Atomic physics branch of physics that studies the structure and state of atoms. A. f. originated in the late 19th and early 20th centuries. In the 10s. 20th century It was found that the atom consists of a nucleus and electrons connected by electrical forces. At the first stage of the development A. f. also covered issues related to the structure of the atomic nucleus. In the 30s. it turned out that the nature of the interactions that take place in the atomic nucleus is different than in the outer shell of the atom, and in the 40s. nuclear physics emerged as an independent field of science. In the 50s. elementary particle physics, or high-energy physics, spun off from it. Prehistory of atomic physics: the doctrine of atoms in the 17th-19th centuries. The idea of the existence of atoms as indivisible particles of matter arose in antiquity; The ideas of atomism were first expressed by the ancient Greek thinkers Democritus and Epicurus. In the 17th century they were revived by the French philosopher P. Gassendi and the English chemist R. Boyle. The ideas about atoms that prevailed in the 17th and 18th centuries were poorly defined. Atoms were considered absolutely indivisible and immutable solid particles, different types of which differ from each other in size and shape. Combinations of atoms in one order or another form various bodies, the movements of atoms determine all phenomena occurring in matter. I. Newton, M. V. Lomonosov and some other scientists believed that atoms can interlock into more complex particles - “corpuscles”. However, atoms were not assigned specific chemical and physical properties. Atomistics still had an abstract, natural-philosophical character. At the end of the 18th - beginning of the 19th centuries. as a result of the rapid development of chemistry, the basis for the quantitative development of atomic science was created. The English scientist J. Dalton for the first time (1803) began to consider the atom as the smallest particle of a chemical element, which differs from the atoms of other elements in its mass. According to Dalton, the main characteristic of an atom is its atomic mass. Chemical compounds are a collection of "composite atoms" containing certain (characteristic for a given complex substance) number of atoms of each element. All chemical reactions are just rearrangements of atoms into new complex particles. Based on these provisions, Dalton formulated his law of multiple ratios (see. Multiple ratios law). The studies of the Italian scientists A. Avogadro (1811) and, in particular, S. Cannizzaro (1858) drew a clear line between the atom and the molecule. In the 19th century along with the chemical properties of atoms, their optical properties were studied. It was found that each element has a characteristic optical spectrum; spectral analysis was discovered (German physicists G. Kirchhoff and R. Bunsen, 1860). Thus, the atom appeared as a qualitatively unique particle of matter, characterized by strictly defined physical and chemical properties. But the properties of the atom were considered eternal and inexplicable. It was believed that the number of types of atoms (chemical elements) was random and that there was no connection between them. However, it gradually became clear that there are groups of elements that have the same chemical properties - the same maximum valency, and similar laws of change (when moving from one group to another) of physical properties - melting point, compressibility, etc. In 1869, D. I. Mendeleev discovered the periodic system of elements (See Periodic system of elements). He showed that as the atomic mass of the elements increases, their chemical and physical properties repeat periodically ( rice. one
and 2
). The periodic system proved the existence of a connection between different types of atoms. The conclusion was that the atom has a complex structure that changes with atomic mass. The problem of revealing the structure of the atom has become the most important in chemistry and physics (for more details, see Atomism).
The emergence of atomic physics. The most important developments in science, from which atomic physics originated, were the discoveries of the electron and radioactivity. When studying the passage of electric current through highly rarefied gases, rays were discovered emitted by the cathode of a discharge tube (cathode rays) and having the property of deflecting in transverse electric and magnetic fields. It turned out that these beams consist of rapidly flying negatively charged particles called electrons. In 1897, the English physicist J. J. Thomson measured the charge ratio e of these particles to their mass m. Metals have also been found to emit electrons when heated strongly or illuminated with short wavelength light (see Thermionic emission, Photoelectron emission). From this it was concluded that electrons are part of any atoms. It further followed from this that neutral atoms must also contain positively charged particles. Positively charged atoms - ions - were indeed discovered in the study of electrical discharges in rarefied gases. The idea of an atom as a system of charged particles explained, according to the theory of the Dutch physicist H. Lorenz, a ,
the very possibility of radiation by an atom of light (electromagnetic waves): electromagnetic radiation occurs when intraatomic charges fluctuate; this was confirmed by studying the effect of a magnetic field on atomic spectra (see the Zeeman phenomenon). It turned out that the ratio of the charge of intraatomic electrons to their mass e/m, found by Lorentz in his theory of the Zeeman phenomenon is exactly equal to the value e/m for free electrons obtained in Thomson's experiments. The theory of electrons and its experimental confirmation gave indisputable proof of the complexity of the atom. The idea of the indivisibility and intransmutability of the atom was finally refuted by the works of the French scientists M. Sklodowska-Curie and P. Curie. .
