In chemistry for grade 11 (O.S. Gabrielyan, 2007),
task №9
to the chapter " § 1. Basic information about the structure of the atom».

The discovery of the complex structure of the atom is the most important stage in the development of modern physics. In the process of creating a quantitative theory of atomic structure, which made it possible to explain atomic systems, new ideas were formed about the properties of microparticles, which are described by quantum mechanics.

The idea of ​​atoms as indivisible smallest particles of substances, as noted above, arose in ancient times (Democritus, Epicurus, Lucretius). In the Middle Ages, the doctrine of atoms, being materialistic, did not receive recognition. By the beginning of the 18th century. atomic theory is gaining increasing popularity. By this time, the works of the French chemist A. Lavoisier (1743-1794), the great Russian scientist M.V. Lomonosov and the English chemist and physicist D. Dalton (1766-1844) proved the reality of the existence of atoms. However, at this time the question of the internal structure of atoms did not even arise, since atoms were considered indivisible.

A major role in the development of atomic theory was played by the outstanding Russian chemist D.I. Mendeleev, who in 1869 developed the periodic system of elements, in which for the first time the question of the unified nature of atoms was raised on a scientific basis. In the second half of the 19th century. It has been experimentally proven that the electron is one of the main parts of any substance. These conclusions, as well as numerous experimental data, led to the fact that at the beginning of the 20th century. The question of the structure of the atom seriously arose.

The existence of a natural connection between all chemical elements, clearly expressed in Mendeleev’s periodic system, suggests that the structure of all atoms is based on a common property: they are all closely related to each other.

However, until the end of the 19th century. In chemistry, the metaphysical conviction prevailed that the atom is the smallest particle of simple matter, the final limit of the divisibility of matter. During all chemical transformations, only molecules are destroyed and created again, while atoms remain unchanged and cannot be split into smaller parts. Various assumptions about the structure of the atom have not been confirmed by any experimental data for a long time.

data. Only at the end of the 19th century. discoveries were made that showed the complexity of the structure of the atom and the possibility of transforming some atoms into others under certain conditions. Based on these discoveries, the doctrine of the structure of the atom began to develop rapidly.

The first indirect evidence of the complex structure of atoms was obtained from the study of cathode rays generated during an electrical discharge in highly rarefied gases. The study of the properties of these rays led to the conclusion that they are a stream of tiny particles carrying a negative electrical charge and flying at a speed close to the speed of light. Using special techniques, it was possible to determine the mass of cathode particles and the magnitude of their charge, and to find out that they do not depend either on the nature of the gas remaining in the tube, or on the substance from which the electrodes are made, or on other experimental conditions. Moreover, cathode particles are known only in a charged state and cannot be stripped of their charges and converted into electrically neutral particles: electric charge is the essence of their nature. These particles, called electrons, were discovered in 1897 by the English physicist J. Thomson.

The study of the structure of the atom practically began in 1897-1898, after the nature of cathode rays as a stream of electrons was finally established and the charge and mass of the electron were determined. Thomson proposed the first model of the atom, presenting the atom as a clump of matter with a positive electrical charge, in which so many electrons are interspersed that it turns it into an electrically neutral formation. In this model, it was assumed that, under the influence of external influences, electrons could oscillate, i.e., move at an accelerated rate. It would seem that this made it possible to answer questions about the emission of light by atoms of matter and gamma rays by atoms of radioactive substances.

Thomson's model of the atom did not assume positively charged particles inside an atom. But how then can we explain the emission of positively charged alpha particles by radioactive substances? Thomson's atomic model did not answer some other questions.

In 1911, the English physicist E. Rutherford, while studying the movement of alpha particles in gases and other substances, discovered a positively charged part of the atom. Further more thorough studies showed that when a beam of parallel rays passes through layers of gas or a thin metal plate, no longer parallel rays emerge, but somewhat diverging ones: alpha particles are scattered, i.e., they deviate from the original path. The deflection angles are small, but there are always a small number of particles (about one in several thousand) that are deflected very strongly. Some particles are thrown back as if they had encountered an impenetrable barrier. These are not electrons - their mass is much less than the mass of alpha particles. Deflection can occur when colliding with positive particles whose mass is of the same order as the mass of alpha particles. Based on these considerations, Rutherford proposed the following diagram of the structure of the atom.

At the center of the atom there is a positively charged nucleus, around which electrons rotate in different orbits. The centrifugal force arising during their rotation is balanced by the attraction between the nucleus and the electrons, as a result of which they remain at certain distances from the nucleus. Since the mass of an electron is negligible, almost the entire mass of an atom is concentrated in its nucleus. The share of the nucleus and electrons, the number of which is relatively small, accounts for only an insignificant part of the total space occupied by the atomic system. The diagram of the structure of the atom proposed by Rutherford, or, as they usually say, the planetary model of the atom, easily explains the phenomena of deflection of alpha particles. Indeed, the size of the nucleus and electrons is extremely small compared to the size of the entire atom, which is determined by the orbits of the electrons farthest from the nucleus, so most alpha particles fly through atoms without noticeable deflection. Only in cases where the alpha particle comes very close to the nucleus does electrical repulsion cause it to deviate sharply from its original path. Thus, the study of the scattering of alpha particles laid the foundation for the nuclear theory of the atom.

