Question 1
Basic concepts and laws of chemistry: an atom is the smallest particle of a chemical element, neutral in charge and the carrier of its properties.
A molecule is the smallest particle of a substance, neutral in charge and the carrier of its properties.
An equivalent is the amount of a substance that interacts with 1 mole of an H atom in exchange reactions or with 1e in redox processes.
Boyle - Mariotte, Gay - Lussac, Avagadro
Richter's law of equivalents - the masses of substances bound by one interaction are directly proportional to the masses of their equivalents.
Quantum mechanical model of the structure of the atom: Bohr-Rutherford model: the center of the atom is the nucleus, which consists of protons and neutrons z - the charge of the atomic nucleus, which determines the atom’s type of chemical element, the serial number of the element in the periodic system, determines the number e of a neutral atom.
N – determines the isotopic composition of the atom.
The dimensions of an atom are determined by the dimensions of its electron shell.
The shell includes
Quantum numbers and types of electronic orbitals: using quantum numbers, you can describe the characteristics of the electron shell, n is the main quantum number, which determines: the number of the quantum layer or level, the capacity of the quantum layer and its energy, the number of sublevels within the level.
Sublevels are described by an orbital quantum number.
The Pauli principle: in an atom there cannot be 2 e having the same set of 4 quantum numbers.
The superscript numbers show combinations of magnetic and spin numbers. 1. The largest value of n determines the period number, outer layer 2. The sum of e on the outer layer determines group 3. The s and p sublevels form the main subgroups. The sublevel being populated defines the subgroup.
Hund's rule regulates permitted models.
Vacant orbitals at the sublevel are initially filled with single-electron clouds with the same spin orientation. Klechkovsky's rule: e sublevels are populated in the direction of increasing the sum of the main and orbital numbers.
For the same values of the sum n and l, the first to be populated is p.sl
Covalent bond: CS base, 2nd cloud for 2 atoms.
1. Each particle or atom provides a one-electron cloud for communication, provided that the clouds of 2 atoms are antiparallel.
2.Implemented due to the 2nd cloud of 1 particle and the vacant orbital of the second particle.
Characteristics: 1.Communication energy. 2. Bond length. 3. Saturability or maximum covalency. 4. Direction of communication. 5.Polarity of connection gender, non-gender.
6. Communication frequency.
Properties of K compounds: hard, brittle, soluble in polar solvents, high boiling and melting temperatures, electrical conductivity.
Ionic bond: when e bonds are completely transferred to a more electronegative atom. The mechanism consists of the formation of ions and the formation of a crystal lattice by the ions. Truly ionic - compounds with 87% ionicity.
Properties: hard, brittle, soluble in polar solvents, high boiling and melting temperatures, electrical conductivity.
Metallic bond: characteristic of elemental metals and occurs to a limited extent in nature. It is characterized by a metal crystal lattice in the nodes of which metal ion atoms are located, and the interstices are occupied by chemical bonds.
Properties of M bonds: chemical properties: the ability to lose valence e, that is, reducing properties. Physical properties of malleability, plasticity. Heat and electrical conductivity.
Complex compounds: higher order compounds that include a complex, highly stable particle - a complex ion. CI and higher sphere ions are connected by electrostatic interaction. The complexing agent and the legends are linked by a covalent bond through a donor-acceptor mechanism.
Characteristics: the complexing agent is an acceptor and also provides a certain number of orbitals, which is called the coordination number.
Legends are characterized by the amount of dentinity.
Dissociation: 1.ionization or primary dissociation, 2.Secondary dissociation occurs to a negligible extent along the covalent bond.
Classification of complex compounds: Classes of inorganic compounds
Reactions of complex compounds: 1.CS participate in metabolic processes with the preservation of the complex ion.
2. Destruction of the CI is possible if a more stable particle is formed.
= STUDENT CITY = Freshman's notebook
SEMESTER 1 EXAM
Exam program for the course "Fundamentals of Inorganic and Experimental Chemistry"
1st semester, JNF, 2011/2012 academic year
Chemical balance. Signs of true balance. Equilibrium constants in homogeneous and heterogeneous systems. Equilibrium concentrations of reagents and products and the concept of their calculation.
Le Chatelier's principle and the shift of chemical equilibrium with changes in temperature, pressure, concentrations of reagents and products.
Redox reactions(OVR). The degree of oxidation of atoms and its change in ORR. Typical oxidizing agents and reducing agents. Substances with oxidizing and reducing functions. The role of the environment in OVR. Drawing up ORR equations using the method of electron-ion half-reactions.
Standard electrochemical potential as a characteristic of the redox properties of substances in aqueous solution. Criterion for the direction of OVR under standard conditions. Solving calculation problems.
General properties of solutions. Solvent and solute. Concentrated and diluted solutions. Saturated, unsaturated and supersaturated solution and methods for their preparation. Solubility. Thermal effect of dissolution. Diagrams (polytherms) of solubility. Dependence of the solubility of gases and crystalline substances in liquid solvents on temperature.
Solutions of electrolytes and non-electrolytes. Ostwald's law of dilution.
