Oxidation is the process of donating electrons, with an increase in the degree of oxidation.
At oxidation substances as a result of recoil electrons it increases oxidation state. atoms oxidizable substance are called donors electrons and atoms oxidizing agent - acceptors electrons.
In some cases, when oxidized, the molecule of the original substance can become unstable and break up into more stable and smaller constituents (see. free radicals). Moreover, some of the atoms of the resulting molecules have a higher degree of oxidation than the same atoms in the original molecule.
The oxidizing agent, accepting electrons, acquires reducing properties, turning into a conjugated reducing agent:
oxidizer+ e − ↔ conjugated reducing agent.
Recovery
Recovery the process of adding electrons to an atom of a substance is called, while its oxidation state goes down.
When recovering atoms or ions join electrons. At the same time, there is a decrease oxidation states 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 − ↔ conjugated oxidizer.
An unbound, free electron is the strongest reducing agent.
Redox reactions are reactions in which reactants gain or donate electrons. An oxidizing agent is a particle (ion, molecule, element) that attaches electrons and at the same time passes from a higher oxidation state to a lower one, i.e. is being restored. A reducing agent is a particle that donates electrons and passes from a lower oxidation state to a higher one, i.e. oxidized.
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-healing) - reactions in which atoms with an intermediate oxidation state are converted into an equimolar mixture of atoms with higher and lower oxidation states, for example:
Cl 2 + H 2 O → HClO + HCl
Reproportionation (comproportionation) - reactions 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 are transferred from one atom, molecule, or ion to another. The process of donating electrons is oxidation. When oxidized, the oxidation state increases:
The process of electron addition is recovery. When reduced, the oxidation state decreases:
The atoms or ions that gain electrons in this reaction are oxidizing agents, and those that donate electrons are reducing agents.
Redox reactions (electrode potential)
Electrons can act as chemical reactants, and the half-reaction is practically used in devices called galvanic cells.
An example of an electrode is a plate of crystalline zinc immersed in a solution of zinc sulfate. After the plate is immersed, 2 processes take place. As a result of the first process, the plate acquires a negative charge, after 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 in various electrochemical measurements and in galvanic cells. A hydrogen electrode (HE) is a plate or wire made of metal that absorbs gaseous gases well. hydrogen(usually used platinum or palladium), saturated with hydrogen (at atmospheric pressure) and immersed in water solution, containing hydrogen ions. The plate potential depends on [ specify ] on the concentration of H + ions in the solution. The electrode is a standard against which the electrode potential of the determined chemical reaction is measured. At a hydrogen pressure of 1 atm, a proton concentration in the solution of 1 mol/l and a temperature of 298 TO the SE potential is taken equal to 0 V. When assembling a galvanic cell from SE and the electrode to be determined, the following reaction proceeds reversibly on the surface of platinum:
2Н + + 2e − = H 2
that is, either recovery hydrogen, or oxidation- it depends on the potential of the reaction taking place on the electrode being determined. By measuring the EMF of a galvanic electrode under standard conditions (see above), determine standard electrode potential defined 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 solubility product, to determine the rate constants of some electrochemical reactions.
Nernst equation
The dependence of the redox potential corresponding to the reduction half-reaction of the permanganate ion in an acidic medium (and, as already noted, simultaneously the half-reaction of the oxidation of the Mn 2+ cation to the permanganate ion in an acidic medium) on the factors listed above that determine 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 this particle in the half-reaction equation, n is the number of electrons accepted by the oxidizing agent, R is the universal gas constant, T- temperature, F is the Faraday number.
