The international unit of atomic mass is equal to 1/12 of the mass of the 12C isotope, the main isotope of natural carbon.
1 amu = 1/12 m (12C) = 1.66057 10-24 g
Relative atomic mass (Ar) is a dimensionless quantity equal to the ratio of the average mass of an atom of an element (taking into account the percentage of isotopes in nature) to 1/12 of the mass of a 12C atom.
The average absolute mass of an atom (m) is equal to the relative atomic mass times the amu.
(Mg) = 24.312 1.66057 10-24 = 4.037 10-23 g
Relative molecular mass (Mr) is a dimensionless quantity that shows how many times the mass of a molecule of a given substance is greater than 1/12 the mass of a 12C carbon atom.
Mg = mg / (1/12 ma(12C))
mr is the mass of a molecule of a given substance;
ma(12C) is the mass of the 12C carbon atom.
Mg = Σ Ar(e). The relative molecular mass of a substance is equal to the sum of the relative atomic masses of all elements, taking into account the indices.
Mg(B2O3) = 2 Ar(B) + 3 Ar(O) = 2 11 + 3 16 = 70
Mg(KAl(SO4)2) = 1 Ar(K) + 1 Ar(Al) + 1 2 Ar(S) + 2 4 Ar(O) =
1 39 + 1 27 + 1 2 32 + 2 4 16 = 258
The absolute mass of a molecule is equal to the relative molecular mass times the amu. The number of atoms and molecules in ordinary samples of substances is very large, therefore, when characterizing the amount of a substance, a special unit of measurement is used - the mole.
Amount of substance, mol. Means certain number structural elements (molecules, atoms, ions). Denoted by ν, measured in moles. A mole is the amount of a substance containing as many particles as there are atoms in 12 g of carbon. Avogadro diQuaregna number (NA). The number of particles in 1 mole of any substance is the same and equals 6.02 1023. (Avogadro’s constant has a dimension of mol-1).
How many molecules are there in 6.4 g of sulfur? The molecular weight of sulfur is 32 g/mol. We determine the amount of g/mol of substance in 6.4 g of sulfur:
ν(s) = m(s) / M(s) = 6.4 g / 32 g/mol = 0.2 mol
Let's determine the number of structural units (molecules) using Avogadro's constant NA
N(s) = ν(s) NA = 0.2 6.02 1023 = 1.2 1023
Molar mass shows the mass of 1 mole of a substance (denoted M).
The molar mass of a substance is equal to the ratio of the mass of the substance to the corresponding amount of the substance.
The molar mass of a substance is numerically equal to its relative molecular mass, however, the first quantity has the dimension g/mol, and the second is dimensionless.
M = NA m(1 molecule) = NA Mg 1 a.m.u. = (NA 1 amu) Mg = Mg
This means that if the mass of a certain molecule is, for example, 80 amu. (SO3), then the mass of one mole of molecules is equal to 80 g. Avogadro’s constant is a proportionality coefficient that ensures the transition from molecular relationships to molar ones. All statements regarding molecules remain valid for moles (with replacement, if necessary, of amu by g). For example, the reaction equation: 2Na + Cl2 → 2NaCl, means that two sodium atoms react with one chlorine molecule or that the same thing, two moles of sodium react with one mole of chlorine.
Stoichiometry. Law of conservation of mass of substances. The law of constancy of the composition of substances of molecular structure. Avogadro's law and consequences from it.
Stoichiometry(from Old Greekστοιχειον “element” + μετρειν “measure”) - section chemistry about the ratios of reagents in chemical reactions.
Allows you to theoretically calculate the required volumes reagents.
Law of Constancy of Composition was discovered by the French scientist Louis Jeanne Prousteau in 1799 and is formulated:
Any pure substance has a constant qualitative and quantitative composition, regardless of its location in nature and the method of production in industry.
For example: H 2 O a) high-quality composition– elements H and O
b) quantitative composition – two hydrogen atoms H, one oxygen atom O.
Water can be obtained:
1. 2H 2 + O 2 = 2H 2 O - reaction of the compound.
2. Cu(OH) 2 t°C H 2 O + CuO – decomposition reaction.
3. HCl + NaOH = H 2 O + NaCl – neutralization reaction.
The meaning of the law of constancy of composition:
· Based on the law, the concepts of “chemical compound” and “mixture of substances” were differentiated
· Based on the law, various practical calculations can be made.
Law of conservation of mass of matter was discovered by M.V. Lomonosov in 1748 and is formulated.
Basic laws of chemistry
The branch of chemistry that considers the quantitative composition of substances and quantitative relationships (mass, volume) between reacting substances is called stoichiometry. In accordance with this, calculations of quantitative relationships between elements in compounds or between substances in chemical reactions are called stoichiometric calculations. They are based on the laws of conservation of mass, constancy of composition, multiple ratios, as well as gas laws - volumetric ratios and Avogadro. The listed laws are considered to be the basic laws of stoichiometry.