Sparingly soluble strong electrolytes and solubility product (SP). Calculations using PR values. Conditions for precipitation and dissolution of sediments. Shift of phase equilibria in saturated solutions of sparingly soluble strong electrolytes.
Basic concepts of proton theory acids and bases. Protic solvents and their ionic products. Acid and base in proton theory. Constants of acidity and basicity and the relationship between them. Ampholytes.
Shift of protolytic equilibria under the influence of temperature, protolyte concentration (dilution) and with the introduction of the same ions of protolysis products. The degree of protolysis and pH in solutions close to infinite dilution.
Ionic product of water. Hydrogen and hydroxide indicators of medium acidity. pH scale for aqueous solutions.
Solvolysis and hydrolysis. Irreversible hydrolysis of binary compounds. Reversible hydrolysis of salts. Shift in hydrolysis equilibria.
Calculations of pH values and the degree of protolysis in the case of strong and weak acids and bases, as well as ampholytes.
The structure of atoms and the Periodic Law. Hydrogen atom. Multielectron atoms. The main thing is the orbital, magnetic and spin quantum numbers. Atomic orbitals, electronic levels and sublevels.
Principle of minimum energy, Hund's rule and Pauli's principle. The order in which electrons occupy atomic orbitals. Klechkovsky's rule. Electronic formulas and energy diagrams of atoms of elements.
Periodic table of chemical elements by D. I. Mendeleev. Periods and groups. Sections s-, p-, d- And f- elements.
Chemical bond. Ionic and covalent bonds. Basic concepts of the valence bond method. Overlap of electron orbitals; sigma, pi and delta binding. Multiple connections. The idea of hybridization and the geometry of molecules.
Polarity of bonds and polarity of molecules. Dipole moment of a chemical bond and dipole moment of a molecule.
The concept of the molecular orbital method. Hydrogen bonding and intermolecular interaction.
Required knowledge for students to receive a positive grade in the 1st semester exam
1. Symbols chemical elements and their names. Sections s-, p-, d- And f- elements in the Periodic Table.2. Nomenclature inorganic substances (formulas and names contained in the lecture course, laboratory practice and homework).
3. Electronic configurations atoms by their coordinates (group number, period number) in the Periodic System.
4. Main, orbital and magnetic quantum numbers, the connection between them and the number of energy levels, sublevels and atomic orbitals.
5. Definition type of hybridization atomic orbitals and prediction of the geometric shape of AB type particles X(molecules or ions), where A, B are atoms s- And p- elements.
6. Equilibrium constant. Acidity and basicity constants. Le Chatelier's principle to shift chemical equilibrium.
7. Solubility inorganic substances. Product of solubility. The condition for precipitation and its dissolution.
8. Drawing up reaction equations following types:
* exchange reactions in aqueous solution (molecular and ionic equation)
* redox reactions in aqueous solution (molecular and ionic equation, selection of coefficients by the method of electron-ionic half-reactions)
* protolytic reactions involving water as a solvent
* reactions of hydrolysis of salts, hydrolysis of binary compounds.
9. Composition of solutions:
* mass fraction
* molarity (molar concentration of solute)
10. Acidic, alkaline and neutral environments aqueous solutions. Hydrogen index (pH). pH scale for aqueous solutions.
What Students Need to Know About the Inorganic Chemistry Written Exam
# The exam starts at 9.00 in room K-2. For students with a cumulative grade in general chemistry for 1 semester from 15 to 24 points, the exam begins at 9.30. Students of the specified category have the right to choose the type of ticket to take the exam: basic level (maximum score 50 points) or tickets reproductive level (maximum score 24 points).
# Students without grade books are not allowed to take the exam. If a student is not admitted to the exam due to lack of credits or for other reasons, the department can accept an exam from him only with written permission (admission) from the dean's office.
# Time to complete written work in the exam from 9.00 to 12.00(from 9.30 to 12.30). During the exam, you are allowed to use reference tables for inorganic chemistry (issued by the teacher on duty) and a microcalculator. Students receive paper for written work from the teacher on duty along with the exam card.
# During the exam not allowed use a mobile phone, electronic notebook, laptop computer. Student leaving the audience during the exam is possible only with the permission of the teacher on duty and in all cases entails a change in the exam card.
# Announcement of results exam - on the day of the exam, at 15.00 at the Department of Inorganic Chemistry. Issuance of test books - at 15.00, only personally to each student.
# Examination ticket includes 6 questions on the following topics:
1. Chemical equilibrium;
2. General properties of solutions, solubility product;
3. Redox reactions;
4. Protolytic equilibria, hydrolysis;
5. The structure of the atom and the Periodic Law;
6. Chemical bonding and molecular structure.
## 2, 3 or 4 ticket question represents calculation problem one of the types studied in the 1st semester.
## calculation problem is accompanied by additional questions, not required to answer a satisfactory or good grade (in italics, surrounded by a box).
## To receive a positive rating (“satisfactory”) you must give correct answers to all six questions(see "Required knowledge of students to receive a positive grade"). Answers to questions must be clear, clear, justified, chemically literate (including the correct representation of formulas, equations of chemical reactions, the use of modern symbols of physical and chemical quantities, the derivation of calculation formulas when solving problems, etc.).
