Part I
1. The position of metals (M) in the Periodic Table of D. I. Mendeleev.
Conditional diagonal from B to At through elements of A groups: IV → V → VI. On the diagonal and above it are non-metals, and below it are metals.
Only M consist of B groups. In total, out of 110 elements, 88 elements are classified as metals.
Group IA are alkali metals.
Group IIA are alkaline earth metals.
2. Features of the structure of M atoms:
1) the number e in the outer layer of the atom is 1-3;
2) R atom – large sizes.
3. The relativity of dividing elements into M and NM (give examples):
1) gray tin – NM, white tin – M.
2) graphite is NM, but electrically conductive.
3) Cr, Zn, Al – M, but amphoteric.
4. A metal chemical bond is connection in metals and alloys between atom-ions through socialized e.
General scheme for the formation of a metallic bond:
5. Fill out the table “Structure and properties of metals.”
6. Write down the signs by which you can distinguish plates made:
a) from aluminum and copper – color, density, electrical and thermal conductivity
b) from lead and aluminum - color, density, melting point
c) from silver and graphite - color, shape, electrical conductivity.
7. Using the pictures, fill in the blanks to create a sequence: name of metal(s), properties(s), area(s) of application.
a) cast iron battery - cast iron, thermal conductivity, strength, wear resistance. In the economy, everyday life, metallurgy.
b) aluminum foil - aluminum, easy to roll out, plasticity, high electrical and thermal conductivity, corrosion resistance. In the food industry, production of alloys.
c) steel buttons and paper clips – steel, “soft” steel, elastic, easy to bend, does not rust, strong and hard. In all sectors of the national economy.
d) metal support - iron (steel), strong, solid, not exposed to the environment. In all sectors of the national economy.
e) domes – gold, inert, appearance. Used in construction - rolling, in jewelry.
f) thermometer - mercury (liquid metal), expands when heated, in medical thermometers. Obtaining alloys for gold mining. Lamps.
8. Fill out the “Classification of Metals” table.
9. Alloy is is a homogeneous metallic material consisting of a mixture of two or more chemical elements with a predominance of metallic components.
10. Ferrous alloys:
11. Fill out the table “Alloys and their components.”
12. Write the names of the alloys from which the objects shown in the pictures can be made.
a) steel
b) cupronickel
c) duralumin
d) bronze
e) bronze
e) cast iron
Part II
1. Metal atoms having in the outer layer:
a) 5e – Sb (antimony), Bi (bismuth)
b) 6e – Po (polonium)
Why?
They are located in 5 and 6 groups respectively
2. Metal atom having 3e in the outer layer, - boron.
Why?
It is located in group 3.
3. Fill out the table “Atomic structure and chemical bonding.”
4. Eliminate the “extra element.”
4) Si
5. Which of the following groups of elements contains only metals?
There is no right answer
6. What physical property is not common to all metals?
3) solid state of aggregation under standard conditions
7. Which statement is true?
4) metal atoms and metals - simple substances exhibit only reducing properties.
8. All elements of the main subgroups are metals if they are located in the Periodic Table below the diagonal:
3) boron - astatine
9. The number of electrons in the outer electronic level of a metal atom located in the main subgroup of the Periodic Table cannot be equal to:
1. Position of metals in the table of elements
Metals are located mainly in the left and lower parts of the PSHE. These include:
2. Structure of metal atoms
Metal atoms usually have 1-3 electrons in their outer energy level. Their atoms have a large radius and easily give up valence electrons, i.e. exhibit restorative properties.
3. Physical properties of metals
Changes in the electrical conductivity of a metal when it is heated and cooled
Metal connection - this is the bond that free electrons carry out between cations in a metal crystal lattice.
4. Obtaining metals
1. Reduction of metals from oxides with coal or carbon monoxide
Me x O y + C = CO 2 + Me or Me x O y + CO = CO 2 + Me
2. Roasting of sulfides followed by reduction
Stage 1 – Me x S y +O 2 =Me x O y +SO 2
Stage 2 -Me x O y + C = CO 2 + Me or Me x O y + CO = CO 2 + Me
3 Aluminothermy (reduction with a more active metal)
Me x O y + Al = Al 2 O 3 + Me
4. Hydrothermy - for the production of high purity metals
Me x O y + H 2 = H 2 O + Me
5. Reduction of metals by electric current (electrolysis)
1) Alkali and alkaline earth metals obtained in industry by electrolysis molten salts (chlorides):
2NaCl – melt, elect. current. → 2 Na + Cl 2
CaCl 2 – melt, elect. current.→ Ca+Cl2
hydroxide melts:
4NaOH – melt, elect. current.→ 4 Na + O 2 + 2 H 2 O
2) Aluminum in industry it is obtained by electrolysis aluminum oxide melt I in Na 3 AlF 6 cryolite (from bauxite):
2Al 2 O 3 – melt in cryolite, electr. current.→ 4 Al + 3 O 2
3) Electrolysis of aqueous salt solutions use to obtain metals of intermediate activity and inactive:
2CuSO 4 +2H 2 O – solution, elect. current.→ 2 Cu + O 2 + 2 H 2 SO 4
5. Finding metals in nature
The most common metal in the earth's crust is aluminum. Metals are found both in compounds and in free form.
1. Active – in the form of salts (sulfates, nitrates, chlorides, carbonates)
2. Moderate activity – in the form of oxides, sulfides ( Fe 3 O 4 , FeS 2 )
3. Noble – in free form ( Au, Pt, Ag)CHEMICAL PROPERTIES OF METALS
General chemical properties of metals are presented in the table:
ASSIGNMENT TASKS
No. 1. Finish equations practicable reactions, name the reaction products
Li+ H 2 O =
Cu + H2O =
Al + H 2 O =
Ba + H2O =
Mg + H2O =
Ca+HCl=
Na + H 2 SO 4 (K) =
Al + H 2 S=
Ca + H3PO4 =
HCl + Zn =
H 2 SO 4 (k)+ Cu=
H 2 S + Mg =
HCl + Cu =
HNO 3 (K)+ С u =
H2S+Pt=
H3PO4 + Fe =
HNO 3 (p)+ Na=
No. 2. Complete the CRM, arrange the coefficients using the electronic balance method, indicate the oxidizing agent (reducing agent):
Al + O 2 =
Li + H 2 O =
Na + HNO 3 (k) =
Mg + Pb(NO 3) 2 =
Ni + HCl =
Ag + H 2 SO 4 (k) =No. 3. Insert missing characters instead of dots (<, >or =)
Core charge | Li…Rb | Na…Al | Ca…K |
Number of energy levels | Li…Rb | Na…Al | Ca…K |
Number of outer electrons | Li…Rb | Na…Al | Ca…K |
Atomic radius | Li…Rb | Na…Al | Ca…K |
Restorative properties | Li…Rb | Na…Al | Ca…K |
No. 4. Complete the CRM, arrange the coefficients using the electronic balance method, indicate the oxidizing agent (reducing agent):
K+ O 2 =
Mg+ H 2 O =
Pb+ HNO 3 (p) =
Fe+ CuCl 2 =
Zn + H 2 SO 4 (p) =
Zn + H 2 SO 4 (k) =
No. 5. Solve test problems
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1.Select a group of elements that contains only metals: A) Al, As, P; B) Mg, Ca, Si; B ) K, Ca, Pb 2. Select a group that contains only simple substances - non-metals: A) K 2 O, SO 2, SiO 2; B) H 2, Cl 2, I 2; B )Ca, Ba, HCl; 3. Indicate the common features in the structure of the K and Li atoms: A) 2 electrons in the last electron layer; B) 1 electron in the last electron layer; C) the same number of electronic layers. 4. Calcium metal exhibits the following properties: A) oxidizing agent; B) reducing agent; C) an oxidizing agent or a reducing agent, depending on the conditions. 5. The metallic properties of sodium are weaker than those of - A) magnesium; B) potassium; C) lithium. 6. Inactive metals include: A) aluminum, copper, zinc; B) mercury, silver, copper; C) calcium, beryllium, silver. 7. What is the physical property is not common to all metals: A) electrical conductivity, B) thermal conductivity, B) solid state of aggregation under normal conditions, D) metallic shine |
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Part B. The answer to the tasks in this part is a set of letters that should be written down Match. With an increase in the ordinal number of an element in the main subgroup of group II of the Periodic System, the properties of the elements and the substances they form change as follows: |
Metals make up the majority of chemical elements. Each period of the periodic table (except for the 1st) of chemical elements begins with metals, and with increasing number of the period there are more and more of them. If in the 2nd period there are only 2 metals (lithium and beryllium), in the 3rd - 3 (sodium, magnesium, aluminum), then already in the 4th - 13, and in the 7th - 29.
Metal atoms are similar in the structure of the outer electron layer, which is formed by a small number of electrons (usually no more than three).
This statement can be illustrated by the examples of Na, aluminum A1 and zinc Zn. When drawing up diagrams of the structure of atoms, you can optionally create electronic formulas and give examples of the structure of elements of long periods, for example zinc.
Due to the fact that the electrons of the outer layer of metal atoms are weakly bound to the nucleus, they can be “given” to other particles, which is what happens in chemical reactions:
The property of metal atoms to give up electrons is their characteristic chemical property and indicates that metals exhibit reducing properties.
