What kind of flame do metals burn with? Chemistry for the curious

LET'S LOOK BEHIND THE SCENES

To formulate the laws of the ongoing processes, we can limit ourselves to considering cations and exclude anions, since they themselves do not participate in the reaction. (However, the rate of deposition is affected by the type of anions.) If, for simplicity, we assume that both the released and dissolved metals are divalent, then we can write:

Me 1 + Me 2 2+ => Me 1 2+ + Me 2

Moreover, for the first experiment Me 1 = Fe, Me 2 = Cu. So, the process consists of the exchange of charges (electrons) between atoms and ions of both metals. If we separately consider (as intermediate reactions) the dissolution of iron or the precipitation of copper, we obtain:

Fe => Fe 2+ + 2е -
Сu 2+ + 2е - => Сu

Now consider the case when a metal is immersed in water or in a salt solution, with a cation of which exchange is impossible due to its position in the stress series. Despite this, the metal tends to go into solution in the form of an ion. In this case, the metal atom gives up two electrons (if the metal is divalent), the surface of the metal immersed in the solution becomes negatively charged relative to the solution, and a double electric layer is formed at the interface. This potential difference prevents further dissolution of the metal, so that the process soon stops. If two different metals are immersed in a solution, they will both charge, but the less active one will be somewhat weaker, due to the fact that its atoms are less prone to losing electrons. Let's connect both metals with a conductor. Due to the potential difference, a flow of electrons will flow from the more active metal to the less active one, which forms the positive pole of the element. A process occurs in which the more active metal goes into solution, and cations from the solution are released on the more noble metal.

The essence of a galvanic cell

Let us now illustrate with several experiments the somewhat abstract reasoning given above (which, moreover, represents a gross simplification).

First, fill a 250 ml beaker to the middle with a 10% solution of sulfuric acid and immerse not too small pieces of zinc and copper in it. We solder or rivet copper wire to both electrodes, the ends of which should not touch the solution.

As long as the ends of the wire are not connected to each other, we will observe the dissolution of zinc, which is accompanied by the release of hydrogen. Zinc, as follows from the voltage series, is more active than hydrogen, so the metal can displace hydrogen from the ionic state. An electrical double layer is formed on both metals. The easiest way to detect the potential difference between the electrodes is with a voltmeter. Immediately after connecting the device to the circuit, the arrow will indicate approximately 1 V, but then the voltage will quickly drop. If you connect a small light bulb that consumes 1 V to the element, it will light up - at first quite strongly, and then the glow will become weak.

Based on the polarity of the device terminals, we can conclude that the copper electrode is the positive pole. This can be proven without a device by considering the electrochemistry of the process. Let's prepare a saturated solution of table salt in a small beaker or test tube, add about 0.5 ml of an alcohol solution of the phenolphthalein indicator and immerse both electrodes closed with wire into the solution. A faint reddish color will be observed near the negative pole, which is caused by the formation of sodium hydroxide at the cathode.

In other experiments, one can place various pairs of metals in a cell and determine the resulting voltage. For example, magnesium and silver will give a particularly large potential difference due to the significant distance between them and a series of voltages, while zinc and iron, on the contrary, will give a very small one, less than a tenth of a volt. By using aluminum, we will not receive practically any current due to passivation.

All these elements, or, as electrochemists say, circuits, have the disadvantage that when measuring current, the voltage across them drops very quickly. Therefore, electrochemists always measure the true value of voltage in a de-energized state using the method voltage compensation, that is, comparing it with the voltage of another current source.

Let us consider the processes in the copper-zinc element in a little more detail. At the cathode, zinc goes into solution according to the following equation:

Zn => Zn 2+ + 2e -

Hydrogen ions of sulfuric acid are discharged at the copper anode. They attach electrons coming through the wire from the zinc cathode and as a result, hydrogen bubbles are formed:

2Н + + 2е - => Н 2

After a short period of time, the copper will be covered with a thin layer of hydrogen bubbles. In this case, the copper electrode will turn into a hydrogen one, and the potential difference will decrease. This process is called polarization electrode. The polarization of the copper electrode can be eliminated by adding a little potassium dichromate solution to the cell after the voltage drop. After this, the voltage will increase again, as potassium dichromate will oxidize hydrogen to water. Potassium bichromate acts in this case as depolarizer

In practice, galvanic circuits are used whose electrodes are not polarized, or circuits whose polarization can be eliminated by adding depolarizers.