Law of Conservation of Mass- the law of physics, according to which the mass of a physical system is conserved during all natural and artificial processes. In its historical, metaphysical form, according to which matter is uncreated and indestructible, the law has been known since ancient times. Later, a quantitative formulation appeared, according to which the measure of the amount of a substance is weight (later mass). The law of conservation of mass has historically been understood as one of the formulations law of conservation of matter. One of the first to formulate it was the ancient Greek philosopher Empedocles (5th century BC): nothing can come from nothing, and in no way can what exists be destroyed. Later, a similar thesis was expressed by Democritus, Aristotle and Epicurus (as retold by Lucretius Cara). With the advent of the concept of mass as a measure amount of substance, proportional to weight, the formulation of the law of conservation of matter was clarified: mass is an invariant (conserved), that is, during all processes the total mass does not decrease or increase(weight, as Newton already assumed, is not an invariant, since the shape of the Earth is far from an ideal sphere). Until the creation of microworld physics, the law of conservation of mass was considered true and obvious. I. Kant declared this law a postulate of natural science (1786). Lavoisier in his “Elementary Textbook of Chemistry” (1789), gives a precise quantitative formulation of the law of conservation of mass of matter, but does not declare it to be something new and important law, but simply mentions it in passing as a well-known and long-established fact. For chemical reactions, Lavoisier formulated the law as follows: nothing happens either in artificial processes or in natural ones, and one can put forward the position that in every operation [chemical reaction] there is the same amount of matter before and after, that the quality and quantity of the principles remained the same, only displacements and regroupings occurred.
In the 20th century, two new properties of mass were discovered: 1. The mass of a physical object depends on its internal energy. When external energy is absorbed, the mass increases, and when it is lost, it decreases. It follows that mass is conserved only in an isolated system, that is, in the absence of energy exchange with the external environment. The change in mass is especially noticeable when nuclear reactions. But even during chemical reactions that are accompanied by the release (or absorption) of heat, the mass is not conserved, although in this case the mass defect is negligible; 2. Mass is not an additive quantity: the mass of a system is not equal to the sum of the masses of its components. In modern physics, the law of conservation of mass is closely related to the law of conservation of energy and is fulfilled with the same limitation - the exchange of energy between the system and the external environment must be taken into account.
Law of Constancy of Composition(J.L. Proust, 1801-1808) - any specific chemically pure compound, regardless of the method of its preparation, consists of the same chemical elements, and the ratios of their masses are constant, and the relative numbers of their atoms are expressed in integers. This is one of the basic laws of chemistry. The law of constant composition is true for daltonides (compounds of constant composition) and is not true for berthollides (compounds of variable composition). However, for the sake of simplicity, the composition of many Berthollides is written as constant.
Law of Multiples discovered in 1803 by J. Dalton and interpreted by him from the standpoint of atomism. This is one of the stoichiometric laws of chemistry: if two elements form more than one compound with each other, then the masses of one of the elements per the same mass of the other element are related as whole numbers, usually small.
Mol. Molar mass
In the International System of Units (SI), the unit of quantity of a substance is the mole.
Mole- this is the amount of a substance containing as many structural units (molecules, atoms, ions, electrons, etc.) as there are atoms in 0.012 kg of the carbon isotope 12 C.
Knowing the mass of one carbon atom (1.933 × 10 -26 kg), we can calculate the number of N A atoms in 0.012 kg of carbon
N A = 0.012/1.933×10 -26 = 6.02×10 23 mol -1
6.02×10 23 mol -1 is called Avogadro's constant(designation N A, dimension 1/mol or mol -1). It shows the number of structural units in a mole of any substance.
Molar mass– a value equal to the ratio of the mass of a substance to the amount of substance. It has the dimension kg/mol or g/mol. It is usually designated M.
In general, the molar mass of a substance, expressed in g/mol, is numerically equal to the relative atomic (A) or relative molecular mass (M) of this substance. For example, the relative atomic and molecular masses of C, Fe, O 2, H 2 O are respectively 12, 56, 32, 18, and their molar masses are respectively 12 g/mol, 56 g/mol, 32 g/mol, 18 g /mol.
It should be noted that mass and quantity of a substance are different concepts. Mass is expressed in kilograms (grams), and the amount of a substance is expressed in moles. There are simple relationships between the mass of a substance (m, g), the amount of substance (ν, mol) and the molar mass (M, g/mol)
m = νM; ν = m/M; M = m/v.
Using these formulas, it is easy to calculate the mass of a certain amount of a substance, or determine the number of moles of a substance in a known mass, or find the molar mass of a substance.
Relative atomic and molecular masses
In chemistry, they traditionally use relative rather than absolute mass values. Since 1961, the atomic mass unit (abbreviated a.m.u.), which is 1/12 of the mass of a carbon-12 atom, that is, the isotope of carbon 12 C, has been adopted as a unit of relative atomic masses since 1961.
Relative molecular weight(M r) of a substance is a value equal to the ratio of the average mass of a molecule of the natural isotopic composition of the substance to 1/12 of the mass of a carbon atom 12 C.
The relative molecular mass is numerically equal to the sum of the relative atomic masses of all atoms that make up the molecule, and is easily calculated using the formula of the substance, for example, the formula of the substance is B x D y C z, then
M r = xA B + yA D + zA C.
Molecular mass has the dimension a.m.u. and is numerically equal to the molar mass (g/mol).
Gas laws
The state of a gas is completely characterized by its temperature, pressure, volume, mass and molar mass. The laws that connect these parameters are very close for all gases, and absolutely accurate for ideal gas , in which there is completely no interaction between particles, and whose particles are material points.
The first quantitative studies of reactions between gases belonged to the French scientist Gay-Lussac. He is the author of laws on thermal expansion gases and the law of volumetric relations. These laws were explained in 1811 by the Italian physicist A. Avogadro. Avogadro's Law - one of the important basic principles of chemistry, which states that “ in equal volumes various gases taken at the same temperature and pressure contain the same number of molecules».
Consequences from Avogadro's law:
1) the molecules of most simple atoms are diatomic (H 2, O 2, etc.);
2) same number molecules of different gases under the same conditions occupy the same volume.
3) when normal conditions one mole of any gas occupies a volume equal to 22.4 dm 3 (l). This volume is called molar volume of gas(V o) (normal conditions - t o = 0 °C or
T o = 273 K, P o = 101325 Pa = 101.325 kPa = 760 mm. rt. Art. = 1 atm).