As a result of the study of radioactivity, it was established (F. Soddy) ,
that atoms undergo transformations of two types. Having emitted an α-particle (a helium ion with a positive charge of 2 e), an atom of a radioactive chemical element turns into an atom of another element located 2 cells to the left in the periodic system, for example, a polonium atom into a lead atom. Having emitted a β-particle (electron) with a negative charge - e, an atom of a radioactive chemical element turns into an atom of an element located 1 cell to the right, for example, a bismuth atom into a polonium atom. The mass of an atom formed as a result of such transformations sometimes turned out to be different from the atomic weight of the element into whose cell it fell. From this followed the existence of varieties of atoms of the same chemical element with different masses; these varieties were later called isotopes (i.e., occupying the same place in the periodic table). So, the ideas about the absolute identity of all atoms of a given chemical element turned out to be wrong. The results of the study of the properties of the electron and radioactivity made it possible to build specific models of the atom. In the model proposed by Thomson in 1903, the atom was represented as a positively charged sphere, in which negative electrons, insignificant in size (compared to the atom), are interspersed ( rice. 3
). They are held in the atom due to the fact that the attractive forces of their distributed positive charge are balanced by the forces of their mutual repulsion. The Thomson model provided a well-known explanation for the possibility of emission, scattering, and absorption of light by an atom. When the electrons are displaced from the equilibrium position, an "elastic" force arises, striving to restore the equilibrium; this force is proportional to the displacement of the electron from the equilibrium position and, therefore, to the dipole moment (See Dipole moment)
atom. Under the action of the electric forces of the incident electromagnetic wave, the electrons in the atom oscillate with the same frequency as the electric intensity in the light wave; the oscillating electrons, in turn, emit light of the same frequency. This is how electromagnetic waves are scattered by atoms of matter. By the degree of attenuation of the light beam in the thickness of the substance, you can find out the total number of scattering electrons, and knowing the number of atoms per unit volume, you can determine the number of electrons in each atom. Creation by Rutherford of the planetary model of the atom. Thomson's model of the atom turned out to be unsatisfactory. On its basis, it was not possible to explain the completely unexpected result of the experiments of the English physicist E. Rutherford and his collaborators H. Geiger and E. Marsden on the scattering of α-particles by atoms. In these experiments, fast α-particles were used for direct probing of atoms. Passing through matter, α-particles collide with atoms. With each collision, the α-particle, flying through the electric field of the atom, changes the direction of motion - it experiences scattering. In the overwhelming majority of scattering events, the deviations of the α-particles (scattering angles) were very small. Therefore, during the passage of a beam of α-particles through a thin layer of matter, only a slight blurring of the beam occurred. However, a very small proportion of α-particles were deflected through angles greater than 90°. This result could not be explained on the basis of the Thomson model, because the electric field in a "solid" atom is not strong enough to deflect a fast and massive α-particle through a large angle. To explain the results of experiments on the scattering of α-particles, Rutherford proposed a fundamentally new model of the atom, reminiscent in structure of the solar system and called the planetary one. It has the following form. In the center of the atom is a positively charged nucleus, the dimensions of which (Atomic Physics10 -12 cm) are very small compared to the size of an atom (Atomic Physics10 -8 cm), and the mass is almost equal to the mass of the atom. Electrons move around the nucleus, like planets around the sun; the number of electrons in an uncharged (neutral) atom is such that their total negative charge compensates (neutralizes) the positive charge of the nucleus. Electrons must move around the nucleus, otherwise they would fall on it under the influence of attractive forces. The difference between the atom and the planetary system is that in the latter, gravitational forces act, and in the atom, electrical (Coulomb) forces. Near the nucleus, which can be considered as a point positive charge, there is a very strong electric field. Therefore, flying near the nucleus, positively charged α-particles (helium nuclei) experience a strong deflection (see Fig. rice. 4
). Later it was found out (G. Moseley) that the charge of the nucleus increases from one chemical element to another by an elementary unit of charge equal to the electron charge (but with a positive sign). Numerically, the charge of the nucleus of an atom, expressed in units of elementary charge e, is equal to the ordinal number of the corresponding element in the periodic system. To test the planetary model, Rutherford and his collaborator Charles Darwin calculated the angular distribution of α-particles scattered by a point nucleus, the center of Coulomb forces. The result obtained was verified experimentally by measuring the number of α-particles scattered at different angles. The results of the experiment exactly matched the theoretical calculations, thus brilliantly confirming Rutherford's planetary model of the atom. However, the planetary model of the atom ran into fundamental difficulties. According to classical electrodynamics, a charged particle moving with acceleration continuously radiates electromagnetic energy. Therefore, electrons, moving around the nucleus, i.e., accelerated, would have to continuously lose energy to radiation. But at the same time, in a tiny fraction of a second, they would lose all their kinetic energy and fall into the core. Another difficulty, also associated with radiation, was as follows: if we accept (in accordance with classical electrodynamics) that the frequency of the light emitted by an electron is equal to the frequency of oscillations of an electron in an atom (i.e., the number of revolutions it makes in its orbit in one second) or has a multiple of it, then the emitted light, as the electron approaches the nucleus, would have to continuously change its frequency, and the spectrum of the light emitted by it should be continuous. But this is contrary to experience. An atom emits light waves of well-defined frequencies, typical for a given chemical element, and is characterized by a spectrum consisting of separate spectral lines - a line spectrum. A number of regularities were experimentally established in the line spectra of elements, the first of which was discovered by the Swiss scientist I. Balmer (1885) in the spectrum of hydrogen. The most general pattern - the combination principle - was found by the Austrian scientist W. Ritz (1908). This principle can be formulated as follows: for the atoms of each element, one can find a sequence of numbers T 1 ,T 2 ,T 3 ,... - so-called. spectral terms such that the frequency v each spectral line of a given element is expressed as the difference of two terms: v = T k - T i .