As an example of a non-polarizable element, consider the Daniel element, which was often used in the past as a current source. This is also a copper-zinc element, but both metals are immersed in different solutions. The zinc electrode is placed in a porous clay cell filled with dilute (about 20%) sulfuric acid. The clay cell is suspended in a large glass containing a concentrated solution of copper sulfate, and at the bottom there is a layer of copper sulfate crystals. The second electrode in this vessel is a cylinder made of copper sheet.

This element can be made from a glass jar, a commercially available clay cell (in extreme cases, we use a flower pot, closing the hole in the bottom) and two electrodes of suitable size.

During operation of the cell, zinc dissolves to form zinc sulfate, and metallic copper is released at the copper electrode. But at the same time, the copper electrode is not polarized and the element produces a voltage of about 1 V. Actually, theoretically, the voltage at the terminals is 1.10 V, but when collecting current we measure a slightly smaller value, due to electrical resistance cells.

If we do not remove the current from the element, we need to remove the zinc electrode from the sulfuric acid solution, because otherwise it will dissolve to form hydrogen.

A diagram of a simple cell that does not require a porous partition is shown in the figure. The zinc electrode is located in glass jar at the top, and copper - near the bottom. The entire cell is filled with a saturated solution of table salt. Place a handful of copper sulfate crystals at the bottom of the jar. The resulting concentrated copper sulfate solution will mix with the table salt solution very slowly. Therefore, when the cell operates, copper will be released on the copper electrode, and zinc will dissolve in the form of sulfate or chloride in the upper part of the cell.

Nowadays they are used almost exclusively for batteries. dry elements, which are more convenient to use. Their ancestor is the Leclanche element. The electrodes are a zinc cylinder and a carbon rod. The electrolyte is a paste that mainly consists of ammonium chloride. Zinc dissolves in the paste, and hydrogen is released on the coal. To avoid polarization, the carbon rod is dipped into a linen bag containing a mixture of coal powder and pyrolusite. The carbon powder increases the electrode surface, and the pyrolusite acts as a depolarizer, slowly oxidizing the hydrogen. True, the depolarizing ability of pyrolusite is weaker than that of the previously mentioned potassium bichromate. Therefore, when current is received in dry elements, the voltage quickly drops, they “tire” due to polarization. Only after some time does the oxidation of hydrogen occur with pyrolusite. Thus, the elements “rest” if no current is passed for some time. Let's check this on a flashlight battery to which we connect a light bulb. We connect a voltmeter parallel to the lamp, that is, directly to the terminals. At first, the voltage will be about 4.5 V. (Most often, such batteries have three cells connected in series, each with a theoretical voltage of 1.48 V.) After some time, the voltage will drop and the glow of the light bulb will weaken. Based on the voltmeter readings, we can judge how long the battery needs to rest.

Special place occupied by regenerating cells known as batteries. They undergo reversible reactions and can be recharged after the cell has been discharged by connecting to an external DC source.

Currently, lead-acid batteries are the most common; The electrolyte in them is dilute sulfuric acid, into which two lead plates are immersed. The positive electrode is coated with lead peroxide PbO 2 ( modern name- lead dioxide), negative represents metallic lead. The voltage at the terminals is approximately 2.1 V. When discharging, lead sulfate is formed on both plates, which again turns into metallic lead and lead peroxide when charging.

chemical element of group III periodic table, atomic number 13, relative atomic mass 26.98. In nature, it is represented by only one stable nuclide 27 Al. A number of radioactive isotopes of aluminum have been artificially obtained, the longest-living ones 26 Al has a half-life of 720 thousand years. Aluminum in nature. 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 main oxide AlO(OH) and hydroxide Al(OH)

3 . Largest deposits bauxite is found in Australia, Brazil, Guinea and Jamaica; industrial production is also carried out in other countries. Alunite (alum stone) is also rich in aluminum (Na,K) 2 SO 4 Al 2 (SO 4) 3 4Al(OH) 3, nepheline (Na,K) 2 O Al 2 O 3 2SiO 2 . 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 and kaolin, which is used to make porcelain and earthenware. see also BOXITE.

