REDOX REACTIONS
Reactions in which a change in the oxidation states of the atoms of the elements that make up the reacting compounds occurs are called redox.
Oxidation state(c.o.) is the charge of an element in a compound, calculated based on the assumption that the compound consists of ions. The 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 oxidation state of oxygen in compounds is usually –2. Exceptions are the 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 have a constant oxidation state (+1); beryllium Be and magnesium Mg (+2); alkaline earth metals Ca, Sr, Ba (+2); fluorine (–1).
5. The algebraic sum of the oxidation states of elements in a neutral molecule is equal to zero, in a complex ion - the charge of the ion.
As an example, let's calculate the oxidation state 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 losing electrons by an atom, molecule or ion, accompanied by an increase in the oxidation state. Recovery – the process of adding electrons, 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 that occur simultaneously.
Oxidizing agents are called substances (atoms, ions or molecules) that gain electrons during a reaction, restorers – 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 a 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, NH 3). If a 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, exhibit reducing properties in the presence of strong oxidizing agents, and when interacting with active reducing agents they are oxidizing agents.
When composing equations for 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 the oxidizing agent and the reducing agent. Write the reaction products.
2. Draw up equations for the processes of oxidation and reduction. Select the multipliers (main coefficients) so that the number of electrons given up during oxidation is equal to the number of electrons accepted 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 sulfuric acid is most often used, and to create an alkaline environment, solutions of sodium or potassium hydroxides are used.
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 reaction discussed above belongs to this type. TO intramolecular reactions include 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 Cl +5 + 6ē → Cl - ½2 Cl +5 - oxidizing agent
oxidation 2O -2 - 4ē → O 2 0 ½3 O -2 - reducing agent
In reactions disproportionation(auto-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
Lesson type. Acquiring new knowledge.
Lesson objectives.Educational. Introduce students to a new classification of chemical reactions based on changes in the oxidation states of elements - oxidation-reduction reactions (ORR); teach students to arrange coefficients using the electronic balance method.
Developmental. Continue the development of logical thinking, ability to analyze and compare, and develop interest in the subject.
Educational. To form the scientific worldview of students; improve work skills.
Methods and methodological techniques. Story, conversation, demonstration of visual aids, independent work of students.
Equipment and reagents. Reproduction with the image of the Colossus of Rhodes, algorithm for arranging coefficients using the electronic balance method, table of typical oxidizing and reducing agents, crossword puzzle; Fe (nail), NaOH, CuSO 4 solutions.
DURING THE CLASSES
Introductory part
(motivation and goal setting)
Teacher. In the 3rd century. BC. On the island of Rhodes, a monument was built in the form of a huge statue of Helios (the Greek god of the Sun). The grandiose design and perfect 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 we know 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 an 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. About 50 years after its completion, the Colossus collapsed. During the earthquake it broke at the level of the knees.
Scientists believe that the real reason for the fragility of this miracle was metal corrosion. And the corrosion process is based on redox reactions.
Today in the lesson you will learn about redox reactions; learn about the concepts of “reducing agent” and “oxidizing agent”, about the processes of reduction and oxidation; learn to place coefficients in equations of redox reactions. Write down the date and 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 for the reactions performed. In each equation, arrange the oxidation states of the elements in the formulas of the starting substances and reaction products.
The student writes reaction equations on the board and assigns 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 - for copper and iron.
Teacher. The second reaction is a redox reaction. Try to define redox reactions.
Student. Reactions that result in changes in the oxidation states of the elements that make up the reactants and reaction products are called redox reactions.
Students write down in their notebooks, under the teacher’s dictation, 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 we 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 we can see, the oxidation state has decreased. Therefore, copper accepts 2 electrons.
Iron donates electrons, it is a reducing agent, and the process of transferring electrons is called oxidation.
Copper accepts electrons, it is an oxidizing agent, and the process of adding electrons is called reduction.
Let us write down the diagrams of these processes:
So, give a definition of the concepts “reducing agent” and “oxidizing agent”.
Student. Atoms, molecules or ions that donate electrons are called reducing agents.
