Oxidation is the process of donating electrons, with an increase in the oxidation state.

At oxidation recoil electrons it increases oxidation state. Atoms oxidizable substances are called donors electrons, and atoms oxidizer - acceptors electrons.

In some cases, during oxidation, the molecule of the starting substance may become unstable and decay into more stable and smaller constituents (see. Free radicals). Moreover, some of the atoms of the resulting molecules have a higher oxidation state than the same atoms in the original molecule.

The oxidant, accepting electrons, acquires reducing properties, turning into a conjugated reducing agent:

oxidizing agent+ e − ↔ conjugate reductant.

Recovery

Restoration is called the process of attaching electrons to an atom of a substance, while its oxidation state goes down.

When recovering atoms or ions attach electrons... At the same time, there is a decrease oxidation state element... Examples: recovery oxides metals to free metals with hydrogen, carbon, other substances; recovery organic acids v aldehydes and alcohols; hydrogenation fat and etc.

The reducing agent, donating electrons, acquires oxidizing properties, turning into a conjugated oxidizing agent:

reducing agent - e − ↔ conjugate oxidant.

An unbound, free electron is the strongest reducer.

Redox reactions Are reactions in which reactants add or donate electrons. An oxidant is a particle (ion, molecule, element) that attaches electrons and passes from a higher oxidation state to a lower one, i.e. is recovering. A reducing agent is a particle that donates electrons and passes from a lower oxidation state to a higher one, i.e. oxidizes.

Intermolecular - reactions in which oxidizing and reducing atoms are in the molecules of different substances, for example:

H 2 S + Cl 2 → S + 2HCl

Intramolecular - reactions in which oxidizing and reducing atoms are in the molecules of the same substance, for example:

2H 2 O → 2H 2 + O 2

Disproportionation (self-oxidation-self-reduction) - reactions in which atoms with an intermediate oxidation state are converted into an equimolar mixture of atoms with a higher and lower oxidation states, for example:

Cl 2 + H 2 O → HClO + HCl

Reproportionation is a reaction in which one oxidation state is obtained from two different oxidation states of the same element, for example:

NH 4 NO 3 → N 2 O + 2H 2 O

Oxidation, reduction

In redox reactions, electrons from one atom, molecule or ion are transferred to another. The process of donating electrons is oxidation. With oxidation, the oxidation state increases:

The process of electron attachment is reduction. Upon reduction, the oxidation state decreases:

The atoms or ions that attach electrons in this reaction are oxidizing agents, and those that donate electrons are reducing agents.

Redox reactions (electrode potential)

Electrons can act as chemical reagents, and half-reaction is practically used in devices called galvanic cells.

An example of an electrode is a plate and crystalline zinc immersed in a zinc sulfate solution. After the immersion of the plate, 2 processes take place. As a result of the first process, the plate acquires a negative charge; some time after immersion in the solution, the velocities equalize and equilibrium occurs. And the plate acquires some electrical potential.

Measure the electrode potential relative to the potential of standard hydrogen.

Copper-hydrogen electrode- an electrode used as reference electrode with various electrochemical measurements and in galvanic cells... A hydrogen electrode (HE) is a metal plate or wire that absorbs gaseous hydrogen(usually use platinum or palladium) saturated with hydrogen (at atmospheric pressure) and immersed in water solution containing hydrogen ions... The potential of the plate depends on [ clarify ] on the concentration of H + ions in the solution. The electrode is a reference against which the electrode potential of the determined chemical reaction is measured. At a hydrogen pressure of 1 atm., A concentration of protons in a solution of 1 mol / L and a temperature of 298 TO the SE potential is taken to be 0 V. When assembling a galvanic cell from an SE and an electrode to be determined, the reaction reversibly proceeds on the platinum surface:

2H + + 2e - = H 2

that is, either recovery hydrogen or its oxidation- it depends on the potential of the reaction proceeding on the determined electrode. By measuring the EMF of a galvanic electrode under standard conditions (see above) determine standard electrode potential determined chemical reaction.

SE is used to measure the standard electrode potential of an electrochemical reaction, to measure concentration(activity) of hydrogen ions, as well as any other ions... VE is also used to determine the product of solubility, to determine the rate constants of some electro chemical reactions.

Nernst equation

The dependence of the redox potential corresponding to the half-reaction of the reduction of the permanganate ion in an acidic medium (and, as already noted, at the same time the half-reaction of the oxidation of the Mn 2+ cation to the permanganate ion in an acidic medium) on the above factors determining it is quantitatively described by the Nernst equation

Each of the concentrations under the sign of the natural logarithm in the Nernst equation is raised to the power corresponding to the stoichiometric coefficient of the given particle in the half-reaction equation, n- the number of electrons accepted by the oxidizer, R- universal gas constant, T- temperature, F Is the Faraday number.

Measure the redox potential in the reaction vessel during the course of the reaction, i.e. under nonequilibrium conditions, it is impossible, since when measuring the potential, electrons must be transferred from the reducing agent to the oxidizing agent not directly, but through the metal conductor connecting the electrodes. In this case, the electron transfer rate (current strength) must be kept very low due to the application of an external (compensating) potential difference. In other words, the measurement of electrode potentials is possible only under equilibrium conditions, when direct contact between the oxidizing agent and the reducing agent is excluded. Therefore, square brackets in the Nernst equation denote, as usual, the equilibrium (under measurement conditions) particle concentrations. Although the potentials of redox pairs during the course of the reaction cannot be measured, they can be calculated by substituting the current ones into the Nernst equation, i.e. responding this moment concentration time. If we consider the change in potential as the reaction proceeds, then first these are the initial concentrations, then the concentrations that depend on time, and, finally, after the termination of the reaction, equilibrium. As the reaction proceeds, the oxidant potential calculated by the Nernst equation decreases, while the potential of the reducing agent corresponding to the second half-reaction, on the contrary, increases. When these potentials equalize, the reaction stops and the system comes to a state of chemical equilibrium.

25. Complex compounds are compounds that exist both in the crystalline state and in solution, the feature of which is the presence of a central atom surrounded by ligands. Complex compounds can be considered as complex compounds of a higher order, consisting of simple molecules capable of independent existence in solution. Werner's coordination theory in each complex compound distinguishes between the inner and outer spheres. The central atom with the surrounding ligands form the inner sphere of the complex. It is usually enclosed in square brackets. Everything else in complex compound constitutes the outer sphere and is written outside square brackets. A certain number of ligands are located around the central atom, which is determined by the coordination number. The number of coordinated ligands is most often 6 or 4. The ligand occupies a coordination position near the central atom. Coordination changes the properties of both the ligands and the central atom. Often, coordinated ligands cannot be detected by chemical reactions that are characteristic of them in a free state. The more tightly bound particles of the inner sphere are called a complex (complex ion). Attraction forces act between the central atom and the ligands (formed covalent bond by exchange and (or) donor – acceptor mechanism), between ligands - repulsive forces. If the charge of the inner sphere is 0, then the outer coordination sphere is absent. The central atom (complexing agent) is the atom or ion that occupies a central position in the complex compound. The role of a complexing agent is most often performed by particles with free orbitals and a sufficiently large positive charge of the nucleus, and therefore can be electron acceptors. These are cations of transition elements. The strongest complexing agents are elements of groups IВ and VIIIВ. Rarely, neutral atoms of d-elements and atoms of non-metals in various oxidation states act as complexing agents. The number of free atomic orbitals provided by the complexing agent determines its coordination number. The value of the coordination number depends on many factors, but usually it is equal to twice the charge of the complexing ion. Ligands are ions or molecules that are directly bound to the complexing agent and are donors of electron pairs. These are electronically abundant systems with free and mobile electron pairs that can be electron donors. Compounds of p-elements exhibit complexing properties and act as ligands in a complex compound. Ligands can be atoms and molecules (protein, amino acids, nucleic acids, carbohydrates). According to the number of bonds formed by ligands with a complexing agent, ligands are divided into mono-, bi-, and polydentate ligands. The above ligands - molecules and anions are monodentate, since they are donors of one electron pair. Bidentate ligands include molecules or ions containing two functional groups capable of donating two electron pairs. The charge of the inner sphere of the complex compound is the algebraic sum of the charges of the particles that form it. Complex compounds having an ionic outer sphere in solution undergo dissociation into a complex ion and ions of the outer sphere. They behave in dilute solutions like strong electrolytes: dissociation occurs instantly and almost completely. SO4 = 2+ + SO42-. If there are hydroxide ions in the outer sphere of the complex compound, then this compound is a strong base.