Occasionally, the extremely hard (second only to diamond) mineral corundum crystalline Al oxide is found

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, cannot occur in nature in a free state, but in 1978 in rocks Siberian platform native aluminum was discovered 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 heating aluminum halides - chloride, bromide, fluoride - 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 valency, 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 were used as a mordant when dyeing fabrics. Latin name, probably goes back to the Greek “halme” brine, brine. 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 1427 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, in Lately 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 insignificant 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 ashes, in order to use only those substances and methods that were available in antiquity. 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 low cost 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 ion 3+ with a small radius is very prone to complex formation, forming a variety of cationic and anionic complexes: AlCl4 , AlF 6 3 , 3+ , Al(OH) 4 , Al(OH) 6 3 , AlH 4and 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, aluminum pan or a spoon are quite stable in air and do not collapse even 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 0.51 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 a thermite mixture, the combustion of which 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 can be reduced to metals2 O 3, CoO, NiO, MoO 3, V 2 O 5, SnO 2, CuO, a number of other oxides. When reducing with aluminum Cr2 O 3, Nb 2 O 5, Ta 2 O 5, SiO 2, TiO 2, ZrO 2, B 2 O 3The heat of reaction is not sufficient 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) can create a thick film 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 3CuCl reaction occurs easily

2 + 2Al ® 2AlCl 3 + 3Cu, 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 a semi-precious spinel stone, usually colored with impurities in a wide variety of colors.

The reaction of aluminum with halogens occurs rapidly. If a thin aluminum wire is introduced into a test tube with 1 ml of bromine, then through a short time the aluminum catches fire 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 are acidic 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 AlN nitride, with sulfur at 200° C (Al sulfide is formed 2 S 3 ), with phosphorus at 500° C (phosphide AlP is formed). When boron is added to molten aluminum, borides of the composition AlB are formed 2 and AlB 12 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 hydride, stable in the absence of moisture at room temperature 3 prepared in a solution of anhydrous ether: Al Cl 3 + LiH ® AlH 3 + 3LiCl. With an excess of LiH, salt-like lithium aluminum hydride LiAlH is formed 4 a very strong reducing agent used in organic syntheses. It decomposes instantly 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 a 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). Made of gold, aluminum and precious stones was made in the 1850s as a rattle 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. Made from precious aluminum top part(in the form of a pyramid) of the Washington Monument in the US capital. 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 book published in 1890 Encyclopedic Dictionary Brockhaus and 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 did outstanding discovery, which made him famous throughout the world.

As a student at the age of 16, Hall heard from his teacher, F. F. Jewett, that if someone could develop a cheap way to produce aluminum, then this person would not only do a great service to humanity, but also make money 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 long hours without a break; and when he could leave the setup for a while, he would rush across our entire long house to play a little... I knew that playing with such

charm and feeling, he constantly thinks 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. Subsequently, Jewett 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, all his life." Hall bequeathed the majority of his inheritance, $5 million, to the college. 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 960970° C, an alumina solution (technical Al

2 O 3 ) in molten cryolite Na 3 AlF 6 , which is partly mined in the form of a mineral, and partly specially synthesized. 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 currentsup 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 because atmospheric air contaminated 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's hard to imagine without him modern life. 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; they make it out of it electric wires, structural details, foil for food products, 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 for the manufacture of mirrors by metal deposition in a vacuum.

In aircraft and mechanical engineering, during manufacturing building structures, use much harder aluminum alloys. 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 widespread - 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 diluted 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 with 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 was 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 in

As a fuel back in 1924, it was proposed by the domestic scientist and inventor F.A. Tsander. According to his plan, aluminum elements can be used spaceship 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; this is 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 fireproof and abrasive (emery) material, raw material for the production of ceramics. It is also used to make laser materials, watch bearings, jewelry stones(artificial rubies). Calcined aluminum oxide adsorbent for purification of gases and liquids and a series catalyst 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 2O, it forms oxide-hydroxide flakes, which, settling, capture and also sorb suspended impurities in the water and even microorganisms on the surface. 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 binders, including Portland cement. Yttrium aluminum garnet (YAG) YAlO 3 laser material. Aluminum nitride refractory material for electric furnaces. Synthetic zeolites (they belong to aluminosilicates) adsorbents in chromatography and catalysts. Organoaluminum compounds (for example, triethylaluminum) components of Ziegler catalysts Natta, 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 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

It is not difficult to guess that the shade of the flame is determined by chemicals, burning in it, in the event that exposure to high temperature releases individual atoms of combustible substances, coloring the fire. To determine the effect of substances on the color of fire, various experiments were carried out, which we will discuss below.

Since ancient times, alchemists and scientists have tried to find out what substances burn, depending on the color that the flame acquires.