Measure the redox potential in the reaction vessel while the reaction is running, i.e. under non-equilibrium 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 a metal conductor connecting the electrodes. In this case, the electron transfer rate (current strength) must be maintained 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) concentrations of particles. Although the potentials of the redox couples during the course of the reaction cannot be measured, they can be calculated by substituting the current ones into the Nernst equation, i.e. responsible present moment concentration time. If the change in potential is considered as the reaction proceeds, then first these are the initial concentrations, then the time-dependent concentrations, and, finally, after the termination of the reaction, the equilibrium ones. As the reaction progresses, the oxidizing agent potential calculated from the Nernst equation decreases, while the reductant potential corresponding to the second half-reaction, on the contrary, increases. When these potentials are equalized, the reaction stops and the system enters 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. By Werner's coordination theory distinguishes between the inner and outer spheres in each complex compound. The central atom with its surrounding ligands form the inner sphere of the complex. It is usually enclosed in square brackets. Everything else in a complex compound is the outer sphere and is written in square brackets. Around the central atom is placed a certain number of ligands, which is determined by the coordination number. The number of coordinated ligands is most often 6 or 4. The ligand occupies a coordination site 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 the free state. More firmly bound particles of the inner sphere are called a complex (complex ion). Attractive forces act between the central atom and ligands (formed covalent bond according to the exchange and (or) donor-acceptor mechanism), between ligands - repulsive forces. If the charge of the inner sphere is 0, then there is no outer coordination sphere. The central atom (complexing agent) is an atom or ion that occupies a central position in a complex compound. The role of a complexing agent is most often performed by particles that have free orbits 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 IB and VIIIB. Rarely, neutral atoms of d-elements and non-metal atoms in various degrees of oxidation - 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 associated with the complexing agent and are donors of electron pairs. These are electron-rich systems that have free and mobile electron pairs and 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 being a donor of two electron pairs. The charge of the inner sphere of a complex compound is the algebraic sum of the charges of its constituent particles. Complex compounds having an ionic outer sphere undergo dissociation in solution into a complex ion and outer sphere ions. They behave in dilute solutions as strong electrolytes: dissociation proceeds 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. 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 group IA is ns1 At the external level, 1 electron. Far from the nucleus. Low ionization potentials. Atoms give up 1 electron. Metallic means are pronounced. 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) Stored under a layer of liquid hydrocarbons (benzene, gasoline, kerosene) 2) Oxidizers. Easily oxidize alkali metals to halides, sulfides, phosphides. Li Na K Rb Cs increase in the metallic radius decrease in ionization energy decrease in electronegativity decrease in melting and boiling points Use of sodium and potassium 1. Preparation of 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 I A and I B groups of the periodic system Alkali 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 are formed, called alkalis. IN Periodic system they immediately follow the inert gases, so the peculiarity of the structure of alkali metal atoms is that they contain one electron in the outer energy level: their electronic configuration is 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 of this subgroup are silver-white (except 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. Sea water contains sodium chloride NaCl, and the soil contains potassium salts - sylvin 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, roentgenium Rg also 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 mint metals is not officially applied to the 11th group of elements, since 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. Also characteristic are high values of density, melting and boiling points, high thermal and electrical conductivity. A feature of the elements of the subgroup is the presence of a filled pre-outer -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 component 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 widely distributed, for example water H20. The composition of natural combustible gas is mainly a combination of carbon and hydrogen - methane CH4 - Hydrogen is also found in many organic substances. 1) If you ignite hydrogen (after checking for purity, see below) and lower the 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 detonating gas. If an explosion occurs in a glass vessel, then its fragments may be injured.
hurt those around you. Therefore, before igniting hydrogen, it is necessary to check its purity. To do this, collect hydrogen 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 out with an explosion. When working with hydrogen, safety regulations must be observed. 2) If, for example, when heated, a jet 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 away oxygen from copper atoms. The reduction process is the opposite of the oxidation process. Substances that take away oxygen are referred to as reducing agents. 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 with an explosion with a slight activation of it. The interaction with Br2 and I2 proceeds more slowly. For the reaction to proceed with hydrogen, it is sufficient to activate a small fraction of the reactants by means of illumination or heating. Activated particles interact with non-activated ones, forming HX and new activated particles, which continue the process, and the reaction of two activated particles according to the main reaction ends with the formation of a product. 4) Oxidation reactions. When hydrogen is heated with metals I and II of the 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, unlimitedly 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 an unstable crystalline hydrate H2O2 2H2O. Hydrogen peroxide has oxidizing as well as reducing properties. It oxidizes nitrites to nitrates, releases iodine from metal iodides, breaks down unsaturated compounds at the site of double bonds. Hydrogen peroxide reduces salts of gold and silver, as well as oxygen when reacting with an aqueous solution of potassium permanganate in an acidic environment. When H2O2 is reduced, H2O or OH- is formed, for example: H2O2 + 2KI + H2SO4 = I2 + K2SO4 + 2H2O chemical analysis to determine the content of H2O2: 5H2O2 + 2KMnO4 + 3H2SO4 → 5O2 + 2MnSO4 + K2SO4 + 8H2O Oxidation of organic compounds with hydrogen peroxide (for example, sulfides and thiols) is advisable to carry out in an acetic acid medium.
29. general characteristics of the properties of elements and their compounds 2 a groups. physical and chemical properties, application. There are 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 with opposite spins, with the expenditure of the necessary energy, one element from the s state passes into the p state. These Metals, but they are less active than alkaline. most distributed in nature Mg Ca Be, found in the form of the mineral Be3AL2 (SiO3) 6 Production method: electrolysis of melted chlorides Physical properties: light metals, but harder alkali metals. Chemical properties: 1 In air, the surface of Be and Mg is covered with an oxide film. 2. interaction with nitrogen at high temperature 3. does not interact with water Be 4. displaces hydrogen from acids (except nitric acid). become. Calcium and its hydride are also used to obtain hard-to-recover metals such as chromium, thorium and uranium. Alloys of calcium with lead are used in batteries and bearing alloys. Calcium granules are also used to remove traces of air from electrovacuum devices.
