A substance can be in three states of aggregation: solid, liquid and gaseous. Molecular physics is a branch of physics that studies the physical properties of bodies in various states of aggregation based on their molecular structure.
Thermal movement- random (chaotic) movement of atoms or molecules of a substance.
FUNDAMENTALS OF MOLECULAR KINETIC THEORY
Molecular kinetic theory is a theory that explains thermal phenomena in macroscopic bodies and the properties of these bodies based on their molecular structure.
Basic principles of molecular kinetic theory:
- matter consists of particles - molecules and atoms, separated by spaces,
- these particles move chaotically,
- particles interact with each other.
MASS AND SIZES OF MOLECULES
The masses of molecules and atoms are very small. For example, the mass of one hydrogen molecule is approximately 3.34 * 10 -27 kg, oxygen - 5.32 * 10 -26 kg. Mass of one carbon atom m 0C =1.995*10 -26 kg
Relative molecular (or atomic) mass of a substance Mr is the ratio of the mass of a molecule (or atom) of a given substance to 1/12 of the mass of a carbon atom: (atomic mass unit).
The amount of a substance is the ratio of the number of molecules N in a given body to the number of atoms in 0.012 kg of carbon N A:
Mole- the amount of a substance containing as many molecules as there are atoms in 0.012 kg of carbon.
The number of molecules or atoms in 1 mole of a substance is called Avogadro's constant:
Molar mass- mass of 1 mole of substance:
The molar and relative molecular mass of a substance are related by the relationship: M = M r * 10 -3 kg/mol.
SPEED OF MOLECULES
Despite the random nature of the movement of molecules, their distribution of velocities has the character of a certain pattern, which called Maxwell's distribution.
The graph characterizing this distribution is called the Maxwell distribution curve. It shows that in a system of molecules at a given temperature there are very fast and very slow, but most of the molecules move at a certain speed, which is called the most probable. As the temperature increases, this most likely rate increases.
IDEAL GAS IN MOLECULAR KINETIC THEORY
Ideal gas is a simplified gas model in which:
- gas molecules are considered material points,
- molecules do not interact with each other
- molecules colliding with obstacles experience elastic interactions.
In other words, the movement of individual molecules of an ideal gas obeys the laws of mechanics. Real gases behave like ideal gases at sufficiently high rarefaction, when the distances between molecules are many times larger than their sizes.
The basic equation of molecular kinetic theory can be written as
Speed called the mean square speed.
TEMPERATURE
Any macroscopic body or group of macroscopic bodies is called thermodynamic system.
Thermal or thermodynamic equilibrium- a state of a thermodynamic system in which all its macroscopic parameters remain unchanged: volume, pressure do not change, heat exchange does not occur, there are no transitions from one state of aggregation to another, etc. Under constant external conditions, any thermodynamic system spontaneously goes into a state of thermal equilibrium.
Temperature- a physical quantity characterizing the state of thermal equilibrium of a system of bodies: all bodies of the system that are in thermal equilibrium with each other have the same temperature.
Absolute zero temperature- the limiting temperature at which the pressure of an ideal gas at constant volume must be equal to zero or the volume of an ideal gas at constant pressure must be equal to zero.
Thermometer- a device for measuring temperature. Typically, thermometers are calibrated on the Celsius scale: the crystallization temperature of water (ice melting) corresponds to 0°C, its boiling point - 100°C.
Kelvin introduced the absolute temperature scale, according to which zero temperature corresponds to absolute zero, the unit of temperature on the Kelvin scale is equal to the degree Celsius: [T] = 1 K(Kelvin).
Relationship between temperature in energy units and temperature in Kelvin:
Where k= 1.38*10 -23 J/K - Boltzmann's constant.
Relationship between the absolute scale and the Celsius scale:
T = t + 273
Where t- temperature in degrees Celsius.
The average kinetic energy of the chaotic movement of gas molecules is proportional to the absolute temperature:
Mean square speed of molecules
Taking into account equality (1), the basic equation of molecular kinetic theory can be written as follows:
EQUATION OF STATE OF AN IDEAL GAS
Let a gas of mass m occupy a volume V at a temperature T and pressure R, A M- molar mass of the gas. By definition, the concentration of gas molecules is: n = N/V, Where N-number of molecules.
Let's substitute this expression into the basic equation of molecular kinetic theory:
Size R is called the universal gas constant, and the equation written in the form
called the ideal gas equation of state or the Mendeleev-Clapeyron equation. Normal conditions - gas pressure is equal to atmospheric ( R= 101.325 kPa) at ice melting temperature ( T = 273,15TO).
1. Isothermal process
The process of changing the state of a thermodynamic system at a constant temperature is called isothermal.
If T =const, then
Boyle-Mariotte Law
For a given mass of gas, the product of the gas pressure and its volume is constant if the gas temperature does not change: p 1 V 1 =p 2 V 2 at T = const
A graph of a process occurring at a constant temperature is called an isotherm.