Atoms, molecules or ions that gain electrons are called oxidizing agents.
Teacher. How can we define the processes of reduction and oxidation?
Student. Reduction is the process by which an atom, molecule, or ion gains electrons.
Oxidation is the process of transfer of electrons by an atom, molecule or ion.
Students write down definitions from dictation in a notebook and draw.
Remember!
Donate electrons and oxidize.
Take electrons - recover.
Teacher. Oxidation is always accompanied by reduction, and vice versa, reduction is always associated with oxidation. The number of electrons given up by the reducing agent is equal to the number of electrons gained by the oxidizing agent.
To select coefficients in the equations of redox reactions, two methods are used - electronic balance and electron-ion balance (half-reaction method).
We will consider only the electronic balance method. To do this, we use an algorithm for arranging coefficients using the electronic balance method (designed on a piece 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 processes of oxidation and reduction:
Fe 2 O 3 + CO Fe + CO 2.
We will use the algorithm for arranging coefficients using the electronic balance method.
3. Let’s write down the elements that change oxidation states:
4. Let’s create 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 starting materials nor the reaction products are charged. We equalize the number of given and received electrons by selecting the least common multiple (LCM) and additional factors:
6. The resulting multipliers 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 made on a piece of Whatman paper is hung.
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 fuel combustion. Redox processes accompany the cycles of substances in nature.
Did you know that approximately 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 (humans produce 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 Earth there are an average of 100 lightning flashes 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 nitrogen oxide (IV) reacts 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.
Exercise . Arrange the coefficients in the given reaction schemes using the electronic balance method, indicate the oxidizing agent, the reducing agent, the processes of oxidation and reduction.
Solution
1. Let's determine the oxidation states of elements:
2. Let us emphasize the symbols of elements whose oxidation states change:
3. Let’s write down the elements that have changed their oxidation states:
4. Let’s create electronic equations (determine the number of given and received electrons):
5. The number of electrons given and received is the same.
6. Let’s transfer the coefficients from the electronic circuits to the reaction diagram:
Next, students are asked to independently arrange the coefficients using the electronic balance method, determine the oxidizing agent, the reducing agent, and indicate the processes of oxidation and reduction in other processes occurring in nature.
The other two reaction equations (with coefficients) have the form:
The correctness of tasks is checked using an overhead projector.
Final part
The teacher asks students to solve a crossword puzzle based on the material they have studied. The result of the work is submitted for verification.
Having solved crossword, you will learn that the substances KMnO 4, K 2 Cr 2 O 7, O 3 are strong ... (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. To select coefficients in the equations of redox reactions, use the electronic... method.
7. According to the diagram, aluminum gave up ... an electron.
8. In reaction:
H 2 + Cl 2 = 2HCl
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 –….
Homework assignment. According to the textbook by O.S. Gabrielyan “Chemistry-8” § 43, p. 178–179, ex. 1, 7 in writing.
Task (for home). The designers of the first spaceships and submarines were faced with a problem: how to maintain a constant air composition on the ship and space stations? Get rid of excess carbon dioxide and replenish oxygen? A solution has been 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.
On a space mission, every gram of cargo counts. Calculate the supply of potassium superoxide that must be taken on a space flight if the flight lasts 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 giant statue stood in a port city on an island in the Aegean Sea, off the coast of modern-day Turkey, where the humid Mediterranean air is laden 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., Voskoboynikova N.P., Yashukova A.V. Teacher's handbook. 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.
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 redox reaction schemes 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 me 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) KS1 (solid) 3) NaOH (p-p) 4) alcohol
A4 The interaction of solutions of barium chloride and sulfuric acid corresponds to the abbreviated ionic equation
1)H+ + SG=HC1
2)Ba2+ + SO42- =BaSO4
3) CO32- + 2H+ = H2O + CO2
4) Ba2+ + CO3- = BaCO3
A5 The reaction between solutions of silver nitrate and hydrochloric acid proceeds to completion, 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(II) hydroxide
3) phosphoric acid and calcium oxide
4) silicic acid and sodium hydroxide
A10 The oxidation process corresponds to the diagram
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 mole 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 testing form under the number of the corresponding task, without changing the order of the numbers.