The flames of gas water heaters and stoves, available in all houses and apartments, have a blue tint. When burned, this shade is produced by carbon, carbon monoxide. The yellow-orange color of the flame of a fire that is lit in the forest, or of household matches, is due to the high content of sodium salts in natural wood. Largely thanks to this - red. The flame of a gas stove burner will acquire the same color if you sprinkle it with ordinary table salt. When copper burns, the flame will be green. I think you have noticed that when you wear a ring or chain made of ordinary copper that is not coated with a protective compound for a long time, the skin becomes green. The same thing happens during the combustion process. If the copper content is high, a very bright green light occurs, almost identical to white. This can be seen if you sprinkle copper shavings on a gas burner.

Many experiments have been carried out involving the common gas burner and various minerals. In this way their composition was determined. You need to take the mineral with tweezers and place it in the flame. The color that fire takes on can indicate the various impurities present in the element. A green flame and its shades indicate the presence of copper, barium, molybdenum, antimony, and phosphorus. Boron produces a blue-green color. Selenium gives the flame a blue tint. The flame is colored red in the presence of strontium, lithium and calcium, and violet - potassium. The yellow-orange color is produced when sodium burns.

Studies of minerals to determine their composition are carried out using a Bunsen burner. The color of its flame is even and colorless; it does not interfere with the course of the experiment. Bunsen invented the burner in the mid-19th century.

He came up with a method that allows one to determine the composition of a substance by the shade of the flame. Scientists had tried to conduct similar experiments before him, but they did not have a Bunsen burner, the colorless flame of which did not interfere with the progress of the experiment. He placed various elements on a platinum wire into the burner fire, since when this metal is added, the flame does not become colored. At first glance, the method seems good; labor-intensive chemical analysis can be dispensed with. You just need to bring the element to the fire and see what it consists of. But substances in their pure form can be found extremely rarely in nature. They usually contain large quantities of various impurities that change the color of the flame.

Bunsen tried to highlight colors and shades using various methods. For example, using colored glass. Suppose, if you look through blue glass, it will not be visible yellow, in which the fire is colored when burning the most common sodium salts. Then the lilac or crimson shade of the desired element becomes distinguishable. But even such tricks led to the correct determination of the composition of a complex mineral in very rare cases. This technology could not achieve more.

Nowadays, such a torch is used only for soldering.

Aluminum burning

Aluminum burning in air

Unlike magnesium, single aluminum particles do not ignite when heated in air or water vapor to 2100 K. Burning magnesium particles were used to ignite aluminum. The latter were placed on the surface of the heating element, and the aluminum particles were placed on the tip of the needle at a distance of 10-4 m above the former.

When aluminum particles are ignited, ignition occurs in the vapor phase, and the intensity of the glow zone that appears around the particle increases slowly. Stationary combustion is characterized by the existence of a glow zone, which does not change its size until the metal is almost completely burned out. The ratio of the sizes of the glow zone and the particle is 1.6-1.9. In the glow zone, small oxide droplets are formed, which merge upon collision.

The residue after combustion of the particle is a hollow shell containing no metal inside. The dependence of the burning time of a particle on its size is expressed by the formula (symmetrical combustion).

Combustion of aluminum in water vapor

Ignition of aluminum in water vapor occurs heterogeneously. The hydrogen released during the reaction contributes to the destruction of the oxide film; in this case, liquid aluminum oxide (or hydroxide) is sprayed in the form of droplets with a diameter of up to 10-15 microns. Such destruction of the oxide shell is periodically repeated. This suggests that a significant fraction of the metal burns on the surface of the particle.

At the beginning of combustion, the ratio rsv /r 0 equals 1.6-1.7. During the combustion process, the particle size decreases, and the gs/?o ratio increases to 2.0-3.0. The burning rate of an aluminum particle in water vapor is almost 5 times greater than in air.

Combustion of aluminum-magnesium alloys

Combustion of aluminum-magnesium alloys in air

Ignition of particles of aluminum-magnesium alloys variable composition in air, oxygen-argon mixtures, water vapor and carbon dioxide, as a rule, proceeds similarly to the ignition of magnesium particles. The onset of ignition is preceded by oxidative reactions occurring on the surface.

The combustion of aluminum-magnesium alloys differs significantly from the combustion of both aluminum and magnesium and strongly depends on the ratio of components in the alloy and on the parameters of the oxidizing environment. The most important feature of the combustion of alloy particles is the two-stage process (Fig. 2.6). At the first stage, the particle is surrounded by a set of torches, forming a non-uniform zone of luminescence of the reaction products. Comparing the nature and size of the luminous zone surrounding the alloy particle during the first stage of combustion with the nature and size of the luminous zone around the burning magnesium particle (see Fig. 2.4), we can conclude that at this stage, mainly magnesium burns out of the particle.