4) one mole of any substance and an atom of any element, regardless of conditions and state of aggregation contains the same number of molecules. This Avogadro's number (Avogadro's constant) - it has been experimentally established that this number is equal to
N A = 6.02213∙10 23 (molecules).
Thus: for gases 1 mol – 22.4 dm 3 (l) – 6.023∙10 23 molecules – M, g/mol;
for substance 1 mol – 6.023∙10 23 molecules – M, g/mol.
Based on Avogadro's law: at the same pressure and the same temperatures, the masses (m) of equal volumes of gases are related as their molar masses (M)
m 1 /m 2 = M 1 /M 2 = D,
where D is the relative density of the first gas relative to the second.
According to law of R. Boyle – E. Mariotte , at a constant temperature, the pressure produced by a given mass of gas is inversely proportional to the volume of the gas
P o /P 1 = V 1 /V o or PV = const.
This means that as pressure increases, the volume of gas decreases. This law was first formulated in 1662 by R. Boyle. Since the French scientist E. Marriott was also involved in its creation, in countries other than England this law is called double name. He is special case ideal gas law(describing a hypothetical gas that ideally obeys all the laws of gas behavior).
By J. Gay-Lussac's law : at constant pressure, the volume of gas changes in direct proportion absolute temperature(T)
V 1 /T 1 = V o /T o or V/T = const.
The relationship between gas volume, pressure and temperature can be expressed general equation, combining the Boyle-Mariotte and Gay-Lussac laws ( combined gas law )
PV/T = P o V o /T o,
where P and V are the pressure and volume of gas at a given temperature T; P o and V o - pressure and volume of gas under normal conditions (n.s.).
Mendeleev-Clapeyron equation(equation of state of an ideal gas) establishes the relationship between the mass (m, kg), temperature (T, K), pressure (P, Pa) and volume (V, m 3) of a gas with its molar mass (M, kg/mol)
where R is the universal gas constant, equal to 8,314 J/(mol K). In addition, the gas constant has two more values: P – mmHg, V – cm 3 (ml), R = 62400 ;
P – atm, V – dm 3 (l), R = 0.082.
Partial pressure(lat. partialis- partial, from lat. pars- part) - the pressure of an individual component of the gas mixture. The total pressure of a gas mixture is the sum of the partial pressures of its components.
The partial pressure of a gas dissolved in a liquid is the partial pressure of the gas that would be formed in the gas formation phase in a state of equilibrium with the liquid at the same temperature. The partial pressure of a gas is measured as the thermodynamic activity of the gas molecules. Gases will always flow from an area of high partial pressure to an area of lower pressure; and the greater the difference, the faster the flow will be. Gases dissolve, diffuse and react according to their partial pressure and are not necessarily dependent on the concentration in the gas mixture. The law of addition of partial pressures was formulated in 1801 by J. Dalton. At the same time, the correct theoretical justification, based on the molecular kinetic theory, was made much later. Dalton's laws - two physical laws that determine the total pressure and solubility of a mixture of gases and are formulated by him early XIX century:
The law on the solubility of the components of a gas mixture: at a constant temperature, the solubility in a given liquid of each of the components of the gas mixture located above the liquid is proportional to their partial pressure
Both Dalton's laws are strictly satisfied for ideal gases. For real gases, these laws are applicable provided that their solubility is low and their behavior is close to that of an ideal gas.
Law of equivalents
The amount of an element or substance that interacts with 1 mole of hydrogen atoms (1 g) or replaces this amount of hydrogen in chemical reactions is called equivalent of a given element or substance(E).
Equivalent mass(M e, g/mol) is the mass of one equivalent of a substance.
The equivalent mass can be calculated from the composition of the compound if the molar masses (M) are known:
1) M e (element): M e = A/B,
where A is the atomic mass of the element, B is the valence of the element;
2) M e (oxide) = M / 2n (O 2) = M e (ele.) + M e (O 2) = M e (element) + 8,
where n(O 2) is the number of oxygen atoms; M e (O 2) = 8 g/mol - equivalent mass of oxygen;
3) Me (hydroxide) = M/n (on-) = Me (element) + Me (OH -) = Me (element) + 17,
where n (he-) is the number of OH - groups; M e (OH -) = 17 g/mol;
4) M e (acids) = M/n (n+) = M e (H +) + M e (acid residue) = 1 + M e (acid residue),
where n (n+) is the number of H + ions; M e (H +) = 1 g/mol; M e (acid residue) – equivalent mass acid residue;
5) Me (salts) = M/n me In me = Me (element) + Me (acid residue),
where n me is the number of metal atoms; In me - the valency of the metal.
When solving some problems containing information about the volumes of gaseous substances, it is advisable to use the value of the equivalent volume (V e).
Equivalent volume is the volume occupied under given conditions
1 equivalent of a gaseous substance. So for hydrogen at no. the equivalent volume is 22.4 1/2 = 11.2 dm 3, for oxygen - 5.6 dm 3.
According to the law of equivalents: the masses (volumes) of substances m 1 and m 2 reacting with each other are proportional to their equivalent masses (volumes)
m 1 /M e1 = m 2 /M e2.
If one of the substances is in a gaseous state, then
m/M e = V o /V e.
If both substances are in gaseous state
V o1 /V e 1 = V o2 /V e2.
Periodic law And
Atomic structure
The periodic law and the periodic system of elements served as a powerful impetus for research into the structure of the atom, which changed the understanding of the laws of the universe and led to practical implementation ideas for using nuclear energy.