You are offered a list of interrelated concepts:
A) acid
B) hydrochloric acid
B) oxygen-free acid
D) strong electrolyte
Write down the letters that represent the concepts in a table so that a chain can be traced from a particular concept to the most general one.
Transfer the resulting sequence of letters to the testing form without changing the order of the letters.
Oxidation is the process of losing electrons, with an increase in the degree of oxidation.
At oxidation substances resulting from recoil electrons it increases oxidation state. Atoms the substance being oxidized is called donors electrons, and atoms oxidizing agent - acceptors electrons.
In some cases, during oxidation, the molecule of the original substance may become unstable and break down into more stable and smaller components (see Fig. Free radicals). In this case, some of the atoms of the resulting molecules have a higher oxidation state than the same atoms in the original molecule.
The oxidizing agent, accepting electrons, acquires reducing properties, turning into a conjugate reducing agent:
oxidizer+ e − ↔ conjugate reducing agent.
Recovery
Restoration is the process of adding electrons to an atom of a substance, while its oxidation state decreases.
Upon recovery atoms or ions attach electrons. At the same time there is a decrease oxidation states element. Examples: recovery oxides metals to free metals using 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 conjugate oxidizing agent:
reducing agent - e − ↔ conjugate oxidizer.
An unbound, free electron is the strongest reducing agent.
Oxidation-reduction reactions are reactions in which reactants gain or donate electrons. An oxidizing agent is a particle (ion, molecule, element) that adds electrons and moves from a higher oxidation state to a lower one, i.e. is being restored. A reducing agent is a particle that donates electrons and moves from a lower oxidation state to a higher one, i.e. oxidizes.
Intermolecular - reactions in which oxidizing and reducing atoms are located in molecules of different substances, for example:
N 2 S + Cl 2 → S + 2HCl
Intramolecular - reactions in which oxidizing and reducing atoms are located in molecules of the same substance, for example:
2H 2 O → 2H 2 + O 2
Disproportionation (auto-oxidation-self-reduction) - 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:
N.H. 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 losing electrons is oxidation. During oxidation, the oxidation state increases:
The process of adding electrons is reduction. During reduction, the oxidation state decreases:
Atoms or ions that gain electrons in a given reaction are oxidizing agents, and those that donate electrons are reducing agents.
Redox reactions (electrode potential)
Electrons can act as chemical reagents, 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 occur. As a result of the first process, the plate acquires a negative charge; after some time after immersion in the solution, the velocities are equalized and equilibrium occurs. And the plate acquires some electrical potential.
The electrode potential is measured relative to the potential of standard hydrogen.
Copper-hydrogen electrode- 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 gas 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 [ specify ] on the concentration of H + ions in solution. The electrode is a standard against which the electrode potential of the chemical reaction being determined 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 potential of the SE is taken equal to 0 V. When assembling a galvanic cell from the SE and the electrode being determined, the following reaction occurs reversibly on the surface of platinum:
2Н + + 2e − = H 2
that is, either happens recovery hydrogen or its oxidation- it depends on the potential of the reaction occurring at the electrode being determined. By measuring the emf of a galvanic electrode under standard conditions (see above), one determines standard electrode potential determined chemical reaction.
HE is used to measure the standard electrode potential of an electrochemical reaction, to measure concentrations(activity) of hydrogen ions, as well as any other ions. VE is also used to determine the solubility product and to determine the rate constants of some electrochemical reactions.
Nernst equation
The dependence of the redox potential corresponding to the half-reaction of reduction of the permanganate ion in an acidic medium (and, as already noted, at the same time the half-reaction of 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 a given particle in the half-reaction equation, n– number of electrons accepted by the oxidizer, R– universal gas constant, T- temperature, F– Faraday number.