Group IA includes lithium, sodium, potassium, rubidium, cesium and francium. These elements are called alkaline elements, and sometimes hydrogen is also included in the IA group. Thus, this group includes elements of each of the 7 periods. The general valence electronic formula of the elements of the IA group is ns1 At the external level 1 electron. Far from the nucleus. Low ionization potentials. Atoms give 1 electron. Metallic means are sharply expressed. Metallic properties increase with increasing serial number. Physical properties: Metals are soft, light, fusible with good electrical conductivity, have a large negative value of electrical potentials. Chemical properties: 1) Store under a layer of liquid hydrocarbons (benzene, gasoline, kerasin) 2) Oxidants. Easily oxidize alkali metals to halides, sulfides, phosphides. Li Na K Rb Cs Increase in metal radius Decrease in ionization energy Decrease in electronegativity Decrease in melting and boiling points Application of sodium and potassium 1. Obtaining peroxides. 2. An alloy of sodium and potassium - a coolant in nuclear power plants. 3. Obtaining organometallic compounds.

27. General comparative characteristics of elements and their compounds of IA and IB groups of the periodic system Alkaline metals are elements of the 1st group of the periodic table of chemical elements (according to the outdated classification - elements of the main subgroup of group I): lithium Li, sodium Na, potassium K, rubidium Rb, cesium Cs and francium Fr. When alkali metals are dissolved in water, soluble hydroxides called alkalis are formed. V Periodic table they immediately follow inert gases, therefore, the peculiarity of the structure of atoms of alkali metals is that they contain one electron at the external energy level: their electronic configuration ns1. It is obvious that the valence electrons of alkali metals can be easily removed, because it is energetically favorable for the atom to donate an electron and acquire the configuration of an inert gas. Therefore, all alkali metals are characterized by reducing properties. This is confirmed by the low values ​​of their ionization potentials (the ionization potential of the cesium atom is one of the lowest) and electronegativity (EO). All metals in this subgroup have a silvery-white color (except for silver-yellow cesium), they are very soft, they can be cut with a scalpel. Lithium, sodium and potassium are lighter than water and float on its surface, reacting with it. Alkali metals occur naturally in the form of compounds containing singly charged cations. Many minerals contain metals of the main subgroup of group I. For example, orthoclase, or feldspar, consists of potassium aluminosilicate K2, a similar mineral containing sodium - albite - has the composition Na2. V sea ​​water contains sodium chloride NaCl, and the soil contains potassium salts - sylvinite KCl, sylvinite NaCl KCl, carnallite KCl MgCl2 6H2O, polyhalite K2SO4 MgSO4 CaSO4 2H2O. Copper subgroup - chemical elements of the 11th group of the periodic table of chemical elements (according to the outdated classification - elements of the secondary subgroup of group I). The group includes transition metals from which coins are traditionally made: copper Cu, silver Ag and gold Au. Based on the structure of the electronic configuration, the Rg X-ray belongs to the same group, but it does not fall into the "coin group" (it is a short-lived transactinide with a half-life of 3.6 sec). The name coin metals does not officially apply to element group 11, as other metals such as aluminum, lead, nickel, stainless steel and zinc are also used to make coins. All elements of the subgroup are relatively chemically inert metals. High values ​​of density, melting and boiling points, high thermal and electrical conductivity are also characteristic. A feature of the elements of the subgroup is the presence of a filled pre-external sublevel, achieved due to the electron hopping from the ns sublevel. The reason for this phenomenon is the high stability of the completely filled d-sublevel. This feature determines the chemical inertness of simple substances, their chemical inactivity, therefore gold and silver are called noble metals 28. Hydrogen. general characteristics... Reaction with oxygen, halogens, metals, oxides. Hydrogen peroxide, its redox properties Hydrogen is the most common chemical element in the Universe. It is the main constituent of the Sun, as well as many stars. In the earth's crust, the mass fraction of hydrogen is only 1%. However, its compounds are widespread, for example, H20 water. The composition of natural combustible gas is mainly a compound of carbon with hydrogen - methane CH4 - Hydrogen is also found in many organic substances. 1) If you ignite hydrogen (after checking for purity, see below) and lower a tube with burning hydrogen into a vessel with oxygen, then water droplets form on the walls of the vessel: Hydrogen without impurities burns quietly. However, a mixture of hydrogen with oxygen or air explodes. The most explosive mixture, consisting of two volumes of hydrogen and one volume of oxygen, is an explosive gas. If an explosion occurs in a glass vessel, then its fragments can be injured.

hurt others. Therefore, before igniting hydrogen, it is necessary to check it for purity. For this, hydrogen is collected in a test tube, which is brought upside down to the flame. If the hydrogen is pure, then it burns quietly, with a characteristic "p-groin" sound. If hydrogen contains an admixture of air, then it burns with an explosion. When working with hydrogen, the safety regulations must be followed. 2) If, for example, during heating, a stream of hydrogen is passed over copper (II) oxide, then a reaction occurs, as a result of which water and metallic copper are formed: In this reaction, a reduction process occurs, since hydrogen takes oxygen away from copper atoms. The reduction process is the opposite of the oxidation process. Substances that take oxygen away are reductants. The processes of oxidation and reduction are mutually related (if one element is oxidized; then the other is reduced, and vice versa). 3) Halogens react with hydrogen, forming HX, and with fluorine and chlorine, the reaction proceeds explosively with a slight activation. The interaction with Br2 and I2 proceeds more slowly. For the reaction with hydrogen to proceed, it is sufficient to activate a small fraction of the reagents using illumination or heating. The activated particles interact with the unactivated ones, forming HX and new activated particles, which continue the process, and the reaction of the two activated particles according to the main reaction ends with the formation of a product. 4) Oxidation reactions. When heating hydrogen with metals of I and II main subgroups: 2Na + H2 (300 ° C) ® 2NaH; Ca + H2 (500-700 ° C) ® CaH2. Hydrogen peroxide (hydrogen peroxide), H2O2 is the simplest representative of peroxides. Colorless liquid with a "metallic" taste, soluble in water, alcohol and ether. Concentrated aqueous solutions are explosive. Hydrogen peroxide is a good solvent. It is released from water in the form of unstable crystalline hydrate H2O2 2H2O. Hydrogen peroxide has oxidizing as well as reducing properties. It oxidizes nitrites to nitrates, releases iodine from metal iodides, splits unsaturated compounds at the site of double bonds. Hydrogen peroxide reduces gold and silver salts, as well as oxygen by reacting with an aqueous solution of potassium permanganate in an acidic medium. During the reduction of H2O2, H2O or OH- is formed, for example: H2O2 + 2KI + H2SO4 = I2 + K2SO4 + 2H2O Under the action of strong oxidants, H2O2 exhibits reducing properties, releasing free oxygen: O22− - 2e− → O2 The reaction of KMnO4 with H2O2 is used in chemical analysis to determine the content of H2O2: 5H2O2 + 2KMnO4 + 3H2SO4 → 5O2 + 2MnSO4 + K2SO4 + 8H2O It is advisable to oxidize organic compounds with hydrogen peroxide (for example, sulfides and thiols) in an acetic acid medium.