2. Isobaric process
The process of changing the state of a thermodynamic system at constant pressure is called isobaric.
Gay-Lussac's Law
The volume of a given mass of gas at constant pressure is directly proportional to the absolute temperature:
If a gas, having a volume V 0, is under normal conditions: and then, at constant pressure, goes into a state with temperature T and volume V, then we can write
Having designated
we get V=V 0 T
The coefficient is called the temperature coefficient of volumetric expansion of gases. The graph of a process occurring at constant pressure is called isobar.
3.Isochoric process
The process of changing the state of a thermodynamic system at a constant volume is called isochoric. If V = const, That
Charles's Law
The pressure of a given mass of gas at constant volume is directly proportional to the absolute temperature:
If a gas, having a volume V 0, is under normal conditions:
and then, maintaining volume, goes into a state with temperature T and pressure R, then we can write
The graph of a process occurring at constant volume is called isochore.
Example. What is the pressure of compressed air in a 20 liter cylinder at 12°C if the mass of this air is 2 kg?
From the equation of state of an ideal gas
Let's determine the pressure value.
All substances in nature are made up of very small particles called molecules. These particles in matter constantly interact with each other. They cannot be seen with the naked eye. We will consider the concept, basic properties and characteristics of molecules in the article.
Molecules are particles that have a neutral electrical charge and consist of varying numbers of atoms. Their number, as a rule, is always more than two, and these atoms are connected to each other by a covalent bond. For the first time the existence of molecules became known in France. For this we must give credit to the physicist Jean Perrin, who made this great discovery in 1906. The composition of the molecule is constant. She does not change it throughout her entire existence. The structure of this small particle depends on the physical properties of the substance it forms.
We hope that from this article you received a lot of useful and interesting information about molecules. Now you know exactly what kind of particle it is, and have an idea of its composition, basic properties and how scientists in the field of chemistry study molecules.
Many experiments show that molecular size very small. The linear size of a molecule or atom can be found in various ways. For example, using an electron microscope, photographs of some large molecules are obtained, and using an ion projector (ion microscope) you can not only study the structure of crystals, but determine the distance between individual atoms in a molecule.
Using the achievements of modern experimental technology, it was possible to determine the linear dimensions of simple atoms and molecules, which are about 10-8 cm. The linear dimensions of complex atoms and molecules are much larger. For example, the size of a protein molecule is 43 * 10 -8 cm.
To characterize atoms, the concept of atomic radii is used, which makes it possible to approximately estimate interatomic distances in molecules, liquids or solids, since atoms do not have clear boundaries in size. That is atomic radius- this is the sphere in which the bulk of the electron density of the atom is contained (at least 90...95%).
The size of the molecule is so small that it can only be imagined using comparisons. For example, a water molecule is as many times smaller than a large apple as the apple is smaller than the globe.
Mole of substance
The masses of individual molecules and atoms are very small, so in calculations it is more convenient to use relative rather than absolute mass values.
Relative molecular weight(or relative atomic mass) of a substance M r is the ratio of the mass of a molecule (or atom) of a given substance to 1/12 of the mass of a carbon atom.
M r = (m 0) : (m 0C / 12)
where m 0 is the mass of a molecule (or atom) of a given substance, m 0C is the mass of a carbon atom.
The relative molecular (or atomic) mass of a substance shows how many times the mass of a molecule of a substance is greater than 1/12 of the mass of the carbon isotope C12. Relative molecular (atomic) mass is expressed in atomic mass units.
Atomic mass unit– this is 1/12 of the mass of the carbon isotope C12. Accurate measurements showed that the atomic mass unit is 1.660 * 10 -27 kg, that is
1 amu = 1.660 * 10 -27 kg
The relative molecular mass of a substance can be calculated by adding the relative atomic masses of the elements that make up the substance's molecule. The relative atomic mass of chemical elements is indicated in the periodic table of chemical elements by D.I. Mendeleev.
In the periodic system D.I. Mendeleev for each element is indicated atomic mass, which is measured in atomic mass units (amu). For example, the atomic mass of magnesium is 24.305 amu, that is, magnesium is twice as heavy as carbon, since the atomic mass of carbon is 12 amu. (this follows from the fact that 1 amu = 1/12 the mass of the carbon isotope, which makes up the majority of the carbon atom).
Why measure the mass of molecules and atoms in amu if there are grams and kilograms? Of course, you can use these units of measurement, but it will be very inconvenient for writing (too many numbers will have to be used in order to write down the mass). To find the mass of an element in kilograms, you need to multiply the atomic mass of the element by 1 amu. Atomic mass is found according to the periodic table (written to the right of the letter designation of the element). For example, the weight of a magnesium atom in kilograms would be:
m 0Mg = 24.305 * 1 a.u.m. = 24.305 * 1.660 * 10 -27 = 40.3463 * 10 -27 kg
The mass of a molecule can be calculated by adding the masses of the elements that make up the molecule. For example, the mass of a water molecule (H 2 O) will be equal to:
m 0H2O = 2 * m 0H + m 0O = 2 * 1.00794 + 15.9994 = 18.0153 a.m. = 29.905 * 10 -27 kg
Mole equal to the amount of substance in a system that contains the same number of molecules as there are atoms in 0.012 kg of carbon C 12. That is, if we have a system with any substance, and in this system there are as many molecules of this substance as there are atoms in 0.012 kg of carbon, then we can say that in this system we have 1 mole of substance.