Measure the redox potential in the reaction vessel during the reaction, 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 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, measuring 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 redox pairs during a reaction cannot be measured, they can be calculated by substituting the current ones into the Nernst equation, i.e. concentrations corresponding to a given point in time. If the change in potential as the reaction proceeds is considered, 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 proceeds, the potential of the oxidizing agent calculated using the Nernst equation decreases, and the potential of the reducing agent corresponding to the second half-reaction, on the contrary, increases. When these potentials are equalized, the reaction stops and the system returns to a state of chemical equilibrium.
25. Complex compounds are compounds that exist both in the crystalline state and in solution, the peculiarity 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. According Werner's coordination theory distinguishes between an internal and an external sphere 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 the complex compound constitutes the outer sphere and is written outside square brackets. A certain number of ligands are placed 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 site near the central atom. Coordination changes the properties of both the ligands and the central atom. Often coordinated ligands cannot be detected using chemical reactions characteristic of them in the free state. More tightly bound particles of the inner sphere are called a complex (complex ion). There are attractive forces between the central atom and the ligands (a covalent bond is formed by an exchange and (or) donor-acceptor mechanism), and repulsive forces between the ligands. If the charge of the inner sphere is 0, then there is no outer coordination sphere. 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 orbitals and a sufficiently large positive nuclear charge, and therefore can be electron acceptors. These are cations of transition elements. The most powerful complexing agents are elements of groups IB and VIIIB. Rarely, neutral atoms of d-elements and atoms of non-metals in varying 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 is usually 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 complex-forming properties and act as ligands in the complex compound. Ligands can be atoms and molecules (protein, amino acids, nucleic acids, carbohydrates). Based on the number of bonds formed by the ligands with the 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 3rd order of the inner sphere of a complex compound is the algebraic sum of the charges of the particles that form it. Complex compounds that have 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 occurs instantly and almost completely. SO4 = 2+ + SO42-. If there are hydroxide ions in the outer sphere of a complex compound, then this compound is a strong base.
Group IA includes lithium, sodium, potassium, rubidium, cesium and francium. These elements are called alkali elements. Sometimes hydrogen is also included in group IA. Thus, this group includes elements from each of the 7 periods. The general valence electronic formula of group IA elements is ns1. At the outer level there is 1 electron. Very far from the nucleus. Low ionization potentials. Atoms donate 1 electron. Metallic means are pronounced. Metallic properties increase with increasing atomic number. Physical properties: Metals are soft, light, fusible with good electrical conductivity, and have a large negative electrical potential. Chemical properties: 1) Stored under a layer of liquid hydrocarbons (benzene, gasoline, kerasin) 2) Oxidizing agents. 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. Preparation of peroxides. 2. An alloy of sodium and potassium - a coolant in nuclear power plants. 3. Preparation of organometallic compounds.
27. General comparative characteristics of elements and their compounds of groups I A and I B 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 called alkalis are formed. In the Periodic Table they immediately follow the noble gases, so the peculiarity of the structure of alkali metal atoms is that they contain one electron in the outer energy level: their electron configuration is ns1. Obviously, the valence electrons of alkali metals can be easily removed because it is energetically favorable for the atom to give up an electron and acquire the inert gas configuration. 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 are silvery-white in color (except silvery-yellow cesium), they are very soft, and 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 in nature 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 soil contains potassium salts - sylvite KCl, sylvinite NaCl KCl, carnallite KCl MgCl2 6H2O, polyhalite K2SO4 MgSO4 CaSO4 2H2O. The copper subgroup is the chemical elements of group 11 of the periodic table of chemical elements (according to the outdated classification, elements of the secondary subgroup of group I). The group includes the 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 seconds). The name coin metals is not officially applied to group 11 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. They are also characterized by high density values, melting and boiling points, and high thermal and electrical conductivity. A feature of the elements of the subgroup is the presence of a filled pre-external -sublevel, achieved due to 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, which is why 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 consists mainly of 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 droplets of water form on the walls of the vessel: Hydrogen without impurities burns calmly. However, a mixture of hydrogen with oxygen or air explodes. The most explosive mixture is one consisting of two volumes of hydrogen and one volume of oxygen—detonating gas. If an explosion occurs in a glass vessel, its fragments may damage