Ernest Rutherford and Niels Bohr played a special role in the creation of the modern theory of atomic structure. In 1911, Rutherford, conducting experiments on bombarding metal foil with α-particles, found out that heavy particles, later called nuclei, are located inside atoms, and he also proposed a planetary model of the structure of the atom. And in 1913, Bohr put forward the first quantum theory of the atom.

So, an atom is a complex microsystem consisting of elementary particles. It consists of a positively charged nucleus and negatively charged electrons. The carrier of the positive charge of the nucleus is the proton. The nuclei of atoms of all elements, except the light isotope of hydrogen, include protons and neutrons.

The nucleus is the fundamental basis of the atom, determines the individuality of the elements, and consists of nucleons.

The atomic radii are 0.05 – 0.30 nm. The mass of an electron is much smaller than the mass of nucleons, so the mass of an atom is approximately equal to the mass of the nucleus.

The properties of the nucleus are determined mainly by its composition. The number of protons - the charge of the nucleus, characterizes the atom's belonging to a given chemical element. Another important characteristic of the nucleus is the mass number A:

A = Z + N, where N is the total number of neutrons; Z is the total number of protons.

Atoms with different numbers of protons and neutrons, but with the same A are called isobars.

Atoms with the same number of protons, and therefore the same nuclear charge, are called isotopes .

Atoms with the same number of neutrons are called isotones.

Based on the determination of isotopes, a more modern definition of a chemical element can be given. Chemical element is a type of atom characterized by a certain nuclear charge.

In chemistry, the atomic nucleus is considered to be a point that has a positive charge + Z and a mass expressed by the mass number A. However, the nucleus is a particle that has its own structure. According to modern theory, the atomic nucleus has a shell structure. Protons and neutrons independently occupy nuclear layers and sublayers, just as is observed for electrons in the electron shell of an atom.

A comparison shows that the mass of the nucleus is always less than the arithmetic sum of the masses of the protons and neutrons that make up its composition. The difference between these quantities is called mass defect. The mass defect corresponds to the energy that is released during the formation of a nucleus from free protons and neutrons, and can be calculated from Einstein's relation E = mc 2. This binding energy nucleons in the nucleus, it is millions of times higher than the binding energy of atoms in a molecule. Therefore, during chemical transformations of substances, atomic nuclei do not change.

It should be noted that the nature of nuclear forces has not been fully clarified; they act at very short distances (about 10 -15 meters) and bind individual protons and neutrons, forming large nuclei. Currently, about 300 stable and over 1400 radioactive nuclei are known.

As we have already found out, the nuclei of atoms are not affected during chemical reactions, therefore the chemical properties of elements mainly depend on the number and location of electrons in their atoms.

Electron is a negatively charged extremely light particle that was discovered in 1897 during experiments with cathode rays.

In 1913, Niels Bohr developed a model of the atom. His theory of atomic structure was based on two provisions called Bohr's postulates.

First postulate: electrons move around the nucleus of an atom in circular orbits. The closer the electron's orbit is to the nucleus, the smaller the energy reserve of the atom, since as the distance between two electric charges of different signs decreases, their total energy decreases.

Bohr suggested that in an atom electrons can only move in certain orbits (meaning the distance from the nucleus), which are called permitted . Each electron orbit corresponds to a certain energy of the atom. An electron can move from orbit to orbit only in jumps. The energy of the atom also changes abruptly; the energy is said to be quantized.

Second postulate States that the angular momentum (momentum mv) of an electron in its orbit is of a quantum nature and is equal to an integer multiple of h/2π(h – Planck’s constant, h = 6.63 ∙ 10 -34 J∙ s).

Introduction

The discovery of the complex structure of the atom is the most important stage in the development of modern physics. The first information about the structure of the atom was obtained by studying the processes of passage of electric current through liquids. In the thirties of the XIX century. The experiments of the outstanding physicist M. Faraday suggested that electricity exists in the form of separate unit charges. The discovery of the spontaneous decay of atoms of some elements, called radioactivity, became direct evidence of the complexity of the structure of the atom.

In 1902, English scientists Ernest Rutherford and Frederick Soddy proved that during radioactive decay, a uranium atom turns into two atoms - a thorium atom and a helium atom. This meant that atoms were not immutable, indestructible particles.

The development of research into radioactive radiation, on the one hand, and quantum theory, on the other, led to the creation of the Rutherford-Bohr quantum model of the atom. But the creation of this model was preceded by attempts to construct a model of the atom based on the concepts of classical electrodynamics and mechanics. In 1904, publications appeared on the structure of the atom, some of which belonged to the Japanese physicist Hantaro Nagaoka, others to the English physicist D.D. Thomson.

