized not only by free oxygen, but also by bound oxygen, which is part of water or carbon dioxide. When aluminum “burns” in water, 8800 kJ is released per 1 kg of products; this is 1.8 times less than during combustion of metal in pure oxygen, but 1.3 times more than during combustion in air. This means that instead of dangerous and expensive compounds, simple water can be used as an oxidizer for such fuel. The idea of ​​using aluminum as a fuel was proposed back in 1924 by the domestic scientist and inventor F.A. Tsander. According to his plan, it is possible to use aluminum elements of a spacecraft as additional fuel. This bold project has not yet been practically implemented, but most currently known solid rocket fuels contain metallic aluminum in the form of fine powder. Adding 15% aluminum to the fuel can increase the temperature of combustion products by a thousand degrees (from 2200 to 3200 K); The rate of flow of combustion products from the engine nozzle also increases noticeably - the main energy indicator that determines the efficiency of rocket fuel. In this regard, only lithium, beryllium and magnesium can compete with aluminum, but all of them are much more expensive than aluminum.

Aluminum compounds are also widely used. Aluminum oxide is a refractory and abrasive (emery) material, a raw material for the production of ceramics. It is also used to make laser materials, watch bearings, and jewelry stones (artificial rubies). Calcined aluminum oxide is an adsorbent for purifying gases and liquids and a catalyst for a number of organic reactions. Anhydrous aluminum chloride is a catalyst in organic synthesis (Friedel-Crafts reaction), the starting material for the production of high-purity aluminum. Aluminum sulfate is used for water purification; reacting with the calcium bicarbonate it contains:

Al 2 (SO 4) 3 + 3Ca(HCO 3) 2 ® 2AlO(OH) + 3CaSO 4 + 6CO 2 + 2H 2 O, it forms oxide-hydroxide flakes, which, settling, capture and also sorb on the surface those in suspended impurities and even microorganisms in water. In addition, aluminum sulfate is used as a mordant for dyeing fabrics, tanning leather, preserving wood, and sizing paper. Calcium aluminate is a component of cementitious materials, including Portland cement. Yttrium aluminum garnet (YAG) YAlO 3 is a laser material. Aluminum nitride is a refractory material for electric furnaces. Synthetic zeolites (they belong to aluminosilicates) are adsorbents in chromatography and catalysts. Organoaluminum compounds (for example, triethylaluminum) are components of Ziegler-Natta catalysts, which are used for the synthesis of polymers, including high-quality synthetic rubber.

Ilya Leenson

Literature:

Tikhonov V.N. Analytical chemistry of aluminum. M., “Science”, 1971
Popular library of chemical elements. M., “Science”, 1983
Craig N.C. Charles Martin Hall and his Metal. J.Chem.Educ. 1986, vol. 63, no. 7
Kumar V., Milewski L. Charles Martin Hall and the Great Aluminum Revolution. J.Chem.Educ., 1987, vol. 64, no. 8



1. Does not interact with H2.

2. How an active metal reacts with almost all non-metals without heating if the oxide film is removed.

4Al + 3O 2 → 2Al 2 O 3

2Al + 3Cl 2 → 2AlCl 3

Al + P → AlP

3. Reacts with H2O:

Aluminum is an active metal with a high affinity for oxygen. In air it becomes covered with a protective film of oxide. If the film is destroyed, the aluminum actively interacts with water.

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

4. With dilute acids:

2Al + 6HCl → 2AlCl 3 + 3H 2

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

It does not react with concentrated HNO 3 and H 2 SO 4 under normal conditions, but only when heated.

5. With alkalis:

2Al + 2NaOH 2NaAlO 2 + 3H 2

Aluminum forms complexes with aqueous solutions of alkalis:

2Al + 2NaOH + 10 H 2 O = 2Na + - + 3H 2

or Na,

Na3, Na2– hydroxoaluminates. The product depends on the alkali concentration.