№31 Alkaline earth metals - chemical elements 2nd group of the main subgroup, except for beryllium and magnesium: calcium, strontium, barium And radium. They belong to the 2nd group of elements according to the new classification IUPAC. So named because they oxides- "lands" (in the terminology alchemists) - report water alkaline reaction. salt alkaline earth metals, except for radium, are widely distributed 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 highest oxidation state: CO2, SO3, P2O5, CrO3, Mn2O7 and amphoteric, for example, ZnO, Al2O3, Cr2O3. Oxides are obtained by burning simple and complex substances, as well as by 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 react 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. Acid oxides interact with both acids and alkalis: ZnO + 2HCl = ZnCl2 + H2O
ZnO + 2NaOH = Na2ZnO2 + H2O
Hydroxides ( hydroxides) - compounds of oxides of chemical elements with a vault. Hydroxides of almost all chemical elements are known; some of them are found in nature in the form of minerals. Alkali metal hydroxides are called alkalis. Depending on whether the corresponding oxide is basic, acidic or amphoteric, one distinguishes accordingly:
basic hydroxides (grounds) - hydroxides that exhibit basic properties (for example, calcium hydroxide Ca (OH) 2, potassium hydroxide KOH, sodium hydroxide NaOH, 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 acid properties(for example, aluminum hydroxide Al (OH) 3, zinc hydroxide Zn (OH) 2).
Carbonates and hydrocarbonates - salts and esters carbonic acid (H 2 CO 3). Among the salts, normal carbonates are known (with the CO 3 2− anion) and acidic or bicarbonates(With anion HCO 3 −).
Chemical properties
When heated, acidic carbonates turn into normal carbonates:
With strong heating, 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: Calcium, magnesium, barium carbonates, etc. are used in construction, 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 buffer substances , regulating the constancy of the reaction blood .
Silicates and aluminosilicates are a large group 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 igneous (pyroxenes, feldspars ), they are also characteristic of pegmatites (mica, tourmaline, beryl, etc.) and skarns (garnets, wollastonite). Widely distributed in metamorphic rocks - shales 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 (the main subgroup), as well as scandium, yttrium, lanthanum and lanthanides, actinium and actinides (side subgroup).
At the outer 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, compounds are typical only with an oxidation state of +3. The elements of the gallium subgroup (gallium, indium, thallium) also have three electrons in the outer electronic level, forming the s 2 p 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 are enhanced.
The elements of the scandium subgroup also have three electrons in the outer electronic level. However, these elements are transitional d-elements, the electronic configuration of their valence layer is d 1 s 2 . These electrons donate all three elements rather easily. The elements of the lanthanide subgroup have a distinctive configuration of the outer electronic level: the 4f level builds up in them and the d level disappears. Starting with cerium, all elements, except for gadolinium and lutetium, have an electronic configuration of the outer 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 of group III have a very strong affinity for oxygen, and the formation of their oxides is accompanied by the release of a large amount of heat.
Elements of the III group find the most diverse application.
33. Physical properties. Aluminum is a silvery-white light metal that melts at 660°C. Very ductile, 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 light.
Chemical properties. Aluminum is a very active metal. In a series of voltages, it is after the 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 interact vigorously with oxygen and water vapor in the air, turning into a loose mass - aluminum hydroxide:
4 Al + 3 O 2 + 6 H 2 O \u003d 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 release of hydrogen:
2 Al + 6 H 2 O \u003d 2 Al (OH) 3 + 3 H 2
Aluminum dissolves well in dilute sulfuric and hydrochloric acids:
2 Al + 6 Hcl \u003d 2 AlCl 3 + 3 H 2
2 Al + 3 H 2 SO 4 \u003d 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 \u003d 8 Al (NO 3) 3 + 3 N 2 O + 15 H 2 O
Concentrated nitric acid passivates aluminum.