By the time the periodic law was discovered, ideas about molecules and atoms had just begun to be established. Moreover, the atom was considered not only the smallest, but also an elementary (that is, indivisible) particle. Direct proof of the complexity of the structure of the atom was the discovery of the spontaneous disintegration of atoms of some elements, called radioactivity. In 1896, the French physicist A. Becquerel discovered that materials containing uranium illuminate a photographic plate in the dark, ionize the gas, and cause fluorescent substances to glow. Later it turned out that not only uranium has this ability. P. Curie and Marie Sklodowska-Curie discovered two new radioactive elements: polonium and radium.
He suggested calling cathode rays discovered by W. Crookes and J. Stoney in 1891 electrons- like elementary particles of electricity. J. Thomson in 1897, studying the flow of electrons, passing it through electric and magnetic fields, established the value of e/m - the ratio of the electron charge to its mass, which led the scientist R. Millikan in 1909 to establish the value of the electron charge q = 4.8∙10 -10 electrostatic units, or 1.602∙10 -19 C (Coulomb), and accordingly to the electron mass –
9.11∙10 -31 kg. Conventionally, the charge of an electron is considered as a unit of negative electric charge and assign it a value (-1). A.G. Stoletov proved that electrons are part of all atoms found in nature. Atoms are electrically neutral, that is, they generally have no electrical charge. This means that atoms must contain positive particles in addition to electrons.
Thomson and Rutherford models
One of the hypotheses about the structure of the atom was put forward in 1903 by J.J. Thomson. He believed that an atom consists of a positive charge, evenly distributed throughout the entire volume of the atom, and electrons oscillating within this charge, like the seeds in a “watermelon” or “raisin pudding.” To test Thomson's hypothesis and more accurately determine internal structure atom in 1909-1911 E. Rutherford, together with G. Geiger (later the inventor of the famous Geiger counter) and students performed original experiments.
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Planetary model of the structure of the atom
The essence of the planetary model of atomic structure can be summarized in the following statements:
1. At the center of the atom there is a positively charged nucleus, occupying an insignificant part of the space inside the atom;
2. All the positive charge and almost all the mass of the atom are concentrated in its nucleus (the mass of the electron is 1/1823 amu);
3. Electrons rotate around the nucleus. Their number is equal to the positive charge of the nucleus.
This model turned out to be very clear and useful for explaining many experimental data, but it immediately revealed its shortcomings. In particular, an electron, moving around a nucleus with acceleration (it is acted upon by a centripetal force), should, according to electromagnetic theory, continuously emit energy. This would cause the electron to spiral around the nucleus and eventually fall onto it. There was no evidence that atoms are continuously disappearing, which means that E. Rutherford’s model is somehow wrong.
Moseley's Law
X-rays were discovered in 1895 and intensively studied in subsequent years; their use for experimental purposes began: they are indispensable for determining the internal structure of crystals and the serial numbers of chemical elements. G. Moseley managed to measure the charge of the atomic nucleus using X-rays. It is in the charge of the nucleus that the main difference between the atomic nuclei of different elements lies. G. Moseley named the charge of the nucleus element serial number. Unit positive charges were later called protons(1 1 r).
X-ray radiation depends on the structure of the atom and is expressed Moseley's law: the square roots of the reciprocal values of the wavelengths are linearly dependent on the serial numbers of the elements. Mathematical expression of Moseley's law:
,
where l is the wavelength of the maximum peak in the X-ray spectrum; a and b are constants that are the same for similar lines of a given series of X-rays. Serial number(Z) is the number of protons in the nucleus. But it was only by 1920 that the name “ proton"and its properties were studied. The charge of a proton is equal in magnitude and opposite in sign to the charge of an electron, that is, 1.602 × 10 -19 C, and conventionally (+1), the mass of a proton is 1.67 × 10 -27 kg, which is approximately 1836 times greater than the mass of an electron . Thus, the mass of a hydrogen atom, consisting of one electron and one proton, practically coincides with the mass of a proton, denoted by 1 1 p.
For all elements, the atomic mass is more than the amount masses of electrons and protons included in their composition. The difference in these values arises due to the presence in atoms of another type of particles called neutrons(1 o n), which were discovered only in 1932 by the English scientist D. Chadwick. Neutrons are almost equal in mass to protons, but lack an electrical charge. The sum of the number of protons and neutrons contained in the nucleus of an atom is called mass number of an atom. The number of protons is equal to the atomic number of the element, the number of neutrons is equal to the difference between the mass number (atomic mass) and the atomic number of the element. The nuclei of all atoms of a given element have the same charge, that is, they contain the same number of protons, but the number of neutrons can be different. Atoms that have the same nuclear charge, and therefore identical properties, but different number neutrons, and therefore different mass numbers are called isotopes ("izos" - equal, "topos" - place ). Each isotope is characterized by two values: the mass number (put down at the top left of the chemical symbol of the element) and the serial number (put down at the bottom left of the chemical sign of the element). For example, an isotope of carbon with a mass number of 12 is written as follows: 12 6 C or 12 C, or in the words: “carbon-12”. Isotopes are known for all chemical elements. Thus, oxygen has isotopes with mass numbers 16, 17, 18: 16 8 O, 17 8 O, 18 8 O. Potassium isotopes: 39 19 K, 40 19 K, 41 19 K. It is the presence of isotopes that explains those rearrangements that in D.I. did his time Mendeleev. Note that he did this only on the basis of the properties of substances, since the structure of atoms was not yet known. Modern science confirmed the rightness of the great Russian scientist. Thus, natural potassium is formed mainly by atoms of its light isotopes, and argon - by heavy ones. Therefore, the relative atomic mass of potassium is less than that of argon, although serial number(atomic nuclear charge) potassium is greater.