For a hydrogen atom, the term T n = R/n 2 , where n- an integer that takes a value n= 1, 2, 3,..., a R- so-called. Rydberg constant (see Rydberg constant).
Thus, within the framework of Rutherford's model of the atom, the stability of the atom with respect to radiation and the line spectra of its radiation could not be explained. On its basis, the laws of thermal radiation and the laws of photoelectric phenomena that arise when radiation interacts with matter could not be explained. It turned out to be possible to explain these laws on the basis of completely new - quantum - concepts, first introduced by the German physicist M. Planck (1900). To derive the law of energy distribution in the spectrum of thermal radiation - the radiation of heated bodies - Planck suggested that the atoms of matter emit electromagnetic energy (light) in the form of separate portions - light quanta, the energy of which is proportional to v(radiation frequency): E = hv where h- a constant characteristic of quantum theory and called the Planck constant (See Planck constant). In 1905, A. Einstein gave a quantum explanation of photoelectric phenomena, according to which the quantum energy hv goes to extract an electron from the metal - work function R - and to communicate to him kinetic energy T kin; hv = R+ Tkin. At the same time, Einstein introduced the concept of light quanta as a special kind of particles; these particles subsequently received the name Photon ov.
It turned out to be possible to resolve the contradictions of Rutherford's model only by abandoning a number of the usual ideas of classical physics. The most important step in the construction of the theory of the atom was made by the Danish physicist N. Bohr (1913). Bohr's postulates and the Bohr model of the atom. In the basis of the quantum theory of the atom, Bohr put 2 postulates characterizing those properties of the atom that did not fit into the framework of classical physics. These postulates of Bohr can be formulated as follows: 1. Existence of stationary states. An atom does not radiate and is stable only in some stationary (invariant in time) states corresponding to a discrete (discontinuous) series of "permissible" energy values E 1 , E 2 , E 3 ,
E 4 ,... Any change in energy is associated with a quantum (jump-like) transition from one stationary state to another. 2. Condition of radiation frequencies (quantum transitions with radiation). Upon transition from one stationary state with energy E i into another with energy E k an atom emits or absorbs light of a certain frequency v in the form of a radiation quantum (photon) hv, according to the ratio hv=E i - E k . When emitted, an atom passes from a state of higher energy E i to a state of lower energy E k , upon absorption, on the contrary, from a state with a lower energy E k to a higher energy state E i .
Bohr's postulates immediately make it possible to understand the physical meaning of the Ritz combination principle (see above); ratio comparison hv = E i - E k and v = T k - T i
shows that the spectral terms correspond to stationary states, and the energy of the latter must equal (up to a constant term) E i = -hT i , E k = -hT k .
When light is emitted or absorbed, the energy of the atom changes, this change is equal to the energy of the emitted or absorbed photon, i.e., the law of conservation of energy takes place. The line spectrum of an atom is the result of the discreteness of the possible values of its energy. Bohr applied classical (Newtonian) mechanics to determine the allowed values of the energy of an atom - the quantization of its energy - and to find the characteristics of the corresponding stationary states. “If we want to make a visual representation of stationary states in general, we have no other means, at least now, except for ordinary mechanics,” Bohr wrote in 1913 (“Three articles on spectra and the structure of atoms”, M.-L., 1923, p. 22). For the simplest atom - a hydrogen atom, consisting of a nucleus with a charge + e(proton) and an electron with a charge - e, Bohr considered the motion of an electron around the nucleus in circular orbits. Comparing the energy of an atom E with spectral terms T n \u003d R / n 2 for the hydrogen atom, found with great accuracy from the frequencies of its spectral lines, he obtained the possible values of the energy of the atom E n= -hT n \u003d -hR / n 2(where n= 1, 2, 3,...). They correspond to circular orbits of radius a n \u003d a 0 n 2, where a 0 = 0.53 10 -8 cm - Bohr radius - the radius of the smallest circular orbit (at n= 1). Bohr calculated the revolution frequencies v electron around the nucleus in circular orbits depending on the energy of the electron. It turned out that the frequencies of the light emitted by the atom do not coincide with the frequencies of revolution v n , as required by classical electrodynamics, but are proportional, according to the relation hv=E i - E k , the energy difference of an electron in two possible orbits. To find the relationship between the frequency of an electron's orbit and the radiation frequency, Bohr made the assumption that the results of quantum and classical theories should coincide at low radiation frequencies (for long wavelengths; such a coincidence takes place for thermal radiation, the laws of which were derived by Planck). He equated for big n transition frequency v = (E n+1 - E n)/ h circulation frequency v n in orbit with given n and calculated the value of the Rydberg constant R, which coincided with great accuracy with the value R, found from experience, which confirmed Bohr's assumption. Bohr also succeeded not only in explaining the spectrum of hydrogen, but also convincingly showing that some of the spectral lines that were attributed to hydrogen belong to helium. Bohr's assumption that the results of quantum and classical theories should coincide in the limiting case of low radiation frequencies represented the original form of the so-called. the principle of conformity. Later, Bohr successfully applied it to find the intensities of the lines of the spectrum. As the development of modern physics has shown, the correspondence principle turned out to be very general (see Correspondence principle) .