Since aluminum oxide and hydroxide have amphoteric
properties, aluminum is easily soluble in aqueous solutions of all alkalis, except for ammonium hydroxide:
2 Al + 6 KOH + 6 H 2 O \u003d 2 K 3 [Al (OH) 6] + 3 H 2
Powdered aluminum readily reacts 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 \u003d 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 \u003d Al 4 C 3 (aluminum carbide)
Aluminum sulfide 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 \u003d 2 Al (OH) 3 + 3 H 2 S
Aluminum easily takes away 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 \u003d 9 Fe + 4 Al 2 O 3
The process of reducing metals from their oxides with aluminum is called aluminothermy. Aluminothermy is used in the production of 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 fuse, a mixture of fine aluminum powder and magnetic iron ore Fe 3 O 4 (termite) is ignited, the reaction proceeds spontaneously with the mixture heated to 3500 ° C. Iron at this temperature is in a molten state.
Receipt. For the first time, aluminum was obtained by reduction from aluminum chloride with sodium metal:
AlCl 3 + 3 Na = 3 NaCl + Al
At present, it is obtained by electrolysis of molten salts in electrolytic baths (Fig. 46). The electrolyte is a melt containing 85-90% cryolite - a complex salt of 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 used very widely. Foil is made from it, which is 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: the 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 withstand heating up to 1000 ° C without being corroded. Aluminum alloys, which are distinguished by their great lightness and strength, are used in the production of heat exchangers, 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 transparent crystals of corundum of a beautiful shape and color. Corundum dyed red with chromium compounds is called ruby, and blue dyed with titanium and iron compounds is called sapphire. Ruby and sapphire are precious stones. Currently, they are quite easily obtained artificially.
Bor-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 the 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.
Receipt
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 metallothermy (more often reduction 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). It has brittleness and semiconductor properties (wide-gap semiconductor).
Chemical properties
In many physical and chemical properties, the non-metal boron resembles silicon.
Chemical boron is rather 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, 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 alkali or alkaline earth metal borides with acid:
When heated strongly, boron exhibits reducing properties. He is able, for example, to restore silicon or phosphorus from 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 acid, sulfuric acid, and aqua regia, boron dissolves to form boric acid.
Boron oxide is a typical acidic 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 anion BO 3 3-), but tetraborates, for example:
Application
Elementary Boron
Boron (in the form of fibers) serves as a strengthening agent for many composite materials.
Also, boron is often used in electronics to change the type of conductivity. silicon.
Boron is used in metallurgy as a microalloying element, which significantly increases the hardenability of steels.
34.haracharacteristics 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 show 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 oxygen subgroup, sulfur and with halogens.The oxidation state +4 is obtained by the transition of the 1s electron to a free p-orbital.
With an increase in the radius of an atom, the strength of the bond between the outer electrons and the nucleus decreases. Non-metallic properties decrease, and metal ones increase. (melting and boiling points decrease, etc.)
Carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb) - elements of group 4 of the main subgroup of PSE. On the outer electron layer, the atoms of these elements have 4 electrons: ns 2 np 2. In the subgroup, with an increase in the ordinal number of the element, the atomic radius increases, non-metallic properties weaken, and metallic properties increase: carbon and silicon are non-metals, germanium, tin, lead are metals.
General characteristics. Carbon and silicon
The carbon subgroup, which includes carbon, silicon, germanium, tin and lead, is the main subgroup of group 4 of the Periodic system.
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. When moving from top to bottom in the group, non-metallic properties are weakened, and 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 of +2 and +4. When elucidating the mechanism of reactions, the carbon isotope 13 is sometimes used C (labeled atom method). Therefore, it is useful to know that the abundance of carbon isotopes: 12 C - 98.89% and 13 C - 1.11%. If we confine ourselves to listing 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.
From the compounds of elements (E) of the subgroup of carbon with hydrogen, we consider compounds of the EN 4 type. With an increase in the charge of the nucleus of the atom E, the stability of hydrides decreases.
In the transition from C to Pb, the stability of compounds with an oxidation state of +4
decreases, and with +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 is contained in fossil coals: from 92% - in anthracite, up to 80% - in brown coal. In a bonded state, carbon is found in carbides: CaCO 3 chalk, limestone and marble, MgCO 3 CaCO 3 - dolomite,
MgCO 3 - magnesite. The air contains carbon in the form of carbon dioxide (0.03% by volume). Carbon is also contained in compounds dissolved in sea water.
Carbon is found in plants and animals, and is found in oil and natural gas.
In reactions with active non-metals, carbon is easily oxidized:
2 C + O 2 \u003d 2 CO,
C + 2 F 2 = CF 4 .