29. general characteristics of the properties of elements and their compounds of group 2 a. Physical and chemical properties, application. Includes s-elements. Be Mg Ca Br Ra Sr With the exception of Be, they are polyisotopic. Atoms of elements at the outer level have 2 S elements each with opposite spins with the expenditure of the necessary energy, one element from the s state passes into the p state. These are Metals, but they are less active than alkaline ones. most distributed in nature Mg Ca Be, found in the form of the mineral Be3AL2 (SiO3) 6 Production method: electrolysis of molten chlorides Physical properties: light metals, but harder than alkaline metals. Chem.Sv-va: 1In air, the surface of Be and Mg is covered with an oxide film. 2. at high temperature interacts with nitrogen 3. does not interact with water. become. Calcium and its hydride are also used to produce hard-to-reduce metals such as chromium, thorium and uranium. Calcium lead alloys are used in batteries and bearing alloys. Calcium granules are also used to remove traces of air from vacuum equipment.

No. 31 Alkaline earth metals - chemical elements 2nd group of the main subgroup, except for beryllium and magnesium: calcium, strontium, barium and radium... Belong to the 2nd group of elements according to the new classification IUPAC... Named so because their oxides- "land" (in the terminology alchemists) - report water alkaline reaction. Salt alkaline earth metals, except for radium, are widespread in nature in the form minerals.

Oxides- substances whose molecules consist of atoms of two elements, one of which is oxygen. Oxides are divided into basic, formed from metal atoms, for example, K2O, Fe2O3, CaO; acidic - formed by atoms of non-metals and some metals in their the highest degree oxidation: СО2, SO3, P2O5, CrO3, Mn2O7 and amphoteric, for example, ZnO, Al2O3, Cr2O3. Oxides are obtained from the combustion of simple and complex substances, as well as from the decomposition of complex substances (salts, bases, acids).

Chemical properties of oxides: 1. Oxides of alkali and alkaline earth metals interact with water, forming soluble bases - alkalis (NaOH, KOH, Ba (OH) 2). Na2O + H2O = 2NaOH

Most acidic oxides interact with water to form acids: CO2 + H2O = H2CO3

2. Some oxides interact with basic oxides: CO2 + CaO = CaCO3

3. Basic oxides interact with acids: BaO + 2HCl = BaCl2 + H2O

4. Acidic oxides interact with both acids and alkalis: ZnO + 2HCl = ZnCl2 + H2O

ZnO + 2NaOH = Na2ZnO2 + H2O

Hydroxides ( hydroxide) - compounds of oxide-chemical elements with water. Hydroxides of almost all chemical elements are known; some of them occur naturally as minerals. Alkali metal hydroxides are called alkalis. Depending on whether the corresponding oxide is basic, acidic or amphoteric, one can respectively distinguish between:

basic hydroxides (foundations) - hydroxides showing basic properties (for example, calcium hydroxideCa (OH) 2, potassium hydroxideKOH, sodium hydroxideNaOH, etc.);

acid hydroxides (oxygenated acids) - hydroxides exhibiting acidic properties (for example, nitric acid HNO 3, sulfuric acid H 2 SO 4, sulfurous acid H 2 SO 3, etc.)

amphoteric hydroxides, exhibiting, depending on the conditions, either basic or acidic properties (for example, aluminum hydroxide Al (OH) 3, zinc hydroxide Zn (OH) 2).

Carbonates and Hydrocarbonates - Salts and Ethers carbonic acid (H 2 CO 3). Among the salts, normal carbonates (with the CO 3 2− anion) and acidic or hydrocarbonates(with anion NSO 3 -).

Chemical properties

When heated, acidic carbonates transform into normal carbonates:

When heated strongly, normal carbonates decompose into oxides and carbon dioxide:

Carbonates react with acids stronger than carbonic (almost all known acids, including organic ones) with the release of carbon dioxide:

Application: Carbonates of calcium, magnesium, barium, etc. are used in the construction business, in the chemical industry, optics, etc. It is widely used in technology, industry and everyday life. soda (Na 2 CO 3 and NaHCO 3). Acid carbonates play an important physiological role, being buffering substances regulating the constancy of the reaction blood .

Silicates and aluminosilicates represent a large group of minerals ... They are characterized by a complex chemical composition and isomorphic substitutions of some elements and complexes of elements by others. The main chemical elements that make up silicates are Si , O , Al , Fe 2+, Fe 3+, Mg , Mn , Ca , Na , K , and Li , B , Be , Zr , Ti , F , H , in the form of (OH) 1− or H 2 O, etc.

Origin (genesis ): Endogenous, mainly magmatic (pyroxenes, feldspars ), they are also typical for pegmatites (mica, tourmaline, beryl, etc.) and skarn (garnets, wollastonite). Widespread in metamorphic rocks - shale and gneisses (garnets, disthene, chlorite). Silicates of exogenous origin are products of weathering or alteration of primary (endogenous) minerals (kaolinite, glauconite, chrysocolla)

No. 32. Group III includes boron, aluminum, gallium, indium, thallium (main subgroup), as well as scandium, yttrium, lanthanum and lanthanides, anemones and actinides (a secondary subgroup).

At the external electronic level of the elements of the main subgroup, there are three electrons each (s 2 p 1). They easily donate these electrons or form three unpaired electron due to the transition of one electron to the p-level. For boron and aluminum, only compounds with an oxidation state of +3 are characteristic. The elements of the gallium subgroup (gallium, indium, thallium) also have three electrons at the outer electronic level, forming the s 2 р 1 configuration, but they are located after the 18-electron layer. Therefore, unlike aluminum, gallium has clearly non-metallic properties. These properties in the series Ga, In, Tl weaken, and the metallic properties increase.

The elements of the scandium subgroup also have three electrons at the outer electronic level. However, these elements belong to transition d-elements, the electronic configuration of their valence layer is d 1 s 2. All three elements give up these electrons quite easily. Elements of the lanthanide subgroup have a distinctive configuration of the external electronic level: the 4f level is built up in them and the d level disappears. Starting with cerium, all elements, except for gadolinium and lutetium, have the electronic configuration of the external electronic level 4f n 6s 2 (gadolinium and lutetium have 5d 1 electrons). The number n varies from 2 to 14. Therefore, s- and f-electrons take part in the formation of valence bonds. Most often, the oxidation state of lanthanides is +3, less often +4.

The electronic structure of the valence layer of actinides in many respects resembles the electronic structure of the valence layer of lanthanides. All lanthanides and actinides are typical metals.

All elements III group have a very strong affinity for oxygen, and the formation of their oxides is accompanied by the release of a large amount of heat.

Group III elements find a wide variety of applications.

33. Physical properties. Aluminum is a silvery white light metal that melts at 660 ° C. Very flexible, easily drawn into wire and rolled into sheets: it can be used to make foil with a thickness of less than 0.01 mm. Aluminum has a very high thermal and electrical conductivity. Its alloys with various metals are strong and lightweight.

Chemical properties. Aluminum is a very active metal. In the series of voltages, it comes after alkali and alkaline earth metals. However, it is quite stable in air, since its surface is covered with a very dense oxide film, which protects the metal from contact with air. If the protective oxide film is removed from the aluminum wire, then aluminum will begin to vigorously interact with oxygen and water vapor in the air, turning into a loose mass - aluminum hydroxide:

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

This reaction is accompanied by the release of heat.