In Bohr's theory of the atom, the quantization of energy, i.e., finding its possible values, turned out to be a special case of the general method for finding "allowed" orbits. According to quantum theory, such orbits are only those for which the angular momentum of an electron in an atom is equal to an integer multiple h/2π. Each permitted orbit corresponds to a certain possible value of the energy of an atom (see Atom).
The main provisions of the quantum theory of the atom - Bohr's 2 postulates - were comprehensively confirmed experimentally. Particularly clear confirmation was given by the experiments of the German physicists J. Frank and G. Hertz (1913-16). The essence of these experiences is as follows. A stream of electrons whose energy can be controlled enters a vessel containing mercury vapor. The electrons are given energy, which gradually increases. As the energy of the electrons increases, the current in the galvanometer included in the electrical circuit increases; when the electron energy turns out to be equal to certain values (4.9; 6.7; 10.4 ev), the current drops sharply ( rice. 5
). At the same time, it can be found that mercury vapor emits ultraviolet rays of a certain frequency. The facts presented allow only one interpretation. As long as the electron energy is less than 4.9 ev, electrons do not lose energy when colliding with mercury atoms - collisions are elastic in nature. When the energy turns out to be equal to a certain value, namely 4.9 ev, electrons transfer their energy to mercury atoms, which then emit it in the form of ultraviolet light quanta. The calculation shows that the energy of these photons is equal to exactly the energy that the electrons lose. These experiments proved that the internal energy of an atom can only have certain discrete values, that the atom absorbs energy from the outside and emits it at once in whole quanta, and that, finally, the frequency of the light emitted by the atom corresponds to the energy lost by the atom. Further development of A. f. showed the validity of Bohr's postulates not only for atoms, but also for other microscopic systems - for molecules and for atomic nuclei. These postulates should be regarded as firmly established experimental quantum laws. They constitute that part of Bohr's theory, which was not only preserved during the further development of quantum theory, but also received its substantiation. The situation is different with Bohr's model of the atom, which is based on the consideration of the motion of electrons in an atom according to the laws of classical mechanics with the imposition of additional quantization conditions. This approach made it possible to obtain a number of important results, but was inconsistent: the quantum postulates were artificially attached to the laws of classical mechanics. A consistent theory was created in the 20s. 20th century Quantum mechanics . Its creation was prepared by the further development of the model representations of Bohr's theory, during which its strengths and weaknesses became clear. Development of the model theory of the Bohr atom. A very important result of Bohr's theory was the explanation of the spectrum of the hydrogen atom. A further step in the development of the theory of atomic spectra was made by the German physicist A. Sommerfeld. Having developed the quantization rules in more detail, based on a more complex picture of the motion of electrons in an atom (along elliptical orbits) and taking into account the screening of an external (so-called valence) electron in the field of the nucleus and internal electrons, he was able to explain a number of regularities in the spectra of alkali metals. Bohr's theory of the atom also shed light on the structure of the so-called. characteristic spectra of x-rays. The X-ray spectra of atoms, as well as their optical spectra, have a discrete line structure characteristic of a given element (hence the name). Investigating the characteristic X-ray spectra of various elements, the English physicist G. Moseley discovered the following pattern: the square roots of the frequencies of the emitted lines increase uniformly from element to element throughout the Mendeleev periodic system in proportion to the atomic number of the element. It is interesting that Moseley's law fully confirmed the correctness of Mendeleev, who in some cases violated the principle of placing elements in the table according to increasing atomic weight and put some heavier elements ahead of lighter ones. On the basis of Bohr's theory, it was possible to give an explanation of the periodicity of the properties of atoms. In a complex atom, electron shells are formed, which are successively filled, starting from the innermost, with certain numbers of electrons (the physical reason for the formation of shells became clear only on the basis of the Pauli principle, see below). The structure of the outer electron shells is periodically repeated, which causes the periodic repetition of the chemical and many physical properties of elements located in the same group of the periodic system. On the basis of Bohr's theory, the German chemist W. Kossel (1916) explained the chemical interaction in the so-called. heteropolar molecules. However, not all questions of the theory of the atom could be explained on the basis of model representations of Bohr's theory. It did not cope with many problems of the theory of spectra, it allowed to obtain only the correct values of the frequencies of the spectral lines of the hydrogen atom and hydrogen-like atoms, while the intensities of these lines remained unexplained; Bohr had to apply the correspondence principle to explain the intensities. In the transition to explaining the motions of electrons in atoms more complex than the hydrogen atom, Bohr's model theory was at an impasse. Already a helium atom, in which 2 electrons move around the nucleus, did not lend itself to theoretical interpretation based on it. Difficulties in this case were not limited to quantitative discrepancies with experience. The theory turned out to be powerless in solving such a problem as the combination of atoms into a molecule. Why