When characterizing the physical properties of metals, their general properties should be noted: electrical conductivity, thermal conductivity, metallic luster, plasticity, which are determined by a single type of chemical bond - metallic and metallic crystal lattice. Their feature is the presence of freely moving socialized electrons between ion-atoms located at the nodes of the crystal lattice.
When characterizing chemical properties, it is important to confirm the conclusion that in all reactions metals exhibit the properties of reducing agents, and to illustrate this by writing the reaction equations. Particular attention should be paid to the interaction of metals with acids and salt solutions, and it is necessary to refer to a number of metal voltages (a number of standard electrode potentials).
Examples of the interaction of metals with simple substances (non-metals):
With salts (Zn in the voltage series is to the left of Cu): Zn + CuC12 = ZnCl2 + Cu!
Thus, despite the wide variety of metals, they all have common physical and chemical properties, which is explained by the similarity in the structure of atoms and the structure of simple substances.
Introduction
Metals are simple substances that, under normal conditions, have characteristic properties: high electrical and thermal conductivity, the ability to reflect light well (which causes their shine and opacity), and the ability to take the desired shape under the influence of external forces (plasticity). There is another definition of metals - these are chemical elements characterized by the ability to donate external (valence) electrons.
Of all the known chemical elements, about 90 are metals. Most inorganic compounds are metal compounds.
There are several types of classification of metals. The most clear is the classification of metals in accordance with their position in the periodic table of chemical elements - chemical classification.
If in the “long” version of the periodic table we draw a straight line through the elements boron and astatine, then metals will be located to the left of this line, and non-metals to the right of it.
From the point of view of atomic structure, metals are divided into intransition and transition. Non-transition metals are located in the main subgroups of the periodic table and are characterized by the fact that in their atoms the electronic levels s and p are sequentially filled. Non-transition metals include 22 elements of the main subgroups a: Li, Na, K, Rb, Cs, Fr, Be, Mg, Ca, Sr, Ba, Ra, Al, Ga, In, Tl, Ge, Sn, Pb, Sb, Bi, Po.
Transition metals are located in secondary subgroups and are characterized by the filling of d- or f-electron levels. The d-elements include 37 metals of secondary subgroups b: Cu, Ag, Au, Zn, Cd, Hg, Sc, Y, La, Ac, Ti, Zr, Hf, Rf, V, Nb, Ta, Db, Cr, Mo , W, Sg, Mn, Tc, Re, Bh, Fe, Co, Ni, Ru, Rh, Pd, Os, Ir, Pt, Hs, Mt.
The f-elements include 14 lanthanides (Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Du, Ho, Er, Tm, Yb, Lu) and 14 actinides (Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr).
Among the transition metals, rare earth metals (Sc, Y, La and lanthanides), platinum metals (Ru, Rh, Pd, Os, Ir, Pt), transuranium metals (Np and elements with higher atomic mass) are also distinguished.
In addition to the chemical one, there is also, although not generally accepted, a long-established technical classification of metals. It is not as logical as the chemical one - it is based on one or another practically important feature of the metal. Iron and alloys based on it are classified as ferrous metals, all other metals are classified as non-ferrous. There are light (Li, Be, Mg, Ti, etc.) and heavy metals (Mn, Fe, Co, Ni, Cu, Zn, Cd, Hg, Sn, Pb, etc.), as well as groups of refractory metals (Ti, Zr , Hf, V, Nb, Ta, Cr, Mo, W, Re), precious (Ag, Au, platinum metals) and radioactive (U, Th, Np, Pu, etc.) metals. In geochemistry, trace (Ga, Ge, Hf, Re, etc.) and rare (Zr, Hf, Nb, Ta, Mo, W, Re, etc.) metals are also distinguished. As you can see, there are no clear boundaries between the groups.
Historical reference
Despite the fact that the life of human society without metals is impossible, no one knows exactly when and how people first began to use them. The most ancient writings that have reached us tell of primitive workshops in which metal was smelted and products were made from it. This means that man mastered metals before writing. When excavating ancient settlements, archaeologists find tools of labor and hunting that people used in those distant times - knives, axes, arrowheads, needles, fish hooks and much more. The older the settlements, the cruder and more primitive the products of human hands were. The most ancient metal products were found during excavations of settlements that existed about 8 thousand years ago. These were mainly jewelry made of gold and silver and arrowheads and spears made of copper.
The Greek word "metallon" originally meant mines, hence the term "metal". In ancient times, it was believed that there were only 7 metals: gold, silver, copper, tin, lead, iron and mercury. This number correlated with the number of planets known at that time - the Sun (gold), the Moon (silver), Venus (copper), Jupiter (tin), Saturn (lead), Mars (iron), Mercury (mercury) (see figure). According to alchemical ideas, metals originated in the bowels of the earth under the influence of the rays of the planets and gradually improved, turning into gold.
Man first mastered native metals - gold, silver, mercury. The first artificially produced metal was copper, then it was possible to master the production of an alloy of copper by nighting - bronze and only later - iron. In 1556, the book of the German metallurgist G. Agricola “On Mining and Metallurgy” was published in Germany - the first detailed guide to obtaining metals that has come down to us. True, at that time lead, tin and bismuth were still considered varieties of the same metal. In 1789, the French chemist A. Lavoisier, in his manual on chemistry, gave a list of simple substances, which included all the then known metals - antimony, silver, bismuth, cobalt, tin, iron, manganese, nickel, gold, platinum, lead, tungsten and zinc. As chemical research methods developed, the number of known metals began to increase rapidly. In the 18th century 14 metals were discovered in the 19th century. - 38, in the 20th century. - 25 metals. In the first half of the 19th century. Satellites of platinum were discovered, and alkali and alkaline earth metals were obtained by electrolysis. In the middle of the century, cesium, rubidium, thallium and indium were discovered by spectral analysis. The existence of metals predicted by D.I. Mendeleev on the basis of his periodic law (these are gallium, scandium and germanium) was brilliantly confirmed. Discovery of radioactivity at the end of the 19th century. led to the search for radioactive metals. Finally, by the method of nuclear transformations in the mid-20th century. radioactive metals that do not exist in nature, in particular transuranium elements, were obtained.
Physical and chemical properties of metals.
All metals are solids (except mercury, which is liquid under normal conditions); they differ from non-metals in a special type of bond (metallic bond). Valence electrons are weakly bound to a particular atom, and inside every metal there is a so-called electron gas. Most metals have a crystalline structure, and the metal can be thought of as a “rigid” crystal lattice of positive ions (cations). These electrons can more or less move around the metal. They compensate for the repulsive forces between cations and, thereby, bind them into a compact body.
All metals are highly electrically conductive (i.e., they are conductors as opposed to non-metals that are dielectrics), especially copper, silver, gold, mercury, and aluminum; The thermal conductivity of metals is also high. A distinctive property of many metals is their ductility (malleability), as a result of which they can be rolled into thin sheets (foil) and drawn into wire (tin, aluminum, etc.), however, there are also quite brittle metals (zinc, antimony, bismuth).
In industry, they often use not pure metals, but mixtures of them called alloys. In an alloy, the properties of one component usually successfully complement the properties of the other. Thus, copper has low hardness and is unsuitable for the manufacture of machine parts, while alloys of copper and zinc, called brass, are already quite hard and are widely used in mechanical engineering. Aluminum has good ductility and sufficient lightness (low density), but is too soft. Based on it, an alloy of ayuralum (duralumin) containing copper, magnesium and manganese is prepared. Duralumin, without losing the properties of its aluminum, acquires high hardness and is therefore used in aircraft technology. Alloys of iron with carbon (and additives of other metals) are the well-known cast iron and steel.
Metals vary greatly in density: for lithium it is almost half that of water (0.53 g/cm3), and for osmium it is more than 20 times higher (22.61 g/cm3). Metals also differ in hardness. The softest are alkali metals; they can be easily cut with a knife; The hardest metal - chromium - cuts glass. There is a great difference in the melting points of metals: mercury is a liquid under normal conditions, cesium and gallium melt at the temperature of the human body, and the most refractory metal, tungsten, has a melting point of 3380 ° C. Metals whose melting point is above 1000 °C are classified as refractory metals, and those below are referred to as fusible metals. At high temperatures, metals are capable of emitting electrons, which are used in electronics and thermoelectric generators to directly convert thermal energy into electrical energy. Iron, cobalt, nickel and gadolinium, after placing them in a magnetic field, are able to permanently maintain a state of magnetization.
Metals also have some chemical properties. Metal atoms relatively easily give up valence electrons and become positively charged ions. Therefore, metals are reducing agents. This, in fact, is their main and most general chemical property.
Obviously, metals as reducing agents will react with various oxidizing agents, which may include simple substances, acids, salts of less active metals and some other compounds. Compounds of metals with halogens are called halides, with sulfur - sulfides, with nitrogen - nitrides, with phosphorus - phosphides, with carbon - carbides, with silicon - silicides, with boron - borides, with hydrogen - hydrides, etc. Many of these compounds found important applications in new technology. For example, metal borides are used in radio electronics, as well as in nuclear engineering as materials for regulating and protecting against neutron radiation.
Under the influence of concentrated oxidizing acids, a stable oxide film is also formed on some metals. This phenomenon is called passivation. Thus, in concentrated sulfuric acid, metals such as Be, Bi, Co, Fe, Mg, and Nb are passivated (and do not react with it), and in concentrated nitric acid the metals Al, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb, Th and U.