hurt others. 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-pang” sound. If hydrogen contains an admixture of air, it burns explosively. When working with hydrogen, safety regulations must be followed. 2) If, for example, when heating, a stream of hydrogen is passed over copper (II) oxide, a reaction occurs as a result of which water and metallic copper are formed: In this reaction, a reduction process occurs, since hydrogen removes oxygen from the copper atoms. The reduction process is the opposite of the oxidation process. Substances that take away oxygen are classified 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 explosively with slight activation. The interaction with Br2 and I2 occurs more slowly. For a reaction with hydrogen to occur, it is sufficient to activate a small fraction of the reagents using light or heat. Activated particles interact with non-activated ones, forming HX and new activated particles, which continue the process, and the reaction of two activated particles in the main reaction ends with the formation of a product. 4) Oxidation reactions. When heating hydrogen with metals of the 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. A colorless liquid with a “metallic” taste, infinitely 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 into nitrates, releases iodine from metal iodides, and breaks down unsaturated compounds at the site of double bonds. Hydrogen peroxide reduces gold and silver salts, 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 When exposed to strong oxidizing agents, H2O2 exhibits reducing properties, releasing free oxygen: O22− - 2e− → O2 The reaction of KMnO4 with H2O2 is used in chemical analysis to determine the H2O2 content: 5H2O2 + 2KMnO4 + 3H2SO4 → 5O2 + 2MnSO4 + K2SO4 + 8H2O It is advisable to carry out the oxidation of organic compounds with hydrogen peroxide (for example, sulfides and thiols) in acetic acid.
29. general characteristics of the properties of elements and their compounds of group 2. 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 with opposite spins; with the expenditure of the necessary energy, one element goes from the s state to 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 Method of preparation: electrolysis of molten chlorides Physical properties: light metals, but harder than alkali metals. Chemical properties: 1 In air, the surface of Be and Mg is covered with an oxide film. 2.at high temperatures it interacts with nitrogen 3.does not interact with water Be 4.displaces hydrogen from acids (except nitric). Application: The main use of calcium metal is its use as a reducing agent in the production of metals, especially nickel, copper and stainless steel become. Calcium and its hydride are also used to produce difficult-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 devices
No. 31 Alkaline earth metals - chemical elements 2nd group of the main subgroup, except beryllium and magnesium: calcium, strontium, barium And radium. Belong to the 2nd group of elements according to the new classification IUPAC. So named because they oxides- “earth” (according to the terminology alchemists) - report water alkaline reaction. Salts alkaline earth metals, except 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 ones, 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 produced by the combustion of 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. Acidic oxides react with both acids and alkalis: ZnO + 2HCl = ZnCl2 + H2O
ZnO + 2NaOH = Na2ZnO2 + H2O
Hydroxides ( hydroxides) - compounds of oxides of chemical elements. 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, a distinction is made accordingly:
basic hydroxides (grounds) - hydroxides exhibiting 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 acidHNO3, sulfuric acidH2SO4, sulfurous acidH2SO3, 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 esters carbonic acid (H 2 CO 3). Among the salts, normal carbonates (with the anion CO 3 2−) 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 acid (almost all known acids, including organic ones) to release carbon dioxide:
Application: Calcium, magnesium, barium carbonates, etc. are used in construction, the chemical industry, optics, etc. They are 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 represent a broad 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 igneous (pyroxenes, feldspars ), they are also typical for pegmatites (mica, tourmaline, beryl, etc.) and skarnov (garnets, wollastonite). Widely distributed 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, actinium and actinides (side subgroup).
At the outer electronic level of the elements of the main subgroup there are three electrons (s 2 p 1). They easily give up these electrons or form three unpaired electrons due to the transition of one electron to the p-level. Boron and aluminum are characterized by compounds 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 metallic properties increase.
Elements of the scandium subgroup also have three electrons in the outer electronic level. However, these elements belong to the 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 outer electronic level: their 4f level is built up and the d level disappears. Starting with cerium, all elements except gadolinium and lutetium have the electronic configuration of the outer electron 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 actinide valence layer is in many ways similar to the electronic structure of the lanthanide valence layer. 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.