A correct, complete and reasonable answer to the additional question serves as the basis for an excellent assessment of the work.
Written examination work is graded in points in the following way:
41-50 points - “excellent”
31-40 points - “good”
21-30 points - “satisfactory”
0-20 points - “unsatisfactory”
n1.doc
2. Atomic-molecular teaching of chemistry.The main provisions were formulated by Lomonosov in the form of a capsular theory of the structure of matter - all substances consist of the smallest particles of capsules (molecules) having the same composition as the whole substance, and being in continuous motion. Chemical element - a type of atom with the same positive nuclear charge. Atom – the smallest particle of a chemical element that is the carrier of its properties. An atom is an electrically neutral microsystem that obeys the laws of quantum physics and consists of a positively charged nucleus and negatively charged electrons. Molecule – the smallest particle of a substance that determines its properties and is capable of independent existence. Atoms are combined into a molecule using chemical bonds, in the formation of which mainly external (valence) electrons take part.
In 1911, Rutherford carried out experiments to clarify the structure of the atom. In 1913, the simplest planetary model of the “hydrogen atom” of Bohr-Rutherford appeared.
This model is currently the generally accepted “official” model of the atom.
advantage is simplicity. According to this model, the atom was supposed to consist of a compact positive nucleus and an electron rotating around it in “stationary circular orbits”. These shortcomings are simply striking:
1) an electron around an atom, according to the solution to the problem of body motion in a central field, cannot move along circular trajectories. The trajectories should be elliptical. But elliptical trajectories are impossible in such a model
N. Bor An atom can only be in special stationary states, each of which has a specific energy. In a stationary state, an atom does not emit electromagnetic waves.
The emission and absorption of energy by an atom occurs during an abrupt transition from one stationary state to another. Advantages:
Explained the discreteness of the energy states of hydrogen-like atoms.
Bohr's theory approached the explanation of intra-atomic processes from a fundamentally new position and became the first semi-quantum theory of the atom. Flaws
Could not explain the intensity of the spectral lines.
Valid only for hydrogen-like atoms and does not work for atoms following it in the periodic table.
3.B1924 G. French physicist Louis de Broglie proposed the idea that matter has both wave and corpuscular properties. According to de Broglie's equation (one of the basic equations of quantum mechanics),
i.e., a particle with mass m moving with speed v corresponds to a wave of length?; h is Planck's constant. For any particle with mass m and known velocity v, the de Broglie wavelength can be calculated. De Broglie's idea was experimentally confirmed in 1927, when both wave and corpuscular properties of electrons were discovered. In 1927, the German scientist W. Heisenberg proposed the uncertainty principle, according to which for microparticles it is impossible to simultaneously accurately determine both the coordinate of the particle X and the px component of the pulse along the x axis. An atom with more than one electron is a complex system of electrons interacting with each other moving in the field of the nucleus. Nevertheless, it turns out that in an atom it is possible, with good accuracy, to introduce the concept of the states of each electron separately as stationary states of electron motion in some effective centrally symmetric field created by the nucleus together with all other electrons. For different electrons in an atom, these fields are, generally speaking, different, and they must all be determined simultaneously, since each of them depends on the states of all other electrons. Such a field is called self-consistent. Since a self-consistent field is centrally symmetric, each state of an electron is characterized by a certain value of its orbital momentum /. The states of an individual electron at a given / are numbered (in increasing order of their energy) using the principal quantum number n, running through the value n = / +1, /+2, ...; this choice of numbering order is established in accordance with that adopted for the hydrogen atom. But the sequence of increasing energy levels with different / in complex atoms, generally speaking, differs from that occurring in the hydrogen atom.
4. Principles of filling orbitals.
1. Pauli principle. There cannot be two electrons in an atom whose values of all quantum numbers (n, l, m, s) would be the same, i.e. Each orbital can contain no more than two electrons (with opposite spins).
Khar-kakov. St.
Light energy, light length, saturation, direction.
12.VS method.
Implied. Images Elect. Density through the socialization of electrons located on the outside. Electron. Level.
Flaws
Could not explain the vapor magnetic properties of some compounds. (O at t -220 becomes liquid, which is attracted by a magnet)
Creatures Mol. Ions (He 2+, H 2+, O 2-)
Provisions
Image. x/s is the result of the transition of electrons from atomic orbitals to new levels having a defined energy. Atom by all. Molecules
After the image. Mol. Orbital - atomic Orb. They lose their individuality.
Each pier Orb. Resp. Defined energy.
Electrons in a molecule are non-localized. In the internuclear spaces of 2 atoms, and find. In the range of nuclear weapons
Hybridization is self-producing. Form and energy leveling process.