Avogadro's constant
Quantity of substanceν is equal to the ratio of the number of molecules in a given body to the number of atoms in 0.012 kg of carbon, that is, the number of molecules in 1 mole of a substance.
ν = N / N A
where N is the number of molecules in a given body, N A is the number of molecules in 1 mole of the substance of which the body consists.
N A is Avogadro's constant. The amount of a substance is measured in moles.
Avogadro's constant is the number of molecules or atoms in 1 mole of a substance. This constant was named after the Italian chemist and physicist Amedeo Avogadro (1776 – 1856).
1 mole of any substance contains the same number of particles.
N A = 6.02 * 10 23 mol -1
Molar mass is the mass of a substance taken in the amount of one mole:
μ = m 0 * N A
where m 0 is the mass of the molecule.
Molar mass is expressed in kilograms per mole (kg/mol = kg*mol -1).
Molar mass is related to relative molecular mass by:
μ = 10 -3 * M r [kg*mol -1 ]
The mass of any quantity of substance m is equal to the product of the mass of one molecule m 0 by the number of molecules:
m = m 0 N = m 0 N A ν = μν
The amount of a substance is equal to the ratio of the mass of the substance to its molar mass:
ν = m/μ
The mass of one molecule of a substance can be found if the molar mass and Avogadro's constant are known:
m 0 = m / N = m / νN A = μ / N A
A more accurate determination of the mass of atoms and molecules is achieved by using a mass spectrometer - a device in which a beam of charged particles is separated in space depending on their charge mass using electric and magnetic fields.
For example, let's find the molar mass of a magnesium atom. As we found out above, the mass of a magnesium atom is m0Mg = 40.3463 * 10 -27 kg. Then the molar mass will be:
μ = m 0Mg * N A = 40.3463 * 10 -27 * 6.02 * 10 23 = 2.4288 * 10 -2 kg/mol
That is, 2.4288 * 10 -2 kg of magnesium “fits” in one mole. Well, or about 24.28 grams.
As we can see, the molar mass (in grams) is almost equal to the atomic mass indicated for the element in the periodic table. Therefore, when indicating the atomic mass, they usually do this:
The atomic mass of magnesium is 24.305 amu. (g/mol).
The content of the article
MOLECULE STRUCTURE(molecular structure), the relative arrangement of atoms in molecules. During chemical reactions, atoms in the molecules of the reactants are rearranged and new compounds are formed. Therefore, one of the fundamental chemical problems is to clarify the arrangement of atoms in the original compounds and the nature of the changes during the formation of other compounds from them.
The first ideas about the structure of molecules were based on an analysis of the chemical behavior of a substance. These ideas became more complex as knowledge about the chemical properties of substances accumulated. The application of the basic laws of chemistry made it possible to determine the number and type of atoms that make up the molecule of a given compound; this information is contained in the chemical formula. Over time, chemists realized that a single chemical formula is not enough to accurately characterize a molecule, since there are isomer molecules that have the same chemical formulas but different properties. This fact led scientists to believe that the atoms in a molecule must have a certain topology, stabilized by the bonds between them. This idea was first expressed in 1858 by the German chemist F. Kekule. According to his ideas, a molecule can be depicted using a structural formula, which indicates not only the atoms themselves, but also the connections between them. Interatomic bonds must also correspond to the spatial arrangement of atoms. The stages of development of ideas about the structure of the methane molecule are shown in Fig. 1. The structure corresponds to modern data G: the molecule has the shape of a regular tetrahedron, with a carbon atom in the center and hydrogen atoms at the vertices.
Such studies, however, did not say anything about the size of the molecules. This information became available only with the development of appropriate physical methods. The most important of these turned out to be X-ray diffraction. From X-ray scattering patterns on crystals, it became possible to determine the exact position of atoms in a crystal, and for molecular crystals it was possible to localize atoms in an individual molecule. Other methods include diffraction of electrons as they pass through gases or vapors and analysis of the rotational spectra of molecules.
All this information gives only a general idea of the structure of the molecule. The nature of chemical bonds allows us to study modern quantum theory. And although the molecular structure cannot yet be calculated with sufficiently high accuracy, all known data on chemical bonds can be explained. The existence of new types of chemical bonds has even been predicted.
Simple covalent bond.