The further to the left a metal is located in this row, the greater the reducing properties it has, i.e., it is easier to oxidize and go into solution as a cation, but it is more difficult to reduce from the cation to a free state.
One non-metal, hydrogen, is placed in the voltage series, since this makes it possible to determine whether this metal will react with acids - non-oxidizing agents in an aqueous solution (more precisely, be oxidized by hydrogen cations H +). For example, zinc reacts with hydrochloric acid, since in the series of voltages it is to the left (before) hydrogen. On the contrary, silver is not transferred into solution by hydrochloric acid, since it is in the voltage series to the right (after) hydrogen. Metals behave similarly in dilute sulfuric acid. Metals in the voltage series after hydrogen are called noble (Ag, Pt, Au, etc.)
An undesirable chemical property of metals is their electrochemical corrosion, i.e. the active destruction (oxidation) of the metal upon contact with water and under the influence of oxygen dissolved in it (oxygen corrosion). For example, corrosion of iron products in water is widely known.
Particularly corrosive-dangerous can be the place of contact of two dissimilar metals - contact corrosion. A galvanic couple occurs between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the voltage series (Fe), to the less active metal (Sn, Cu), and the more active metal is destroyed (corroded).
It is because of this that the tinned surface of cans (iron coated with tin) rusts when stored in a humid atmosphere and handled carelessly (iron quickly collapses after even a small scratch appears, allowing the iron to come into contact with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, since even if there are scratches, it is not the iron that corrodes, but the zinc (a more active metal than iron).
The corrosion resistance of a given metal increases when it is coated with a more active metal or when they are fused; Thus, coating iron with chromium or making alloys of iron and chromium eliminates corrosion of iron. Chromed iron and steels containing chromium (stainless steels) have high corrosion resistance.
General methods for obtaining metals:
Electrometallurgy, i.e., the production of metals by electrolysis of melts (for the most active metals) or solutions of their salts;
Pyrometallurgy, i.e., the reduction of metals from their ores at high temperatures (for example, the production of iron using the blast furnace process);
Hydrometallurgy, i.e., the separation of metals from solutions of their salts by more active metals (for example, the production of copper from a CuSO 4 solution by replacing zinc, iron
or aluminum).
In nature, metals are sometimes found in free form, for example native mercury, silver and gold, and more often in the form of compounds (metal ores). The most active metals, of course, are present in the earth's crust only in bound form.
Lithium (from the Greek Lithos - stone), Li, a chemical element of subgroup Ia of the periodic table; atomic number 3, atomic mass 6, 941; refers to alkali metals.
The lithium content in the earth's crust is 6.5-10 -3% by mass. It is found in more than 150 minerals, of which about 30 are lithium minerals. The main minerals are spodumene LiAl, lepidolite KLi 1.5 Al 1.5 (F.0H) 2 and petalite (LiNa). The composition of these minerals is complex; many of them belong to the class of aluminosilicates, which are very common in the earth’s crust. Promising sources of raw materials for lithium production are brines (brine) of salt-bearing deposits and groundwater. The largest deposits of lithium compounds are located in Canada, the USA, Chile, Zimbabwe, Brazil, Namibia and Russia.
Interestingly, the mineral spodumene occurs in nature in the form of large crystals weighing several tons. A needle-shaped crystal 16 m long and weighing 100 tons was found at the Etta mine in the USA.
The first information about lithium dates back to 1817. The Swedish chemist A. Arfvedson, while analyzing the mineral petalite, discovered an unknown alkali in it. Arfvedson's teacher J. Berzelius gave it the name “lithion” (from the Greek litheos - stone), because unlike potassium and sodium hydroxides, which were obtained from plant ash, a new alkali was discovered in the mineral. He also named the metal that is the “base” of this alkali, lithium. In 1818, the English chemist and physicist G. Davy obtained lithium by electrolysis of LiOH hydroxide.
Properties. Lithium is a silvery-white metal; m.p. 180.54 °C, bp. 1340 "C; the lightest of all metals, its density is 0.534 g/cm - it is 5 times lighter than aluminum and almost half as light as water. Lithium is soft and ductile. Lithium compounds paint the flame a beautiful carmine red color. This very sensitive method is used in qualitative analysis for lithium detection.
Configuration of the outer electron layer of the lithium atom 2s 1 (s-element). In compounds it exhibits an oxidation state of +1.
Lithium is first in the electrochemical series of voltages and displaces hydrogen not only from acids, but also from water. However, many of lithium's chemical reactions are less vigorous than those of other alkali metals.
Lithium practically does not react with air components in the complete absence of moisture at room temperature. When heated in air above 200 °C, Li 2 O oxide forms as the main product (only traces of Li 2 O 2 peroxide are present). In humid air it produces predominantly Li 3 N nitride; at air humidity above 80% it produces LiOH hydroxide and Li 2 CO 3 carbonate. Lithium nitride can also be obtained by heating the metal in a stream of nitrogen (lithium is one of the few elements that directly combines with nitrogen): 6Li + N 2 = 2Li 3 N
Lithium easily alloys with almost all metals and is highly soluble in mercury. Directly combines with halogens (with iodine when heated). At 500 °C it reacts with hydrogen, forming LiH hydride, when interacting with water - LiOH hydroxide, with dilute acids - lithium salts, with ammonia - LiNH2 amide, for example:
2Li + H 2 = 2LiH
2Li + 2H 2 O = 2LiOH + H 2
2Li + 2НF = 2LiF + Н 2
2Li + 2NH 3 = 2LiNH 2 + H 2
LiH hydride - colorless crystals; used in various fields of chemistry as a reducing agent. When interacting with water, it releases a large amount of hydrogen (2820 l of H2 are obtained from 1 kg of LiH):
LiH + H 2 O = LiOH + H 2
This makes it possible to use LiH as a source of hydrogen for filling balloons and rescue equipment (inflatable boats, belts, etc.), as well as as a kind of “warehouse” for storing and transporting flammable hydrogen (in this case, it is necessary to protect LiH from the slightest traces of moisture).
Mixed lithium hydrides are widely used in organic synthesis, for example lithium aluminum hydride LiAlH 4 - a selective reducing agent. It is obtained by reacting LiH with aluminum chloride AlCl3
LiOH hydroxide is a strong base (alkali), its aqueous solutions destroy glass and porcelain; Nickel, silver and gold are resistant to it. LiOH is used as an additive to the electrolyte of alkaline batteries, which increases their service life by 2-3 times and capacity by 20%. Based on LiOH and organic acids (especially stearic and palmitic), frost- and heat-resistant greases (lithols) are produced to protect metals from corrosion in the temperature range from -40 to +130 "C.
Lithium hydroxide is also used as a carbon dioxide absorbent in gas masks, submarines, airplanes and spacecraft.
Receipt and application. The raw material for lithium production is its salts, which are extracted from minerals. Depending on the composition, minerals are decomposed with sulfuric acid H 2 SO 4 (acid method) or by sintering with calcium oxide CaO and its carbonate CaCO3 (alkaline method), with potassium sulfate K 2 SO 4 (salt method), with calcium carbonate and its chloride CaCl (alkali-salt method). Using the acid method, a solution of Li 2 SO 4 sulfate is obtained [the latter is freed from impurities by treating with calcium hydroxide Ca(OH) 2 and soda Na 2 Co 3 ]. The cake formed by other methods of mineral decomposition is leached with water; in this case, with the alkaline method, LiOH goes into the solution, with the salt method - Li 2 SO 4, with the alkali-salt method - LiCl. All these methods, except alkaline, provide for the production of the finished product in the form of Li 2 CO 3 carbonate. which is used directly or as a source for the synthesis of other lithium compounds.
Lithium metal is produced by electrolysis of a molten mixture of LiCl and potassium chloride KCl or barium chloride BaCl 2 with further purification from impurities.
Interest in lithium is huge. This is due, first of all, to the fact that it is a source of industrial production of tritium (a heavy hydrogen nuclide), which is the main component of a hydrogen bomb and the main fuel for thermonuclear reactors. The thermonuclear reaction takes place between the nuclide 6 Li and neutrons (neutral particles with mass number 1); reaction products - tritium 3 H and helium 4 He:
6 3 Li + 1 0 n= 3 1 H + 4 2 He
Large quantities of lithium are used in metallurgy. An alloy of magnesium with 10% lithium is stronger and lighter than magnesium itself. Aluminum and lithium alloys - scleron and aeron, containing only 0.1% lithium, in addition to lightness, have high strength, ductility, and increased resistance to corrosion; they are used in aviation. The addition of 0.04% lithium to lead-calcium bearing alloys increases their hardness and reduces the coefficient of friction.
Lithium halides and carbonate are used in the production of optical, acid-resistant and other special glasses, as well as heat-resistant porcelain and ceramics, various glazes and enamels.
Fine lithium particles cause chemical burns to wet skin and eyes. Lithium salts irritate the skin. When working with lithium hydroxide, precautions must be taken as when working with sodium and potassium hydroxides.
Sodium (from Arabic, natrun, Greek nitron - natural soda, a chemical element of subgroup Ia of the periodic system; atomic number 11, atomic weight 22.98977; belongs to the alkali metals. In nature, it is found in the form of one stable nuclide 23 Na.