do 2 neutral hydrogen atoms combine to form a hydrogen molecule? How to explain the nature of valency in general? What binds the atoms of a solid? These questions remained unanswered. Within the framework of the Bohr model, it was impossible to find an approach to their solution. Quantum mechanical theory of the atom. The limitations of Bohr's model of the atom were rooted in the limitations of classical ideas about the motion of microparticles. It became clear that for the further development of the theory of the atom, it is necessary to critically reconsider the basic ideas about the motion and interaction of microparticles. The unsatisfactory nature of a model based on classical mechanics with the addition of quantization conditions was clearly understood by Bohr himself, whose views exerted a great influence on the further development of algebraic functions. The beginning of a new stage in the development of A. f. was the idea expressed by the French physicist L. de Broglie (1924) about the dual nature of the motion of micro-objects, in particular the electron (see De Broglie waves). This idea became the starting point of quantum mechanics (see Quantum Mechanics), created in 1925–26 by the works of W. Heisenberg and M. Born (Germany), E. Schrödinger (Austria), and P. Dirac (England), and developed on its basis modern quantum mechanical theory of the atom. The ideas of quantum mechanics about the motion of an electron (microparticles in general) differ radically from the classical ones. According to quantum mechanics, an electron does not move along a trajectory (orbit), like a solid ball; The motion of an electron also has certain features characteristic of the propagation of waves. On the one hand, an electron always acts (for example, in collisions) as a single whole, as a particle with an indivisible charge and mass; at the same time, electrons with a certain energy and momentum propagate like a plane wave with a certain frequency (and a certain wavelength). Electron energy E how particles are related to frequency v electron wave ratio: E=hv, and its momentum R - with wavelength λ
ratio: p = h/λ.
The stable motions of an electron in an atom, as shown by Schrödinger (1926), are in some respects analogous to standing waves (See standing waves) ,
whose amplitudes are different at different points. At the same time, in the atom, as in an oscillatory system, only some “selected” movements are possible with certain values of energy, angular momentum and projection of the electron momentum in the atom. Each stationary state of an atom is described using some wave function (See Wave function) ,
which is a solution of a wave equation of a special type - the Schrödinger equation; wave function corresponds to the "electron cloud", which characterizes (on average) the distribution of the electron charge density in the atom (see Atom ,
there on rice. 3
projections of the "electron clouds" of the hydrogen atom are shown). In the 20-30s. Approximate methods were developed for calculating the distribution of the electron charge density in complex atoms, in particular the Thomas-Fermi method (1926, 1928). This value and the associated value of the so-called. atomic factor (See atomic factor)
important in the study of electron collisions with atoms, as well as their scattering of x-rays. On the basis of quantum mechanics, it was possible to correctly calculate the energies of electrons in complex atoms by solving the Schrödinger equation. Approximate methods for such calculations were developed in 1928 by D. Hartree (England) and in 1930 by V. A. Fok (USSR). Studies of atomic spectra fully confirmed the quantum mechanical theory of the atom. It turned out that the state of an electron in an atom essentially depends on its Spin a -
own mechanical moment of momentum. An explanation was given for the action of external electric and magnetic fields on the atom (see Stark phenomenon (See Stark effect), Zeeman phenomenon). An important general principle related to the electron spin was discovered by the Swiss physicist W. Pauli (1925) (see Pauli principle), according to this principle, only one electron can be in each electronic state in an atom; if this state is already occupied by some electron, then the next electron, entering the atom, is forced to occupy another state. On the basis of the Pauli principle, the filling numbers of electron shells in complex atoms were finally established, which determine the periodicity of the properties of elements. Based on quantum mechanics, the German physicists W. Geytler and F. London (1927) gave the theory of the so-called. homeopolar chemical bond of two identical atoms (for example, hydrogen atoms in the H 2 molecule), which cannot be explained within the framework of the Bohr model of the atom. Important applications of quantum mechanics in the 30s. and later there were studies of bound atoms that make up a molecule or crystal. The states of an atom that is part of a molecule are essentially different from the states of a free atom. The atom also undergoes significant changes in a crystal under the action of an intracrystalline field, the theory of which was first developed by H. Bethe (1929). Investigating these changes, one can establish the nature of the interaction of the atom with its environment. The largest experimental achievement in this area is A. f. was the discovery by E. K. Zavoisky in 1944 of electron paramagnetic resonance (See Electron paramagnetic resonance) ,
which made it possible to study the various bonds of atoms with the environment. Modern atomic physics. The main sections of modern A. f. are the theory of the atom, atomic (optical) spectroscopy, X-ray spectroscopy, radio spectroscopy (it also investigates the rotational levels of molecules), and the physics of atomic and ion collisions. Different sections of spectroscopy cover different ranges of radiation frequencies and, accordingly, different ranges of photon energies. While X-ray spectroscopy studies the radiation of atoms with photon energies up to hundreds of thousands of electrons. ev, radio spectroscopy deals with very small quanta - up to quanta less than 10 -6 ev.