The hydrogen molecule H2 consists of two identical atoms. According to physical measurements, the bond length - the distance between the nuclei of hydrogen atoms (protons) - is 0.70 Å (1 Å = 10 -8 cm), which corresponds to the radius of the hydrogen atom in the ground state, i.e. in a state of minimal energy. The formation of bonds between atoms can only be explained on the assumption that their electrons are localized mainly between the nuclei, forming a cloud of negatively charged bonding particles and holding together positively charged protons.
Let us consider two hydrogen atoms in the ground state, i.e. state in which their electrons are at 1 s-orbitals. Each of these electrons can be thought of as a wave, and the orbital as a standing wave. As the atoms approach each other, the orbitals begin to overlap (Fig. 2), and, as in the case of ordinary waves, interference occurs - the superposition of waves (wave functions) in the overlap region. If the signs of the wave functions are opposite, then during interference the waves destroy each other (destructive interference), and if they are the same, then they add up (constructive interference). When hydrogen atoms come together, two outcomes are possible, depending on whether the wave functions are in phase (Fig. 2, A) or in antiphase (Fig. 2, b). In the first case, constructive interference will occur, in the second - destructive interference, and two molecular orbitals will appear; one of them is characterized by high density in the region between the nuclei (Fig. 2, V), for the other – low (Fig. 2, G) is actually a node with zero amplitude separating the nuclei.
Thus, when hydrogen atoms come closer and interact 1 s-orbitals form two molecular orbitals, and two electrons must fill one of them. Electrons in atoms always strive to occupy the most stable position - the one in which their energy is minimal. For the orbital shown in Fig. 2, V, there is a high density in the region between the nuclei, and each electron that occupies this orbital will most of the time be located near positively charged nuclei, i.e. its potential energy will be small. On the contrary, the orbital shown in Fig. 2, G, the maximum density occurs in the regions located to the left and right of the nuclei, and the energy of the electrons located in this orbital will be high. So electrons have less energy when they occupy an orbital V, and this energy is even less than what they would have if the atoms were infinitely distant from each other. Since there are only two electrons in this case, both of them can occupy a more energetically favorable orbital if their spins are antiparallel (Pauli principle). Therefore, the energy of a system consisting of two hydrogen atoms decreases as the atoms approach each other, and in order to then remove the atoms from each other, energy will be required equal to the energy of formation of a stable hydrogen molecule H2. Note that a necessary condition for the existence of a hydrogen molecule is the preferential localization of electrons between nuclei in accordance with what we have already said above. Molecular orbital V is called a bonding orbital, and the orbital G- loosening.
Let us now consider the approach of two helium atoms (atomic number 2). Here too there is overlap 1 s-orbitals leads to the formation of two molecular orbitals, one of which corresponds to a lower and the other to a higher energy. This time, however, 4 electrons must be placed in the orbitals, 2 electrons from each helium atom. The low-energy bonding orbital can only be filled by two of them, the other two must occupy the high-energy orbital G. The decrease in energy due to the favorable location of the first pair is approximately equal to the increase in energy due to the unfavorable location of the second pair. Now bringing the atoms closer together does not provide any gain in energy, and molecular helium He 2 is not formed. This can be conveniently illustrated using a diagram (Fig. 3); the different orbitals on it are represented as energy levels in which electrons can reside. The latter are indicated by arrows pointing up and down to distinguish the direction of the spins. Two electrons can occupy the same orbital only if their spins are antiparallel.
These general principles are followed in the formation of molecules from atoms. As soon as two atoms get so close that their atomic orbitals (AO) begin to overlap, two molecular orbitals (MO) appear: one bonding, the other antibonding. If each AO has only one electron, both of them can occupy a bonding MO with lower energy than the AO and form a chemical bond. Bonds of this type, now called covalent, have long been known to chemists (the idea of a covalent bond formed the basis of the octet theory of bonding, formulated by the American physical chemist G. Lewis in 1916). Their formation was explained by the sharing of a pair of electrons by interacting atoms. According to modern concepts, the bond strength depends on the degree of overlap of the corresponding orbitals. All of the above suggests that bonds between atoms can be formed by sharing not only two, but also one or three electrons. However, they will be weaker than ordinary covalent bonds for the following reasons. When a one-electron bond is formed, the energy of only one electron decreases, and in the case of a bond formed as a result of the sharing of three electrons, the energy of two of them decreases, and the third, on the contrary, increases, compensating for the decrease in the energy of one of the first two electrons. As a result, the resulting three-electron bond turns out to be twice as weak as an ordinary covalent bond.
The sharing of one and three electrons occurs during the formation of the molecular hydrogen ion H 2 + and the HHe molecule, respectively. In general, bonds of this type are rare, and the corresponding molecules are highly reactive.
Valence. Donor-acceptor bonds.