Nagaoka presented the structure of the atom as similar to the structure of the solar system: the role of the Sun is played by the positively charged central part of the atom, around which “planets” - electrons - move in established ring-shaped orbits. At slight displacements, electrons excite electromagnetic waves.

In a Thomson atom, positive electricity is “distributed” over a sphere in which electrons are embedded. In the simplest hydrogen atom, the electron is located in the center of a positively charged sphere. In multielectron atoms, electrons are arranged in stable configurations calculated by Thomson. Thomson considered each such configuration to determine the chemical properties of atoms. He made an attempt to theoretically explain the periodic system of elements by D.I. Mendeleev. Bohr later pointed out that since this attempt, the idea of ​​dividing the electrons in an atom into groups became the starting point.

But it soon turned out that new experimental facts refute Thomson’s model and, on the contrary, testify in favor of the planetary model. These facts were discovered by Rutherford. First of all, it should be noted the discovery of the nuclear structure of the atom.

The basis of the modern theory of the electronic structure of atoms was the planetary model of the atom by Niels Bohr.

The purpose of the abstract: to reflect the process of evolution of ideas about the structure of atoms using the examples of the models of Ernest Rutherford and Niels Bohr.

Objectives of the abstract: study, analyze, generalize the ideas about the structure of atoms expressed by E. Rutherford and N. Bohr, draw conclusions about the most correct assumption, from the point of view of modern physics. In the process of work, various types of sources were used: textbooks by S. Kh. Karpenkov and T.I. Trofimova, intended for higher education. They tell in accessible language the history of the origin and development of knowledge about the structure of the atom. This approach is driven by the desire to study the problem in all its complexity.

Atomic structure

An atom (from the Greek atomos - indivisible) is a particle of a substance of microscopic size and very low mass (microparticle), the smallest part of a chemical element, which is the carrier of its properties. Each element corresponds to a certain type of atom, designated by the symbol of the element (for example, a hydrogen atom H, an iron atom Fe; a mercury atom Hg; a uranium atom U).

According to modern concepts, an atom is a complex system consisting of a positively charged nucleus and electrons surrounding the nucleus.

The nucleus is the central part of an atom, in which almost the entire mass of the atom and its positive electric charge are concentrated. All atomic nuclei consist of elementary particles: protons and neutrons, which are considered two charge states of one particle - the nucleon. A proton has a positive electric charge, equal in absolute value to the charge of an electron. A neutron has no electrical charge.

The electrons surrounding the nucleus of an atom are negatively charged microparticles with a mass of ~ 5∙10 -4 atomic mass unit and a charge of -1.6 ∙ 10 -19 K (-1). Since the mass of an electron is negligible compared to the mass of a proton or neutron, the mass of an atom is practically equal to the mass of its nucleus, i.e. the sum of the masses of protons and neutrons. The number of electrons in an atom is equal to the number of positively charged protons that make up the nucleus.

The dimensions of an atom as a whole are determined by the dimensions of its electron shell and are large compared to the dimensions of the atomic nucleus. The electron shells of an atom do not have a strictly defined boundary; the values ​​of atomic sizes depend to a greater or lesser extent on the methods of their determination and are very diverse

In 1911, the English scientist Ernest Rutherford came up with a “planetary” model of the atom, according to which in the center of the atom Rutherford placed a tiny but very dense nucleus in which almost the entire mass of the atom was concentrated, and the electrons revolved around it in certain orbits, like planets around the Sun .

Then it turned out that each electron moves around the nucleus so quickly that it not only cannot be examined with the most powerful microscope, but it is even impossible to imagine it as a point moving along a certain trajectory. The electron is, as it were, “smeared” in space and forms an electron cloud, and the shape of the electron cloud can be different.

At the moment, four forms of electron clouds are known: s-electrons (spherical shape of the electron cloud); p-electrons (electron cloud shape - dumbbell or figure eight); d-electrons; f electrons.

Rutherford and Bohr models

In 1911, the English physicist Ernest Rutherford, while studying the movement of alpha particles in gases and other substances, discovered a positively charged part of the atom. Further more thorough studies showed that when a beam of parallel rays passes through layers of gas or a thin metal plate, no longer parallel rays emerge, but somewhat diverging ones: alpha particles are scattered, i.e., they deviate from the original path. The deflection angles are small, but there are always a small number of particles (about one in several thousand) that are deflected very strongly. Some particles are thrown back as if they had encountered an impenetrable barrier. These are not electrons - their mass is much less than the mass of alpha particles. Deflection can occur when colliding with positive particles whose mass is of the same order as the mass of alpha particles. Based on these considerations, Rutherford proposed a nuclear (planetary) model of the structure of the atom.