Purified from the protective oxide film, aluminum interacts with water with the evolution of hydrogen:

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

Aluminum dissolves well in dilute sulfuric and hydrochloric acids:

2 Аl + 6 НСl = 2 AlСl 3 + 3 Н 2

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

Diluted nitric acid cold passivates aluminum, but when heated, aluminum dissolves in it with the release of nitrogen monoxide, nitrogen hemioxide, free nitrogen or ammonia, for example:

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

Concentrated nitric acid passivates aluminum.

Since aluminum oxide and hydroxide are amphoteric

properties, aluminum easily dissolves in aqueous solutions of all alkalis, except for ammonium hydroxide:

2 Al + 6 KOH + 6 H 2 O = 2 K 3 [Al (OH) 6] + 3 H 2

Powdered aluminum easily interacts with halogens, oxygen and all non-metals. To start the reactions, heating is necessary, then they proceed very intensively and are accompanied by the release of a large amount of heat:

2 Al + 3 Br 2 = 2 AlBr 3 (aluminum bromide)

4 Al + 3 O 2 = 2 Al 2 O 3 (aluminum oxide)

2 Al + 3 S = Al 2 S 3 (aluminum sulfide)

2 Al + N 2 = 2 AlN (aluminum nitride)

4 Al + 3 C = Al 4 C 3 (aluminum carbide)

Aluminum sulphide can only exist in solid form. In aqueous solutions, it undergoes complete hydrolysis with the formation of aluminum hydroxide and hydrogen sulfide:

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

Aluminum easily removes oxygen and halogens from oxides and salts of other metals. The reaction is accompanied by the release of a large amount of heat:

8 Al + 3 Fe 3 O 4 = 9 Fe + 4 Al 2 O 3

The process of reducing metals from their oxides with aluminum is called aluminothermy. Alumothermy is used to obtain some rare metals that form a strong bond with oxygen (niobium, tantalum, molybdenum, tungsten, etc.), as well as for welding rails. If, with the help of a special igniter, set fire to a mixture of fine aluminum powder and magnetic iron ore Fe 3 O 4 (thermite), then the reaction proceeds spontaneously with the mixture warming up to 3500 ° C. Iron at this temperature is in a molten state.

Receiving. For the first time, aluminum was obtained by reduction from aluminum chloride with metallic sodium:

AlCl 3 + 3 Na = 3 NaCl + Al

Currently, it is obtained by electrolysis of molten salts in electrolytic baths (Fig. 46). The electrolyte is a melt containing 85-90% cryolite - complex salt 3NaF · AlF 3 (or Na 3 AlF 6) and 10-15% alumina - aluminum oxide Al 2 O 3. This mixture melts at about 1000 ° C.

Application. Aluminum is widely used. It is used to make foil used in radio engineering and for packaging food products. Steel and cast iron products are coated with aluminum in order to protect them from corrosion: products are heated to 1000 ° C in a mixture of aluminum powder (49%), aluminum oxide (49%) and aluminum chloride (2%). This process is called aluminizing.

Aluminized products can withstand heating up to 1000 ° C without corroding. Slavs of aluminum, distinguished by their great lightness and strength, are used in the production of heat exchange devices, in aircraft construction and mechanical engineering.

Aluminum oxide Al 2 O 3. It is a white substance with a melting point of 2050 ° C. In nature, aluminum oxide occurs in the form of corundum and alumina. Sometimes there are beautifully shaped transparent crystals of corundum and ocher. Corundum colored red with chromium compounds is called ruby, and blue colored corundum with titanium and iron compounds is called sapphire. Ruby and sapphire are precious stones. Nowadays, they are quite easily obtained artificially.

Boron-element the main subgroup of the third group, the second period periodic table of chemical elements D. I. Mendeleev, with atomic number 5. Indicated by the symbol B(Borium). In a free state boron- colorless, gray or red crystalline or dark amorphous substance. More than 10 allotropic modifications of boron are known, the formation and mutual transitions of which are determined by the temperature at which boron was obtained.

Receiving

The purest boron is obtained by pyrolysis of borohydrides. Such boron is used for the production of semiconductor materials and fine chemical syntheses.

1. Method of metallothermia (more often restoration with magnesium or sodium):

2. Thermal decomposition of boron bromide vapor on a hot (1000-1200 ° C) tantalum wire in the presence of hydrogen:

Physical properties

Extremely hard substance (second only to diamond, carbon nitride, boron nitride (borazon), boron carbide, boron-carbon-silicon alloy, scandium-titanium carbide). Possesses fragility and semiconducting properties (wide-gap semiconductor).

Chemical properties

In many physical and chemical properties, the nonmetal boron resembles silicon.

Chemical boron is quite inert and at room temperature interacts only with fluorine:

When heated, boron reacts with other halogens to form trihalides, with nitrogen forms boron nitride BN, with phosphorus- phosphide BP, with carbon - carbides of various compositions (B 4 C, B 12 C 3, B 13 C 2). When heated in an oxygen atmosphere or in air, boron burns out with a large release of heat, the oxide B 2 O 3 is formed:

Boron does not directly interact with hydrogen, although a fairly large number of borohydrides (boranes) of various compositions are known, obtained by treating borides of alkali or alkaline earth metals with acid:

Boron exhibits reducing properties under strong heating. He is able, for example, to restore silicon or phosphorus of their oxides:

This property of boron can be explained by the very high strength of chemical bonds in boron oxide B 2 O 3.

In the absence of oxidizing agents, boron is resistant to the action of alkali solutions. In hot nitric, sulfuric acids and aqua regia, boron dissolves to form boric acid.

Boron oxide - typical acid oxide... It reacts with water to form boric acid:

When boric acid interacts with alkalis, salts are formed not of boric acid itself - borates (containing the BO 3 3- anion), but tetraborates, for example:

Application

Elementary boron

Boron (in the form of fibers) serves as a reinforcing agent in many composites.

Boron is also often used in electronics to change the type of conductivity. silicon.

Boron is used in metallurgy as a microalloying element that significantly increases the hardenability of steels.

34.haracharacteristic of elements of group 4A. Tin, lead.

(addition)

The group includes 5 elements: two non-metals - carbon and silicon, located in the second and third periods of the Mendeleev system and 3 metals - germanium (intermediate between non-metals and metals, tin and lead, located at the end of large periods - IV, V, VI All these elements are characterized by the fact that they have 4 electrons at the external energy level, and therefore they can exhibit an oxidation state from +4 to -4. These elements form gaseous compounds with hydrogen: CH4, Si H4, Sn H4, PbH4. When heated in air, they combine with elements of the subgroup oxygen, sulfur and halogens.The oxidation state +4 is obtained when the 1s-electron passes to the free p-orbital.

With an increase in the radius of the atom, the strength of the bond between the outer electrons and the nucleus decreases. Non-metallic properties decrease, while metallic properties increase. (melting and boiling points decrease, etc.)

Carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb) are elements of the 4th group of the main subgroup of the PSE. On the outer electron layer, the atoms of these elements have 4 electrons: ns 2 np 2. In a subgroup, with an increase in the ordinal number of an element, the atomic radius increases, the non-metallic properties weaken, and the metallic ones intensify: carbon and silicon are non-metals, germanium, tin, lead are metals.

General characteristics. Carbon and silicon

The subgroup of carbon, which includes carbon, silicon, germanium, tin and lead, is the main subgroup of group 4 of the Periodic Table.