13. MO method
An improved version of the valence bond method. Based on principles. 1. Chemical bonds between atoms are carried out through one or more electron pairs. 2. When a common electron pair is formed, the electron clouds overlap. The stronger the overlap, the stronger the chemical bond. 3. When a common electron pair is formed, the electron spins must be antiparallel. 4. Only unpaired electrons of atoms can participate in the formation of common electron pairs. Paired electrons must be separated to form bonds. 5. When a covalent bond is formed from a certain number of electron clouds of two atoms, the same number of electron clouds of a molecule belonging to both atoms is formed. 6. When electron clouds combine, their mutual overlap with the formation of bonding clouds of a molecule and mutual repulsion with the formation of loosening clouds of a molecule are possible. 7. The filling of the orbitals of a molecule with electrons occurs in accordance with the principles of minimum energy and Pauli (An atom cannot have 2 electrons that have the same values of all 4 quantum numbers. No more than 2 electrons can be located in one orbital). 8. A bond is formed when the number of electrons in bonding orbitals is greater than in antibonding orbitals. Properties of covalent bonds. It is durable. Has the property of saturation. Has directionality in space.
14.chem. thermodynamics studies energy. Changes.under consideration processes in the state Equilibrium p-I either did not begin or ended and the flows into the outside. There are no environments.
Thermodyne. A system is a macroscopic body isolated from the mental environment. or physical shells.
By number of phases:
Homogeneous (all components of the system are in one phase)
Heterogeneous (chemical reactions occur in different phase sections)
According to the nature of the interaction with the environment. Wednesday:
Open (exchange of things and energy), Closed (exchange of energy), Isolated (no exchange)
All vehicles are characterized by parameters: pressure, tempo, volume, mass. Thermodyne. Studies the transition of the system. From one composition. In the other - process: Equilibrium any chemical. district in composition Equilibrium, Stationary.
Isobaric(constant Pressure), Isochoric(constant Volume), Isothermal(constant Temperature)
Vehicle energy: E = K + P + delta U (internal)
Chem. thermodyne Based on 2 laws
Law. Save Energy - change in ext. Energy Syst. Def. Quantity of heat released and work done
Standard enthalpy is the enthalpy of that solution in which 1 mole of a substance is formed from simple substances that are stable. At std. Terms.
15.First law of thermodynamics
Enthalpy – state function equal to the internal energy of the system + expansion work. . At constant pressure
1 law-thermal effect p-i = thermal. Ef. Reverse p-i, but opposite in sign. (The greater the heat. Effect of formation of a complex substance, the more stable it is.)
16.Hess's law - heat. Ef. Chem. p-i does not depend on the path along which it flows, but depends on the initial and final state. syst.
Consequence
-cheating Enthalpies chem. r-i does not depend on the number of int. stages
High selectivity
The ability to regulate catalytic properties.
24.
Chemical
equilibrium
– the state of the system in which the rates of forward and reverse reactions are equal.
Reversible-protek. Not completely and the products of such r-th mutual. from images. ref. in-in.
Irreversible r-i- leaked. until the end, until complete consumption. ref. in-in and product. r-i (image of sediment, gas, water)
Constant
chemical equilibrium reaction = the product of the concentrations of reaction products, taken to the powers of their stoichiometric coefficients in the reaction equation, divided by the product of the concentrations of the starting substances, taken to the powers of the stoichiometric coefficients
25.
the process proceeds spontaneously in the forward direction if the potential decreases, therefore the equilibrium constant is greater than 1. Concentration of products > concentration of starting substances. If on the contrary, then there was practically no reaction. When the temperature increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, towards the exothermic reaction. As the pressure increases, the equilibrium shifts in the direction of the reaction that occurs with a decrease in the volume of gaseous substances; when the pressure decreases, in the direction of the reaction that occurs with an increase in volume. As the concentration of the starting substances increases, the equilibrium shifts towards the direct reaction.
Le Chatelier-Brown principle . If an external influence is exerted on a system in equilibrium, then the equilibrium shifts in a direction that weakens this influence
26. Solutions - solid, liquid, gas - homogeneous system. image. growth, growth and product. Their interaction
The solvent is a component that does not change its aggregate. comp. with images. solutions
Concentration - quantity of solution. in units volume or mass of the ras-ra or rast-la.
27.
Solubility is the ability of a substance to form homogeneous systems with other substances - solutions in which the substance is found in the form of individual atoms, ions, molecules or particles.
The growth process is complex physical and chemical. yavl., one of the physical. phenomena processes Diffusion solution. in the growth of this process of spontaneous movement. The force of diffusion is Warm. Movement
The reasons for the difference are an increase in entropy and the speed of the solution. depends on the rate of diffusion.
Rule of earthenware phases
28.
dissolution of gases into liquids. ectotherm process. (when gases disintegrate into liquids.
Henry's Law:
Gas mass at a given temp. And this volume is liquid. directly proportional partial pressure gas
Dalton's Law:
The growth of each of the gas components of the mixture at constant. Temp., directly proportional. partial pressure liquid component and does not depend on the general pressure. mixtures and individual component.
Sechenov's Law:
In the presence of electrolytes, gas grows into liquid. decrease
29.Collegiate name saints depending on the concentration. raster, but not dependent. from their chem. comp.
Pressure
rich
pair
liquids
called the pressure that is established above the liquid when the rate of evaporation of the liquid = the rate of condensation of vapor into liquid. 1
law
Raoul.