4Al + 3O 2 → 2Al 2 O 3

Al 2 O 3 (alumina) occurs in nature in the form of the mineral corundum (close to diamond in hardness). Precious stones ruby ​​and sapphire are also Al 2 O 3, colored with impurities of iron and chromium

Aluminium oxide– amphoteric. When it is fused with alkalis, salts of meta-aluminum acid HAlO 2 are obtained. For example:

Also interacts with acids

White gelatinous sediment aluminum hydroxide dissolves in acids

Al(OH) 3 + 3HCl = AlCl 3 + 3 H 2 O,

and in excess of alkali solutions, exhibits amphoteric properties

Al(OH) 3 + NaOH + 2H 2 O = Na

When fused with alkalis, aluminum hydroxide forms salts of meta-aluminum or ortho-aluminum acids

Al(OH) 3 Al 2 O 3 + H 2 O

Aluminum salts are highly hydrolyzed. Aluminum salts and weak acids are converted into basic salts or undergo complete hydrolysis:

AlCl 3 + HOH ↔ AlOHCl 2 + HCl

Al +3 + HOH ↔ AlOH +2 + H + pH>7 occurs in stage I, but when heated it can also occur in stage II.

AlOHCl 2 + HOH ↔ Al(OH) 2 Cl + HCl

AlOH +2 + HOH ↔ Al(OH) 2 + + H +

During boiling, stage III can also occur

Al(OH) 2 Cl + HOH ↔ Al(OH) 3 + HCl

Al(OH) 2 + + HOH ↔ Al(OH) 3 + H +

Aluminum salts are highly soluble.

AlCl 3 – aluminum chloride is a catalyst in oil refining and various organic syntheses.

Al 2 (SO 4) 3 ×18H 2 O - aluminum sulfate is used to purify water from colloidal particles captured by Al (OH) 3 formed during hydrolysis and reduction of hardness

Al 2 (SO 4) 3 + Ca(HCO 3) 2 = Al(OH) 3 + CO 2 + CaSO 4 ↓

In the leather industry, it serves as a mordant for crumbling cotton fabrics - KAl(SO 4) 2 × 12H 2 O - potassium aluminum sulfate (potassium alum).

The main use of aluminum is the production of alloys based on it. Duralumin is an alloy of aluminum, copper, magnesium and manganese.

Silumin – aluminum and silicon.

Their main advantage is low density and satisfactory resistance to atmospheric corrosion. The hulls of artificial Earth satellites and spacecraft are made from aluminum alloys.

Aluminum is used as a reducing agent in metal smelting (aluminothermy)

Cr 2 O 3 + 2 Al t = 2Cr + Al 2 O 3.

Also used for thermite welding of metal products (a mixture of aluminum and iron oxide Fe 3 O 4) called thermite gives a temperature of about 3000 ° C.

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 its outer electron layer: 2 on the 3s and 1 on the 3p sublevel. Due to this structure, it is characterized by reactions as a result of which the aluminum atom loses three electrons from the outer level and acquires an oxidation state of +3. Aluminum is a highly reactive metal and exhibits very strong reducing properties.

Interaction of aluminum with simple substances

with oxygen

When absolutely pure aluminum comes into contact with air, aluminum atoms located in the surface layer instantly interact with oxygen in the air and form a thin, tens of atomic layers thick, durable oxide film of the composition Al 2 O 3, 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 = 2Al 2 O 3

with halogens

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

2Al + 3I 2 =2AlI 3

Aluminum also reacts with bromine, which is a dark brown liquid, without heating. Simply add a sample of aluminum to liquid bromine: a violent reaction immediately begins, releasing a large amount of heat and light:

2Al + 3Br 2 = 2AlBr 3

The reaction between aluminum and chlorine occurs when heated aluminum foil or fine aluminum powder is added to 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 o 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 reacts 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 reacts 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 prevents aluminum from oxidizing in air. The same protective oxide film makes aluminum inert towards water. When removing the protective oxide film from the surface by methods such as treatment with aqueous solutions of alkali, ammonium chloride or mercury salts (amalgiation), aluminum begins to react vigorously with water to form aluminum hydroxide and hydrogen gas:

with metal oxides

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

2AI + Fe 2 O 3 = 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. with almost 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 (diluted) = Al 2 (SO 4) 3 + 3H 2