The most important task of A. f. - detailed definition of all characteristics of states of an atom. We are talking about determining the possible values of the energy of an atom - its energy levels, the values of the moments of momentum and other quantities that characterize the state of the atom. Fine and hyperfine structures of energy levels are studied (see Atomic Spectra) ,
changes in energy levels under the influence of electric and magnetic fields - both external, macroscopic, and internal, microscopic. Of great importance is such a characteristic of the states of the atom as the lifetime of an electron at the energy level. Finally, much attention is paid to the mechanism of excitation of atomic spectra. The areas of phenomena studied by different sections of the AF overlap. X-ray spectroscopy by measuring the emission and absorption of X-rays makes it possible to determine mainly the binding energies of internal electrons with the nucleus of an atom (ionization energy), the distribution of the electric field inside the atom. Optical spectroscopy studies the sets of spectral lines emitted by atoms, determines the characteristics of the energy levels of the atom, the intensities of the spectral lines and the lifetimes of the atom in excited states associated with them, the fine structure of the energy levels, their displacement and splitting in electric and magnetic fields. Radio spectroscopy investigates in detail the width and shape of spectral lines, their hyperfine structure, shift and splitting in a magnetic field, and, in general, intra-atomic processes caused by very weak interactions and influences of the medium. Analysis of the results of collisions of fast electrons and ions with atoms makes it possible to obtain information about the distribution of the electron charge density ("electron cloud") inside the atom, about the excitation energies of the atom, and ionization energies. The results of a detailed study of the structure of atoms find the widest application not only in many branches of physics, but also in chemistry, astrophysics and other fields of science. Based on the study of the broadening and shift of spectral lines, one can judge the local (local) fields in the medium (liquid, crystal) that cause these changes, and the state of this medium (temperature, density, etc.). Knowing the distribution of the electron charge density in an atom and its changes during external interactions makes it possible to predict the type of chemical bonds that an atom can form, the behavior of an ion in a crystal lattice. Information about the structure and characteristics of the energy levels of atoms and ions is extremely important for quantum electronics devices.
The special theory of relativity (SRT) is based on two postulates:
- The principle of relativity: in any inertial reference frames, all physical phenomena under the same initial conditions proceed in the same way, i.e. no experiments carried out in a closed system of bodies can reveal whether the body is at rest or moves uniformly and rectilinearly.
- The principle of constancy of the speed of light: in all inertial frames of reference the speed of light in vacuum is the same and does not depend on the speed of the moving light source.
Equal to the postulates of SRT, the position of SRT on the limiting nature of the speed of light in vacuum matters: the speed of any signal in nature cannot exceed the speed of light in vacuum: c= 3∙10 8 m/s. When objects move at a speed comparable to the speed of light, various effects are observed, described below.
1. Relativistic length contraction.
The length of a body in the reference frame where it is at rest is called its own length. L 0 . Then the length of the body moving with speed V in the inertial reference frame decreases in the direction of motion to a length:
where: c is the speed of light in vacuum, L 0 is the length of the body in a fixed frame of reference (the length of a body at rest), L is the length of the body in the frame of reference moving with the speed V(length of a body moving at a speed V). Thus, body length is relative. The reduction of bodies is noticeable only at speeds comparable to the speed of light.
2. Relativistic lengthening of the event time.
The duration of a phenomenon occurring at a certain point in space will be the smallest in that inertial frame of reference, relative to which this point is stationary. This means that clocks moving relative to an inertial frame of reference run slower than stationary clocks and show a longer time interval between events. Relativistic time dilation becomes noticeable only at speeds comparable to the speed of light, and is expressed by the formula:
Time τ 0 , measured by a clock resting relative to the body, is called the proper time of the event.