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 plastic, easily drawn into wire and rolled into sheets: foil with a thickness of less than 0.01 mm can be made from it. Aluminum has 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 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, the 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.
Aluminum, cleared of the protective oxide film, reacts with water to release hydrogen:
2 Al + 6 H 2 O = 2 Al(OH) 3 + 3 H 2
Aluminum dissolves well in dilute sulfuric and hydrochloric acids:
2 Al + 6 HCl = 2 AlCl 3 + 3 H 2
2 Al + 3 H 2 SO 4 = Al 2 (SO 4) 3 +3 H 2
Dilute nitric acid passivates aluminum in the cold, but when heated, aluminum dissolves in it, releasing 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 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 intensely 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 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 = 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. 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 you use a special fuse to 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 heating to 3500 ° C. Iron at this temperature is in a molten state.
Receipt. Aluminum was first obtained by reduction from aluminum chloride with sodium metal:
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 a temperature of about 1000 °C.
Application. Aluminum is used very widely. It is used to make foil used in radio engineering and for packaging food products. Steel and cast iron products are coated with aluminum 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 can withstand heating up to 1000 °C without corroding. Aluminum alloys, distinguished by their great lightness and strength, are used in the production of heat exchangers, 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 transparent corundum crystals of beautiful shape and color are found. Corundum colored red by chromium compounds is called ruby, and blue colored by titanium and iron compounds is called sapphire. Ruby and sapphire are precious stones. Currently, they are quite easily obtained artificially.
Bor-element main subgroup of the third group, 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- a 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 the boron was obtained.
Receipt
The purest boron is obtained by pyrolysis of borohydrides. This boron is used for the production of semiconductor materials and fine chemical syntheses.
1. Metallothermy method (usually reduction with magnesium or sodium):
2. Thermal decomposition of boron bromide vapor on hot (1000-1200°C) tantalum wire in the presence of hydrogen:
Physical properties
An 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 nonmetal boron resembles silicon.
The chemical boron is quite inert and at room temperature only reacts 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 with a large release of heat, forming the oxide B 2 O 3:
Boron does not interact directly 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:
When heated strongly, boron exhibits restorative properties. It is capable, for example, of restoring 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 alkali solutions. In hot nitric and sulfuric acids and in 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 BO 3 3- anion), but tetraborates, for example:
Application
Elemental boron
Boron (in the form of fibers) serves as a strengthening agent for many composite materials.
Boron is also often used in electronics to change the type of conductivity silicon.
Boron is used in metallurgy as a micro-alloying 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 periodic system and 3 metals - germanium (intermediate between non-metals and metals, tin and lead, located at the end of large periods - IV, V, VI What is characteristic of all these elements is that they have 4 electrons in the outer energy level. And therefore they can exhibit an oxidation state from +4 to -4. These elements form gaseous compounds with hydrogen: CH4, SiH4, SnH4, PbH4. at when heated in air, they combine with elements of the oxygen subgroup, sulfur and with halogens.The oxidation state +4 is obtained when the 1s electron passes to a free p-orbital.
As the radius of the atom increases, the strength of the bond between the outer electrons and the nucleus decreases. Non-metallic properties decrease, and metallic properties increase. (melting and boiling points decrease, etc.)
Carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb) are 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 a subgroup, as the atomic number of an element increases, 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 Table.
There are 4 electrons in 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 a group, non-metallic properties weaken and metallic ones increase, so 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 nonmetals (N, O, S) and halogens they exhibit oxidation states +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 is: 12 C - 98.89% and 13 C - 1.11%. If we limit ourselves to listing isotopes whose abundance 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 allotropes
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 increasing charge of the nucleus of the E atom, the stability of hydrides decreases.
When moving from C to Pb, the stability of compounds with an oxidation state of +4
decreases, and with +2 it increases. The acidic character of EO 2 oxides decreases, and the basic character increases of EO oxides.