Even in ancient times, sodium compounds were known - table salt (sodium chloride) NaCl, caustic alkali (sodium hydroxide) NaOH and soda (sodium carbonate) Na 2 CO3. The last substance was called “nitron” by the ancient Greeks; This is where the modern name of the metal comes from - “sodium”. However, in the UK, USA, Italy, France, the word sodium (from the Spanish word “soda”, which has the same meaning as in Russian) is retained.
The production of sodium (and potassium) was first reported by the English chemist and physicist G. Davy at a meeting of the Royal Society in London in 1807. He managed to decompose the caustic alkalis KOH and NaOH using electric current and isolate previously unknown metals with extraordinary properties. These metals oxidized very quickly in air and floated on the surface of the water, releasing hydrogen from it.
Prevalence in nature. Sodium is one of the most common elements in nature. Its content in the earth's crust is 2.64% by weight. In the hydrosphere it is contained in the form of soluble salts in an amount of about 2.9% (with a total concentration of salts in sea water of 3.5-3.7%). The presence of sodium has been established in the solar atmosphere and interstellar space. In nature, sodium is found only in the form of salts. The most important minerals are halite (rock salt) NaCl, mirabilite (Glauber's salt) Na 2 SO 4 *10H 2 O, thenardite Na 2 SO 4, Chelyan nitrate NaNO 3, natural silicates, for example albite Na, nepheline Na
Russia is exceptionally rich in rock salt deposits (for example, Solikamsk, Usolye-Sibirskoye, etc.), large deposits of the mineral trona in Siberia.
Properties. Sodium is a silver-white, fusible metal, mp. 97.86 °C, bp. 883.15 °C. This is one of the lightest metals - it is lighter than water, density 0.99 g/cm 3 at 19.7 ° C). Sodium and its compounds color the burner flame yellow. This reaction is so sensitive that it reveals the presence of the slightest traces of sodium everywhere (for example, in indoor or outdoor dust).
Sodium is one of the most active elements of the periodic table. The outer electron layer of the sodium atom contains one electron (3s 1 configuration, sodium is an s-element). Sodium easily gives up its only valence electron and therefore always exhibits an oxidation state of +1 in its compounds.
In air, sodium is actively oxidized, forming Na 2 O oxide or Na 2 O 2 peroxide, depending on conditions. Therefore, sodium is stored under a layer of kerosene or mineral oil. Reacts vigorously with water, displacing hydrogen:
2Na + H 2 0 = 2NaOH + H 2
This reaction occurs even with ice at a temperature of -80 ° C, and with warm water or at the contact surface it occurs with an explosion (it’s not for nothing that they say: “If you don’t want to become a freak, don’t throw sodium into the water”).
Sodium reacts directly with all non-metals: at 200 ° C it begins to absorb hydrogen, forming a very hygroscopic hydride NaH; with nitrogen in an electric discharge produces Na 3 N nitride or NaN 3 azide; in an atmosphere of fluorine it ignites; in chlorine burns at temperature; reacts with bromine only when heated:
2Na + H 2 = 2NaH
6Na + N 2 = 2Na 3 N or 2Na + 3Na 2 = 2NaN 3
2Na+ С1 2 = 2NaСl
At 800-900 °C, sodium combines with carbon, forming Na 2 C 2 carbide; when triturated with sulfur it gives Na 2 S sulfide and a mixture of polysulfides (Na 2 S 3 and Na 2 S 4)
Sodium easily dissolves in liquid ammonia, the resulting blue solution has metallic conductivity, with gaseous ammonia at 300-400 °C or in the presence of a catalyst when cooled to -30 °C gives the amide NaNH 2.
Sodium forms compounds with other metals (intermetallic compounds), such as silver, gold, cadmium, lead, potassium and some others. With mercury it produces amalgams NaHg 2, NaHg 4, etc. The most important are liquid amalgams, which are formed by the gradual introduction of sodium into mercury located under a layer of kerosene or mineral oil.
Sodium forms salts with dilute acids.
Receipt and application. The main method for producing sodium is electrolysis of molten table salt. In this case, chlorine is released at the anode, and sodium is released at the cathode. To reduce the melting point of the electrolyte, other salts are added to table salt: KCl, NaF, CaCl 2. Electrolysis is carried out in electrolyzers with a diaphragm; anodes are made of graphite, cathodes are made of copper or iron.
Sodium can be obtained by electrolysis of the molten NaOH hydroxide, and small quantities can be obtained by the decomposition of NaN 3 azide.
Metallic sodium is used to restore pure metals from their compounds - potassium (from KOH), titanium (from TiCl 4), etc. An alloy of sodium with potassium is a coolant for nuclear reactors, since alkali metals do not absorb neutrons well and therefore do not prevent the fission of uranium nuclei. Sodium vapor, which has a bright yellow glow, is used to fill gas-discharge lamps used to illuminate highways, marinas, train stations, etc. Sodium is used in medicine: the artificially obtained nuclide 24 Na is used for radiological treatment of some forms of leukemia and for diagnostic purposes.
The use of sodium compounds is much more extensive.
Peroxide Na 2 O 2 - colorless crystals, yellow technical product. When heated to 311-400 °C, it begins to release oxygen, and at 540 °C it rapidly decomposes. A strong oxidizing agent, due to which it is used for bleaching fabrics and other materials. In air, it absorbs CO 2, releasing oxygen and forming carbonate 2Na 2 O 2 +2CO 2 = 2Na 2 Co 3 +O 2). The use of Na 2 O 2 for air regeneration in enclosed spaces and insulating breathing devices (submarines, insulating gas masks, etc.) is based on this property.
NaOH hydroxide; the outdated name is caustic soda, the technical name is caustic soda (from the Latin caustic - caustic, burning); one of the strongest foundations. The technical product, in addition to NaOH, contains impurities (up to 3% Na 2 CO3 and up to 1.5% NaCl). A large amount of NaOH is used for the preparation of electrolytes for alkaline batteries, the production of paper, soap, paints, cellulose, and is used to purify petroleum and oils.
Among sodium salts, Na 2 CrO 4 chromate is used - in the production of dyes, as a mordant for dyeing fabrics and as a tanning agent in the leather industry; Na 2 SO 3 sulfite is a component of fixers and developers in photography; hydrosulfite NaHSO 3 - bleaches fabrics, natural fibers, used for canning fruits, vegetables and plant feed; Na 2 S 2 O 3 thiosulfate - for removing chlorine when bleaching fabrics, as a fixative in photography, an antidote for poisoning with mercury compounds, arsenic, etc., an anti-inflammatory agent; chlorate NaClO 3 - an oxidizing agent in various pyrotechnic compositions; Na 5 P 3 O 10 triphosphate is an additive to synthetic detergents for water softening.
Sodium, NaOH and its solutions cause severe burns to the skin and mucous membranes.
In appearance and properties, potassium is similar to sodium, but more reactive. Reacts vigorously with water and causes hydrogen to ignite. It burns in air, forming orange superoxide CO2. At room temperature it reacts with halogens, and with moderate heating - with hydrogen and sulfur. In humid air it quickly becomes covered with a layer of KOH. Store potassium under a layer of gasoline or kerosene.
The greatest practical applications are potassium compounds - KOH hydroxide, KNO 3 nitrate and K 2 CO 3 carbonate.
Potassium hydroxide KOH (technical name - caustic potassium) - white crystals that spread in moist air and absorb carbon dioxide (K 2 CO 3 and KHCO 3 are formed). Very soluble in water with a high exo-effect. The aqueous solution is highly alkaline.
Potassium hydroxide is produced by electrolysis of a KCl solution (similar to the production of NaOH). The initial potassium chloride KCl is obtained from natural raw materials (minerals sylvite KCl and carnallite KMgCl 3 6H 2 0). KOH is used for the synthesis of various potassium salts, liquid soap, dyes, as an electrolyte in batteries.
Potassium nitrate KNO 3 (mineral potassium nitrate) - white crystals, very bitter in taste, low melting point (t pl = 339 ° C). Highly soluble in water (no hydrolysis). When heated above the melting point, it decomposes into potassium nitrite KNO 2 and oxygen O 2 and exhibits strong oxidizing properties. Sulfur and charcoal ignite upon contact with molten KNO 3, and the C + S mixture explodes (combustion of “black powder”):
2КNO 3 + ЗС (coal) + S=N 2 + 3CO 2 + K 2 S
Potassium nitrate is used in the production of glass and mineral fertilizers.
Potassium carbonate K 2 CO 3 (technical name - potash) is a white hygroscopic powder. It is very soluble in water, strongly hydrolyzes at the anion and creates an alkaline environment in solution. Used in making glass and soap.
The production of K 2 CO 3 is based on the reactions:
K 2 SO 4 + Ca(OH) 2 + 2CO = 2K(HCOO) + CaSO 4
2К(НСОО) + O 2 = К 2 С0 3 + Н 2 0 + С0 2
Potassium sulfate from natural raw materials (minerals kainite KMg(SO 4)Cl ZN 2 0 and schoenite K 2 Mg(SO 4) 2 * 6H 2 0) is heated with slaked lime Ca(OH) 2 in a CO atmosphere (under a pressure of 15 atm) , potassium formate K(HCOO) is obtained, which is calcined in a stream of air.
Potassium is a vital element for plants and animals. Potassium fertilizers are potassium salts, both natural and their processed products (KCl, K 2 SO 4, KNO 3); high content of potassium salts in plant ash.