Rice. 2.6. Combustion of an alloy particle 30% A1 + 70% Mg at normal atmospheric pressure in a mixture containing 15% O by volume 2and 85% Ar:

1, 2 – magnesium burnout; 3-6 – aluminum burnout

A feature of the first stage of alloy combustion is the constancy of particle size and flame zone. This means that the liquid drop of the alloy is contained within a solid oxide shell. The oxide film is dominated by magnesium oxide. Through film defects, magnesium flows out, burning in a vapor-phase diffusion flame.

At the end of the first stage, the occurrence of heterogeneous reactions increases, as evidenced by the appearance of foci of bright luminescence on the surface of the particle. The heat released during heterogeneous reactions contributes to heating the particle to the melting point of the oxide and the beginning of the second stage of combustion.

At the second stage of combustion, the particle is surrounded by a uniform, brighter glow zone, which decreases as the metal burns out. The homogeneity and sphericity of the flame zone indicates that the oxide film on the surface of the particle is molten. Diffusion of the metal through the film is ensured by the low diffusion resistance of the liquid oxide. The size of the flame zone significantly exceeds the particle size, which indicates combustion of the metal in the vapor phase. Comparison of the nature of the second stage of combustion with the known pattern of aluminum combustion indicates a great similarity; it is likely that aluminum burns at this stage of the process. As it burns out, the size of the flame and, consequently, the burning drop decrease. Burnt particle long time glows.

Changing the size of the glow zone of a particle burning in accordance with the described mechanism is complex (Fig. 2.7). After ignition the value r St. /r 0 quickly (in -0.1 ms) reaches the maximum value (section ab). Further, during the main time of the first stage of combustion, the ratio r St/ r 0 remains constant (section bv). When the magnesium burnout ends, r cv/ r 0 is reduced to a minimum (point G), and then, with the beginning of aluminum combustion, it increases (section gd). Finally, but as aluminum burns out r St. /r 0 decreases monotonically (section de) before final value, corresponding to the size of the formed oxide.

Rice. 2.7.:

1 – alloy 30% Al + 70% Mg, air; 2 – alloy 30% A1 + 70% Mg, mixture 15% O2 + 85% Ar; 3 – alloy 50% A1 + 50% Mg, air

The mechanism and parameters of the combustion process of aluminum-magnesium alloys significantly depend on the composition of the alloy. With a decrease in magnesium content in the alloy, the size of the glow zone during the first stage of combustion and the duration of this stage decreases. When the magnesium content in the alloy is less than 30%, the process remains a two-stage process, but becomes intermittent. At the end of the first stage, the glow zone is reduced to the size of the particle itself, the combustion process stops, and aluminum burns out only after the particle is re-ignited. The particles that do not ignite again are hollow, porous oxide shells containing droplets of unburned aluminum inside.

The dependence of the burning time of particles on their initial diameter is expressed by the following empirical formulas:

Combustion of aluminum-magnesium alloys in mixtures of oxygen with argon, in water vapor and in carbon dioxide.

The nature of combustion of particles of aluminum-magnesium alloys in oxygen-argon mixtures is the same as in air. With a decrease in oxygen content, the size of the glow zone during magnesium burnout noticeably decreases. The dependence of the combustion time of particles of the 50% Al + 50% Mg alloy on the particle size and oxygen content in the mixture in volume percent is expressed by the formula

The combustion of alloys in water vapor is significantly different (Fig. 2.8). The oxide film formed during the first stage is destroyed by hydrogen, and the particle takes on the appearance of coral. The aluminum remaining in the coral ignites only 1-10 ms after the end of the first stage. Such intermittency of the process is typical for alloys of any composition.

Rice. 2.8. Combustion of aluminum-magnesium alloy particles (50:50) spherical(A) and wrong(b) forms in water vapor at normal atmospheric pressure:

1 – initial particle; 2 – particle before ignition; 3 – magnesium burnout; 4 – aluminum burnout; 5 – coral formed after the particle

When aluminum-magnesium alloys burn in carbon dioxide, only magnesium burns out of the particle, after which the combustion process stops.