All of the above assumes that atoms can form as many covalent bonds as their orbitals are occupied by one electron, but this is not always the case. [In the accepted scheme for filling an AO, the number of the shell is first indicated, then the type of orbital, and then, if there is more than one electron in the orbital, their number (superscript). So, record (2 s) 2 means that on s-orbitals of the second shell contain two electrons.] A carbon atom in the ground state (3 R) has an electronic configuration (1 s) 2 (2s) 2 (2p x)(2 p y), while two orbitals are not filled, i.e. contain one electron each. However, divalent carbon compounds are very rare and are highly reactive. Typically, carbon is tetravalent, and this is due to the fact that for its transition to excited 5 S-state (1 s) 2 (2s) (2p x)(2 p y)(2 p z) With four unfilled orbitals, very little energy is needed. Energy costs associated with transition 2 s-electron to free 2 R-orbital, are more than compensated by the energy released during the formation of two additional bonds. For the formation of unfilled AOs, it is necessary that this process be energetically favorable. Nitrogen atom with electronic configuration (1 s) 2 (2s) 2 (2p x)(2 p y)(2 p z) does not form pentavalent compounds, since the energy required for the transfer of 2 s-electron for 3 d-orbital to form a pentavalent configuration (1 s) 2 (2s)(2p x)(2 p y)(2 p z)(3 d), is too big. Similarly, boron atoms with the usual configuration (1 s) 2 (2s) 2 (2p) can form trivalent compounds when in an excited state (1 s) 2 (2s)(2p x)(2 p y), which occurs during transition 2 s-electron for 2 R-AO, but does not form pentavalent compounds, since the transition to the excited state (1 s)(2s)(2p x)(2 p y)(2 p z), due to the transfer of one of 1 s-electrons to a higher level requires too much energy. The interaction of atoms with the formation of a bond between them occurs only in the presence of orbitals with close energies, i.e. orbitals with the same principal quantum number. The relevant data for the first 10 elements of the periodic table are summarized below. The valence state of an atom is the state in which it forms chemical bonds, for example state 5 S for tetravalent carbon.
| VALENCE STATES AND VALENCES THE FIRST TEN ELEMENTS OF THE PERIODIC TABLE |
|||
| Element | Ground state | Normal valence state | Regular valency |
| H | (1s) | (1s) | 1 |
| He | (1s) 2 | (1s) 2 | 0 |
| Li | (1s) 2 (2s) | (1s) 2 (2s) | 1 |
| Be | (1s) 2 (2s) 2 | (1s) 2 (2s)(2p) | 2 |
| B | (1s) 2 (2s) 2 (2p) | (1s) 2 (2s)(2p x)(2 p y) | 3 |
| C | (1s) 2 (2s) 2 (2p x)(2 p y) | (1s) 2 (2s)(2p x)(2 p y)(2 p z) | 4 |
| N | (1s) 2 (2s) 2 (2p x)(2 p y)(2 p z) | (1s) 2 (2s) 2 (2p x)(2 p y)(2 p z) | 3 |
| O | (1s) 2 (2s) 2 (2p x) 2 (2 p y)(2 p z) | (1s) 2 (2s) 2 (2p x) 2 (2 p y)(2 p z) | 2 |
| F | (1s) 2 (2s) 2 (2p x) 2 (2 p y) 2 (2 p z) | (1s) 2 (2s) 2 (2p x) 2 (2 p y) 2 (2 p z) | 1 |
| Ne | (1s) 2 (2s) 2 (2p x) 2 (2 p y) 2 (2 p z) 2 | (1s) 2 (2s) 2 (2p x) 2 (2 p y) 2 (2 p z) 2 | 0 |
These patterns are manifested in the following examples:
All of the above applies only to neutral atoms. Ions and corresponding atoms have different numbers of electrons; ions can have the same valence as other atoms with the same number of electrons. Thus, N + and B – ions have the same number of electrons (six) as a neutral carbon atom, and accordingly they are tetravalent. Ammonium ions NH 4 + and boron hydride BH 4 – form complex salts and are similar in their electronic configuration to methane CH 4.
Let us now assume that the molecules of ammonia NH 3 and boron trifluoride BF 3 are brought closer to each other. When an electron transfers from a nitrogen atom to a boron atom, we obtain two ions, NH 3 + and BF 3 –, each with an unoccupied orbital, which can lead to the formation of a covalent bond. The H 3 N–BF 3 molecule is an electronic analogue of 1,1,1-trifluoroethane H 3 C–CF 3 . Bonds formed as a result of interatomic electron transfer followed by the formation of a covalent bond are called donor-acceptor.
Geometry of molecules. Hybridization.