Potassium is the ninth most abundant element in the earth's crust. Contained only in bound form in minerals, sea water (up to 0.38 g of K + ions in 1 l), plants and living organisms (inside cells). The human body contains = 175 g of potassium, the daily requirement reaches ~4 g. The radioactive isotope 40 K (an admixture to the predominant stable isotope 39 K) decays very slowly (half-life 1 10 9 years), it, along with the isotopes 238 U and 232 Th, makes a large contribution to the geothermal reserve of our planet (internal heat of the earth's interior) .
From (lat. Cuprum), Cu, a chemical element of subgroup 16 of the periodic table; atomic number 29, atomic mass 63.546 belongs to the transition metals. Natural copper is a mixture of nuclides with mass numbers 63 (69.1%) and 65 (30.9%).
Prevalence in nature. The average copper content in the earth's crust is 4.7-10~ 3% by mass.
In the earth's crust, copper is found both in the form of nuggets and in the form of various minerals. Copper nuggets, sometimes of considerable size, are covered with a green or blue coating and are unusually heavy compared to stone; the largest nugget weighing about 420 tons was found in the USA in the Great Lakes region (picture). The vast majority of copper is present in rocks in the form of compounds. More than 250 minerals containing copper are known. Of industrial importance are: chalcopyrite (copper pyrite) CuFeS 2, covellite (copper indigo) Cu 2 S, chalcocite (copper luster) Cu 2 S, cuprite Cu 2 O, malachite CuCO3*Cu(OH) 2 and azurite 2CuCO3*Cu(OH ) 2 . Almost all copper minerals are brightly and beautifully colored, for example, chalcopyrite has a gold shimmer, copper luster has a steel-blue color, azurite has a deep blue with a glassy luster, and pieces of covellite shimmer in all the colors of the rainbow. Many of the copper minerals are semi-precious and precious stones; Malachite and turquoise CuA1 6 (PO 4) 4 (OH) 8 *5H 2 O are very highly valued. The largest deposits of copper ores are located in North and South America (mainly in the USA, Canada, Chile, Peru, Mexico), Africa (Zambia, South Africa), Asia (Iran, Philippines, Japan). In Russia, there are copper ore deposits in the Urals and Altai.
Copper ores are usually polymetallic: in addition to copper, they contain Fe, Zn, Pb, Sn, Ni, Mo, Au, Ag, Se, platinum metals, etc.
Historical reference. Copper has been known since time immemorial and is one of the “magnificent seven” of the oldest metals used by mankind - gold, silver, copper, iron, tin, lead and mercury. According to archaeological data, copper was known to people already 6,000 years ago. It turned out to be the first metal that replaced stone for ancient man in primitive tools. This was the beginning of the so-called. the Copper Age, which lasted about two thousand years. Axes, knives, maces, and household items were forged from copper and then smelted. According to legend, the ancient blacksmith god Hephaestus forged a shield of pure copper for the invincible Achilles. The stones for the 147-meter Cheops pyramid were also quarried and hewn with copper tools.
The ancient Romans exported copper ore from the island of Cyprus, hence the Latin name for copper - “cuprum”. The Russian name "copper" is apparently related to the word "smida", which in ancient times meant "metal".
In the ores mined on the Sinai Peninsula, ores sometimes contained an admixture of tin, which led to the discovery of an alloy of copper and tin - bronze. Bronze turned out to be more fusible and harder than copper itself. The discovery of bronze marked the beginning of the long Bronze Age (4th-1st millennium BC).
Properties. Copper is a red metal. T.pl. 1083 "C, boiling point 2567 °C, density 8.92 g/cm. This is a ductile, malleable metal; leaves 5 times thinner than tissue paper can be rolled out of it. Copper reflects light well, conducts heat and electricity well, second only to silver
The configuration of the outer electronic layers of the copper atom is 3d 10 4s 1 (d-element). Although copper and alkali metals are in the same group I, their behavior and properties are very different. Copper is similar to alkali metals only in its ability to form monovalent cations. When forming compounds, a copper atom can lose not only its outer s-electron, but one or two d-electrons from the previous layer, exhibiting a higher oxidation state. For copper, the oxidation state +2 is more typical than +1.
Metallic copper is low-active and stable in dry and clean air. In humid air containing CO 2, a greenish film of Cu(OH) 2* CuCO3, called patina, forms on its surface. Patina gives products made from copper and its alloys a beautiful “antique” look; a continuous coating of patina, in addition, protects the metal from further destruction. When copper is heated in pure and dry oxygen, the formation of black oxide CuO occurs; heating above 375°C leads to red Cu 2 O oxide. At normal temperatures, copper oxides are stable in air.
In the series of voltages, copper is to the right of hydrogen, and therefore it does not displace hydrogen from water and does not in oxygen-free acids. Copper can dissolve in acids only when it is simultaneously oxidized, for example in nitric acid or concentrated sulfuric acid:
3Сu + 8НNO 3 = 3Сu(NO 3) 2 + 2NO + 4Н 2 O
Cu + 2H 2 S0 4 = CuSO 4 + SO 2 + 2H 2 O
Fluorine, chlorine and bromine react with copper to form the corresponding dihalides, for example:
Cu + Cl 2 = CuCl 2
When heated copper powder reacts with iodine, Cu(I) iodide or copper monoiodide is obtained:
2Сu +I 2 = 2СuI
Copper burns in sulfur vapor, forming monosulfide CuS. It does not interact with hydrogen under normal conditions. However, if copper samples contain microimpurities of Cu 2 O oxide, then in an atmosphere containing hydrogen, methane or carbon monoxide, copper oxide is reduced to metal:
Сu 2 O+ Н 2 = 2Сu + Н 2 O
Cu 2 O+ CO = 2Cu + CO 2
The released water vapor and CO 2 cause cracks to appear, which sharply worsens the mechanical properties of the metal (“hydrogen disease”). Monivalent copper salts - CuCl chloride, Cu 2 SO3 sulfite, Cu 2 S sulfide and others - as a rule, are poorly soluble in water. For divalent copper, there are salts of almost all known acids; the most important of them are CuSO 4 sulfate, CuCl 2 chloride, Cu(NO3) nitrate. All of them are highly soluble in water, and when separated from it they form crystalline hydrates, for example CuCl 2 *2H 2 O, Cu(NO3) 2 *6H 2 O, Cu80 4 -5H 2 0. The color of the salts is from green to blue, because the Cu ion in water is hydrated and is in the form of a blue aqua ion [Cu(H 2 O) 6 ] 2+, which determines color of solutions of divalent copper salts.
One of the most important copper salts - sulfate - is obtained by dissolving the metal in heated dilute sulfuric acid while blowing air:
2Сu + 2Н 2 SO 4 + O 2 = 2СuSO 4 + 2Н 2 O
Anhydrous sulfate is colorless; adding water, it turns into copper sulfate CuSO 4 -5H 2 O - azure blue transparent crystals. Due to the property of copper sulfate to change color when moistened, it is used to detect traces of water in alcohols, ethers, gasolines, etc.
When divalent copper salt interacts with alkali, a voluminous blue precipitate is formed - Cu(OH) 2 hydroxide. It is amphoteric: it dissolves in concentrated alkali to form a salt in which copper is present as an anion, for example:
Cu(OH) 2 + 2KOH = K 2 [Cu(OH) 4 ]
Unlike alkali metals, copper is characterized by a tendency to form complexes - Cu and Cu 2+ ions in water can form complex ions with anions (Cl -, CN -), neutral molecules (NH 3) and some organic compounds. These complexes are usually brightly colored and highly soluble in water.
Receipt and application. Back in the 19th century. copper was smelted from ores containing at least 15% metal. Currently, rich copper ores are practically exhausted, so copper ch. arr. obtained from sulfide ores containing only 1-7% copper. Metal smelting is a long and multi-stage process.
After flotation treatment of the original ore, the concentrate containing iron and copper sulfides is placed in reverberatory copper smelting furnaces heated to 1200 °C. The concentrate melts, forming the so-called. matte containing molten copper, iron and sulfur, as well as solid silicate slags that float to the surface. The smelted matte in the form of CuS contains about 30% copper, the rest is iron sulfide and sulfur. The next stage is the transformation of matte into the so-called. blister copper, which is carried out in horizontal converter furnaces purged with oxygen. FeS is oxidized first; To bind the resulting iron oxide, quartz is added to the converter - this forms an easily separated silicate slag. Then CuS is oxidized, turning into metallic copper, and SO 2 is released:
CuS + O 2 = Cu + SO 2
After removing SO 2 with air, the blister copper remaining in the converter, containing 97-99% copper, is poured into molds and then subjected to electrolytic purification. To do this, blister copper ingots, shaped like thick boards, are suspended in electrolysis baths containing a solution of copper sulfate with the addition of H 2 SO 4. Thin sheets of pure copper are also suspended in the same baths. They serve as cathodes, and blister copper castings serve as anodes. During the passage of current, copper dissolves at the anode, and its release occurs at the cathode:
Cu - 2е = Cu 2+
Сu 2+ + 2е = Сu
Impurities, including silver, gold, platinum, fall to the bottom of the bath in the form of a silt-like mass (sludge). The recovery of precious metals from the sludge usually pays for this entire energy-intensive process. After such refining, the resulting metal contains 98-99% copper.