There are 4 electrons on the outer electron shell of the atoms of these elements and their electronic configuration in general can be written as follows: ns 2 np 2, where n is the number of the period in which the chemical element is located. Moving from top to bottom in the group, the non-metallic properties are weakened, and the metallic ones increase, therefore carbon and silicon are non-metals, and tin and lead exhibit the properties of typical metals. Forming covalent polar bonds with hydrogen atoms, C and Si exhibit a formal oxidation state of -4, and with more active non-metals (N, O, S) and halogens, they exhibit oxidation states +2 and + 4. When clarifying the reaction mechanism, the carbon isotope 13 is sometimes used. C (tracer method). Therefore, it is useful to know that the abundance of carbon isotopes: 12 С - 98.89% and 13 С - 1.11%. If we restrict ourselves to the enumeration of isotopes, the prevalence of which is more than 0.01%, then silicon has 3 such isotopes, germanium has 5, tin has 10, and lead has 4 stable isotopes.

Under normal conditions, carbon can exist in the form of two allotropic

modifications: diamond and graphite; ultrapure crystalline silicon

Semiconductor.

Among the compounds of elements (E) of the carbon subgroup with hydrogen, we consider compounds of the EN 4 type. With an increase in the charge of the nucleus of the E atom, the stability of hydrides decreases.

On going from C to Pb, the stability of compounds with an oxidation state of +4

decreases, and from +2 - increases. For EO 2 oxides, the acidic character decreases, while for EO oxides, the basic character increases.

Carbon

Carbon occurs naturally in the form of diamond and graphite. It contains in fossil coals: from 92% - in anthracite, up to 80% - in brown coal. In a coherent state, carbon is found in carbides: CaCO 3 chalk, limestone and marble, MgCO 3 CaCO 3 - dolomite,

MgCO 3 - magnesite. In the air, carbon is contained in the form of carbon dioxide (0.03% by volume). Carbon is also contained in compounds dissolved in seawater.

Carbon is found in plants and animals, oil and natural gas.

In reactions with active non-metals, carbon is easily oxidized:

2 C + O 2 = 2 CO,

C + 2 F 2 = CF 4.

Carbon can also exhibit reducing properties when interacting with complex substances:

C + 2 CuO = 2 Cu + CO 2,

C + 2 H 2 SO 4 (conc) = CO 2 + 2 SO 2 + H 2 O,

2 C + BaSO 4 = BaS + 2 CO 2.

In reactions with metals and less active non-metals, carbon is an oxidizing agent: 2C + H 2 = C 2 H 2,

2 C + Ca CaC 2,

3 C + 4 Al = Al 4 C 3.

Aluminum carbide is a true carbide: by all four valence bonds, each carbon atom is bonded to a metal atom. Calcium carbide is an acetylenide because there is a triple bond between the carbon atoms. Therefore, when aluminum carbides interact with water, methane is released, and when calcium carbide interacts with water, acetylene is released.

Al 4 C 3 + 12H 2 O = 4Al (OH) 3 + 3CH 4,

CaC 2 + 2H 2 O = Ca (OH) 2 + C 2 H 2.

Bituminous coal is used as a fuel and is used to produce synthesis gas. Electrodes are made of graphite, graphite rods are used as a moderator

neutrons in nuclear reactors... Diamonds are used to make cutting tools, abrasives, cut diamonds (diamonds) are precious stones.

Silicon

Silicon is found in nature only in a bound form in the form of silica SiO2 and various silicic acid salts (silicates). It is the second most abundant chemical element (after oxygen) in the earth's crust (27.6%).

In 1811, the French J.L. Gay-Lussac and L.J. Tener obtained a brown-brown substance (silicon) by the reaction:

SiF 4 + 4 K = 4 KF + Si

and only in 1824 the Swede J. Berzelius, having obtained silicon by the reaction:

K 2 SiF 6 + 4 K = 6 KF + Si,

proved that it is a new chemical element. Now silicon is obtained from silica:

SiO 2 + 2 Mg = Si + 2 MgO,

3SiO 2 + 4Al = Si + 2Al 2 O 3,

reducing it with magnesium or carbon. It turns out also when decomposing silane:

SiH 4 = Si + 2 H 2.

In reactions with non-metals, silicon can be oxidized (i.e. Si is a reducing agent):

Si + O 2 = SiO 2,

Si + 2 F 2 = SiF 4,

Silicon is soluble in alkalis:

Si + 2 NaOH + H 2 O = Na 2 SiO 3 + 2 H 2,

insoluble in acids (except for hydrofluoric acid).

In reactions with metals, silicon exhibits oxidizing properties:

2 Mg + Si = Mg 2 Si.

When magnesium silicide is decomposed with hydrochloric acid, a silane is obtained:

Mg 2 Si + 4 HCl = 2MgCl 2 + SiH 4.

Silicon is used to obtain many alloys based on iron, copper

and aluminum. The addition of silicon to steel and cast iron improves their mechanical properties. Large additions of silicon impart acid resistance to iron alloys.

Ultrapure silicon is a semiconductor used in microchips and solar cells.

Oxygen compounds. Getting, properties and application

Carbon oxides

Carbon monoxide (II) (CO - carbon monoxide)

CO is a colorless and odorless poisonous gas, poorly soluble in water.

Receiving

In the laboratory, CO is obtained by decomposition of formic or oxalic acid (in the presence of concentrated H 2 SO 4):

HCOOH = CO + H 2 O,

H 2 C 2 O 4 = CO + CO 2 + H 2 O

or by heating zinc dust with calcium carbonate:

CaCO 3 + Zn = CaO + ZnO + CO.

In the factory, CO is produced by passing air or carbon dioxide through hot coal:

2C + O 2 = 2CO,

Properties

Poisonous action carbon monoxide is caused by the fact that the affinity of hemoglobin for carbon monoxide is greater than for oxygen. In this case, carboxyhemoglobin is formed and thereby the transfer of oxygen in the body is blocked.

Carbon (II) oxide is easily oxidized, burns in air with the release of a large amount of heat:

2 CO + O 2 = 2 CO 2 + 577 kJ / mol.

CO reduces many metals from their oxides:

FeO + CO = Fe + CO 2,

CuO + CO = Cu + CO 2.

CO easily enters into addition reactions:

CO + Cl 2 = COCl 2,

CO + NaOH = HCOONa,

Ni + 4 CO = Ni (CO) 4.

In industry, not pure CO is often used, but its various mixtures with other gases. Generator gas is obtained by passing air in a shaft furnace through hot coal:

2 C + O 2 = 2 CO + 222 kJ.

Water gas is obtained by passing water vapor through a hot coal:

C + H 2 O = CO + H 2 - 132 kJ.

The first reaction is exothermic, and the second is heat absorption. If both processes are alternated, then it is possible to maintain the required temperature in the oven. By combining the producer gas and the water gas, a mixed gas is obtained. These gases are used not only as fuel, but also for the synthesis of, for example, methanol:

CO + 2H 2 = CH 3 OH.

Carbon monoxide (IV) (CO 2 - carbon dioxide)

CO 2 is a colorless, odorless, incombustible gas. It is released when animals breathe. Plants absorb CO 2 and release oxygen. Air usually contains 0.03% carbon dioxide. Due to human activities (uncontrolled deforestation,

burning more and more coal, oil and gas), the content of CO 2 in the atmosphere gradually increases, which causes Greenhouse effect and threatens humanity with an ecological catastrophe.

Receiving

In the laboratory, CO 2 is obtained in the Kipp apparatus, acting with hydrochloric acid on marble:

CaCO 3 + 2HCl = CaCl 2 + H 2 O + CO 2.