Carbon can also exhibit reducing properties when interacting with complex substances:
C + 2 CuO \u003d 2 Cu + CO 2,
C + 2 H 2 SO 4 (conc) = CO 2 + 2 SO 2 + H 2 O,
2 C + BaSO 4 \u003d BaS + 2 CO 2.
In reactions with metals and less active non-metals, carbon is an oxidizing agent: 2C + H 2 \u003d C 2 H 2,
2 C + Ca CaC 2 ,
3 C + 4 Al \u003d Al 4 C 3.
Aluminum carbide is a true carbide: with all four valence bonds, each carbon atom is connected to metal atoms. 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 \u003d 4Al (OH) 3 + 3CH 4,
CaC 2 + 2H 2 O \u003d Ca (OH) 2 + C 2 H 2.
Coal is used as a fuel, used to produce synthesis gas. Electrodes are made from 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 occurs in nature only in a bound form in the form of silica SiO2 and various salts of silicic acid (silicates). It is the second most abundant chemical element (after oxygen) in the earth's crust (27.6%).
In 1811, the Frenchmen J.L. Gay-Lussac and L.J. Tener received a brown-brown substance (silicon) by the reaction:
SiF 4 + 4 K = 4 KF + Si
and only in 1824 the Swede J. Berzelius, having received silicon according to 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 \u003d Si + 2 MgO,
3SiO 2 + 4Al \u003d Si + 2Al 2 O 3,
reducing it with magnesium or carbon. It also turns out when the silane is decomposed:
SiH 4 \u003d Si + 2 H 2.
In reactions with non-metals, silicon can be oxidized (i.e., Si is a reducing agent):
Si + O 2 \u003d SiO 2,
Si + 2 F 2 \u003d SiF 4,
Silicon is soluble in alkalis:
Si + 2 NaOH + H 2 O \u003d Na 2 SiO 3 + 2 H 2,
insoluble in acids (except hydrofluoric).
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 \u003d 2MgCl 2 + SiH 4.
Silicon is used to produce many alloys based on iron, copper
and aluminium. The addition of silicon to steel and cast iron improves their mechanical properties. Large additions of silicon give iron alloys acid resistance.
Ultra-pure silicon is a semiconductor and is used to make microcircuits and in the production of solar cells.
oxygen compounds. Getting, properties and application
Oxides of carbon
Carbon monoxide (II) (CO - carbon monoxide)
CO is a poisonous gas, colorless and odorless, poorly soluble in water.
Receipt
In the laboratory, CO is obtained by decomposition of formic or oxalic acid (in the presence of concentrated H 2 SO 4):
HCOOH \u003d CO + H 2 O,
H 2 C 2 O 4 \u003d CO + CO 2 + H 2 O
or by heating zinc dust with calcium carbonate:
CaCO 3 + Zn \u003d CaO + ZnO + CO.
In the factory, CO is produced by passing air or carbon dioxide through hot coal:
2C + O 2 \u003d 2CO,
Properties
The toxic effect of carbon monoxide is due to the fact that the affinity of hemoglobin for carbon monoxide is greater than for oxygen. This forms carboxyhemoglobin and thus blocks the transfer of oxygen in the body.
Carbon monoxide (II) is easily oxidized, burns in air with the release of a large amount of heat:
2 CO + O 2 \u003d 2 CO 2 + 577 kJ / mol.
CO reduces many metals from their oxides:
FeO + CO \u003d Fe + CO 2,
CuO + CO \u003d Cu + CO 2.
CO easily enters into addition reactions:
CO + Cl 2 \u003d COCl 2,
CO + NaOH = HCOONa,
Ni + 4 CO \u003d Ni (CO) 4.
In industry, it is often not pure CO that is 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 \u003d 2 CO + 222 kJ.
Water gas is obtained by passing water vapor through hot coal:
C + H 2 O \u003d CO + H 2 - 132 kJ.
The first reaction is exothermic, and the second is with the absorption of heat. If both processes are alternated, then it is possible to maintain the required temperature in the furnace. By combining generator and 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(IV) monoxide (CO 2 - carbon dioxide)
CO 2 is a colorless, odorless, non-flammable gas. It is released during the respiration of animals. Plants take in CO 2 and release oxygen. The air usually contains 0.03% carbon dioxide. Due to human activities (uncontrolled deforestation,
burning everything more coal, oil and gas), the content of CO 2 in the atmosphere is gradually increasing, which causes a greenhouse effect and threatens humanity with an ecological catastrophe.
Receipt
In the laboratory, CO 2 is obtained in the Kipp apparatus by acting on marble with hydrochloric acid:
CaCO 3 + 2HCl \u003d CaCl 2 + H 2 O + CO 2.
There are many reactions that produce CO2:
KHCO 3 + H 2 SO 4 \u003d KHSO 4 + H 2 O + CO 2,
C + O 2 \u003d CO 2,
2 CO + O 2 \u003d 2 CO 2,
Ca (HCO 3) 2 CaCO 3 Ї + CO 2 + H 2 O,
CaCO 3 \u003d CaO + CO 2,
BaSO 4 + 2 C \u003d 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 CO2 is dissolved in water, carbonic acid is formed:
H 2 O + CO 2 \u003d H 2 CO 3.