“At the center of the atom there is a positively charged nucleus, around which electrons rotate in different orbits. The centrifugal force arising during their rotation is balanced by the attraction between the nucleus and the electrons, as a result of which they remain at certain distances from the nucleus. Since the mass of an electron is negligible, almost the entire mass of an atom is concentrated in its nucleus. The share of the nucleus and electrons, the number of which is relatively small, accounts for only an insignificant part of the total space occupied by the atomic system.”

The diagram of the structure of the atom proposed by Rutherford, or, as they usually say, the nuclear model of the atom, easily explains the phenomena of deflection of alpha particles. Indeed, the size of the nucleus and electrons is extremely small compared to the size of the entire atom, which is determined by the orbits of the electrons farthest from the nucleus, so most alpha particles fly through atoms without noticeable deflection. Only in cases where the alpha particle comes very close to the nucleus does electrical repulsion cause it to deviate sharply from its original path. Thus, the study of the scattering of alpha particles laid the foundation for the nuclear theory of the atom. But, despite consistent reasoning, Rutherford's model could not explain all the properties of atoms. Thus, according to the laws of classical physics, an atom from a positively charged nucleus and electrons revolving in circular orbits should emit electromagnetic waves. “The emission of electromagnetic waves should lead to a decrease in the potential energy reserve in the nucleus-electron system, to a gradual decrease in the radius of the electron’s orbit and the electron’s fall onto the nucleus. However, atoms usually do not emit electromagnetic waves, electrons do not fall on atomic nuclei, that is, atoms are stable.” Attempts to build a model of the atom within the framework of classical physics did not lead to success: Thomson's model was refuted by Rutherford's experiments, while the nuclear model turned out to be unstable electrodynamically and contradicted experimental data. Overcoming the difficulties that arose required the creation of a qualitatively new theory of the atom.

The first attempt to construct a qualitatively new theory of the atom was made in 1913. Danish physicist Niels Bohr. He set himself the goal of linking into a single whole the empirical laws of line spectra, the Rutherford nuclear model of the atom, and the quantum nature of the emission and absorption of light. Bohr based his theory on Rutherford's nuclear model. He suggested that electrons move around the nucleus in circular orbits. Circular motion, even at constant speed, has acceleration. This accelerated movement of charge is equivalent to alternating current, which creates an alternating electromagnetic field in space. Energy is consumed to create this field. The field energy can be created due to the energy of the Coulomb interaction of the electron with the nucleus. As a result, the electron must move in a spiral and fall onto the nucleus. However, experience shows that atoms are very stable formations. It follows from this that the results of classical electrodynamics, based on Maxwell’s equations, are not applicable to intra-atomic processes. It is necessary to find new patterns. Bohr based his theory on two postulates.

Bohr’s first postulate (postulate of stationary states): “in an atom there are stationary (not changing with time) states in which it does not emit energy. Stationary states of an atom correspond to stationary orbits along which electrons move. The movement of electrons in stationary orbits is not accompanied by the emission of electromagnetic waves. In the stationary state of an atom, an electron moving in a circular orbit must have discrete quantum values ​​of angular momentum that satisfy the condition.

Bohr’s second postulate (frequency rule): “when an electron passes from one stationary orbit to another, one photon is emitted (absorbed) with energy = En – Em equal to the difference in the energies of the corresponding stationary states (En and Em are, respectively, the energies of the stationary states of the atom before and after radiation and absorption). At En > Em, photon emission occurs (the transition of an atom from a state with higher energy to a state with lower energy, i.e., the transition of an electron from an orbit more distant from the nucleus to a closer one), at En< Em - его поглощение (переход атома в состояние с большей энергией, т.е. переход атома на более отдалённую от ядра орбиту)».

Bohr's theory brilliantly explained the experimentally observed line spectrum of hydrogen. But the successes of the theory of the hydrogen atom were achieved at the cost of abandoning the fundamental principles of classical mechanics, which has remained unconditionally valid for more than 200 years. Therefore, direct experimental proof of the validity of Bohr's postulates, especially the first - on the existence of stationary states - was of great importance. The second postulate can be considered as a consequence of the law of conservation of energy and the hypothesis about the existence of photons.

German physicists D. Frank and G. Hertz, studying the collision of electrons with gas atoms using the retarding potential method (1913), experimentally confirmed the existence of stationary states and the discreteness of atomic energy values.

Despite the undoubted success of Bohr's concept in relation to the hydrogen atom, for which it turned out to be possible to construct a quantitative theory of the spectrum, it was not possible to create a similar theory for the helium atom next to hydrogen on the basis of Bohr's ideas. Regarding the helium atom and more complex atoms, Bohr's theory allowed us to draw only qualitative (albeit very important) conclusions. The idea of ​​certain orbits along which an electron moves in a Bohr atom turned out to be very conditional. In fact, the movement of electrons in an atom has little in common with the movement of planets in orbit.

Currently, with the help of quantum mechanics, it is possible to answer many questions regarding the structure and properties of atoms of any elements.