Relative decrease in solvent vapor pressure above solution = mole fraction of solute
Solutions
subordinate
this
law
are called
ideal. 2
law
Raoul.
Ebulioscopic.
The increase in boiling point of a nonelectrolyte solution is proportional to the molal concentration of the solute. 
, E-ebullioscopic constant. E = increase in boiling point caused by 1 mole of a substance dissolved in 1000 g of solvent. Cryoscopic.
The decrease in the freezing point of a nonelectrolyte solution is proportional to the molar concentration of the solute. 
,
K-cryoscopic = lowering the freezing point of solutions in which there is 1 mole of dissolved non-electrolyte per 1000 g of solvent.
30.Diffusion and osmosis.
Osmosis is the one-way diffusion of solvent molecules in a solution through a membrane impermeable to the dissolved substance.
reactions,
divided
on
work
concentrations
original
substances
taken
V
degrees
their
stoichiometric.
Let us denote K* by KH 2 O. The quantity is called the ionic product of water. Ionic
work
water= product of the concentration of hydrogen cations and the concentration hydroxide anions. Water dissociation constant 
. Changing the concentrations of protons and hydroxide ions in a solution creates an acidic or alkaline environment. -7 – alkaline,
>10 -7 – acidic. 
.
Hydrogen
indicator (pH) numerically = the decimal logarithm of the concentration of hydrogen cations, taken with the opposite sign. 
, the hydroxide index is calculated similarly 
. For a neutral environment [pH] =7, alkaline - [pH] >7, acidic - [pH]
38. Hydrolysis of salts. Constant and degree of hydrolysis. Hydrolysis– reaction of salt with water to form a weak electrolyte. Accompanied by a change in the pH of the environment. Example Na 2 CO 3 =Na + +CO 3 2- dissociation, CO 3 2- +H 2 O=HCO 3 - +OH - hydrolysis. Hydrolysis consists of the chemical interaction of dissolved salt ions with water molecules, leading to the formation slightly dissociated compounds and changes in the reaction of the environment. Quantitative value characterizing hydrolysis is called the degree of hydrolysis h. Degree hydrolysis– number ratio hydrolyzed salt molecules to the total number of dissolved salt molecules. . Dependence of the degree of hydrolysis. Concentration substances– the greater the dilution, the greater the degree of hydrolysis. Temperature – the higher the temperature, the stronger the hydrolysis. Addition strangers substances– the introduction of substances that give an alkaline reaction, suppresses the hydrolysis of salts with pH > 7 and enhances hydrolysis with pH 7, and vice versa, substances that give an acidic reaction to the environment increase hydrolysis with pH > 7 and suppress with pH 7. nature dissolved substances– the degree of hydrolysis depends on the chemical. the nature of the dissolved salt. There are 3 options.
42.cooking methods:
Without solution (by mixing the shielded quantities of liquids; by adding shielding quantities of solids to the solution)
According to the equation p-i
43.Buffer solutions– solutions that practically do not change their pH value when diluted or added to them in certain quantities of a strong acid or strong base
Buffer capacity. Expressed as the amount of substance equivalent to a strong acid or base that must be added to 1 liter of a buffer solution to shift its pH value by one.
44. Heterogeneous equilibria
At contact solid substance with a solvent, the substance begins to dissolve and upon establishment thermodynamic equilibrium, a saturated solution is formed. When sparingly soluble electrolyte in an aqueous solution saturated relatively sparingly soluble electrolyte.
Product of solubility - product of ion concentration sparingly soluble electrolyte in its saturated solution at constant temperature and pressure. Work solubility-value constant.
A precipitate will form if the ionic product is greater than the solubility product
45.ORP. Redox reactions– such reactions that occur with a change in the oxidation states of the elements that make up the compounds. Oxidation state is the actual charge of an atom in a molecule resulting from redistribution. electron density.
46. Oxidation is the process of losing electrons, leading to an increase in CO. Oxidizing agents: simple substances, atoms that have a large electronegativity (F, O. CE); substances, containing. Elements in max CO; cations ME and N.
Reducing agents: simple substances whose atoms have low EO; uh, you're in the bottom. CO
47.Intermolecular- change CO in different molecules exl.comproportionation(ok, it's the same email but in different COs)
Intramolecular -ism. CO in one molecule
2. Klechkovsky's rule (principle of least energy). In the ground state, each electron is arranged so that its energy is minimal. The smaller the sum (n + l), the lower the energy of the orbital. For a given value (n + l), the orbital with the smaller n has the lowest energy. The energy of the orbitals increases in the series:
1S
3. Hund's rule. An atom in the ground state must have the maximum possible number of unpaired electrons within a certain sublevel.
The state of an atom with the minimum possible energy of electrons in it is called the ground, or unexcited, state. However, if atoms receive energy from the outside (for example, during irradiation, heating), then the electrons of the outer electron layer can “steam apart” and move to free orbitals characterized by higher energy. This state of the atom is called excited.