All atomic orbitals except s, are spherically asymmetric, and the degree of their overlap with the AO of other atoms depends on the mutual orientation of the orbitals. So, R-AO will overlap with the AO of another atom to the greatest extent if the latter is located along its axis (Fig. 4, A). This means that the bonds formed as a result of overlapping AOs must have a specific geometry. Consider the carbon atom in 5 S-condition. It has one electron in three R-orbitals and in the fourth, spherically symmetric s-orbitals. It would seem that the three bonds it forms will be different from the fourth, while R-connections will be located in mutually perpendicular directions along the axes R-AO. In fact, a different, completely symmetrical picture is observed. The easiest way to explain it is as follows. Orbital set (2 s)+(2p x)+(2 p y)+(2 p z) is a certain volume of “orbital space” capable of holding four pairs of electrons. We can obtain an equivalent description of this situation by mixing all the orbitals and dividing their sum into four equal parts, so that each of the resulting mixed or hybrid orbitals contains one pair of electrons. Therefore 5 S-state of carbon can be represented as (1 s) 2 (t 1)(t 2)(t 3)(t 4), where t i– hybrid orbitals, which successfully explains the formation of a symmetrical tetravalent carbon molecule. Let's now consider what happens when mixing R-AO s s-AO. Strengthening one half R-dumbbell interference will invariably be accompanied by a weakening of its other half (Fig. 4, b), resulting in the formation of an asymmetric hybrid orbital (Fig. 4, V). It will effectively overlap with other orbitals oriented in the same direction, forming fairly strong bonds. This is one of the reasons why the carbon atom prefers to form bonds through AO hybridization. But there is another reason. Consider a typical tetravalent carbon compound, such as methane CH4. In it, each hydrogen atom is held near a carbon atom by a pair of shared electrons. These pairs repel each other, and the optimal configuration of the molecule is one in which they are at the maximum possible distance from each other. In this case, the hydrogen atoms will be located at the vertices of a regular tetrahedron, and the carbon atom will be at its center. This geometry can be realized using the so-called. sp 3-hybrid orbitals, each formed by 1/4 of 2 s-AO and one of 2 R-AO. All these orbitals are identical in shape, easily form bonds and are directed from the carbon atom in the center of a regular tetrahedron to its four vertices (Fig. 1, G).
The nitrogen atom could form bonds with only 2 R-AO, the angles between which would be 90°, but the mutual repulsion of pairs of bonding electrons and pairs of non-bonding electrons of the 2nd shell is minimized if “tetrahedral” ones participate in the formation of bonds sp 3 -orbitals. Here, however, another feature emerges. For an N+ ion configuration (1 s) 2 (2s)(2p) 3 and (1 s) 2 (t) 4 , where t – sp 3-hybrid AOs are truly equivalent. Another thing is the neutral nitrogen atom, the 7th electron of which can occupy either 2 s-AO, and then you get the configuration (1 s) 2 (2s)(2p) 4 , or t-AO in configuration (1 s) 2 (t) 5 . Since 2 s-AO is located below 2 p-AO and therefore lower than any sp-hybrid orbital, the first configuration turns out to be energetically more favorable and one would expect that, other things being equal, trivalent nitrogen would prefer the “non-hybridized” configuration. However, the mutual repulsion of electron pairs is apparently sufficient for hybridization to occur, in which the bond angles in a nitrogen compound such as ammonia NH 3 are close to the corresponding angles in a regular tetrahedron, i.e. to 109°. The same applies to divalent oxygen in the composition of the water molecule H 2 O. In all these cases, bonded atoms occupy three (or two) vertices of the tetrahedron, and pairs of lone electrons of the 2nd shell occupy the remaining vertices.
Similar reasoning applies to other typical elements of groups IV, V and VI of the periodic table. Tetravalent elements of group IV (Si, Ge, Sn and Pb) always form tetrahedral structures, but other elements of groups V and VI (P, S, As, Se, Sb, Te) differ from nitrogen and oxygen and form compounds with bond angles, close to 90°. Apparently, due to the larger size of these atoms, the mutual repulsion of the valence electrons is not enough to allow the hybridization observed for N and O.
Bonds involving d-orbitals.
Unlike nitrogen, the phosphorus atom can form five covalent bonds. In the ground state, phosphorus has the configuration (1 s) 2 (2s) 2 (2p) 6 (3s) 2 (3p x)(3 p y)(3 p z) and is trivalent, forming, like nitrogen, compounds of the PF 3 type. However, in this case it is possible to participate 3 s-electrons in the formation of bonds, since d-AO (3 d) have the same principal quantum number. Indeed, pentavalent phosphorus compounds of the PF 5 type are also known, where phosphorus is in the +5 valence state, consistent with the electronic configuration (1 s) 2 (2s) 2 (2p) 6 (3s)(3p x)(3 p y)(3 p z)(3 d); connections in this case are formed as a result sp 3 d-hybridization (i.e. as a result of mixing one s-, three R- and one d-AO). The optimal structure from the point of view of reducing the mutual repulsion of pairs of valence electrons is a triangular bipyramid (Fig. 5, A). Sulfur can be not only divalent, but also tetravalent (SF 4) and hexavalent (SF 6), being in states (1 s) 2 (2s) 2 (2p) 6 (3s) 2 (3p x)(3 p y)(3 p z)(3 d) and (1 s) 2 (2s) 2 (2p) 6 (3s)(3p x)(3 p y)(3 p z)(3 d 1)(3d 2) accordingly. In tetravalent sulfur compounds, the mutual repulsion of electrons of the 3rd shell is optimized by hybridization of the orbitals of all its electrons. The structure of compounds of this type is similar to the structure of PF 5, but one of the vertices of the triangular bipyramid is occupied by a pair of lone electrons of the 3rd shell (Fig. 5, b). In hexavalent sulfur compounds, the mutual repulsion of electrons is minimized when sp 3 d 2 -
hybridization, when all orbitals are equivalent and directed towards the vertices of a regular octahedron (Fig. 5, V).