Modern ideas about the structure of the atom

The dual nature of the electron, which has the properties of not only a particle, but also a wave, was confirmed experimentally in 1927, prompting scientists to create a new theory of the structure of the atom that takes into account both of these properties. The modern theory of atomic structure is based on quantum mechanics.
The duality of the properties of an electron is manifested in the fact that, on the one hand, it has the properties of a particle (has a certain rest mass), and on the other, its movement resembles a wave and can be described by a certain amplitude, wavelength, oscillation frequency, etc. Therefore, one cannot say about any specific trajectory of an electron's movement - one can only judge one or another degree of probability of its being at a given point in space.
Consequently, the electron orbit should be understood not as a specific line of movement of the electron, but as a certain part of the space around the nucleus, within which the probability of the electron being is greatest. In other words, the electron orbit does not characterize the sequence of movement of an electron from point to point, but is determined by the probability of finding an electron at a certain distance from the nucleus. In this regard, the electron is not represented as a material point, but as if “smeared” throughout the entire volume of the atom in the form of a so-called electron cloud, which has areas of condensation and rarefaction of the electric charge. The idea of ​​an electron as some cloud of electric charge is convenient; it quite accurately conveys the behavior of the electron. However, it should be borne in mind that the electron cloud does not have sharply defined boundaries, and even at a great distance from the nucleus there is a possibility of an electron remaining. To characterize the shape of the electron cloud, the concept of an orbital instead of the concept of an orbit was introduced precisely in order not to confuse the motion of an electron with the motion of a body in classical physics. However, when considering the structure of an atom in a simplified manner, the term orbit is sometimes retained, nevertheless remembering the special nature of the movement of the electron in the atom.

Modern ideas about the structure of the atom are subject to the quantum model of the structure of the atom, which takes into account the wave properties of elementary particles. Let us present its main provisions.

The electron has a dual (particle-wave) nature, i.e. behaves both as a particle and as a wave. As a particle, an electron has mass and charge; like a wave, it has the ability to diffraction.

It is impossible for an electron to accurately measure its position and velocity at the same time.

An electron in an atom does not move along certain trajectories, but can be located in any part of the perinuclear space, but the probability of its being in different parts of this space is not the same. The region of space where an electron is most likely to be located is called an orbital.

The nuclei of atoms consist of protons and neutrons, which have a common name - nucleons.

Conclusion

The basis of the modern theory of atomic structure is the planetary model, supplemented and improved. According to this theory, the nucleus of an atom consists of protons (positively charged particles) and neurons (particles without a charge). And around the nucleus electrons (negatively charged particles) move along uncertain trajectories.

This study reflected the process of evolution of ideas about the structure of atoms using the models of Ernest Rutherford and Niels Bohr as an example. The ideas about the structure of atoms expressed by Rutherford and Bohr have been fully studied, analyzed and generalized. From the point of view of modern physics, the most correct assumption about the structure of the atom was made by the Danish scientist Niels Bohr.

Thus, the discoveries of Rutherford and Bohr are fundamental and of great importance for modern physics and for all humanity. The history of science teaches that every time humanity masters the next rung of the ladder leading into the depths of matter, this leads to the discovery of a new, even more powerful type of energy.

Combustion and explosion are associated with the rearrangement of molecules. Intraatomic processes are accompanied by the release of millions of times more energy. An even greater release of energy occurs at the level of elementary particles. What will happen at the next steps? The discoveries of Rutherford and Bohr proved that the atom is not an indivisible particle, and make it possible for modern physics to answer this question.

Bibliography

1. Alekseev I. S. Development of ideas about the structure of the atom. – M.: Nauka, 2000.

2. Bochkarev A.I. Bochkareva T.S., Saksonov S.V. Concepts of modern natural science. - M.: Nauka, 2008.

3. Gorbachev V.V. Concepts of modern natural science, - M.: Alfa-M, 2003.

4. Korenev Yu. M. General and inorganic chemistry, in 3 parts. M.: Moscow University Publishing House, 2002.

5. Kudryavtsev L. S. Course in the history of physics. – M.: Nauka, 2006.

6. Karpenkov S. Kh. Concepts of modern natural science. Textbook for universities. – M.: Academic Project, 2000.

7. Trofimova T. I. Physics course: textbook. - M.: Higher School, 2007.

1. Pieces of matter. Democritus believed that the properties of a particular substance are determined by the shape, mass, and other characteristics of the atoms that form it. So, let’s say, the atoms of fire are sharp, so the fire is capable of burning; the atoms of solids are rough, so they adhere tightly to each other; those of water are smooth, so it is capable of flowing. Even the human soul, according to Democritus, consists of atoms.

2. Corpuscular-kinetic theory of heat. M.V. Lomonosov claims that all substances consist of “corpuscles” - “molecules”, which are “assemblies” of “elements” - “atoms”: “An element is a part of the body that does not consist of any other smaller and bodies different from it... A corpuscle is a collection of elements that form one small mass.” He gives the “element” its contemporary meaning - in the sense of the limit of divisibility of bodies - their last component. The scientist points to its spherical shape. It was M.V. Lomonosov who came up with the idea of ​​the “internal rotational (“rotary”) motion of particles”—the rotation speed is affected by an increase in temperature. With all the costs of such a model, it is important for scientists to give the concept of movement a deeper physical significance.