5.Periodic law. The properties of elements, as well as the structure and properties of their compounds, periodically depend on the charge of the nuclei of their atoms. The atomic number of an element = the charge of its nucleus and the number of electrons. Number of neutrons = atomic mass – atomic number. Each period begins with s - elements (s 1 alkali metal) and ends with p - element (s 2 p 6 inert gas). 1st period contains 2 s-elements. 2-3 contains 2 s - elements and 6 p - elements. In 4-5 d elements are wedged between s and p. Number of electronic levels = period number. For elements of the main subgroups, the number of electrons = group number. In the group from top to bottom, the metallic properties are enhanced. From left to right, non-metallic properties (the ability to accept electrons) are enhanced. Frequency of changes in the properties of s-, p- and d elements.
Atom chem. element consists of 3 main elementary particles: positively charged protons, uncharged neurons and negatively charged electrons. At the center of the atom there is a nucleus consisting of protons and neutrons, and electrons rotate in orbitals around it. Number of electrons = charge of nucleus. Chemical element– a type of atom with a certain nuclear charge. Isotopes- atoms of the same element that have the same nuclear charge but different masses. Isobars – atoms of different elements having different nuclear charges, but the same atomic mass. The modern model is based on 2 fundamental principles of quantum physics. 1. An electron has the properties of both a particle and a wave at the same time. 2. particles do not have strictly defined coordinates and velocities. Energy level(quantum number n) – distance from the nucleus. As n increases, the electron energy increases. The number of energy levels = the number of the period in which the element is located. The maximum number of electrons is determined by N=2n 2. Energy sublevel denoted by the letters s (spherical), p (dumbbell-shaped), d (4 petal rosette), f (more complex). Magnetic quantum number interaction of an electron cloud with external magnetic fields. Spin quantum number is the intrinsic rotation of an electron around its axis .
7. x/s- result of interaction atom drive. to image chem. molecules.
8.energy- necessary for the rupture of x/c or released during the formation of x/c.
Length is the shortest distance between the nuclei of interacting atoms
Saturation-number x/s which can image. Atom of a given element.
Saturation - valence
Focus - strict location x/s in three-dimensional space
9.1.Orientation-interaction. Communication With the presence of 2 or more floors. they say
2. induction - one mol. Polar, the second is not
3. dispersive - associated with the image. Instantaneous dipoles (characteristic for non-pol. Mol.)
10.Inon light-result electrostat. mutual m/y ions. (limiting case of forged field. St.) total electr. A pair refers to only one of the interactions. Atoms.
polarization-phenomenon Spat. Atom deformations found. In the operating area constant or electric Molek. cathode(-) anode(+)
the ability to undergo polarization (polarizability) of an ion, radius.
11.Kov x/s - the process of socialization of electrons is found. On external Energetic Level.
Non-polar (non-difference H2) polar (NSE)
Mechanisms image.
Exchange- into the image x/c participation One electron from each atom
Donor-acceptor- donor(electronic pair) acceptor(orbital)
Dative- variety Donor-acceptance. In which each of the atoms simultaneously appears. Both donor and acceptor
-enthalpy
x/r
=
sum
enthalpy
product image
district
behind
minus
amounts
enthalpy
arr.
Exodus.
thing
1. Subject and tasks of chemistry. Basic concepts and laws of chemistry.
2. Periodic law and the Periodic table of chemical elements D.I. Mendeleev based on ideas about the structure of atoms. The significance of the Periodic Law for the development of science.
3. The structure of atoms of chemical elements and patterns in changes in their properties using the example of: a) elements of the same period; b) elements of one main subgroup.
4. Types of chemical bonds: ionic, metallic, covalent (polar, non-polar); simple and multiple bonds in organic compounds. Types of crystal lattices.
5. Classification of chemical reactions in inorganic chemistry.
6. Classification of chemical reactions in organic chemistry
7. Rate of chemical reactions. Dependence of speed on nature, concentration of reactants, temperature, catalyst.
8. Chemical equilibrium and conditions for its displacement: changes in the concentration of reactants, temperature, pressure.
9. The concept of allotropy. Allotropy of inorganic substances using the example of carbon and oxygen.
10. Disperse systems. Classification, examples. Colloidal solutions. Application in medicine of suspensions, emulsions, aerosols, gels.
11. Solutions. True solutions. Solubility of substances as a physical and chemical phenomenon.. Classification of solutions. Types of concentration.
12.Electrolytic dissociation. Electrolytes and non-electrolytes. Ion exchange reactions. Degree of dissociation.
13. The most important classes of inorganic compounds.
14.Oxides. Higher oxides of chemical elements of the third period. Regularities in changes in their properties in connection with the position of chemical elements in the Periodic Table.
15. Acids, their classification and properties based on ideas about electrolytic dissociation.
16. Bases, their classification and properties based on ideas about electrolytic dissociation.
17. Salts, their composition and names, interaction with metals, acids, alkalis, with each other, taking into account the characteristics of oxidation - reduction and ion exchange reactions.
18. Hydrolysis of salts. Types of hydrolysis.
19. Redox reactions (using the example of the interaction of aluminum with oxides of some metals, concentrated sulfuric acid with copper).
20.Electrolysis of melts and salt solutions.