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

b) 2AI + 6HCl = 2AICl3 + 3H2

with oxidizing acids

-concentrated sulfuric acid

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

- dilute nitric acid

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

8Al + 30HNO 3(dil.) = 8Al(NO 3) 3 +3N 2 O + 15H 2 O

8Al + 30HNO 3(pure dilute) = 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 = 2Na

Al 2 O 3 + 2NaOH = 2NaAlO 2 + H 2 O

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

2Al + 6H 2 O = 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

  • Designation - Al (Aluminium);
  • Period - III;
  • Group - 13 (IIIa);
  • Atomic mass - 26.981538;
  • Atomic number - 13;
  • Atomic radius = 143 pm;
  • Covalent radius = 121 pm;
  • Electron distribution - 1s 2 2s 2 2p 6 3s 2 3p 1 ;
  • melting temperature = 660°C;
  • boiling point = 2518°C;
  • Electronegativity (according to Pauling/according to Alpred and Rochow) = 1.61/1.47;
  • Oxidation state: +3.0;
  • Density (no.) = 2.7 g/cm3;
  • Molar volume = 10.0 cm 3 /mol.

Aluminum (alum) was first obtained in 1825 by the Dane G. K. Ørsted. Initially, before the discovery of an industrial method of production, aluminum was more expensive than gold.

Aluminum is the most abundant metal in the earth's crust (mass fraction is 7-8%), and the third most abundant of all elements after oxygen and silicon. Aluminum is not found in free form in proirod.

The most important natural aluminum compounds:

  • aluminosilicates - Na 2 O Al 2 O 3 2SiO 2 ; K 2 O Al 2 O 3 2SiO 2
  • bauxite - Al 2 O 3 · n H2O
  • corundum - Al 2 O 3
  • cryolite - 3NaF AlF 3


Rice. Structure of the aluminum atom.

Aluminum is a chemically active metal - at its outer electronic level there are three electrons that participate in the formation of covalent bonds when aluminum interacts with other chemical elements (see Covalent bond). Aluminum is a strong reducing agent and exhibits an oxidation state of +3 in all compounds.

At room temperature, aluminum reacts with oxygen contained in the atmospheric air to form a strong oxide film, which reliably prevents the process of further oxidation (corrosion) of the metal, as a result of which the chemical activity of aluminum decreases.

Thanks to the oxide film, aluminum does not react with nitric acid at room temperature, therefore, aluminum cookware is a reliable container for storing and transporting nitric acid.

Physical properties of aluminum:

  • silver-white metal;
  • solid;
  • lasting;
  • easy;
  • plastic (stretched into thin wire and foil);
  • has high electrical and thermal conductivity;
  • melting point 660°C
  • natural aluminum consists of one isotope 27 13 Al

Chemical properties of aluminum:

  • when removing the oxide film, aluminum reacts with water:
    2Al + 6H 2 O = 2Al(OH) 3 + 3H 2;
  • at room temperature it reacts with bromine and chlorine to form salts:
    2Al + 3Br 2 = 2AlCl 3;
  • at high temperatures, aluminum reacts with oxygen and sulfur (the reaction is accompanied by the release of a large amount of heat):
    4Al + 3O 2 = 2Al 2 O 3 + Q;
    2Al + 3S = Al 2 S 3 + Q;
  • at t=800°C reacts with nitrogen:
    2Al + N 2 = 2AlN;
  • at t=2000°C reacts with carbon:
    2Al + 3C = Al 4 C 3;
  • reduces many metals from their oxides - aluminothermy(at temperatures up to 3000°C) tungsten, vanadium, titanium, calcium, chromium, iron, manganese are produced industrially:
    8Al + 3Fe 3 O 4 = 4Al 2 O 3 + 9Fe;
  • reacts with hydrochloric and dilute sulfuric acid to release hydrogen:
    2Al + 6HCl = 2AlCl3 + 3H2;
    2Al + 3H 2 SO 4 = Al 2 (SO 4) 3 + 3H 2;
  • reacts with concentrated sulfuric acid at high temperature:
    2Al + 6H 2 SO 4 = Al 2 (SO 4) 3 + 3SO 2 + 6H 2 O;
  • reacts with alkalis with the release of hydrogen and the formation of complex salts - the reaction occurs in several stages: when aluminum is immersed in an alkali solution, the durable protective oxide film that is on the surface of the metal dissolves; after the film dissolves, aluminum, as an active metal, reacts with water to form aluminum hydroxide, which reacts with alkali as an amphoteric hydroxide:
    • Al 2 O 3 +2NaOH = 2NaAlO 2 +H 2 O - dissolution of the oxide film;
    • 2Al+6H 2 O = 2Al(OH) 3 +3H 2 - interaction of aluminum with water to form aluminum hydroxide;
    • NaOH+Al(OH) 3 = NaAlO 2 + 2H 2 O - interaction of aluminum hydroxide with alkali
    • 2Al+2NaOH+2H 2 O = 2NaAlO 2 +3H 2 - the overall equation for the reaction of aluminum with alkali.

Aluminum connections

Al 2 O 3 (alumina)

Aluminium oxide Al 2 O 3 is a white, very refractory and hard substance (in nature, only diamond, carborundum and borazone are harder).

Alumina properties:

  • does not dissolve in water and reacts with it;
  • is an amphoteric substance, reacting with acids and alkalis:
    Al 2 O 3 + 6HCl = 2AlCl 3 + 3H 2 O;
    Al 2 O 3 + 6NaOH + 3H 2 O = 2Na 3;
  • how an amphoteric oxide reacts when fused with metal oxides and salts to form aluminates:
    Al 2 O 3 + K 2 O = 2KAlO 2.

In industry, alumina is obtained from bauxite. In laboratory conditions, alumina can be obtained by burning aluminum in oxygen:
4Al + 3O 2 = 2Al 2 O 3.

Applications of Alumina:

  • for the production of aluminum and electrical ceramics;
  • as an abrasive and refractory material;
  • as a catalyst in organic synthesis reactions.

Al(OH)3

Aluminum hydroxide Al(OH) 3 is a white crystalline solid that is obtained as a result of an exchange reaction from a solution of aluminum hydroxide - it precipitates as a white gelatinous precipitate that crystallizes over time. This amphoteric compound is almost insoluble in water:
Al(OH) 3 + 3NaOH = Na 3;
Al(OH) 3 + 3HCl = AlCl 3 + 3H 2 O.

  • interaction of Al(OH) 3 with acids:
    Al(OH) 3 +3H + Cl = Al 3+ Cl 3 +3H 2 O
  • interaction of Al(OH) 3 with alkalis:
    Al(OH) 3 +NaOH - = NaAlO 2 - +2H 2 O

Aluminum hydroxide is obtained by the action of alkalis on solutions of aluminum salts:
AlCl 3 + 3NaOH = Al(OH) 3 + 3NaCl.

Production and use of aluminum

Aluminum is quite difficult to isolate from natural compounds by chemical means, which is explained by the high strength of bonds in aluminum oxide; therefore, for the industrial production of aluminum, electrolysis of a solution of alumina Al 2 O 3 in molten cryolite Na 3 AlF 6 is used. As a result of the process, aluminum is released at the cathode, and oxygen is released at the anode:

2Al 2 O 3 → 4Al + 3O 2

The starting raw material is bauxite. Electrolysis occurs at a temperature of 1000°C: the melting point of aluminum oxide is 2500°C - it is not possible to carry out electrolysis at this temperature, so aluminum oxide is dissolved in molten cryolite, and only then the resulting electrolyte is used in electrolysis to produce aluminum.