For CO 2, all those reactions that are characteristic of acid oxides are known:
Na 2 O + CO 2 \u003d Na 2 CO 3,
Ca (OH) 2 + 2 CO 2 \u003d Ca (HCO 3) 2,
Ca (OH) 2 + CO 2 \u003d CaCO 3 + H 2 O.
The ignited Mg continues to burn in carbon dioxide:
CO 2 + 2 Mg \u003d 2 MgO + C.
Carbonic acid is a weak dibasic acid:
H 2 O + CO 2 \u003d H 2 CO 3
H + + HCO 3 - \u003d H + + CO 3 2-
and can displace weaker acids from solutions of their salts:
Na 2 SiO 3 + CO 2 + H 2 O \u003d H 2 SiO 3 + Na 2 CO 3,
KCN + CO 2 + H 2 O \u003d KHCO 3 + HCN.
Salts of carbonic acid. Carbonates and bicarbonates
General methods for obtaining salts are also typical for obtaining salts of carbonic acid:
CaCO 3 + CO 2 + H 2 O \u003d Ca (HCO 3) 2,
Ca (HCO 3) 2 + Ca (OH) 2 \u003d 2 CaCO 3 + 2 H 2 O.
Alkali metal and ammonium carbonates are highly soluble in water and
subject to hydrolysis. All other carbonates are practically insoluble:
Na 2 CO 3 + H 2 O \u003d 2 Na + + OH - + HCO 3 -.
With relatively low heating, hydrocarbons decompose:
Ca (HCO 3) 2 \u003d CaCO 3 + CO 2 + H 2 O.
When carbonates are calcined, metal oxides and CO 2 are obtained:
CaCO 3 \u003d CaO + CO 2.
Carbonates are easily decomposed by stronger (than carbonic) acids:
MgCO 3 + 2HCl \u003d MgCl 2 + CO 2 + H 2 O.
CaCO 3 + 2HCl \u003d CaCl 2 + CO 2 + H 2 O.
When calcining carbonates with sand, SiO 2 displaces a more volatile oxide:
Na 2 CO 3 + SiO 2 \u003d 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 the 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 a raw material for the production of lime.
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.
Receipt
SiO 2 is an acid oxide, which can be obtained by the reactions:
Si + O 2 \u003d SiO 2,
H 2 SiO 3 \u003d SiO 2 + H 2 O,
SiH 4 + 2O 2 \u003d SiO 2 + 2H 2 O.
Properties
When interacting with metals or carbon, SiO 2 can be reduced to silicon
SiO 2 + 2 Mg \u003d Si + 2 MgO,
SiO 2 + 2 C \u003d Si + 2 CO
or give carborundum (SiC) SiO 2 + 3 C \u003d 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 \u003d K 2 SiO 3 + CO 2,
SiO 2 + CaO \u003d CaSiO 3.
Acids do not act on SiO 2 . An exception is hydrofluoric acid:
SiO 2 + 4HF \u003d SiF 4 + 2H 2 O,
SiF 4 + 2HF \u003d 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 \u003d H 2 SiO 3 + 2 NaCl.
When heated, silicic acid decomposes:
H 2 SiO 3 \u003d SiO 2 + H 2 O.
silicates
Many hundreds of silicate minerals are known. They make up 75% of the mass earth's crust. Among them there are a lot of aluminosilicates. Silicates are the main component of cement, glass, concrete and bricks.
Only Na and K silicates are soluble in water. Their aqueous solutions are called "liquid glass". During hydrolysis, these solutions have an alkaline reaction. They are used for the manufacture of acid-resistant cement and concrete.
Lesson type. Acquisition of new knowledge.
Lesson objectives.Educational. To acquaint students with a new classification of chemical reactions on the basis of changes in the oxidation states of elements - with redox reactions (ORD); teach students to arrange coefficients using the electronic balance method.
Developing. Continue development logical thinking, the ability to analyze and compare, the formation of interest in the subject.
Educational. To form a scientific worldview of students; improve work skills.
Methods and methodological techniques. Story, conversation, demonstration of visual aids, independent work students.
Equipment and reagents. Reproduction depicting the Colossus of Rhodes, the algorithm for placing coefficients according to the electronic balance method, a table of typical oxidizing and reducing agents, a crossword puzzle; Fe (nail), solutions of NaOH, СuSO 4 .