3. Relativistic law of addition of velocities.
The law of addition of velocities in Newtonian mechanics contradicts the postulates of SRT and is replaced by a new relativistic law of addition of velocities. If two bodies move towards each other, then their speed of approach is expressed by the formula:
where: V 1 and V 2 - speeds of movement of bodies relative to a fixed frame of reference. If the bodies move in the same direction, then their relative speed:
4. Relativistic increase in mass.
Mass of the moving body m greater than the rest mass of the body m 0:
5. Relationship between energy and body mass.
From the point of view of the theory of relativity, the mass of a body and the energy of a body are practically the same thing. Thus, only the fact of the existence of a body means that the body has energy. Least Energy E 0 the body has in the inertial reference frame relative to which it is at rest and is called the body's own energy (rest energy of the body):
Any change in body energy means a change in body mass and vice versa:
where: ∆ E is the change in body energy, ∆ m is the corresponding change in mass. Total body energy:
where: m- body mass. Total body energy E proportional relativistic mass and depends on the speed of the moving body, in this sense the following relations are important:
By the way, the kinetic energy of a body moving at a relativistic speed can only be calculated using the formula:
From the point of view of the theory of relativity, the law of conservation of rest masses is unfair. For example, the rest mass of an atomic nucleus is less than the sum of the rest masses of the particles in the nucleus. However, the rest mass of a particle capable of spontaneous decay is greater than the sum of its own masses of its constituents.
This does not mean a violation of the law of conservation of mass. In the theory of relativity, the law of conservation of relativistic mass is valid, since in an isolated system of bodies the total energy is preserved, and hence the relativistic mass, which follows from the Einstein formula, so we can talk about a single law of conservation of mass and energy. This does not mean that mass can be converted into energy and vice versa.
There is a relationship between the total energy of the body, rest energy and momentum:
Photon and its properties
Light is a stream of quanta of electromagnetic radiation called photons. Photon is a particle that carries the energy of light. It cannot be at rest, but always moves at a speed equal to the speed of light. A photon has the following characteristics:
1. The energy of photons is equal to:
where: h= 6.63∙10 –34 J∙s = 4.14∙10 –15 eV∙s – Planck’s constant, ν is the frequency of the light, λ is the wavelength of the light, c is the speed of light in vacuum. The energy of a photon in Joules is very small, therefore, for mathematical convenience, it is often measured in an off-system unit - electron volts:
1 eV = 1.6∙10 -19 J.
2. A photon travels in a vacuum at the speed of light. c.
3. A photon has momentum:
4. A photon does not have mass in the usual sense for us (the mass that can be measured on scales, calculated according to Newton's second law, and so on), but in accordance with Einstein's theory of relativity, it has mass as a measure of energy ( E = mc 2). Indeed, any body that has some energy also has mass. If we consider that a photon has energy, then it also has a mass, which can be found as:
5. A photon has no electric charge.
Light has a dual nature. When light propagates, its wave properties appear (interference, diffraction, polarization), and when interacting with matter, corpuscular (photoelectric effect). This dual nature of light is called wave-particle duality.
external photoelectric effect
photoelectric effect- a phenomenon consisting in the appearance of a photocurrent in a vacuum bottle when the cathode is illuminated with monochromatic light of a certain wavelength λ .
When the voltage across the anode is negative, the electric field between the cathode and anode slows down the electrons. Measuring the given retarding voltage at which the photocurrent disappears, it is possible to determine the maximum kinetic energy of photoelectrons escaping from the cathode:
Numerous experimenters have established the following basic laws of the photoelectric effect:
- The photoelectric effect is inertialess. This means that electrons begin to fly out of the metal immediately after the start of irradiation with light.
- The maximum kinetic energy of photoelectrons increases linearly with increasing light frequency ν and does not depend on its intensity.
- For every substance there is a so-called red border photo effect, that is, the lowest frequency ν min (or the longest wavelength λ max) at which the external photoelectric effect is still possible.
- The number of photoelectrons pulled out by light from the cathode in 1 s is directly proportional to the light intensity.
When interacting with matter, a photon transfers all of its energy E = hν one electron. Part of this energy can be dissipated by an electron in collisions with atoms of matter. In addition, part of the electron energy is spent on overcoming the potential barrier at the metal–vacuum interface. To do this, the electron must make work function A out, depending on the properties of the cathode material. The highest kinetic energy that a photoelectron emitted from the cathode can have, in this case, is determined by the energy conservation law:
This formula is called Einstein's equation for the external photoelectric effect. Using the Einstein equation, one can explain all the regularities of the external photoelectric effect. For red border photo effect, according to Einstein's formula, we can get the expression:
Bohr's postulates
Bohr's first postulate (stationary state postulate): an atomic system can only be in special stationary or quantum states, each of which corresponds to a certain number n and energy E n. In stationary states, an atom does not emit or absorb energy.
The state with the lowest energy is assigned the number "1". It's called main. All other states are assigned sequential numbers "2", "3", and so on. They're called excited. An atom can remain in its ground state indefinitely. In the excited state, the atom lives for some time (about 10 ns) and passes into the ground state.