Copper has long been used in construction: the ancient Egyptians built copper water pipes; the roofs of medieval castles and churches were covered with copper sheets, for example the famous royal castle in Elsinore (Denmark) was covered with copper roofing. Coins and jewelry were made from copper. Due to its low electrical resistance, copper is the main metal in electrical engineering: more than half of all copper produced is used in the production of electrical wires for high-voltage transmissions and low-current cables. Even insignificant impurities in copper lead to an increase in its electrical resistance and large losses of electricity.
High thermal conductivity and corrosion resistance make it possible to manufacture copper parts for heat exchangers, refrigerators, vacuum devices, pipelines for pumping oils and fuels, etc. Copper is also widely used in electroplating when applying protective coatings to steel products. So, for example, when nickel or chrome plating steel objects, copper is pre-deposited on them; in this case, the protective coating lasts longer and is more effective. Copper is also used in electroplating (i.e., when replicating products by obtaining a mirror image), for example, in the manufacture of metal matrices for printing banknotes and reproducing sculptural products.
A significant amount of copper is spent on the manufacture of alloys, which it forms with many metals. The main alloys of copper are generally divided into three groups: bronzes (alloys with tin and other metals other than zinc and nickel), brasses (alloys with zinc) and cupro-nickel alloys. There are separate articles about bronzes and brasses in the encyclopedia. The most famous copper-nickel alloys are cupronickel, nickel silver, constantan, manganin; they all contain up to 30-40% nickel and various alloying additives. These alloys are used in shipbuilding, for the manufacture of parts operating at elevated temperatures, in electrical devices, as well as for household metal products instead of silver (cutlery).
Copper compounds have been and are found in various applications. Cupric oxide and sulfate are used for the manufacture of certain types of artificial fiber and for the production of other copper compounds; CuO and Cu 2 O are used for the production of glass and enamels; Сu(NO3) 2 - calico printing; CuCl 2 is a component of mineral paints, a catalyst. Mineral paints containing copper have been known since ancient times; Thus, an analysis of the ancient frescoes of Pompeii and wall paintings in Rus' showed that the composition of the paints included the main copper acetate Cu(OH) 2 * (CH3COO) 2 Cu 2, which served as a bright green paint, called verdigris in Russia .
Copper belongs to the so-called. bioelements necessary for the normal development of plants and animals. In the absence or deficiency of copper in plant tissues, the chlorophyll content decreases, the leaves turn yellow, the plants stop bearing fruit and may die. Therefore, many copper salts are included in copper fertilizers, for example, copper sulfate, copper-potassium fertilizers (copper sulfate mixed with KSD). Copper salts are also used to combat plant diseases. For more than a hundred years, Bordeaux mixture has been used for this purpose, containing basic copper sulfate [Cu(OH) 2 ]3CuSO 4 ; get it by the reaction:
4CuSO 4 + 3Ca(OH) 2 = CuSO 4 * 3Cu(OH) 2 + 3СаSO 4
The gelatinous sediment of this salt covers the leaves well and stays on them for a long time, protecting the plant. Cu 2 O, copper chloroxide 3Cu(OH) 2 *CuCl 2, as well as copper phosphate, borate and copper arsenate have a similar property.
In the human body, copper is part of some enzymes and is involved in the processes of hematopoiesis and enzymatic oxidation; The average copper content in human blood is about 0.001 mg/l. In the organisms of lower animals there is much more copper, for example hemocyanin - the blood pigment of mollusks and crustaceans - contains up to 0.26% copper. The average copper content in living organisms is 2-10 - 4% by weight.
For humans, copper compounds are mostly toxic. Despite the fact that copper is included in some pharmaceuticals, if it enters the stomach with water or food in large quantities, it can cause severe poisoning. People who work for a long time in the smelting of copper and its alloys often develop “copper fever” - the temperature rises, pain occurs in the stomach, and the vital activity of the lungs decreases. If copper salts get into the stomach, before the doctor arrives, you must urgently rinse it and take a diuretic.
Conclusion.
Metals serve as the main structural materials in mechanical engineering and instrument making. They all have common so-called metallic properties, but each element exhibits them in accordance with its position in the periodic system of D.I. Mendeleev, i.e., in accordance with the structural features of its atom.
Metals actively interact with elementary oxidizing agents with high electronegativity (halogens, oxygen, sulfur, etc.) and therefore, when considering the general properties of metal elements, it is necessary to take into account their chemical activity towards non-metals, the types of their compounds and the forms of chemical bonds, since this determines not only metallurgical processes during their production, but also the performance of metals under operating conditions.
Today, when the economy is developing at a rapid pace, there is a need for prefabricated buildings that do not require significant capital investments. This is mainly needed for the construction of shopping pavilions, entertainment centers, and warehouses. With the use of metal structures, such buildings can now not only be easily and quickly erected, but also dismantled with the same ease when the rental period ends or for moving to another location. Moreover, it is not difficult to install communications, heating, and light into such easily erected buildings. Buildings made of metal structures can withstand harsh natural conditions not only in terms of temperature conditions, but also, which is equally important, in terms of seismological activity, where it is not easy or safe to erect brick buildings.
The range of metal structures offered by industry today is easily transportable and can be lifted by any cranes. The connection and installation of such structures can be done either using bolts or welding. The emergence of lightweight metal structures, which are manufactured and supplied in an integrated manner, plays a large positive role in the construction of public buildings in comparison with the construction of buildings made of reinforced concrete, and significantly reduces the time required to complete the work.
Bibliography.
1. Khomchenko G.P. A manual on chemistry for applicants to universities. - 3rd edition - M.: Novaya Volna Publishing House LLC, ONIKS Publishing House CJSC, 1999. - 464 p.
2. A.S. Egorova. Chemistry. A guide for applicants to universities - 2nd edition - Rostov N/D: Phoenix Publishing House, 1999. - 768 p.
3. Frolov V.V. Chemistry: Textbook for mechanical engineering special universities. – 3rd ed., revised. and additional – M.: Higher School, 1986.-543 p.
Supports with his approval the student’s incorrect or not entirely accurate answer. 1.2 Improving school chemical experiments in problem-based learning 1.2.1 Principles for developing a methodological system and the content of experiments in chemistry in a problem-based learning system A characteristic feature of developmental education is the widespread use of a problem-based approach, which includes the creation...
objectively existing relationship between chemical elements. That is why it was called by Mendeleev a “natural” system of elements. The periodic law has no equal in the history of science. Instead of disparate, unrelated substances, science faced a single, harmonious system that united all the chemical elements into one whole. Mendeleev pointed out the path of directed search in chemistry...
Position of metals in the periodic table
If in D.I. Mendeleev’s table we draw a diagonal from boron to astatine, then in the main subgroups under the diagonal there will be metal atoms, and in the secondary subgroups all elements are metals. Elements located near the diagonal have dual properties: in some of their compounds they behave like metals; in some - as non-metals.
Structure of metal atoms
In periods and main subgroups there are regularities in the change in metallic properties.
Many metal atoms have 1, 2 or 3 valence electrons, for example:
Na(+11): 1S2 2S22p6 3S1
Ca (+ 20): 1S2 2S22p6 3S23p63d0 4S2
Alkali metals (group 1, main subgroup): ...nS1.
Alkaline earth (group 2, main subgroup): ...nS2.
The properties of metal atoms are periodically dependent on their location in D.I. Mendeleev’s table.
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a - copper; b - magnesium; c - α-modification of iron
Metal atoms tend to give up their outer electrons. In a piece of metal, ingot or metal product, the metal atoms give up external electrons and send them into this piece, ingot or product, turning into ions. “Detached” electrons move from one ion to another, temporarily recombine with them into atoms, are detached again, and this process occurs continuously. Metals have a crystal lattice, at the nodes of which there are atoms or ions (+); Between them there are free electrons (electron gas). The connection diagram in metal can be displayed as follows:
М0 ↔ nē + Мn+,
atom - ion
Where n is the number of external electrons participating in the bond (y Na - 1 ē, y Ca - 2 ē, y Al - 3 ē).
This type of bond is observed in metals - simple substances - metals and alloys.
A metallic bond is a bond between positively charged metal ions and free electrons in a metal crystal lattice.
A metallic bond has some similarities with a covalent bond, but also some differences, since a metallic bond is based on the sharing of electrons (similarity), all atoms take part in the sharing of these electrons (difference). That is why crystals with a metal bond are plastic, electrically conductive and have a metallic luster. However, in the vapor state, metal atoms are connected to each other by a covalent bond; metal pairs consist of individual molecules (monatomic and diatomic).
General characteristics of metals
The ability of atoms to give up electrons (oxidize) |
← Increasing |
||
Interaction with atmospheric oxygen | Oxidizes quickly at normal temperatures | Oxidizes slowly at normal temperatures or when heated | Do not oxidize |
Interaction with water | At normal temperature, H2 is released and hydroxide is formed | When heated, H2 is released | H2 is not displaced from water |
Interaction with acids | Displaces H2 from dilute acids | Does not displace H2 from dilute acids |
|
React with conc. and dil. HNO3 and conc. H2SO4 when heated | Does not react with acids |
||
Being in nature | Only in connections | In connections and in free form | Mainly in free form |
Methods of obtaining | Electrolysis of melts | Reduction with coal, carbon monoxide (2), aluminothermy, or electrolysis of aqueous salt solutions | |
The ability of ions to gain electrons (recover) | Li K Ca Na Mg Al Mn Zn Cr Fe Ni Sn Pb (H) Cu Hg Ag Pt Au |
||
Increasing → |
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Electrochemical voltage series of metals. Physical and chemical properties of metals
General physical properties of metals
The general physical properties of metals are determined by the metallic bond and the metallic crystal lattice.