Until now, we have considered only those elements of the periodic table that have shells with d-orbitals are either completely filled or completely empty. Let us now dwell on the transition elements in which these shells are not completely filled. The energy of electrons in different orbitals of the 3rd shell increases in the following order: 3 s p d; all orbitals are too far from the 2nd shell orbitals for hybridization to occur. At the same time 3 d-orbitals and orbitals of the 4th shell are energetically close enough so that interaction 3 is possible d-, 4s- and 4 R-orbitals, and transition elements from Sc to Cu can form covalent bonds by hybridizing these orbitals. In all cases where there are two 3 d-orbitals, bond formation occurs through d 2 sp 3-hybridization, while the hybrid orbitals are similar in shape to sp 3 d 2 -orbitals. The elements in compounds of this type are hexavalent, and the molecules of the compounds themselves have the shape of an octahedron (Fig. 5, V). Most of them contain ions, and can be considered to be formed by the interaction of an ion of the central atom with six molecules, each of which has a pair of lone electrons. Covalent bonds with the central ion are called donor-acceptor bonds. A simple example of such a compound is the hexammine ion of trivalent cobalt Co(NH 3) 6 3+. The Co 3+ ion has an electronic configuration (1 s) 2 (2s) 2 (2p) 6 (3s) 2 (3p) 6 (3d 1) 2 (3d 2) 2 (3d 3) 2, and three of his five 3 are fully occupied d-orbitals, and two are 3 d-AO are free. These orbitals can hybridize with 4 s- and 4 R-AO with the formation of six octahedral d 2 sp 3-orbitals; all of them are free and can participate in the formation of acceptor bonds with six ammonia molecules.
A different picture is observed when the central atom has only one free d-orbital. An example is the doubly charged nickel ion Ni 2+, in which the optimal configuration occurs when four bonds are formed using dsp 2 -orbitals. These orbitals lie in the same plane at an angle of 90° to each other.
Multiple connections.
One of the well-known carbon compounds is ethylene C 2 H 4, in which each carbon atom is bonded to only three other atoms. By analogy with boron, we can assume that the optimal geometry will be such that sp 2-hybrid orbitals lie in the same plane. In this case, each carbon atom will have one unused (in sp 2 -hybridization) R-orbital that contains one of the four valence electrons. If all six ethylene atoms lie in the same plane, then the two unused R-AOs overlap with each other as shown in Fig. 6, A. This overlap leads to the formation of a pair of MOs: one binding (Fig. 6, b) and one loosening (Fig. 6, V). Because they each contain only one electron, they can form a low-energy bonding MO. This creates an additional bond between carbon atoms, and the structural formula of ethylene has the form
This new type of bond differs from those formed by overlapping orbitals along the line of bonding of atoms in two respects. The last type of bonds, C–C single bonds, are axially symmetrical and are therefore not affected by the rotation of the groups they connect. On the contrary, overlap R-orbitals depends on whether all six atoms in the ethylene molecule lie in the same plane, since for optimal overlap R-AO must be parallel. Thus, while rotation around a single C–C bond can occur relatively freely, rotation around a double C=C bond is very difficult. Indeed, the ethylene molecule is a rigid, flat structure. The second difference concerns the degree of orbital overlap. Cross overlap R-AO is relatively inefficient, and therefore this type of connection is weak. Therefore, ethylene is chemically more active than saturated compounds that have only single bonds.
S-bonds, and with transverse overlap – p- connections.
The molecules of some compounds, for example acetylene C 2 H 2, contain triple bonds. In them, each carbon atom is connected to its neighbor s- connections formed sp-hybrid orbitals. They are collinear, so four atoms in an acetylene molecule lie on the same straight line. Rest R-AO carbon atoms, when overlapping, form two p- connections.
Aromatic compounds.
The benzene molecule C 6 H 6 is represented as a six-membered ring of carbon atoms, each of which also has a hydrogen atom attached (Fig. 7, A). Since each carbon atom has three neighbors, it can be assumed that the corresponding bonds are formed as a result sp 2-hybridization and lie in the same plane at an angle of 120° to each other. Indeed, the benzene molecule is a flat structure. Unused R-AO carbon atoms can form p-connections (Fig. 7, b), however, for benzene the situation turns out to be more complicated than in the cases considered above, when bonds were formed as a result of overlapping AO pairs. In benzene 2 R-The AO of each carbon atom must overlap equally effectively with 2 R-AO of all neighboring atoms. (Here we can draw an analogy with multiple wave interference by comparing the overlap of orbitals in a benzene molecule with the overlap of waves diffracted by two slits or a diffraction grating.) As a result, for benzene we obtain a set of ring molecular orbitals covering all six carbon atoms (Fig. 7, V). The total energy of the system with such an electron configuration is less than if R-AOs formed ordinary ones in pairs p- connections. Indeed, benzene is more stable and less active than would be expected based on its “classical” structure (Fig. 7, G). All bonds in its molecule are symmetrical, and their lengths are the same, and in strength they occupy an intermediate position between single and double bonds. Other compounds are also known in which p-electrons participate in the formation of “multicenter” MOs and for which similar features of bond lengths and chemical activity are observed. 