21. Non-metals, position in the periodic table of chemical elements D.I. Mendeleev, the structure of their atoms. Redox properties of nonmetals using the example of elements of the oxygen subgroup. . Hydrogen compounds of nonmetals. Regularities in changes in their properties in connection with the position of chemical elements in the Periodic Table D.I. Mendeleev
22. Halogens. General characteristics of halogens. Chlorine. Physico-chemical properties. Hydrochloric acid, its properties. Chlorides.
23. Oxygen subgroup. General characteristics of the VIA subgroup. Sulfur, its physicochemical properties. Sulfur compounds: hydrogen sulfide, sulfur oxides, sulfuric acid and its salts.
24. Nitrogen subgroup.. Nitrogen compounds: ammonia, ammonium salts, nitric acid and its salts.
25. Subgroup of carbon. General characteristics. Carbon. Atomic structure. Allotropic modifications of carbon. Chemical properties. Carbon compounds: oxides, carbonic acid and its salts.
26. Metals, their position in the Periodic Table of Chemical Elements D.I. Mendeleev, the structure of their atoms, metallic bonds. General chemical properties of metals. . Electrochemical voltage series of metals. Displacement of metals from salt solutions by other metals
27. Chemical and electrochemical corrosion of metals. Conditions under which corrosion of metals occurs. Conditions under which corrosion occurs, measures to protect metals and alloys from corrosion
28. General methods of obtaining metals. The practical significance of electrolysis using the example of salts of oxygen-free acids.
29. Alkali metals. General characteristics based on the position in D.I. Mendeleev’s PSHE. Properties of sodium and its compounds. Biological role of sodium and potassium ions.
30. alkaline earth metals. Calcium, its properties. The most important calcium compounds. Biological role of calcium ions.
31. Iron: position in the Periodic Table of Chemical Elements D.I. Mendeleev, atomic structure, possible oxidation states, physical properties, interactions with oxygen, halogens, solutions of acids and salts. Iron alloys.
32. Reasons for the diversity of inorganic and organic substances; relationship of substances.
33 Basic principles of the theory of the chemical structure of organic substances A.M. Butlerov. Chemical structure as the order of connection and mutual influence of atoms in molecules.
34. Isomerism of organic compounds and its types.
35. Saturated hydrocarbons, general formula and chemical structure of homologues of this series. Properties and applications of methane.
36. Unsaturated hydrocarbons of the ethylene series, general formula and chemical structure. Properties and applications of ethylene. Methods for producing ethylene hydrocarbons
37. Acetylene is a representative of hydrocarbons with a triple bond in the molecule. Properties, production and use of acetylene.
38. Aromatic hydrocarbons. Benzene, structural formula, properties and preparation. Application of benzene and its homologues.
39. Natural sources of hydrocarbons: gas, oil, coal and their practical use.
40. Saturated monohydric alcohols, their structure, properties. Preparation and use of ethyl alcohol. Preparation of alcohols from saturated and unsaturated hydrocarbons.
41. Phenol, its chemical structure, properties, preparation and use.
42. Aldehydes, their chemical structure and properties. Preparation and use of formic and acetaldehydes.
43. Limit monobasic carboxylic acids, their structure and properties using acetic acid as an example.
44. Fats, their composition and properties. Fats in nature, transformation of fats in the body. Products of technical processing of fats. The concept of synthetic detergents.
45. Glucose is a representative of monosaccharides, chemical structure, physical and chemical properties, application
46. Starch, occurrence in nature, practical significance, hydrolysis of starch
47. Cellulose, composition of molecules, physical and chemical properties, application. The concept of artificial fibers using the example of acetate fiber.
48. Amino acids, their composition and chemical properties: interaction with hydrochloric acid, alkalis, with each other. Biological role of amino acids and their use.
49. Aniline is a representative of amines; chemical structure and properties; production and practical application.
50. The relationship between the most important classes of organic compounds. Genetic connection.
51. Proteins as biopolymers. Properties and biological functions of proteins.
52. General characteristics of high-molecular compounds: composition, structure, reactions underlying their production (for example, polyethylene or synthetic rubber).
53. Types of synthetic rubbers, their properties and applications.
54. Vitamins. Classification of vitamins. Biological role of vitamins.
55. Enzymes. Classification. Biological role.
56. Hormones. Classification. Biological role.
Related information.
Chemistry tickets for the 10th grade course.
Ticket No. 1
The periodic law and the periodic system of chemical elements of D. I. Mendeleev based on ideas about the structure of atoms. The importance of the periodic law for the development of science.In 1869, D.I. Mendeleev, based on an analysis of the properties of simple substances and compounds, formulated the Periodic Law:
The properties of simple bodies... and compounds of elements are periodically dependent on the size of the atomic masses of the elements.
Based on the periodic law, the periodic system of elements was compiled. In it, elements with similar properties were combined into vertical columns - groups. In some cases, when placing elements in the Periodic Table, it was necessary to disrupt the sequence of increasing atomic masses in order to maintain the periodicity of the repetition of properties. For example, it was necessary to “swap” tellurium and iodine, as well as argon and potassium.