DURING THE CLASSES
Introduction
(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 god of the Sun). The grandiose idea and perfection of execution of the Colossus of Rhodes - one of the wonders of the world - amazed everyone who saw it.
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 Haret and took 12 years to build.
The bronze shell was attached to the iron frame. The hollow statue began to be built from the bottom and, as it grew, it was filled with stones to make it more stable. Approximately 50 years after the completion of construction, the Colossus collapsed. During the earthquake, he broke at the level of his knees.
Scientists believe that the true 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; learn about the concepts of "reducing agent" and "oxidizing agent", about the processes of reduction and oxidation; learn how to arrange the coefficients in the equations of redox reactions. Write in your workbooks the number, the topic of the lesson.
Learning new material
The teacher makes 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 of the starting materials and reaction products.
The student writes 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, but in the second they changed - in 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 up 2 electrons.
Copper has an oxidation state of +2 before the reaction, and 0 after the reaction. As you can see, the oxidation state has decreased. Therefore, copper accepts 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 adding electrons is called reduction.
We write the schemes of these processes:
So, give the definition of the concepts of "reducing agent" and "oxidizing agent".
Student. Atoms, molecules or ions that donate electrons are called reducing agents.
Atoms, molecules, or ions that accept electrons are called oxidizing agents.
Teacher. What is the definition of reduction and oxidation processes?
Student. Recovery is the process of adding electrons to an atom, molecule or ion.
Oxidation is the process by which electrons are transferred by an atom, molecule, or ion.
Students write the definitions in a notebook under dictation and complete the 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 attached by the oxidizing agent.
To select the coefficients in the equations of redox reactions, two methods are used - electron balance and electron-ion balance (half-reaction method).
We will consider only the electronic balance method. To do this, we use the algorithm for arranging the coefficients using the electronic balance method (drawn up on a piece of drawing paper).
EXAMPLE Arrange the coefficients in this reaction scheme using the electron balance method, determine the oxidizing agent and reducing agent, indicate the processes of oxidation and reduction:
Fe 2 O 3 + CO Fe + CO 2.
We will use the algorithm for placing the coefficients using the electronic balance method.
3. Let's write out the elements that change the degree of oxidation:
4. Compose electronic equations, determining the number of given and received electrons:
5. The number of given and received electrons must be the same, because neither the reactants nor the products of the reaction are charged. We equalize the number of given and received electrons by choosing the least common multiple (LCM) and additional factors:
6. The resulting multipliers are coefficients. We transfer the coefficients to the reaction scheme:
Fe 2 O 3 + 3CO \u003d 2Fe + 3CO 2.
Substances that are oxidizing or reducing agents in many reactions are called typical.
A table made on a Whatman sheet is posted.
Teacher. Redox reactions are very common. They are associated not only with corrosion processes, but also with fermentation, decay, photosynthesis, and metabolic processes occurring in a living organism. They can be observed during the combustion of fuel. Redox processes accompany the cycles of substances in nature.
Did you know that about 2 million tons of nitric acid are formed in the atmosphere every day, or
700 million tons per year, and in the form of a weak solution fall to the ground with rain (man produces only 30 million tons of nitric acid per year).
What happens in the atmosphere?
Air contains 78% nitrogen by volume, 21% oxygen and 1% other gases. Under the action of lightning discharges, and an average of 100 lightning flashes on Earth every second, nitrogen molecules interact with oxygen molecules to form nitric oxide (II):
Nitric oxide (II) is easily oxidized by atmospheric oxygen to nitric oxide (IV):
NO + O 2 NO 2 .
The resulting 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 above reaction schemes using the electronic balance method, indicate the oxidizing agent, reducing agent, oxidation and reduction processes.
Solution
1. Let's determine the oxidation states of the elements:
2. We underline the symbols of elements whose oxidation states change:
3. Let's write out the elements that have changed their oxidation states:
4. Compose electronic equations (determine the number of given and received electrons):
5. The number of given and received electrons is the same.
6. Let's transfer the coefficients from electronic circuits to the reaction scheme:
Next, students are invited to independently arrange the coefficients using the electronic balance method, determine the oxidizing agent, reducing agent, indicate the processes of oxidation and reduction in other processes occurring in nature.
The other two reaction equations (with coefficients) are:
Checking the correctness of the tasks is carried out using a codoscope.
Final part
The teacher asks 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 substances KMnO 4, K 2 Cr 2 O 7, O 3 are strong ... (vertically (2)).
Horizontally:
1. What process does the scheme 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 scheme reflect:
6. For the selection of coefficients in the equations of redox reactions, the method of electronic ... is used.
7. According to the diagram, aluminum gave ... an electron.
8. In reaction:
H 2 + Cl 2 \u003d 2HCl
hydrogen H 2 - ....