3. Thomson's model of the atom (Plum pudding model). J. J. Thomson proposed to consider the atom as some positively charged body with electrons enclosed inside it. This model did not explain the discrete nature of the radiation of an atom and its stability. It was finally refuted by Rutherford after his famous experiment on the scattering of alpha particles.

4. Nagaoka's early planetary model of the atom. In 1904, Japanese physicist Hantaro Nagaoka proposed a model of the atom, built by analogy with the planet Saturn. In this model, electrons, united in rings, rotated in orbits around a small positive nucleus. The model turned out to be erroneous, but some of its important provisions were included in Rutherford's model.

5. Bohr-Rutherford planetary model of the atom. In 1911, Ernest Rutherford, after conducting a series of experiments, came to the conclusion that the atom is a kind of planetary system in which electrons move in orbits around a heavy, positively charged nucleus located in the center of the atom (“Rutherford’s atom model”). However, such a description of the atom came into conflict with classical electrodynamics. The fact is that, according to classical electrodynamics, an electron, when moving with centripetal acceleration, should emit electromagnetic waves and, therefore, lose energy. Calculations showed that the time it takes for an electron in such an atom to fall onto the nucleus is absolutely insignificant. To explain the stability of atoms, Niels Bohr had to introduce postulates that boiled down to the fact that an electron in an atom, being in some special energy states, does not emit energy (“Bohr-Rutherford model of the atom”). Bohr's postulates showed that classical mechanics is inapplicable to describe the atom. Further study of atomic radiation led to the creation of quantum mechanics, which made it possible to explain the vast majority of observed facts.

6. Quantum mechanical model of the atom. The modern model of the atom is a development of the planetary model. According to this model, the nucleus of an atom consists of positively charged protons and uncharged neutrons and is surrounded by negatively charged electrons. However, the concepts of quantum mechanics do not allow us to assume that electrons move around the nucleus along any definite trajectories (the uncertainty of the coordinate of an electron in an atom can be comparable to the size of the atom itself). The chemical properties of atoms are determined by the configuration of the electron shell and are described by quantum mechanics. The position of an atom in the periodic table is determined by the electrical charge of its nucleus (that is, the number of protons), while the number of neutrons does not fundamentally affect chemical properties; in this case, there are, as a rule, more neutrons in the nucleus than protons (see: atomic nucleus). If an atom is in a neutral state, then the number of electrons in it is equal to the number of protons. The main mass of the atom is concentrated in the nucleus, and the mass fraction of electrons in the total mass of the atom is insignificant (a few hundredths of a percent of the mass of the nucleus).

The idea of ​​atoms as indivisible smallest particles of matter arose in ancient times, but only in the 18th century, through the works of A. Lavoisier, M.V. Lomonosov and other scientists, the reality of the existence of atoms was proven. But the question of their internal structure did not even arise, and atoms were still considered indivisible particles. In the 19th century, the study of the atomic structure of matter made significant progress. In 1833, while studying the phenomenon of electrolysis, M. Faraday established that the current in an electrolyte solution is the ordered movement of charged particles - ions. Faraday determined the minimum charge of an ion, which was called the elementary electric charge. Its approximate value turned out to be e = 1.60·10 –19 C.

Based on Faraday's research, it was possible to conclude that there are electric charges inside atoms.

A major role in the development of atomic theory was played by the outstanding Russian chemist D.I. Mendeleev, who developed the periodic system of elements in 1869, in which the question of the unified nature of atoms was first raised.

Important evidence of the complex structure of atoms came from spectroscopic studies, which led to the discovery of line spectra of atoms. At the beginning of the 19th century, discrete spectral lines were discovered in the radiation of hydrogen atoms in the visible part of the spectrum. Subsequently, in 1885, I. Balmer established mathematical laws connecting the wavelengths of these lines.

In 1896, A. Becquerel discovered the phenomenon of invisible penetrating radiation emitted by atoms, called radioactivity. In subsequent years, the phenomenon of radioactivity was studied by many scientists (M. Sklodowska-Curie, P. Curie, E. Rutherford, etc.). It was discovered that atoms of radioactive substances emit three types of radiation of different physical nature (alpha, beta and gamma rays). Alpha rays turned out to be a stream of helium ions, beta rays - a stream of electrons, and gamma rays - a stream of hard X-ray quanta.

In 1897, J. Thomson discovered the electron and measured the e/m ratio of the electron's charge to mass. Thomson's experiments confirmed the conclusion that electrons are part of atoms.