The reason is that Mendeleev proposed the periodic law at a time when nothing was known about the structure of the atom.
After the planetary model of the atom was proposed in the 20th century, the periodic law was formulated as follows:
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The properties of chemical elements and compounds are periodically dependent on the charges of atomic nuclei.
The charge of the nucleus is equal to the number of the element in the periodic table and the number of electrons in the electron shell of the atom.
This formulation explained the “violations” of the Periodic Law.
In the Periodic System, the period number is equal to the number of electronic levels in the atom, the group number for elements of the main subgroups is equal to the number of electrons in the outer level.
The reason for the periodic change in the properties of chemical elements is the periodic filling of electronic shells. After filling the next shell, a new period begins. The periodic change of elements is clearly visible in the changes in the composition and properties of the oxides.
Scientific significance of the periodic law. The periodic law made it possible to systematize the properties of chemical elements and their compounds. When compiling the periodic table, Mendeleev predicted the existence of many undiscovered elements, leaving free cells for them, and predicted many properties of undiscovered elements, which facilitated their discovery.
Ticket No. 2
The structure of atoms of chemical elements using the example of elements of the second period and IV-A group of the periodic system of chemical elements by D. I. Mendeleev. Regularities in the changes in the properties of these chemical elements and the simple and complex substances formed by them (oxides, hydroxides) depending on the structure of their atoms.As you move from left to right along a period, the metallic properties of the elements become less and less pronounced. When moving from top to bottom within one group, elements, on the contrary, display increasingly pronounced metallic properties. Elements located in the middle part of the short periods (2nd and 3rd periods) usually have a skeleton covalent structure, and elements from the right part of these periods exist in the form of simple covalent molecules.
Atomic radii change as follows: decrease when moving from left to right along a period; increase as you move from top to bottom along the group. As you move from left to right across a period, electronegativity, ionization energy, and electron affinity increase, reaching a maximum for the halogens. For noble gases, electronegativity is 0. Changes in the electron affinities of elements when moving from top to bottom along the group are not so characteristic, but at the same time the electronegativity of the elements decreases.
In elements of the second period, 2s and then 2p orbitals are filled.
The main subgroup of group IV of the periodic system of chemical elements by D. M. Mendeleev contains carbon C, silicon Si, germanium Ge, tin Sn and lead Pb. The outer electron layer of these elements contains 4 electrons (s 2 p 2 configuration). Therefore, the elements of the carbon subgroup must have some similarities. In particular, their highest oxidation state is the same and is +4.
What causes the difference in the properties of the elements of the subgroup? The difference between the ionization energy and the radius of their atoms. As the atomic number increases, the properties of elements naturally change. Thus, carbon and silicon are typical non-metals, tin and lead are metals. This is manifested primarily in the fact that carbon forms a simple non-metal substance (diamond), and lead is a typical metal.
Germanium occupies an intermediate position. According to the structure of the electron shell of the atom, p-elements of group IV have even oxidation states: +4, +2, – 4. The formula of the simplest hydrogen compounds is EN 4, and the E-H bonds are covalent and equivalent due to the hybridization of s- and p-orbitals with the formation of sp 3 orbitals directed at tetrahedral angles.
The weakening of the characteristics of a non-metallic element means that in the subgroup (C-Si-Ge-Sn-Pb) the highest positive oxidation state +4 becomes less and less characteristic, and the oxidation state +2 becomes more typical. So, if for carbon the most stable compounds are those in which it has an oxidation state of +4, then for lead the compounds in which it exhibits an oxidation state of +2 are most stable.
What can be said about the stability of compounds of elements in the negative oxidation state -4? Compared to non-metallic elements of groups VII-V, p-elements of group IV exhibit the signs of a non-metallic element to a lesser extent. Therefore, for elements of the carbon subgroup, a negative oxidation state is atypical.
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Ticket number 3.
Types of chemical bonds and methods of their formation in inorganic compounds: covalent (polar, nonpolar, simple and multiple bonds), ionic, hydrogen.
^ Covalent bond formed by the overlap of the electron clouds of two atoms. Each atom contributes one unpaired electron to form one chemical bond, which results in the formation of shared electron pair. If a covalent bond is formed between two identical atoms, it is called non-polar.
If a covalent bond is formed between two different atoms, the shared electron pair is shifted to the atom with greater electronegativity (electronegativity is the ability of an atom to attract electrons). In this case, there is polar covalent bond.
A special case of a covalent bond is donor-acceptor bond. For its formation, one atom must have a free orbital at the outer electronic level, and the other must have a pair of electrons. One atom (donor) provides another (acceptor) with its electron pair, as a result it becomes shared and a chemical bond is formed. Example - CO molecule:
^ Ionic bond formed between atoms with very different electronegativity. In this case, one atom gives up electrons and turns into a positively charged ion, and the atom that received electrons turns into a negatively charged one. The ions are held together by electrostatic attractive forces.
^ Hydrogen bond is formed between polar molecules (water, alcohols, ammonia) due to the attraction of opposite charges.
The strength of a hydrogen bond is significantly (~20 times) less than that of an ionic or covalent bond.