Application of aluminum:

  • aluminum alloys are widely used as structural materials in automobile, aircraft, and shipbuilding: duralumin, silumin, aluminum bronze;
  • in the chemical industry as a reducing agent;
  • in the food industry for the production of foil, tableware, packaging material;
  • for making wires, etc.

One of the most common elements on the planet is aluminum. The physical and chemical properties of aluminum are used in industry. You will find everything you need to know about this metal in our article.

Atomic structure

Aluminum is the 13th element of the periodic table. It is in the third period, group III, the main subgroup.

The properties and uses of aluminum are related to its electronic structure. The aluminum atom has a positively charged nucleus (+13) and 13 negatively charged electrons, located at three energy levels. The electronic configuration of the atom is 1s 2 2s 2 2p 6 3s 2 3p 1.

The outer energy level contains three electrons, which determine the constant valence of III. In reactions with substances, aluminum goes into an excited state and is able to give up all three electrons, forming covalent bonds. Like other active metals, aluminum is a powerful reducing agent.

Rice. 1. Structure of the aluminum atom.

Aluminum is an amphoteric metal that forms amphoteric oxides and hydroxides. Depending on the conditions, the compounds exhibit acidic or basic properties.

Physical Description

Aluminum has:

  • lightness (density 2.7 g/cm 3);
  • silver-gray color;
  • high electrical conductivity;
  • malleability;
  • plasticity;
  • melting point - 658°C;
  • boiling point - 2518.8°C.

Tin containers, foil, wire, and alloys are made from metal. Aluminum is used in the manufacture of microcircuits, mirrors, and composite materials.

Rice. 2. Tin containers.

Aluminum is paramagnetic. Metal is attracted to a magnet only in the presence of a magnetic field.

Chemical properties

In air, aluminum quickly oxidizes, becoming covered with an oxide film. It protects the metal from corrosion and also prevents interaction with concentrated acids (nitric, sulfuric). Therefore, acids are stored and transported in aluminum containers.

Under normal conditions, reactions with aluminum are possible only after removing the oxide film. Most reactions occur at high temperatures.

The main chemical properties of the element are described in the table.

Reaction

Description

The equation

With oxygen

Burns at high temperatures releasing heat

4Al + 3O 2 → 2Al 2 O 3

With non-metal

Reacts with sulfur at temperatures above 200°C, with phosphorus - at 500°C, with nitrogen - at 800°C, with carbon - at 2000°C

2Al + 3S → Al 2 S 3 ;

Al + P → AlP;

2Al + N 2 → 2AlN;

4Al + 3C → Al 4 C 3

With halogens

Reacts under normal conditions, with iodine - when heated in the presence of a catalyst (water)

2Al + 3Cl 2 → 2AlCl 3 ;

2Al + 3I 2 → 2AlI 3 ;

2Al + 3Br 2 → 2AlBr 3

With acids

Reacts with dilute acids under normal conditions, with concentrated acids when heated

2Al + 3H 2 SO 4 (diluted) → Al 2 (SO 4) 3 + 3H 2;

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

With alkalis

Reacts with aqueous solutions of alkalis and upon fusion

2Al + 2NaOH + 10H 2 O → 2Na + 3H 2;

2Al + 6KOH → 2KAlO 2 + 2K 2 O + 3H 2

With oxides

Displaces less active metals

2Al + Fe 2 O 3 → 2Fe + Al 2 O 3

Aluminum does not react directly with hydrogen. Reaction with water is possible after removing the oxide film.

Rice. 3. Reaction of aluminum with water.

What have we learned?

Aluminum is an amphoteric active metal with constant valence. It has low density, high electrical conductivity, and plasticity. Attracted by a magnet only in the presence of a magnetic field. Aluminum reacts with oxygen, forming a protective film that prevents reactions with water, concentrated nitric and sulfuric acids. When heated, it reacts with non-metals and concentrated acids, and under normal conditions - with halogens and dilute acids. In oxides it displaces less active metals. Does not react with hydrogen.

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