The atomic mass of an element is equal to the average value of all its natural isotopes, taking into account their abundance. For example, natural chlorine consists of 75.4% isotope with mass number 35 and 24.6% isotope with mass number 37; the average atomic mass of chlorine is 35.453. Atomic masses of elements given in periodic table
DI. Mendeleev, there are average mass numbers of natural mixtures of isotopes. This is one of the reasons why they are different from integer values.
Stable and unstable isotopes. All isotopes are divided into: stable and radioactive. Stable isotopes are not subject to radioactive decay, which is why they are preserved in natural conditions. Examples of stable isotopes are 16 O, 12 C, 19 F. Most natural elements consist of a mixture of two or more stable isotopes. Of all the elements greatest number Tin has stable isotopes (10 isotopes). In rare cases, such as aluminum or fluorine, only one stable isotope occurs in nature, and the remaining isotopes are unstable.
Radioactive isotopes are, in turn, divided into natural and artificial, both of which spontaneously decay, emitting α- or β-particles until a stable isotope is formed. The chemical properties of all isotopes are basically the same.
Isotopes are widely used in medicine and scientific research. Ionizing radiation capable of destroying living tissue. Malignant tumor tissues are more sensitive to radiation than healthy tissues. This makes it possible to treat cancer with γ-radiation (radiation therapy), which is usually obtained using the radioactive isotope cobalt-60. The radiation is directed to the area of the patient’s body affected by the tumor; the treatment session usually lasts several minutes and is repeated for several weeks. During the session, all other parts of the patient's body must be carefully covered with radiation-impermeable material to prevent the destruction of healthy tissue.
In method labeled atoms Radioactive isotopes are used to trace the “route” of an element in the body. Thus, a patient with a diseased thyroid gland is injected with a drug of radioactive iodine-131, which allows the doctor to monitor the passage of iodine through the patient’s body. Since the half-life
iodine-131 is only 8 days, then its radioactivity quickly decreases.
Particularly interesting is the use of radioactive carbon-14 to determine the age of objects of organic origin based on the radiocarbon method (geochronology), developed by the American physical chemist W. Libby. This method was awarded the Nobel Prize in 1960. When developing his method, W. Libby used known fact formation of the radioactive isotope carbon-14 (in the form of carbon monoxide (IV)) in upper layers earth's atmosphere when nitrogen atoms are bombarded by neutrons that make up cosmic rays
14 7 N + 1 0 n → 14 6 C + 1 1 p
Radioactive carbon-14 in turn decays, emitting beta particles and turning back into nitrogen
14 6 C → 14 7 N + 0 -1 β
Atoms of different elements having the same mass numbers (atomic masses) are called isobars. In the periodic table With There are 59 pairs and 6 triplets of isobars. For example, 40 18 Ar 40 19 K 40 20 Ca.
Atoms of different elements that have the same number of neutrons are called isotones. For example, 136 Ba and 138 Xe - they each have 82 neutrons in the nucleus of the atom.
Periodic law and
Covalent bond
In 1907 N.A. Morozov and later in 1916-1918. Americans J. Lewis and I. Langmuir introduced the concept of education chemical bond by a shared electron pair and proposed to denote valence electrons with dots
A bond formed by electrons belonging to two interacting atoms is called covalent. According to the Morozov-Lewis-Langmuir ideas:
1) when atoms interact between them, shared - common - electron pairs are formed that belong to both atoms;
2) due to common electron pairs, each atom in the molecule acquires eight electrons at the external energy level, s 2 p 6;
3) configuration s 2 p 6 is a stable configuration of an inert gas and in the process of chemical interaction each atom strives to achieve it;
4) the number of common electron pairs determines the covalency of the element in the molecule and is equal to the number of electrons in the atom, missing up to eight;
5) the valence of a free atom is determined by the number of unpaired electrons.
Chemical bonds are depicted in different ways:
1) using electrons in the form of dots placed at the chemical symbol of the element. Then the formation of a hydrogen molecule can be shown by the diagram
Н× + Н× ® Н: Н;
2) using quantum cells (orbitals) as placing two electrons with opposite spins in one molecular quantum cell
The arrangement diagram shows that the molecular energy level is lower than the original atomic levels, which means that the molecular state of the substance is more stable than the atomic one;
3) often, especially in organic chemistry, a covalent bond is represented by a dash (for example, H-H), which symbolizes a pair of electrons.
The covalent bond in the chlorine molecule is also carried out using two shared electrons, or an electron pair.
As you can see, each chlorine atom has three lone pairs and one unpaired electron. The formation of a chemical bond occurs due to the unpaired electrons of each atom. Unpaired electrons bond into a shared pair of electrons, also called shared pair.
Valence bond method
Ideas about the mechanism of chemical bond formation using the example of a hydrogen molecule extend to other molecules. The theory of chemical bonding, created on this basis, was called valence bond method (VBC). Key points:
1) a covalent bond is formed as a result of the overlap of two electron clouds with oppositely directed spins, and the resulting common electron cloud belongs to two atoms;
2) the stronger the covalent bond, the more the interacting electron clouds overlap. The degree to which electron clouds overlap depends on their size and density;
3) the formation of a molecule is accompanied by compression of electron clouds and a decrease in the size of the molecule compared to the size of atoms;
4) s- and p-electrons of the external energy level and d-electrons of the pre-external energy level take part in the formation of the bond.
Sigma (s) and pi (p) bonds
In a chlorine molecule, each of its atoms has a complete outer level of eight electrons s 2 p 6, and two of them (electron pair) belong equally to both atoms. The overlap of electron clouds during the formation of a molecule is shown in the figure.