Many reactions can be cited that result in CO 2:

KHCO 3 + H 2 SO 4 = KHSO 4 + H 2 O + CO 2,

C + O 2 = CO 2,

2 CO + O 2 = 2 CO 2,

Ca (HCO 3) 2 CaCO 3 Ї + CO 2 + H 2 O,

CaCO 3 = CaO + CO 2,

BaSO 4 + 2 C = BaS + 2 CO 2,

C + 2 H 2 SO 4 (conc) = CO 2 + 2 SO 2 + 2H 2 O,

C + 4 HNO 3 (conc) = CO 2 + 4 NO 2 + 2 H 2 O.

Properties

When CO 2 dissolves in water, carbonic acid is formed:

H 2 O + CO 2 = H 2 CO 3.

For CO 2, all those reactions are known that are characteristic of acidic oxides:

Na 2 O + CO 2 = Na 2 CO 3,

Ca (OH) 2 + 2 CO 2 = Ca (HCO 3) 2,

Ca (OH) 2 + CO 2 = CaCO 3 + H 2 O.

The ignited Mg continues to burn in carbon dioxide:

CO 2 + 2 Mg = 2 MgO + C.

Carbonic acid is a weak diacid:

H 2 O + CO 2 = H 2 CO 3

H + + HCO 3 - = H + + CO 3 2-

and can displace weaker acids from solutions of their salts:

Na 2 SiO 3 + CO 2 + H 2 O = H 2 SiO 3 + Na 2 CO 3,

KCN + CO 2 + H 2 O = KHCO 3 + HCN.

Carbonic acid salts. Carbonates and hydrocarbons

General methods for the preparation of salts are also typical for the preparation of carbonic acid salts:

CaCO 3 + CO 2 + H 2 O = Ca (HCO 3) 2,

Ca (HCO 3) 2 + Ca (OH) 2 = 2 CaCO 3 + 2 H 2 O.

Alkali metal and ammonium carbonates are readily soluble in water and

susceptible to hydrolysis. All other carbonates are practically insoluble:

Na 2 CO 3 + H 2 O = 2 Na + + OH - + HCO 3 -.

With relatively weak heating, bicarbonates decompose:

Ca (HCO 3) 2 = CaCO 3 + CO 2 + H 2 O.

When carbonates are calcined, metal oxides and CO 2 are obtained:

CaCO 3 = CaO + CO 2.

Carbonates are readily decomposed by stronger (than carbonic) acids:

MgCO 3 + 2HCl = MgCl 2 + CO 2 + H 2 O.

CaCO 3 + 2HCl = CaCl 2 + CO 2 + H 2 O.

When calcining carbonates with sand, SiO2 displaces a more volatile oxide:

Na 2 CO 3 + SiO 2 = Na 2 SiO 3 + CO 2.

Application

Sodium carbonate Na 2 CO 3 (soda ash) and its crystalline hydrate Na 2 CO 3 10H 2 O

(crystalline soda) are used in glass, soap, pulp and paper industries. Sodium bicarbonate NaHCO 3 (baking soda)

applied in Food Industry and in medicine. Limestone is a building stone and raw material for lime production.

Silicon (IV) oxides (SiO 2 )

Silica SiO 2 exists in nature in crystalline (mainly quartz) and amorphous (for example, opal SiO 2 nH 2 O) forms.

Receiving

SiO 2 is an acidic oxide that can be obtained by reactions:

Si + O 2 = SiO 2,

H 2 SiO 3 = SiO 2 + H 2 O,

SiH 4 + 2O 2 = SiO 2 + 2H 2 O.

Properties

When interacting with metals or carbon, SiO 2 can be reduced to silicon

SiO 2 + 2 Mg = Si + 2 MgO,

SiO 2 + 2 C = Si + 2 CO

or give carborundum (SiC) SiO 2 + 3 C = SiC + 2 CO.

When SiO 2 is fused with metal oxides, alkalis and some salts, silicates are formed:

SiO 2 + 2 NaOH = Na 2 SiO 3 + H 2 O,

SiO 2 + K 2 CO 3 = K 2 SiO 3 + CO 2,

SiO 2 + CaO = CaSiO 3.

Acids have no effect on SiO 2. An exception is hydrofluoric acid:

SiO 2 + 4HF = SiF 4 + 2H 2 O,

SiF 4 + 2HF = H 2,

SiO 2 + 6HF = H 2 + 2H 2 O.

Silicic acid H 2 SiO 3 is the simplest of the silicic acid family. Her general formula xSiO 2 yH 2 O. It can be obtained from silicates

Na 2 SiO 3 + 2 HCl = H 2 SiO 3 + 2 NaCl.

When heated, silicic acid decomposes:

H 2 SiO 3 = SiO 2 + H 2 O.

Silicates

Many hundreds of silicate minerals are known. They make up 75% of the mass crust... There are a lot of aluminosilicates among them. Silicates are the main constituents of cement, glass, concrete and brick.

Only Na and K silicates are soluble in water. Their aqueous solutions are called "liquid glass". During hydrolysis, these solutions are alkaline. They are used for the manufacture of acid-resistant cement and concrete.

OXIDATION-REDUCTION REACTIONS

Reactions in which there is a change in the oxidation states of the atoms of the elements that make up the reacting compounds, are called redox.

Oxidation state(s.r.) is the charge of an element in a compound, calculated based on the assumption that the compound consists of ions... Determination of the oxidation state is carried out using the following provisions:

1. The oxidation state of an element in a simple substance, for example, in Zn, Ca, H 2, Br 2, S, O 2, is zero.

2. The degree of oxygen oxidation in compounds is usually –2. Exceptions are peroxides H 2 +1 O 2 –1, Na 2 +1 O 2 –1 and oxygen fluoride O +2 F 2.

3. The oxidation state of hydrogen in most compounds is +1, with the exception of salt-like hydrides, for example, Na +1 H -1.

4. Alkali metals (+1) have a constant oxidation state; beryllium Be and magnesium Mg (+2); alkaline earth metals Ca, Sr, Ba (+2); fluorine (–1).

5. Algebraic sum of degrees oxidation of elements in a neutral molecule is zero, in a complex ion - the charge of the ion.

As an example, we calculate the oxidation state of chromium in the compound К 2 Cr 2 O 7 and nitrogen in the anion (NO 2) -

K 2 +1 Cr 2 NS O 7 –2 2 ∙ (+1) + 2 x + 7 (–2) = 0 x = + 6

(NO 2) - x + 2 (–2) = –1 x = + 3

In redox reactions, electrons from one atom, molecule or ion are transferred to another. Oxidation – the process of giving up electrons by an atom, molecule or ion, accompanied by an increase in the oxidation state. Recovery – the process of electron attachment, accompanied by a decrease in the oxidation state.

-4 -3 -2 -1 0 +1 +2 +3 +4 +5 +6 +7 +8
Recovery process

Oxidation and reduction are interrelated processes occurring simultaneously.

Oxidizing agents are called substances (atoms, ions or molecules) that attach electrons during the reaction, restorers – electron donating substances... Oxidizing agents can be halogen atoms and oxygen, positively charged metal ions (Fe 3+, Au 3+, Hg 2+, Cu 2+, Ag +), complex ions and molecules containing metal atoms in the highest oxidation state (KMnO 4, K 2 Cr 2 O 7, NaBiO 3, etc.), nonmetal atoms in a positive oxidation state (HNO 3, concentrated H 2 SO 4, HClO, HClO 3, KClO 3, NaBrO, etc.).

Typical reducing agents are almost all metals and many non-metals (carbon, hydrogen) in a free state, negatively charged ions of non-metals (S 2-, I -, Br -, Cl -, etc.), positively charged metal ions in the lowest oxidation state (Sn 2+, Fe 2+, Cr 2+, Mn 2+, Cu +, etc.).