Carbon
Carbon occurs naturally in the form of diamond and graphite. Fossil coals contain it: from 92% in anthracite, 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. Carbon in the air is contained 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 nonmetals, 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: each carbon atom is connected to metal atoms by all four valence bonds. 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.
Coal is used as fuel and 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 and abrasives; cut diamonds are precious stones.
Silicon
Silicon occurs in nature only in bound form in the form of silica SiO2 and various silicic acid salts (silicates). It is the second (after oxygen) most abundant chemical element 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 also turns out when silane decomposes:
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 hydrofluoric acid).
In reactions with metals, silicon exhibits oxidizing properties:
2 Mg + Si = Mg 2 Si.
When magnesium silicide is decomposed with hydrochloric acid, silane is obtained:
Mg 2 Si + 4 HCl = 2MgCl 2 + SiH 4.
Silicon is used to produce many alloys based on iron and copper
and aluminum. Adding 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 microchips and in the production of solar cells.
Oxygen compounds. Preparation, properties and applications
Carbon oxides
Carbon (II) monoxide (CO - carbon monoxide)
CO is a poisonous gas, colorless and odorless, poorly soluble in water.
Receipt
In the laboratory, CO is obtained by the 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 a factory setting, CO is produced by passing air or carbon dioxide through hot coal:
2C + O 2 = 2CO,
Properties
The poisonous effect of 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 blocks the transfer of oxygen in the body.
Carbon (II) oxide easily oxidizes and burns in air, releasing 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 readily undergoes addition reactions:
CO + Cl 2 = COCl 2,
CO + NaOH = HCOONa,
Ni + 4 CO = Ni(CO) 4 .
In industry, it is often not pure CO that is used, but various mixtures of it with other gases. Producer gas is produced by passing air through hot coal in a shaft furnace:
2 C + O 2 = 2 CO + 222 kJ.
Water gas is produced by passing water vapor through hot coal:
C + H 2 O = CO + H 2 - 132 kJ.
The first reaction is exothermic, and the second occurs with the absorption of heat. If both processes are alternated, it is possible to maintain the required temperature in the oven. When generator and water gas are combined, a mixed gas is obtained. These gases are used not only as fuel, but also for the synthesis, for example, of methanol:
CO + 2H 2 = CH 3 OH.
Carbon monoxide (IV) (CO 2 - carbon dioxide)
CO 2 is a colorless, non-flammable, odorless gas. It is released when animals breathe. Plants absorb CO 2 and release oxygen. The air usually contains 0.03% carbon dioxide. Due to human activities (uncontrolled deforestation,
burning more and more coal, oil and gas), the CO 2 content in the atmosphere gradually increases, which causes the greenhouse effect and threatens humanity with an environmental disaster.
Receipt
In the laboratory, CO 2 is obtained in a Kipp apparatus by treating marble with hydrochloric acid:
CaCO 3 + 2HCl = CaCl 2 + H 2 O + CO 2.
There are many reactions 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 that are characteristic of acid oxides are known:
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 dibasic acid:
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.
Salts of carbonic acid. Carbonates and bicarbonates
General methods for obtaining salts are also typical for obtaining 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 highly soluble in water and
subject to hydrolysis. All other carbonates are practically insoluble:
Na 2 CO 3 + H 2 O = 2 Na + + OH - + HCO 3 - .
With relatively low heating, hydrocarbonates 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 easily 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, SiO 2 displaces the 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 crystal hydrate Na 2 CO 3 10H 2 O
(crystalline soda) are used in the glass, soap, and pulp and paper industries. Sodium bicarbonate NaHCO 3 (baking soda)
used in the food industry and medicine. Limestone is a building stone and 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 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. The 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. Its general formula is 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 of the earth's crust. Among them there are a lot of aluminosilicates. Silicates are the main constituent of cement, glass, concrete and brick.
Only Na and K silicates are soluble in water. Their aqueous solutions are called “liquid glass”. Upon hydrolysis, these solutions have an alkaline reaction. They are used for the production of acid-resistant cement and concrete.