Scheme of the formation of a chemical bond in molecules of chlorine Cl 2 (a) and hydrogen chloride HCl (b)
A chemical bond for which the line connecting atomic nuclei, is the axis of symmetry of the connecting electron cloud, called sigma (σ)-bond. It occurs when atomic orbitals overlap head-on. Bonds when s-s orbitals overlap in the H 2 molecule; p-p-orbitals in the Cl 2 molecule and s-p-orbitals in the HCl molecule are sigma bonds. “lateral” overlap of atomic orbitals is possible. When overlapping p-electron clouds oriented perpendicular to the bond axis, i.e. along the y- and z-axis, two overlap regions are formed, located on either side of this axis. This covalent bond is called pi (p)-bond. There is less overlap of electron clouds during π bond formation. In addition, the overlap regions lie further from the nuclei than during the formation of a σ bond. Due to these reasons, the π bond has less strength compared to the σ bond. Therefore, the energy of a double bond is less than twice the energy of a single bond, which is always a σ bond. In addition, the σ bond has axial, cylindrical symmetry and is a body of revolution around the line connecting the atomic nuclei. The π bond, on the contrary, does not have cylindrical symmetry.
A single bond is always a pure or hybrid σ bond. A double bond consists of one σ- and one π-bond, located perpendicular to each other. The σ bond is stronger than the π bond. In compounds with multiple bonds, there is always one σ bond and one or two π bonds.
Donor-acceptor bond
Another mechanism for the formation of a covalent bond is also possible - donor-acceptor. In this case, a chemical bond occurs due to the two-electron cloud of one atom and the free orbital of another atom. Let us consider as an example the mechanism of formation of ammonium ion (NH 4 +). In an ammonia molecule, the nitrogen atom has a lone pair of electrons (two-electron cloud)
The hydrogen ion has a free (not filled) 1s orbital, which can be denoted as H + (here the square means a cell). When an ammonium ion is formed, the two-electron cloud of nitrogen becomes common to the nitrogen and hydrogen atoms, that is, it turns into a molecular electron cloud. This means that a fourth covalent bond appears. The process of formation of ammonium ion can be represented by the diagram
The charge of the hydrogen ion becomes common (it is delocalized, i.e., dispersed between all atoms), and the two-electron cloud (lone electron pair) belonging to nitrogen becomes common with H +. In diagrams, the image of cell is often omitted.
An atom that provides a lone pair of electrons is called donor , and the atom that accepts it (that is, provides a free orbital) is called acceptor .
The mechanism of formation of a covalent bond due to the two-electron cloud of one atom (donor) and the free orbital of another atom (acceptor) is called donor-acceptor. The covalent bond formed in this way is called a donor-acceptor or coordination bond.
However, this is not a special type of bond, but only a different mechanism (method) for the formation of a covalent bond. According to the properties, quarter N-H connection in the ammonium ion is no different from the other three.
For the most part, donors are molecules containing N, O, F, Cl atoms associated with atoms of other elements. An acceptor can be a particle that has vacant electronic levels, for example, atoms of d-elements that have unfilled d-sublevels.
Properties of covalent bonds
Link length is the internuclear distance. The shorter the length of a chemical bond, the stronger it is. The bond length in molecules is: HC 3 -CH 3 1.54 ; H 2 C=CH 2
1,33 ; NS≡CH 1.20 .In terms of single bonds, these values increase, and the reactivity of compounds with multiple bonds increases. A measure of bond strength is the bond energy.
Communication energy determined by the amount of energy required to break the bond. It is usually measured in kilojoules per 1 mole of substance. As the bond multiplicity increases, the bond energy increases and its length decreases. Bond energy values in compounds (alkanes, alkenes, alkynes): C-C 344 kJ/mol; C=C 615 kJ/mol; С≡С 812 kJ/mol. That is, the energy of a double bond is less than twice the energy of a single bond, and the energy of a triple bond is less than three times the energy of a single bond, so alkynes are the more reactive of this group of hydrocarbons.
Under saturation understand the ability of atoms to form a limited number of covalent bonds. For example, a hydrogen atom (one unpaired electron) forms one bond, a carbon atom (four unpaired electrons in an excited state) forms no more than four bonds. Due to the saturation of the bonds, the molecules have a certain composition: H 2, CH 4, HCl, etc. However, even with saturated covalent bonds, more complex molecules can be formed by the donor-acceptor mechanism.
Focus covalent bonds determine the spatial structure of molecules, that is, their shape. Let's consider this using the example of the formation of molecules HCl, H 2 O, NH 3.
According to the MBC, a covalent bond occurs in the direction of maximum overlap of the electron orbitals of interacting atoms. When an HCl molecule is formed, the s-orbital of the hydrogen atom overlaps with the p-orbital of the chlorine atom. Molecules of this type have a linear shape.
On external level The oxygen atom has two unpaired electrons. Their orbitals are mutually perpedicular, i.e. are located relative to each other at an angle of 90°. When a water molecule is formed
In chemical literature, a situation has historically developed when the masses of atoms and molecules are expressed through the concepts of atomic weight and molecular weight.
As is known, if a body with mass m (as well as an atom or molecule) moves under the influence of the Earth’s gravity with acceleration g, then the force of gravity of this body is equal, i.e., the force of gravity is proportional to the mass of the body on which it acts.
If the body is at rest, then the weight of the body is equal to the force of gravity acting on it, and in the formula we can consider P as the weight of the body. Consequently, for bodies at rest, their weights are proportional to their masses. However, the acceleration at various points earth's surface different, therefore the weight of the same body (atom, molecule!) will be different here. The body's weight will also decrease as it rises above the Earth's surface.
In conclusion, let us ask ourselves the question: Is the weight of a body (and, accordingly, the weight of an atom, molecule) the same on Earth, on a space orbital station, on the surface of the Moon?