Compounds containing elements in the maximum and minimum oxidation states can be, respectively, or only oxidizing agents (KMnO 4, K 2 Cr 2 O 7, HNO 3, H 2 SO 4, PbO 2), or only reducing agents (KI, Na 2 S, NH 3). If the substance contains an element in an intermediate oxidation state, then, depending on the reaction conditions, it can be both an oxidizing agent and a reducing agent. For example, potassium nitrite KNO 2, containing nitrogen in the oxidation state +3, hydrogen peroxide H 2 O 2, containing oxygen in the oxidation state -1, in the presence of strong oxidants exhibit reducing properties, and when interacting with active reducing agents are oxidizing agents.

When drawing up equations for redox reactions, it is recommended to adhere to the following order:

1. Write the formulas of the starting materials. Determine the oxidation state of elements that can change it, find an oxidizing agent and a reducing agent. Write the reaction products.

2. Make up the equations of oxidation and reduction processes. Select the factors (basic coefficients) so that the number of electrons donated during oxidation is equal to the number of electrons received during reduction.

3. Place the coefficients in the reaction equation.

K 2 Cr 2 +6 O 7 + 3H 2 S -2 + 4H 2 SO 4 = Cr 2 +3 (SO 4) 3 + 3S 0 + K 2 SO 4 + 7H 2 O

oxidizing agent reducing agent medium

oxidation S -2 - 2ē → S 0 ½3

recovery 2Cr +6 + 6ē → 2Cr +3 ½1

The nature of many redox reactions depends on the environment in which they occur. To create an acidic environment, diluted sulfuric acid, to create alkaline - solutions of sodium or potassium hydroxides.

There are three types of redox reactions: intermolecular, intramolecular, disproportionation. Intermolecular redox reactions - these are reactions in which an oxidizing agent and a reducing agent are in different substances ... The reaction discussed above is of this type. TO intramolecular include reactions in which the oxidizing agent and the reducing agent are in the same substance.

2KCl +5 O 3 -2 = 2KCl -1 + 3O 2 0

reduction Сl +5 + 6ē → Cl - ½2 Cl +5 - oxidizing agent

oxidation 2O -2 - 4ē → O 2 0 ½3 O -2 - reducing agent

In reactions disproportionation(self-oxidation - self-healing) molecules of the same substance react with each other as an oxidizing agent and as a reducing agent.

3K 2 Mn +6 O 4 + 2H 2 O = 2KMn +7 O 4 + Mn +4 O 2 + 4KOH

oxidation Mn +6 - ē → Mn +7 ½ 2 Mn +6 - reducing agent

reduction Mn +6 + 2ē → Mn +4 ½ 1 Mn +6 - oxidizing agent

pliz at least something Using the electronic balance method, select the coefficients in the schemes of redox reactions and indicate the oxidation process

and recovery:

1. P + HNO3 + H2O = H3PO4 + NO

2. P + HNO3 = H3PO4 + NO2 + H2O

3. K2Cr2O7 + HCl = Cl2 + KCl + CrCl3 + H20

4. KMnO4 + H2S + H2SO4 = MnSO4 + S + K2SO4 + H2O

5. KMnO4 + HCl = Cl2 + MnCl2 + KCl + H2O

Using the electronic balance method, select the coefficients in the schemes of redox reactions and indicate the oxidation and reduction process:

CuO + NH3 = Cu + N2 + H2O

Ag + HNO3 = AgNO3 + NO + H2O

Zn + HNO3 = Zn (NO3) 2 + N2 + H2O

Cu + H2SO4 = CuSO4 + SO2 + H2O

Help solve: ELECTROLYTIC DISSOCIATION. OXIDATION-REDUCTION REACTIONS

Part A
A2 When studying electrical conductivity various substances using a special device, the students observed the following:

Which of the following was in the glass?
1) sugar (solution)
2) KC1 (tv.) 3) NaOH (p-p) 4) alcohol
A4 The abbreviated ionic equation corresponds to the interaction of solutions of barium chloride and sulfuric acid
1) H + + SG = HC1
2) Ba2 + + SO42- = BaSO4
3) CO32- + 2H + = H2O + CO2
4) Ba2 + + COz2- = BaCOz
A5 The reaction between solutions of silver nitrate and hydrochloric acid proceeds to the end, since
1) both substances are electrolytes
2) silver nitrate is a salt
3) insoluble silver chloride is formed
4) soluble nitric acid is formed

A7 The equation H + + OH = H2O reflects the essence of the interaction

1) hydrochloric acid and barium hydroxide
2) sulfuric acid and copper hydroxide (P)
3) phosphoric acid and calcium oxide
4) silicic acid and sodium hydroxide

A10 The oxidation process corresponds to the scheme
1) S + 6 → S + 4
2) Cu + 2 → Cu0
3) N + 5 → N-3
4) C-4 → C + 4

Part B

B2 Establish a correspondence between the formula of the substance and total ions formed during the complete dissociation of 1 mol of this substance: for each position from the first column, select the corresponding position from the second column, indicated by a number.
FORMULA NUMBER OF IONS (IN MOLES)
A) A1 (NO3) 3 1) 1 B) Mg (NO3) 2 2) 2
B) NaNO3 3) 3 D) Cu (NO3) 2 4) 4
5) 5

Write down the selected numbers in the table under the corresponding letters.

Transfer the answer in the form of a sequence of four numbers to the test form under the number of the corresponding task, without changing the order of the numbers.

You are presented with a list of interrelated concepts:

A) acid
B) hydrochloric acid
B) anoxic acid
D) strong electrolyte
Write down the letters that designate concepts in the table so that a chain is traced from a particular concept to the most general one.

Transfer the resulting sequence of letters to the test form without changing the order of the letters.

Lesson type. Acquisition of new knowledge.

Lesson objectives.Educational. To acquaint students with the new classification of chemical reactions based on changes in the oxidation states of elements - with redox reactions (ORR); to teach students to arrange the coefficients using the electronic balance method.

Developing. Continue the development of logical thinking, the ability to analyze and compare, the formation of interest in the subject.

Educational. To form the scientific outlook of students; improve labor skills.

Methods and methodological techniques. Storytelling, conversation, demonstration of visual aids, independent work students.

Equipment and reagents. Reproduction with the image of the Colossus of Rhodes, an algorithm for arranging coefficients using the electronic balance method, a table of typical oxidants and reducing agents, a crossword puzzle; Fe (nail), solutions of NaOH, CuSO 4.

DURING THE CLASSES

Introductory part

(motivation and goal setting)

Teacher. In the III century. BC. on the island of Rhodes, a monument was built in the form of a huge statue of Helios (among the Greeks - the sun god). The grandiose design and perfection of the execution of the Colossus of Rhodes - one of the wonders of the world - amazed everyone who saw him.

We do not know exactly what the statue looked like, but it is known that it was made of bronze and reached a height of about 33 m. The statue was created by the sculptor Hareth and took 12 years to build.

The bronze shell was attached to an iron frame. The hollow statue began to be built from below and, as it grew, was filled with stones to make it more stable. About 50 years after completion, the Colossus collapsed. During the earthquake, it broke at the knee level.

Scientists believe that the real reason for the fragility of this miracle was the corrosion of the metal. And at the heart of the corrosion process are redox reactions.

Today in the lesson you will get acquainted with redox reactions; you will learn about the concepts of "reducing agent" and "oxidizing agent", about the processes of reduction and oxidation; learn to place the coefficients in the equations of redox reactions. Write down the number, topic of the lesson in your workbooks.

Learning new material

The teacher performs two demonstration experiments: the interaction of copper (II) sulfate with alkali and the interaction of the same salt with iron.