Malleability, ductility
Mechanical action on a metal crystal causes displacement of layers of atoms. Since the electrons in the metal move throughout the crystal, no bond breaking occurs. Plasticity decreases in the series Au, Ag, Cu, Sn, Pb, Zn, Fe. Gold, for example, can be rolled into sheets no more than 0.001 mm thick, which are used for gilding various objects. Aluminum foil appeared relatively recently and before tea and chocolate were forged into tin foil, which was called staniol. However, Mn and Bi do not have ductility: these are brittle metals.
Metallic shine
Metallic luster, which all metals lose in powder except Al And Mg. The most shiny metals are Hg(the famous “Venetian mirrors” were made from it in the Middle Ages), Ag(modern mirrors are now made from it using the “silver mirror” reaction). By color (conventionally) a distinction is made between ferrous and non-ferrous metals. Among the latter, we highlight the precious ones - Au, Ag, Pt. Gold is the metal of jewelers. It was on its basis that Faberge's wonderful Easter eggs were made.
Ringing
Metals ring, and this property is used to make bells (remember the Tsar Bell in the Moscow Kremlin). The most sonorous metals are Au, Ag, Cu. The copper rings with a thick, buzzing ringing - a crimson ringing. This figurative expression is not in honor of the raspberry, but in honor of the Dutch city of Malina, where the first church bells were smelted. In Russia, Russian craftsmen then began to cast bells of even better quality, and residents of cities and towns donated gold and silver jewelry so that the bells cast for churches would sound better. In some Russian pawnshops, the authenticity of gold rings accepted for commission was determined by the ringing of a gold wedding ring suspended on a woman’s hair (a very long and clear high-pitched sound is heard).
Under normal conditions, all metals except mercury Hg are solids. The hardest metal is chromium Cr: it scratches glass. The softest are alkali metals; they can be cut with a knife. Alkali metals are stored with great precautions - Na - in kerosene, and Li - in Vaseline because of its lightness, kerosene - in a glass jar, a jar - in asbestos chips, asbestos - in a tin jar.
Electrical conductivity
The good electrical conductivity of metals is explained by the presence of free electrons in them, which, under the influence of even a small potential difference, acquire directional movement from the negative pole to the positive one. As the temperature increases, the vibrations of atoms (ions) increase, which impedes the directional movement of electrons and thereby leads to a decrease in electrical conductivity. At low temperatures, the oscillatory motion, on the contrary, is greatly reduced and the electrical conductivity increases sharply. Near absolute zero, metals exhibit superconductivity. Ag, Cu, Au, Al, Fe have the highest electrical conductivity; the worst conductors are Hg, Pb, W.
Thermal conductivity
Under normal conditions, the thermal conductivity of metals changes essentially in the same sequence as their electrical conductivity. Thermal conductivity is determined by the high mobility of free electrons and the vibrational motion of atoms, due to which the temperature quickly equalizes in the metal mass. The highest thermal conductivity is for silver and copper, the lowest for bismuth and mercury.
Density
The density of metals is different. The smaller the atomic mass of the metal element and the larger the radius of its atom, the smaller it is. The lightest metal is lithium (density 0.53 g/cm3), the heaviest is osmium (density 22.6 g/cm3). Metals with a density of less than 5 g/cm3 are called light, the rest are called heavy.
The melting and boiling points of metals vary. The most fusible metal is mercury (tbp = -38.9°C), cesium and gallium melt at 29 and 29.8°C, respectively. Tungsten is the most refractory metal (tbp = 3390°C).
The concept of allotropy of metals using the example of tin
Some metals have allotropic modifications.
For example, tin is distinguished into:
· α-tin, or gray tin (“tin plague” - the transformation of ordinary β-tin into α-tin at low temperatures caused the death of R. Scott’s expedition to the South Pole, which lost all its fuel, as it was stored in sealed tanks tin), stable at t<14°С, серый порошок.
· β-tin, or white tin (t = 14 - 161°C) is a very soft metal, but harder than lead, amenable to casting and soldering. Used in alloys, for example to make tinplate (tinned iron).
Electrochemical voltage series of metals and its two rules
The arrangement of atoms in a row according to their reactivity can be represented as follows:
Li, K, Ca, Na, Mg, Al, Mn, Zn, Fe, Ni, Sn, Pb,H2 , Cu, Hg, Ag, Pt, Au.
The position of an element in the electrochemical series shows how easily it forms ions in an aqueous solution, i.e. its reactivity. The reactivity of elements depends on the ability to accept or donate electrons involved in bond formation.
1st rule of voltage series
If a metal is in this series before hydrogen, it is capable of displacing it from acid solutions; if after hydrogen, then not.
For example, Zn, Mg, Al gave a substitution reaction with acids (they are in the voltage range up to H), A Cu no (she is after H).
2nd rule of voltage series
If a metal is in the stress series before the salt metal, then it is able to displace this metal from the solution of its salt.
For example, CuSO4 + Fe = FeSO4 + Cu.
In such cases, the position of the metal before or after hydrogen may not matter, what is important is that the reacting metal precedes the metal forming the salt:
Cu + 2AgNO3 = 2Ag + Cu(NO3)2.
General chemical properties of metals
In chemical reactions, metals are reducing agents (donate electrons).
Interaction with simple substances.
1. Metals form salts with halogens - halides:
Mg + Cl2 = MgCl2;
Zn + Br2 = ZnBr2.
2. Metals form oxides with oxygen:
4Na + O2 = 2 Na2O;
2Cu + O2 = 2CuO.
3. Metals form salts with sulfur - sulfides:
4. With hydrogen, the most active metals form hydrides, for example:
Ca + H2 = CaH2.
5. many metals form carbides with carbon:
Ca + 2C = CaC2.
Interaction with complex substances
1. Metals located at the beginning of the voltage series (from lithium to sodium), under normal conditions, displace hydrogen from water and form alkalis, for example:
2Na + 2H2O = 2NaOH + H2.
2. Metals located in the voltage series up to hydrogen interact with dilute acids (HCl, H2SO4, etc.), as a result of which salts are formed and hydrogen is released, for example:
2Al + 6НCl = 2AlCl3 + 3H2.
3. Metals interact with solutions of salts of less active metals, as a result of which a salt of a more active metal is formed, and the less active metal is released in free form, for example:
CuSO4 + Fe = FeSO4 + Cu.
Metals in nature.
Finding metals in nature.
Most metals occur in nature in the form of various compounds: active metals are found only in the form of compounds; low-active metals - in the form of compounds and in free form; noble metals (Ag, Pt, Au...) in free form.
Native metals are usually found in small quantities as grains or inclusions in rocks. Occasionally there are also quite large pieces of metal - nuggets. Many metals in nature exist in a bound state in the form of chemical natural compounds - minerals. Very often these are oxides, for example iron minerals: red iron ore Fe2O3, brown iron ore 2Fe2O3 ∙ 3Н2О, magnetic iron ore Fe3O4.
Minerals are part of rocks and ores. Rudami are natural formations containing minerals in which metals are found in quantities technologically and economically suitable for the production of metals in industry.
Based on the chemical composition of the mineral included in the ore, oxide, sulfide and other ores are distinguished.
Usually, before obtaining metals from ore, it is pre-enriched - waste rock and impurities are separated, resulting in the formation of a concentrate that serves as raw material for metallurgical production.
Methods for obtaining metals.
Obtaining metals from their compounds is the task of metallurgy. Any metallurgical process is a process of reduction of metal ions with the help of various reducing agents, resulting in the production of metals in free form. Depending on the method of carrying out the metallurgical process, pyrometallurgy, hydrometallurgy and electrometallurgy are distinguished.
Pyrometallurgy- this is the production of metals from their compounds at high temperatures using various reducing agents: carbon, carbon monoxide (II), hydrogen, metals (aluminum, magnesium), etc.
Examples of metal recovery
ZnO + C → Zn + CO2;
carbon monoxide:
Fe2O3 + 3CO → 2Fe + 3CO2;
hydrogen:
WO3 + 3H2 → W + 3H2O;
CoO + H2 → Co + H2O;
aluminum (aluminothermy):
4Al + 3MnO2 → 2Al2O3 + 3Mn;
Cr2O3 + 2Al = 2Al2O3 + 2Cr;
magnesium:
TiCl4 + 2Mg = Ti + 2MgCl2.
Hydrometallurgy- this is the production of metals, which consists of two processes: 1) a natural metal compound is dissolved in an acid, resulting in a solution of the metal salt; 2) this metal is displaced from the resulting solution by a more active metal. For example:
1. 2CuS + 3O2 = 2CuO + 2SO2.
CuO + H2SO4 = CuSO4 + H2O.
2. CuSO4 + Fe = FeSO4 + Cu.
Electrometallurgy- this is the production of metals by electrolysis of solutions or melts of their compounds. Electric current plays the role of a reducing agent in the electrolysis process.