If necessary, you can repeat it physical concepts“weight”, “mass”, etc..
Relative units to express the weight of atoms were first used by Dalton, who defined atomic weight as a number showing how many times an atom of an element is heavier than an atom of another element. As a unit of atomic weights, he proposed the weight of the lightest atom - hydrogen.
More correctly, as shown above, we need to talk about a unit of atomic or molecular masses, therefore, in further presentation, the authors tried to use these concepts everywhere instead of “atomic weight”, “molecular weight”.
Since the atomic masses of elements were calculated from experimental data on the weight ratios in various compounds, and oxygen forms compounds with a much larger number of elements compared to hydrogen, then in subsequent years, until 1961, a fraction of the mass of an atom was adopted as a unit of atomic mass oxygen. This relative unit of measure of the mass of atoms was called the oxygen unit (o.u.).
However, by 1930 it was discovered that in addition to oxygen atoms with a mass of 16 k.u., there are isotopes of oxygen that differ in mass (0.039%) and (0.204%). The chemical properties of oxygen isotopes are the same, but the physical properties, although not very much, differ, therefore the isotopic composition of oxygen in different natural compounds is not the same. For example, the average atomic mass of atmospheric oxygen is 0.00011 atomic units less than the average atomic mass of oxygen from seawater.
Physical and chemical system units of atomic mass. Physicists took part of the mass of an isotope as a unit of atomic mass, while chemists took part of the average mass of an oxygen atom of natural isotopic composition. This led to different values of atomic masses and made it difficult to compare physical and chemical atomic masses, which was ultimately the main reason for the abandonment of the oxygen atomic unit.
In 1961, the International Union of Pure and Applied Chemistry decided to choose a standard unit of atomic mass and move to a unified atomic mass scale. The carbon unit (cu) was chosen as the new standard unit of atomic mass - part of the mass of the carbon isotope. Atomic masses based on the new unit (cu) are equal to the old ones (cu) multiplied by 0.99996, so that changes in previous atomic masses are so small, and this should be especially emphasized, that they do not affect almost all chemical calculations.
Thus, the mass of an atom expressed in carbon units is called atomic mass. Atomic mass shows how many times the mass of an atom of a given element is heavier than the mass of a C10 carbon atom. The mass of molecules is also expressed in carbon units (cu).
The molecular mass of a substance is the mass of its molecule, expressed in carbon units. Molecular mass shows how many times the mass of a molecule of a given substance is heavier than the mass of carbon C12. Therefore, both atomic and molecular masses are relative units of measurement. When writing, they usually do not indicate the dimensions of atomic and molecular masses, remembering that they are expressed in carbon units.
For quantitative calculations, it is convenient to use the following characteristics - gram-atom and gram-molecule.
A gram atom is the number of grams of a substance that is numerically equal to the atomic mass of that element. For example, the atomic mass of sodium is 23 cu. i.e., therefore, G-sodium atom has a mass of 23 g.
The number of grams of a substance, numerically equal to its molecular weight, is called a gram molecule of this substance, or mole. For example, the molecular weight of potassium permanganate is 158 c.u. e., therefore, constitute 1 gram molecule.
The concepts of atomic and molecular masses are fundamentally different from the concepts of gram-atomic and gram-molecular masses. If the values of atomic and molecular masses are relative numbers and show how many times the mass of an atom or molecule is larger than part of an atom of a carbon isotope, then gram-atom and gram-molecule are absolute numbers, showing the number of grams of a substance.
After the discovery of Avogadro's law (see § 5, Chapter IV) "it was proven that the number of molecules (atoms) contained in one gram-molecule (gram-atom) of any substance is the same and equal (Avogadro's number), i.e. . and the mass of a gram molecule is equal to the mass of molecules of a given substance. It is worth emphasizing that molecules (atoms) are contained in 1 mole
(1 g-atom) - any substance in any state of aggregation - solid, liquid, gaseous.
One of the most important characteristics atoms is his weight.
Absolute mass is the mass of an atom, expressed in kilograms (grams).
Absolute atomic mass ( m a volume) value is extremely small. Thus, an atom of the light isotope of hydrogen (protium) has a mass of 1.66 × 10 –27 kg.
m(N) = 1.66 10 –27 kg, m(H) = 1.66 10 –24 g,
an atom of one of the oxygen isotopes has a mass of 2.67 10 –26 kg,
m(O) = 2.67 10 –26 kg, m(ABOUT) = 2.67 10 –23 g,
an atom of the carbon isotope 12 C has a mass of 1.99 10 –26 kg,
m(C) = 1.99 10 –26 kg, m(C) = 1.99 10 –23 g.
In practical calculations it is extremely inconvenient to use such quantities. Therefore, they usually use not the absolute masses of atoms, but the values relative atomic masses.
Relative atomic mass is denoted Ar, index r – the initial letter of the English word relative, which means relative.
The unit used to measure the masses of atoms and molecules is atomic mass unit (a.m.u.).
An atomic mass unit (amu) is 1/12 of the mass of an atom of the carbon isotope 12 C, i.e.
a.e.m. = =
· 1.99 · 10 –26 kg = 
· 1.99 · 10 –23 g.
Relative atomic mass shows how many times the mass of an atom of a given element is greater than 1/12 of the mass of an atom of the carbon isotope 12 C, i.e., an atomic mass unit.
Relative atomic mass is a dimensionless quantity, but it is possible to designate its value in atomic mass units (amu). For example:
Thus, the relative atomic mass value of the element hydrogen is 1.001 or, in round numbers,
Аr(Н) ≈ 1 amu, and oxygen – Аr(O) = 15.999 ≈ 16 amu.