Teacher. Write down the molecular equations of the reactions performed. In each equation, arrange the oxidation states of the elements in the formulas for the starting materials and reaction products.

The student writes down the reaction equations on the board and arranges the oxidation states:

Teacher. Did the oxidation states of the elements change in these reactions?

Student. In the first equation, the oxidation states of the elements did not change, and in the second they changed - for copper and iron.

Teacher. The second reaction is redox. Try to define redox reactions.

Student. Reactions as a result of which the oxidation states of the elements that make up the reactants and reaction products change are called redox reactions.

Students write down in a notebook under the dictation of the teacher the definition of redox reactions.

Teacher. What happened as a result of the redox reaction? Before the reaction, iron had an oxidation state of 0, after the reaction it became +2. As you can see, the oxidation state has increased, therefore, iron gives 2 electrons.

For copper, before the reaction, the oxidation state is +2, after the reaction - 0. As you can see, the oxidation state has decreased. Consequently, copper takes 2 electrons.

Iron donates electrons, it is a reducing agent, and the process of electron transfer is called oxidation.

Copper accepts electrons, it is an oxidizing agent, and the process of attaching electrons is called reduction.

Let's write down the diagrams of these processes:

So, give a definition of the terms "reducing agent" and "oxidizing agent".

Student. The atoms, molecules or ions that donate electrons are called reducing agents.

The atoms, molecules, or ions that attach electrons are called oxidants.

Teacher. What definition can be given to the processes of reduction and oxidation?

Student. Reduction is the process of attaching electrons to an atom, molecule or ion.

Oxidation refers to the transfer of electrons by an atom, molecule or ion.

Students write down definitions under dictation in a notebook and complete a drawing.

Remember!

Donate electrons - oxidize.

Take electrons - recover.

Teacher. Oxidation is always accompanied by reduction, and vice versa, reduction is always associated with oxidation. The number of electrons donated by the reducing agent is equal to the number of electrons donated by the oxidizing agent.

To select the coefficients in the equations of redox reactions, two methods are used - electronic balance and electronic-ion balance (half-reaction method).

We will only consider the electronic balance method. To do this, we use the algorithm for arranging the coefficients using the electronic balance method (drawn up on a sheet of Whatman paper).

EXAMPLE Arrange the coefficients in this reaction scheme using the electronic balance method, determine the oxidizing agent and the reducing agent, indicate the oxidation and reduction processes:

Fe 2 O 3 + CO Fe + CO 2.

Let's use the algorithm for arranging the coefficients using the electronic balance method.

3. Let us write out the elements that change the oxidation state:

4. Let's compose electronic equations, determining the number of given and received electrons:

5. The number of donated and received electrons must be the same, because neither the starting materials nor the reaction products are charged. We equalize the number of electrons given and received by choosing the least common multiple (LCM) and additional factors:

6. The resulting factors are coefficients. Let's transfer the coefficients to the reaction scheme:

Fe 2 O 3 + 3CO = 2Fe + 3CO 2.

Substances that are oxidizing or reducing agents in many reactions are called typical.

A table is hung out on a sheet of Whatman paper.

Teacher. Redox reactions are very common. They are associated not only with corrosion processes, but also fermentation, decay, photosynthesis, metabolic processes occurring in a living organism. They can be observed during fuel combustion. Redox processes accompany the cycles of substances in nature.

Did you know that about 2 million tons of nitric acid is generated in the atmosphere every day, or
700 million tons per year, and in the form of a weak solution falls on the ground with rain (man produces only 30 million tons of nitric acid per year).

What is happening in the atmosphere?

Air contains 78% by volume nitrogen, 21% oxygen and 1% other gases. Under the influence of lightning discharges, and on the Earth an average of 100 lightning flashes every second, the interaction of nitrogen molecules with oxygen molecules occurs with the formation of nitric oxide (II):

Nitric oxide (II) is easily oxidized by atmospheric oxygen to nitric oxide (IV):

NO + O 2 NO 2.

The formed nitric oxide (IV) interacts with atmospheric moisture in the presence of oxygen, turning into nitric acid:

NO 2 + H 2 O + O 2 HNO 3.

All these reactions are redox reactions.

Exercise ... Arrange the coefficients in the given reaction schemes using the electronic balance method, indicate the oxidizing agent, reducing agent, oxidation and reduction processes.

Solution

1. Determine the oxidation state of the elements:

2. We emphasize the symbols of the elements whose oxidation states change:

3. Let us write out the elements that have changed the oxidation states:

4. Let's compose the electronic equations (determine the number of electrons given and received):

5. The number of electrons donated and received is the same.

6. Let's transfer the coefficients from the electronic circuits to the reaction circuit:

Further, students are invited to independently arrange the coefficients using the electronic balance method, to determine the oxidizing agent, reducing agent, to indicate the oxidation and reduction processes in other processes occurring in nature.

The other two reaction equations (with coefficients) are:

Checking the correctness of the tasks is carried out using an overhead scope.

Final part

The teacher invites students to solve a crossword puzzle based on the material studied. The result of the work is submitted for verification.

Having guessed crossword, you will learn that the substances KMnO 4, K 2 Cr 2 O 7, O 3 are strong ... (along the vertical (2)).

Horizontally:

1. What process does the diagram reflect:

3. Reaction

N 2 (g) + 3H 2 (g) 2NH 3 (g) + Q

is redox, reversible, homogeneous,….

4.… carbon (II) is a typical reducing agent.

5. What process does the diagram reflect:

6. For the selection of coefficients in the equations of redox reactions use the method of electronic ....

7. According to the scheme, aluminum gave ... an electron.

8. In reaction:

Н 2 + Сl 2 = 2НCl

hydrogen H 2 -….

9. What type of reactions are always only redox?

10. The oxidation state of simple substances is….

11. In reaction:

reducing agent -….

Home assignment. According to OS Gabrielyan's textbook "Chemistry-8" § 43, p. 178-179, exercise. 1, 7 in writing.

Task (at home). Constructors first spaceships and submarines are faced with a problem: how to maintain a constant air composition on a ship and space stations? Get rid of excess carbon dioxide and replenish oxygen? The solution was found.

Potassium superoxide KO 2, as a result of interaction with carbon dioxide, forms oxygen:

As you can see, this is a redox reaction. Oxygen in this reaction is both an oxidizing agent and a reducing agent.

In a space expedition, every gram of cargo counts. Calculate the supply of potassium superoxide that must be taken on a space flight if the flight is designed for 10 days and if the crew consists of two people. It is known that a person exhales 1 kg of carbon dioxide per day.

(Answer. 64.5 kg KO 2. )

Assignment (increased level of difficulty). Write down the equations of redox reactions that could lead to the destruction of the Colossus of Rhodes. Keep in mind that this gigantic statue stood in a port city on an island in the Aegean Sea, off the coast of modern Turkey, where the humid Mediterranean air is saturated with salts. It was made of bronze (an alloy of copper and tin) and mounted on an iron frame.

Literature

Gabrielyan O.S.... Chemistry-8. M .: Bustard, 2002;
Gabrielyan O.S., Voskoboinikova N.P., Yashukova A.V. Handbook of the teacher. 8th grade. M .: Bustard, 2002;
Cox R., Morris N... Seven wonders of the world. The ancient world, the Middle Ages, our time. M .: BMM AO, 1997;
Small children's encyclopedia. Chemistry. M .: Russian encyclopedic partnership, 2001; Encyclopedia for children "Avanta +". Chemistry. T. 17.M .: Avanta +, 2001;
Khomchenko G.P., Sevastyanova K.I. Redox reactions. M .: Education, 1989.