Thus, based on all the experimental facts known by the beginning of the 20th century, it was possible to conclude that the atoms of matter have a complex internal structure. They are electrically neutral systems, and the carriers of the negative charge of the atoms are light electrons, the mass of which is only a small fraction of the mass of the atoms. The bulk of the mass of atoms is associated with a positive charge.

The first direct experiments to study the internal structure of atoms were carried out by E. Rutherford and his collaborators E. Marsden and H. Geiger in 1909–1911. Rutherford proposed using atomic probing using α-particles, which arise during the radioactive decay of radium and some other elements. The mass of alpha particles is approximately 7300 times the mass of an electron, and the positive charge is equal to twice the elementary charge. In his experiments, Rutherford used α-particles with a kinetic energy of about 5 MeV (the speed of such particles is very high - about 10 7 m/s, but still significantly less than the speed of light). α particles are fully ionized helium atoms. They were discovered by Rutherford in 1899 while studying the phenomenon of radioactivity. Rutherford bombarded atoms of heavy elements (gold, silver, copper, etc.) with these particles. The electrons that make up the atoms, due to their low mass, cannot noticeably change the trajectory of the α particle. Scattering, that is, a change in the direction of motion of α-particles, can only be caused by the heavy, positively charged part of the atom. The diagram of Rutherford's experiment is shown in Fig. 6.1.2.

From a radioactive source enclosed in a lead container, alpha particles were directed onto a thin metal foil. Scattered particles fell on a screen covered with a layer of zinc sulfide crystals, capable of glowing when hit by fast charged particles. Scintillations (flashes) on the screen were observed by eye using a microscope. Observations of scattered α particles in Rutherford's experiment could be carried out at different angles φ to the original direction of the beam. It was found that most α particles pass through a thin layer of metal with little or no deflection. However, a small part of the particles are deflected at significant angles exceeding 30°. Very rare alpha particles (about one in ten thousand) were deflected at angles close to 180°.

This result was completely unexpected even for Rutherford. His ideas found bcm in sharp contradiction with Thomson's model of the atom, according to which the positive charge is distributed throughout the entire volume of the atom. With such a distribution, the positive charge cannot create a strong electric field that can throw α particles back. The electric field of a uniform charged ball is maximum on its surface and decreases to zero as it approaches the center of the ball. If the radius of the ball in which all the positive charge of the atom is concentrated decreased by a factor of n, then the maximum repulsive force acting on an α-particle, according to Coulomb’s law, would increase by a factor of n 2. Consequently, for a sufficiently large value of n, alpha particles could experience scattering at large angles up to 180°. These considerations led Rutherford to the conclusion that the atom is almost empty, and all its positive charge is concentrated in a small volume. Rutherford called this part of the atom the atomic nucleus. This is how the nuclear model of the atom arose. Rice. 6.1.3 illustrates the scattering of an α particle in a Thomson atom and in a Rutherford atom.

Figure 6.1.3. Scattering of an α particle in a Thomson atom (a) and in a Rutherford atom (b)

Thus, the experiments of Rutherford and his colleagues led to the conclusion that at the center of the atom there is a dense, positively charged nucleus, the diameter of which does not exceed 10–14–10–15 m. This nucleus occupies only 10–12 of the total volume of the atom, but contains the entire positive charge and at least 99.95% of its mass. The substance constituting the nucleus of the atom should have been assigned a colossal density of the order of ρ ≈ 10 15 g/cm 3 . The charge of the nucleus must be equal to the total charge of all the electrons that make up the atom. Subsequently, it was possible to establish that if the charge of an electron is taken as one, then the charge of the nucleus is exactly equal to the number of a given element in the periodic table.

The radical conclusions about the structure of the atom that followed from Rutherford's experiments forced many scientists to doubt their validity. Rutherford himself was no exception, publishing the results of his research only in 1911, two years after the first experiments were performed. Based on classical ideas about the movement of microparticles, Rutherford proposed a planetary model of the atom. According to this model, at the center of the atom there is a positively charged nucleus, in which almost the entire mass of the atom is concentrated. The atom as a whole is neutral. Electrons rotate around the nucleus, like planets, under the influence of Coulomb forces from the nucleus (Fig. 6.1.4). Electrons cannot be at rest, since they would fall onto the nucleus.

Danish physicist Niels Bohr (1885–1962) created the first quantum theory of the atom in 1913, linking into a single whole the empirical laws of the line spectra of hydrogen, the Rutherford nuclear model of the atom and the quantum nature of the emission and absorption of light.

Bohr based his theory on three postulates, about which the American physicist L. Cooper remarked: “Of course, it was somewhat arrogant to put forward proposals that contradicted Maxwell’s electrodynamics and Newton’s mechanics, but Bohr was young.”

The first postulate (postulate of stationary states): in an atom, electrons can move only along certain, so-called allowed, or stationary, circular orbits, in which they, despite the presence of acceleration, do not emit electromagnetic waves (therefore, these orbits are called stationary). An electron in each stationary orbit has a certain energy En.