Compounds containing multicenter bonds.
Even in such simple molecules as CH 4, individual molecular orbitals necessarily interact with each other. Therefore, the idea of localized two-center covalent bonds can only be considered as a certain approximation. Typically, however, these interactions are weak because the degree of orbital overlap is small (except p-MO in aromatic and similar compounds). Nevertheless, we cannot rule out the existence of molecules with multiple overlapping AOs responsible for the formation of bonds by sharing electrons with three or more atoms. An example is diborane B 2 H 6, which has six pairs of valence electrons; this is not enough to form the seven bonds needed to create the classical H 3 B–BH 3 structure. H. Longuet-Higgins proposed the structure of diborane, shown in Fig. 8, A. In this structure, the central hydrogen atoms are connected by three-center bonds formed as a result of overlapping sp 3-hybrid orbitals of two boron atoms with 1 s-AO of the hydrogen atom (Fig. 8, b). Four of the six pairs of valence electrons participate in the formation of ordinary s-bonds with “terminal” hydrogen atoms, and two pairs of three-center bonds. A more complex example of a multicenter bond is provided by the dibenzene chromium molecule (Fig. 8, V). The benzene rings in this molecule are connected to the metal atom by complex multicenter orbitals formed by overlapping p-Benzene MO with 3 d-, 4s- and 4 R-AO of the central atom. Other similar compounds are known that have a sandwich-type structure. 
Prospects.
By now, the general principles of the structure of molecules can be considered established. Physicochemical methods have been developed for determining the structure of complex molecules, including biological ones. Progress in two related directions is possible in the near future. We should expect, firstly, an increase in the accuracy of quantum mechanical calculations and, secondly, an improvement in experimental methods for measuring the corresponding molecular parameters.
May contain positively and negatively charged, i.e.; in this case are implemented. In addition to those indicated, there are also weaker interactions between. Repulsive forces act between valence-unbonded bonds.
The development of the doctrine of structure is inextricably linked with success, first of all. The theory of structure, created in the 60s. 19th century the works of A. M. Butlerov, F. A. Kekule, A. S. Cooper and others, made it possible to represent or by structural formulas expressing the sequence of valence in. With the same empirical formula, there can be different structures with different properties (phenomenon). These are, for example, C 5 H 5 OH and (CH 3) 2 O. These compounds differ:
In some cases, isomeric ones quickly transform into one another and a dynamic relationship is established between them (see). Subsequently, J. H. Van't Hoff and independently the French chemist A. J. Le Bel came to an understanding of the spatial arrangement in and to an explanation of the phenomenon. A. Werner (1893) extended the general ideas of the theory of structure to inorganic ones. By the beginning of the 20th century. had a detailed theory based on the study of only their chemical properties. It is remarkable that direct physical research methods, developed later, in the overwhelming majority of cases completely confirmed those established by studying macroscopic quantities, and not individual ones.
Equilibrium internuclear distances r 0 and energies D (at 25° C) of some diatomic
| r 0, Ǻ | r 0 , Ǻ | |||||||
| C-Br……………. | ||||||||
| Cº C……………... | C-I……………… | |||||||
| C-H…………….. | C-S…………….. | |||||||
| C-O…………….. | O-H……………. | |||||||
| C=O……………... | N-H…………….. | |||||||
| C-N…………….. | S-H…………….. |
|
In the vast majority of cases, the total valence in is equal to zero, i.e., they are pairwise saturated. , containing unpaired - (for example, atomic H · · , methyl CH · · 3) are usually unstable, because when they combine with each other, a significant decrease in energy occurs due to the formation of valence bonds. The most effective method for studying the structure is (). Electrical and optical properties. Behavior in an electric field is determined by the basic electrical characteristics - constant and . means a discrepancy between the centers of gravity of positive and negative charges, i.e. electrical asymmetry. Accordingly, those with a center, for example H 2, are deprived of a constant; on the contrary, in HCl they are shifted towards Cl and are equal to 1.03 D (1.03 × 10 -18 CGS units). characterized by the ability of any electron shell to shift under the influence of an electric field, as a result of which an induced one is created. The values of and are found experimentally using dielectric constant measurements. In the case of additivity of properties, it can be represented by the sum of connections (taking into account their direction), the same applies to. Elements with or odd numbers have nuclear spin paramagnetism. Such nuclei are characterized |
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