The values of the relative atomic masses of elements are given in the periodic system of D.I. Mendeleev. These values represent the average value of the mass of an atom of any element, taking into account the isotopes of this element existing in nature and their quantity. For ordinary calculations, rounded values of the relative atomic masses of elements should be used. (see table 4 of the appendix).
Similar to the concepts of absolute atomic mass and relative atomic mass, we can formulate the concepts absolute molecular mass and relative molecular mass.
Absolute molecular mass(m) mol. – molecular mass chemical substance, expressed in kilograms (grams).
Relative molecular weight(Mr) (or just molecular weight) – the mass of a molecule, expressed in atomic mass units.
Knowing the chemical formula of a compound, you can easily determine the value of its molecular mass, which is defined as the sum of the atomic masses of all elements that make up the molecule of the substance.
For example, the relative molecular mass of sulfuric acid Mr(H 2 SO 4) will be the sum of two relative atomic masses of the element hydrogen, one relative atomic mass of the element sulfur and four relative atomic masses of the element oxygen:
Mr(H 2 SO 4) = 2Аr (H) + Аr (S) + 4Аr(O) = 2 1 + 32 + 4 16 = 98.
Thus, the molecular weight of sulfuric acid is 98 or 98 amu.
Molecular weight (relative molecular weight) shows how many times the mass of a molecule of a given substance is greater than 1/12 of the mass of a 12 C carbon atom.
In the above example, the molecular weight of sulfuric acid is 98 amu, that is, a sulfuric acid molecule has a mass 98 times greater than 1/12 of the mass of a 12 C carbon atom .
Atomic-molecular theory. Atom, molecule. Chemical element. Simple and complex substance. Allotropy.
Chemistry- the science of substances, the laws of their transformations (physical and chemical properties) and application. Currently, more than 100 thousand inorganic and more than 4 million organic compounds are known.
Chemical phenomena: Some substances are transformed into others that differ from the original ones in composition and properties, while the composition of the atomic nuclei does not change.
Physical phenomena: the physical state of substances changes (vaporization, melting, electrical conductivity, release of heat and light, malleability, etc.) or new substances are formed with a change in the composition of atomic nuclei.
1. All substances are made up of molecules. Molecule- the smallest particle of a substance that has its chemical properties.
2. Molecules are made up of atoms. Atom- the smallest particle of a chemical element that retains all its chemical properties. Various elements different atoms correspond.
3. Molecules and atoms are in continuous motion; there are forces of attraction and repulsion between them.
Chemical element- this is a type of atom characterized by certain nuclear charges and structure electron shells. Currently, 117 elements are known: 89 of them are found in nature (on Earth), the rest are obtained artificially. Atoms exist in a free state, in compounds with atoms of the same or other elements, forming molecules. The ability of atoms to interact with other atoms and form chemical compounds determined by its structure. Atoms consist of a positively charged nucleus and negatively charged electrons moving around it, forming an electrically neutral system that obeys the laws characteristic of microsystems.
Chemical formula - this is a conventional notation of the composition of a substance using chemical symbols (proposed in 1814 by J. Berzelius) and indices (index is the number at the bottom right of the symbol. Indicates the number of atoms in the molecule). The chemical formula shows which atoms of which elements and in what ratio are connected to each other in a molecule.
Allotropy- the phenomenon of the formation by a chemical element of several simple substances that differ in structure and properties.
Simple substances- molecules consist of atoms of the same element.
Complex substances- molecules consist of atoms of various chemical elements.
The international unit of atomic mass is equal to 1/12 of the mass of the 12 C isotope - the main isotope of natural carbon: 1 amu = 1/12 m (12 C) = 1.66057 10 -24 g
Relative atomic mass (Ar)- a dimensionless quantity equal to the ratio of the average mass of an atom of an element (taking into account the percentage of isotopes in nature) to 1/12 of the mass of a 12 C atom.
Average absolute atomic mass (m) equal to the relative atomic mass times the amu. (1 amu=1.66*10 -24)
Relative molecular weight (Mr)- a dimensionless quantity showing how many times the mass of a molecule of a given substance is greater than 1/12 the mass of a carbon atom 12 C.
Mr = mr / (1/12 mа(12 C))
mr is the mass of a molecule of a given substance;
ma(12 C) - mass of carbon atom 12 C.
Mr = S Ar(e). The relative molecular mass of a substance is equal to the sum of the relative atomic masses of all elements, taking into account formula indices.
The absolute mass of a molecule is equal to the relative molecular mass times the amu. The number of atoms and molecules in ordinary samples of substances is very large, therefore, when characterizing the amount of a substance, a special unit of measurement is used - mole.
Amount of substance, mol. Means a certain number of structural elements (molecules, atoms, ions). Denoted n and measured in moles. A mole is the amount of a substance containing as many particles as there are atoms in 12 g of carbon.
Avogadro di Quaregna number(N A). The number of particles in 1 mole of any substance is the same and equals 6.02 10 23. (Avogadro's constant has the dimension - mol -1).
Molar mass shows the mass of 1 mole of a substance (denoted by M): M = m/n
The molar mass of a substance is equal to the ratio of the mass of the substance to the corresponding amount of the substance.
The molar mass of a substance is numerically equal to its relative molecular mass, however, the first quantity has the dimension g/mol, and the second is dimensionless: M = N A m(1 molecule) = N A Mr 1 amu. = (N A 1 amu) Mr = Mr
Equivalent- is a real or conditional particle of a substance that is equivalent to:
a) one H + or OH - ion in a given acid-base reaction;
b) one electron in a given ORR (redox reaction);
c) one unit of charge in a given exchange reactions,
d) the number of monodentate ligands participating in the complex formation reaction.