Second postulate (frequency rule): an atom emits or absorbs a quantum of electromagnetic energy when an electron moves from one stationary orbit to another:

hv = E 1 – E 2,

where E 1 and E 2 are the electron energy before and after the transition, respectively.

When E 1 > E 2, quantum emission occurs (the transition of an atom from one state with higher energy to a state with lower energy, that is, the transition of an electron from any orbit far away to any orbit close to the nucleus); at E 1< E 2 – поглощение кванта (переход атома в состояние с большей энергией, то есть переход электрона на более удаленную от ядра орбиту).

Convinced that Planck's constant should play a major role in atomic theory, Bohr introduced the third postulate (quantization rule): in stationary orbits the angular momentum of the electron

m en r n = nh, n = 1, 2, 3, …,

where = 1.05 · 10 -34 e = 9.1 · 10 -31 kg – electron mass; r p n is the speed of the electron in this orbit.

Wave-particle duality. De Broglie's hypothesis. Electron diffraction. Wave properties of matter. Corpuscular-wave nature of microparticles. Heisenberg uncertainty relation.

Particle-wave dualism (from the Latin dualis - dual) is the most important universal property of nature, which consists in the fact that each micro-object has both corpuscular and wave characteristics.

For example, an electron, a neutron, a photon, in some conditions, behave like particles that move along classical trajectories and have a certain energy and momentum, and in others, they reveal their wave nature, which is characteristic of the phenomena of interference and diffraction of particles.

The wave-particle duality was previously defined for light. The propagation of light as a stream of photons and the quantum nature of the interaction of light with matter are confirmed by numerous experiments. But a number of optical phenomena (interference, polarization, diffraction) undeniably indicate the wave properties of light.

Classical physics has always clearly distinguished objects that have a wave nature (for example, light and sound), and objects that have a discrete corpuscular structure (for example, systems of material points). One of the most important achievements of modern physics is the conviction that the opposition between wave and quantum properties of light is false. If we consider light as a stream of photons, and photons as quanta of electromagnetic radiation, which at the same time have both wave and corpuscular properties, modern physics can combine antagonistic theories - wave and corpuscular. As a result, the idea of ​​wave-particle duality was created, which underlies modern physics (wave-particle dualism turns out to be the primary principle of quantum mechanics and quantum field theory).

A quantum of light is neither a wave nor a corpuscle in Newton's understanding. Photons are specific microparticles in which energy and momentum (unlike ordinary material points) are expressed using material characteristics - frequency and wavelength.

In 1924, the French scientist Louis de Broglie voiced the hypothesis that wave-particle duality is inherent in each and every type of matter - electrons, protons, atoms, and the quantitative relationships between the wave and corpuscular properties of particles are the same as those previously established for photons. That is, if a particle has energy E and momentum, the absolute value of which is equal to p, then a wave with frequency v=E/h and length is associated with this particle

where h - in this case is Planck's constant.

This is the famous de Broglie formula - one of the most important formulas in the physics of the microworld.

It is worth noting that the de Broglie wavelength decreases with increasing particle mass m and its velocity v: for particles with true .

Thus, a particle with a mass of 1 g, which moves at a speed of 1 m/s, corresponds to a de Broglie wave of length , so small that it is impossible to observe. Therefore, wave properties are unimportant in the mechanics of macroscopic bodies, which is completely consistent with the correspondence principle.

In the 20s of the 20th century, it was established that any particle has a particle-wave nature. According to the theory of L. de Broglie (1924), each particle with momentum corresponds to a wave process with wavelength λ, i.e. λ = h/p. The smaller the particle mass, the longer the wavelength. For elementary particles, W. Heisenberg formulated the uncertainty principle, according to which it is impossible to simultaneously determine the position of a particle in space and its momentum. Consequently, it is impossible to calculate the trajectory of an electron in the field of a nucleus; one can only estimate the probability of its presence in the atom using the wave function ψ, which replaces the classical concept of trajectory. The wave function ψ characterizes the amplitude of the wave depending on the coordinates of the electron, and its square ψ 2 determines the spatial distribution of the electron in the atom. In the simplest version, the wave function depends on three spatial coordinates and makes it possible to determine the probability of finding an electron in atomic space or its orbital. Thus, an atomic orbital (AO) is a region of atomic space in which the probability of finding an electron is greatest.

Wave functions are obtained by solving the fundamental relation of wave mechanics - the Schrödinger equation. An exact solution exists for the hydrogen atom or hydrogen-like ions; various approximations are used for multielectron systems. The surface that limits the probability of finding an electron or electron density to 90–95% is called the boundary surface. The atomic orbital and electron cloud density have the same boundary surface (shape) and the same spatial orientation. The atomic orbitals of an electron, their energy and direction in space depend on four parameters - quantum numbers.

Schrödinger equation. Wave function and its statistical meaning. Application of the Schrödinger equation: a particle in a one-dimensional infinitely deep potential well; passage of particles through a potential barrier.