Combustion of aluminum-magnesium alloys in a high-temperature flame

To study the combustion process of metal particles at high temperatures ah, under a particle mounted on the tip of a needle, a pressed tablet of mixtures of ammonium perchlorate and hexamine, having calculated combustion temperatures of 2500, 2700 and 3100 K, was burned.

Combustion of particles of aluminum-magnesium alloys under these conditions occurs, as a rule, with an explosion. The presence of an explosion is typical for particles of all compositions. As a result of the explosion, a significant luminescence zone is formed, which is a sign of the predominance of vapor-phase combustion. Photographs of a burning particle at the beginning of combustion (Fig. 2.9, A) show that heterogeneous reactions occur over the entire surface of the oxide shell. Due to the heat of heterogeneous reactions, rapid evaporation of the metal occurs (Fig. 2.9, b), promoting rupture of the oxide shell and splashing of the unevaporated drop (Fig. 2.9, V).

Rice. 2.9. Combustion of 95% Al alloy particle with 5% Mg in oxidizing flame (temperature 2700 K):

A– initial stage of combustion; b– stationary combustion; V- splitting up

According to B. G. Lrabey, S. E. Salibekov and Yu. V. Leninsky, crushing of particles of aluminum-magnesium alloys is caused by a very large difference in the boiling temperatures of magnesium and aluminum, as a result of which the boiling of magnesium when the particle is in a high-temperature zone is explosive. and leads to crushing of the remaining aluminum. A temperature of 2500 K is already sufficient for explosive combustion to occur, which is quite natural, since this temperature exceeds the boiling point of both components.

• Arabey B. G., Salibekov S. E., Levinsky Yu. V. Some characteristics of ignition and combustion of metal dust // Powder metallurgy. 1964. No. 3. P. 109-118.

Aluminum - flammable metal, atomic mass 26.98; density 2700 kg/m 3, melting point 660.1 °C; boiling point 2486 °C; calorific value -31087 kJ/kg. Aluminum shavings and dust can ignite under the local action of low-calorie ignition sources (match flame, spark, etc.). When aluminum powder, shavings, and foil interact with moisture, aluminum oxide is formed and a large amount of heat is released, leading to their spontaneous combustion when accumulated in heaps. This process is facilitated by the contamination of these materials with oils. The release of free hydrogen when aluminum dust interacts with moisture facilitates its explosion. The self-ignition temperature of a sample of aluminum dust with a dispersion of 27 microns is 520 °C; smoldering temperature 410 °C; lower concentration limit of flame propagation 40 g/m 3 ; maximum explosion pressure 1.3 MPa; pressure rise rate: average 24.1 MPa/s, maximum 68.6 MPa/s. The maximum oxygen concentration at which ignition of the air suspension by an electric spark is excluded is 3% of the volume. Settled dust is a fire hazard. Self-ignition temperature 320 °C. Aluminum easily reacts at room temperature with aqueous solutions of alkalis and ammonia, releasing hydrogen. Mixing aluminum powder with alkaline aqueous solution may cause an explosion. Reacts vigorously with many metalloids. Aluminum turnings burn, for example, in bromine, forming aluminum bromide. The interaction of aluminum with chlorine and bromine occurs at room temperature, and with iodine - when heated. When heated, aluminum combines with sulfur. If you add aluminum powder to the vapor of boiling sulfur, the aluminum will catch fire. Heavily ground aluminum reacts with halogenated hydrocarbons; the small amount of aluminum chloride present (formed during this reaction) acts as a catalyst, accelerating the reaction, in some cases leading to an explosion. This phenomenon is observed when aluminum powder is heated with methyl chloride, carbon tetrachloride, a mixture of chloroform and carbon tetrachloride to a temperature of about 150 °C.

Aluminum in the form of a compact material does not interact with carbon tetrachloride. Mixing aluminum dust with some chlorinated hydrocarbons and alcohol causes the mixture to spontaneously ignite. A mixture of aluminum powder with copper oxide, silver oxide, lead oxide and especially lead dioxide burns explosively. A mixture of ammonium nitrate, aluminum powder with coal or nitro compounds is an explosive. Extinguishing agents: dry sand, alumina, magnesite powder, asbestos blanket. The use of water and fire extinguishers is prohibited.

Aluminum does not occur in nature in its pure form, because it is very quickly oxidized by atmospheric oxygen to form strong oxide films that protect the surface from further interaction.

Typically, not pure aluminum is used as a structural material, but various alloys based on it, which are characterized by a combination of satisfactory strength, good ductility, very good weldability and corrosion resistance. In addition, these alloys are characterized by high vibration resistance.