9. What type of reactions are always only redox reactions?
10. The oxidation state of simple substances is ....
11. In reaction:
reducer...
Homework assignment. According to O.S. Gabrielyan's textbook "Chemistry-8" § 43, p. 178–179, ex. 1, 7 in writing.
A task (at home). Constructors of the first spaceships and submarines faced a problem: how to maintain a constant air composition on the ship and space stations? Get rid of excess carbon dioxide and replenish oxygen? The solution has been found.
Potassium superoxide KO 2 forms oxygen as a result of interaction with carbon dioxide:
As you can see, this is a redox reaction. Oxygen is both an oxidizing agent and a reducing agent in this reaction.
In a space expedition, every gram of cargo counts. Calculate the supply of potassium superoxide to be taken in 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 KO2. )
Exercise ( elevated level difficulties). Write down the equations for the redox reactions that could have led to the destruction of the Colossus of Rhodes. Keep in mind that this giant statue stood in a port city on an island in the Aegean 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. Moscow: Bustard, 2002;
Gabrielyan O.S., Voskoboynikova N.P., Yashukova A.V. Handbook of the teacher. 8th grade. Moscow: 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.
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 process of oxidation and reduction:
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. REDOX REACTIONSPart A
A2 When studying the electrical conductivity of various substances using a special device, students observed the following:
Which of the following substances was in the glass?
1) sugar (solution)
2) KC1 (tv.) 3) NaOH (p-p) 4) alcohol
A4 The interaction of solutions of barium chloride and sulfuric acid corresponds to a reduced ionic equation
1) H+ + SG=HC1
2) Ba2+ + SO42- \u003d BaSO4
3) CO32- + 2H+ = H2O + CO2
4) Ba2+ + CO32- = BaCO3
A5 Reaction between solutions of silver nitrate and of hydrochloric acid runs to the end, because
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(II) hydroxide
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 a substance and the total number of 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 in the table the selected numbers under the corresponding letters.
Transfer the answer in the form of a sequence of four digits to the test form under the number of the corresponding task, without changing the order of the numbers.
Here is a list of related concepts:
A) acid
B) hydrochloric acid
B) anoxic acid
D) strong electrolyte
Write down the letters that denote the concepts in the table in such a way that a chain can be 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.
REDOX REACTIONS
Reactions in which there is a change in the oxidation states of the atoms of the elements that make up the reacting compounds, called redox.
Oxidation state(c.d.) is the charge of the element in the compound, calculated assuming that the compound is composed of ions. The determination of the degree of oxidation 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 oxidation state of oxygen 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 the oxidation states of elements in a neutral molecule is zero, in a complex ion - the charge of the ion.
As an example, we calculate the degree of oxidation of chromium in the compound K 2 Cr 2 O 7 and nitrogen in the anion (NO 2) -
K 2 +1 Cr 2 X O 7 –2 2∙(+1)+ 2 x + 7 (–2) = 0 x = + 6
(NO 2) - x + 2 (–2) = –1 x = + 3
In redox reactions, electrons are transferred from one atom, molecule, or ion to another. Oxidation – the process of donating electrons by an atom, molecule, or ion, accompanied by an increase in the degree of oxidation. Recovery – the process of adding electrons, accompanied by a decrease in the degree of oxidation.
-4 -3 -2 -1 0 +1 +2 +3 +4 +5 +6 +7 +8
Recovery process
Oxidation and reduction are interrelated processes occurring simultaneously.
Oxidizers called substances (atoms, ions or molecules) that gain electrons during a reaction, reducing agents – substances that donate electrons. 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.), non-metal 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 the free state, negatively charged non-metal ions (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, either 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, NH3). 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 +3 oxidation state, hydrogen peroxide H 2 O 2, containing oxygen in the -1 oxidation state, exhibit reducing properties in the presence of strong oxidizing agents, and when interacting with active reducing agents, they are oxidizing agents.
When compiling the equations of redox reactions, it is recommended to adhere to the following order:
1. Write the formulas of the starting substances. Determine the oxidation state of elements that can change it, find an oxidizing agent and a reducing agent. Write the reaction products.
2. Make equations for the processes of oxidation and reduction. Select multipliers (basic coefficients) so that the number of electrons given away during oxidation is equal to the number of electrons received during reduction.
3. Arrange 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 medium
oxidation S -2 - 2ē → S 0 ½3
reduction 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, dilute water is most often used. 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 the oxidizing agent and the reducing agent are in different substances.. The above reaction is of this type. TO intramolecular include reactions, in which the oxidizing agent and 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 - oxidizer
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 \u003d 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 - oxidizer