General characteristics of group IA metals.
The metals of the main subgroup of the first group (IA-group) include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr). These metals are called alkali because they and their oxides form alkalis when reacting with water.
Alkali metals belong to the s-elements. Metal atoms have one s-electron (ns1) in their outer electron layer.
Potassium, sodium - simple substances
Alkali metals in ampoules:
a - cesium; b - rubidium; c - potassium; g – sodium
Basic information about the elements of group IA
Li lithium | Na sodium | K potassium | Rb rubidium | Cs cesium | Fr France |
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Atomic number | ||||||
Oxidation state | ||||||
Main Natural Compounds | Li2O Al2O3 4SiO2 (spodumene); LiAl(PO4)F, LiAl(PO4)OH (amblygonite) | NaCl (table salt); Na2SO4· 10H2O (Glauber's salt, mirabile); KCl NaCl (sylvinite) | KCl (sylvinite), KCl NaCl (sylvinite); K (potassium feldspar, orthoglaze); KCl MgCl2 6H2O (carnallite) - found in plants | As an isoamorphous impurity in potassium minerals - sylvinite and carnallite | 4Cs2O 4Al2O3 18 SiO2 2H2O (hemi-cyt); satellite of potassium minerals | Actinium α-decay product |
Physical properties
Potassium and sodium - soft silvery metals (cut with a knife); ρ(K) = 860 kg/m3, Tmelt(K) = 63.7°C, ρ(Na) = 970 kg/m3, Tmelt(Na) = 97.8°C. They have high heat and electrical conductivity, color the flame in characteristic colors: K - pale violet, Na - yellow.
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Interaction with complex substances:
1. 2Na + 2H2O → 2NaOH + H2.
2. 2Na + Na2O2 → 2Na2O.
3. 2Na + 2НCl → 2NaCl + Н2.
Pulp and paper industry" href="/text/category/tcellyulozno_bumazhnaya_promishlennostmz/" rel="bookmark">production of paper, artificial fabrics, soap, for cleaning oil pipelines, in the production of artificial fiber, in alkaline batteries.
Finding metal compoundsI.A.groups in nature.
Salts ― NaCl- sodium chloride, NaNO3- sodium nitrate (Chilean saltpeter), Na2СО3- sodium carbonate (soda), NaHCO3- sodium bicarbonate (baking soda), Na2SO4- sodium sulfate, Na2SO4 10H2O- Glauber's salt, KCl- potassium chloride, KNO3- potassium nitrate (potassium nitrate), К2SO4- potassium sulfate, K2CO3- potassium carbonate (potash) - crystalline ionic substances, almost all soluble in water. Sodium and potassium salts exhibit the properties of medium salts:
· 2NaCl(solid) + Н2SO4(conc.) → Na2SO4 + 2НCl;
· KCl + AgNo3 → KNO3 + AgCl ↓;
· Na2СО3 + 2НCl → NaCl + CO2 + Н2О;
· K2СО3 + Н2О ↔ KHCO3 + KOH;
CO32- + H2O ↔ HCO3- + OH - (alkaline medium, pH< 7).
Table salt crystals
Salt mine
Na2СО3 used for the production of paper, soap, glass;
NaHCO3- in medicine, cooking, in the production of mineral waters, in fire extinguishers;
K2CO3- for producing liquid soap and glass;
Potash – potassium carbonate
NaNO3, KNO3, KCl, K2SO4― the most important potassium fertilizers.
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Sea salt contains 90-95% NaCl (sodium chloride) and up to 5% other minerals: magnesium salts, calcium salts, potassium salts, manganese salts, phosphorus salts, iodine salts, etc. All together, over 40 useful elements of the periodic table - all this exists in sea water.
Dead Sea
There is something extraordinary, almost fantastic, about it. In the eastern lands, even the tiniest stream of moisture is a source of life, where gardens bloom and grains ripen. But this water kills all living things.
Many peoples visited these shores: Arabs, Jews, Greeks, Romans; each of them called this huge lake in their own language, but the meaning of the name was the same: dead, rotten, lifeless.
We stood on a deserted shore, the dull appearance of which evoked sadness: a dead land - no grass, no birds. On the other side of the lake, reddish mountains rose steeply from the green water. Bare, wrinkled slopes. It seemed as if some force had torn off their natural cover, and the muscles of the earth were exposed.
We decided to swim, but the water turned out to be cold, we just washed ourselves with thick, flowing water like cool brine. After a few minutes, my face and hands were covered with a white coating of salt, and an unbearably bitter taste remained on my lips, which I could not get rid of for a long time. It is impossible to drown in this sea: the thick water itself holds a person on the surface.
Sometimes fish swim from the Jordan into the Dead Sea. She dies within a minute. We found one such fish washed up on the shore. It was hard as a stick, in a strong salt shell.
This sea can become a source of wealth for the people. After all, this is a giant storehouse of mineral salts.
Each liter of Dead Sea water contains 275 grams of potassium, sodium, bromine, magnesium, and calcium salts. Mineral reserves here are estimated at 43 billion tons. Bromine and potash can be produced extremely cheaply, and there is no limit to the scale of production. The country has huge reserves of phosphates, which are in great demand on the world market, but only a negligible amount is mined.
General characteristics of group IIA elements.
The metals of the main subgroup of the second group (IIA group) include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), radium (Ra). These metals are called alkaline earth metals, since their hydroxides Me(OH)2 have alkaline properties, and their oxides MeO are similar in their refractoriness to the oxides of heavy metals, previously called “earths”.
Alkaline earth metals belong to the s-elements. Metal atoms have two s electrons (ns2) in their outer electron layer.
Basic information about Group IIA elements
Be beryllium | Mg magnesium | Ca calcium | Sr strontium | Ba barium | Ra radium |
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Atomic number | ||||||
Structure of the outer electron shells of atoms |
where n = 2, 3, 4, 5, 6, 7, n is the period number |
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Oxidation state | ||||||
Main Natural Compounds | 3BeO Al2O3 6SiO2 (beryl); Be2SiO4 (phenacite) | 2MgO SO2 (olivine); MgCO3 (magnesite); MgCO3 · CaCO3 (dolomite); MgCl2 KCl 6H2O (carnal lite) | CaCO3 (calcite), CaF2—fluorite, CaO Al2O3 6SiO2 (anorthite); CaSO4 2H2O (gypsum); MgCO3 CaCO3 (dolomite), Сa3(PO4)2 – phosphorite, Сa5(PO4)3Х (Х = F, Cl, OH) – apatite | SrCO3 (stron-cyanite), SrSO4 (celestine) | BaCO3 (baterite) BaSO4 (barite, heavy spar) | As part of uranium ores |
Alkaline earth― light silvery-white metals. Strontium has a golden hue and is much harder than alkali metals. Barium is softer than lead. In air at ordinary temperatures, the surface of beryllium and magnesium is covered with a protective oxide film. Alkaline earth metals actively interact with oxygen in the air, so they are stored under a layer of kerosene or in sealed vessels, like alkali metals.
Calcium is a simple substance
Physical properties
Natural calcium is a mixture of stable isotopes. The most common calcium is 97%). Calcium is a silvery-white metal; ρ = 1550 kg/m3, Tmelt = 839°C. Colors the flame orange-red.
Chemical properties
Interaction with simple substances (non-metals):
1. With halogens: Ca + Cl2 → CaCl2 (calcium chloride).
2. With carbon: Ca + 2C → CaC2 (calcium carbide).
3. With hydrogen: Ca + H2 → CaH2 (calcium hydride).
Salts: CaCO3 Calcium carbonate is one of the most common compounds on Earth: chalk, marble, limestone. The most important of these minerals is limestone. It itself is an excellent building stone; in addition, it is a raw material for the production of cement, slaked lime, glass, etc.
Lime gravel is used to strengthen roads, and powder is used to reduce soil acidity.
Natural chalk represents the remains of shells of ancient animals. It is used as school crayons, in toothpastes, and in the production of paper and rubber.
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Physical properties
Iron is a silver-white or gray metal, hard, with high ductility, thermal and electrical conductivity, refractory; ρ = 7874 kg/m3, Tmelt = 1540°C. Unlike other metals, iron can be magnetized and has ferromagnetism.
Chemical properties
Iron interacts with both simple and complex substances.
Interaction of iron with oxygen
a) when heated (combustion), b) at N. u. (corrosion)
Chemical properties of iron
When n. at. | When heated |
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Reaction | 3FeSO4 + 2K3 = Fe32↓ + 3K2SO4 (turbulen blue - dark blue precipitate). | 1. 4FeCl3 + 3K4 = Fe43↓ + 12KCl (Prussian blue - dark blue precipitate). 2. FeCl3 + 3NH4CNS ⇆ Fe(CNS)3 + 3NH4Cl (fe blood red rhodanide + ammonia). |
Biological role of iron
Biochemists reveal the enormous role of iron in the life of plants, animals and humans. As part of hemoglobin, iron causes the red color of this substance, which, in turn, determines the color of blood. The adult human body contains 3 g of iron, of which 75% is part of hemoglobin, thanks to which the most important biological process - respiration - is carried out. Iron is also necessary for plants. It participates in the oxidative processes of protoplasm, during plant respiration and in the construction of chlorophyll, although it itself is not part of its composition. Iron has long been used in medicine to treat anemia, exhaustion, and loss of strength.