According to Bohr's first postulate, an atom is characterized by a system of energy levels, each of which corresponds to a certain stationary state. The mechanical energy of an electron moving along a closed path around a positively charged nucleus is negative. Therefore, all stationary states correspond to the energy values E n < 0. При E n≥ 0 the electron moves away from the nucleus (ionization occurs). Value | E 1 | called ionization energy. State with energy E 1 is called the ground state of the atom.
Bohr's second postulate (frequency rule): during the transition of an atom from one stationary state with energy E n to another stationary state with energy E m a quantum is emitted or absorbed, the energy of which is equal to the difference between the energies of the stationary states:
hydrogen atom
The simplest of the atoms is the hydrogen atom. It contains a single electron. The nucleus of an atom is a proton - a positively charged particle, the charge of which is equal in absolute value to the charge of an electron. Usually, an electron is at the first (main, unexcited) energy level (an electron, like any other system, tends to a state with a minimum of energy). In this state, its energy is E 1 = -13.6 eV. In the hydrogen atom, the following relations are satisfied that relate the radius of the trajectory of an electron rotating around the nucleus, its speed and energy in the first orbit with similar characteristics in other orbits:
On any orbit in a hydrogen atom, the kinetic ( To) and potential ( P) the electron energies are related to the total energy ( E) by the following formulas:
atomic nucleus
At present, it is firmly established that the atomic nuclei of various elements consist of two particles - protons and neutrons, which are usually called nucleons. A number of notations are introduced to characterize atomic nuclei. The number of protons that make up the atomic nucleus is denoted by the symbol Z and is called the charge number or atomic number (this is the serial number in the periodic table of Mendeleev). The number of neutrons is denoted by the symbol N. The total number of nucleons (that is, protons and neutrons) is called the mass number A, for which the following formula can be written:
Communication energy. mass defect
The most important role in nuclear physics is played by the concept nuclear binding energy. The binding energy of the nucleus is equal to the minimum energy that must be expended for the complete splitting of the nucleus into individual particles. It follows from the law of conservation of energy that the binding energy is equal to the energy that is released during the formation of a nucleus from individual particles.
The binding energy of any nucleus can be determined by accurately measuring its mass. Such measurements show that the mass of any nucleus M i is always less than the sum of the masses of its constituent protons and neutrons: M I< Zm p + N m n. The difference between these masses is called mass defect, and is calculated by the formula:
The mass defect can be determined using the Einstein formula E = mc 2 the energy released during the formation of a given nucleus, that is, the binding energy of the nucleus E St:
But it is more convenient to calculate the binding energy using a different formula (here, the masses are taken in atomic units, and the binding energy is obtained in MeV):
Radioactivity. Law of radioactive decay
Almost 90% of known atomic nuclei are unstable. An unstable nucleus spontaneously transforms into other nuclei with the emission of particles. This property of nuclei is called radioactivity.
Alpha decay. Alpha decay is the spontaneous transformation of an atomic nucleus with the number of protons Z and neutrons N into another (daughter) nucleus containing the number of protons Z - 2 and neutrons N - 2. In this case, α -particle - the nucleus of a helium atom 4 2 He. The general scheme of alpha decay:
Beta decay. During beta decay, an electron (0 –1 e) flies out of the nucleus. Scheme of beta decay:
Gamma decay. Unlike α - and β -radioactivity γ -radioactivity of nuclei is not associated with a change in the internal structure of the nucleus and is not accompanied by a change in charge or mass numbers. As with α - as well as β -decay, the daughter nucleus may be in some excited state and have an excess of energy. The transition of the nucleus from the excited state to the ground state is accompanied by the emission of one or more γ -quanta, the energy of which can reach several MeV.
Law of radioactive decay. Any sample of radioactive material contains a huge number of radioactive atoms. Since radioactive decay is random and does not depend on external conditions, the law of decreasing quantity N(t) undecayed to this point in time t nuclei can serve as an important statistical characteristic of the process of radioactive decay. The law of radioactive decay has the form:
Value T called half-life, N 0 is the initial number of radioactive nuclei at t= 0. The half-life is the main quantity that characterizes the rate of radioactive decay. The shorter the half-life, the more intense the decay.
At α - and β In radioactive decay, the daughter nucleus may also be unstable. Therefore, a series of successive radioactive decays are possible, which end in the formation of stable nuclei.
Nuclear reactions
nuclear reaction- this is the process of interaction of an atomic nucleus with another nucleus or elementary particle, accompanied by a change in the composition and structure of the nucleus and the release of secondary particles or γ -quanta. As a result of nuclear reactions, new radioactive isotopes can be formed that are not found on Earth in natural conditions.
In nuclear reactions, several conservation laws are fulfilled: momentum, energy, angular momentum, charge. In addition to these classical conservation laws, nuclear reactions hold the law of conservation of the so-called baryon charge(that is, the number of nucleons - protons and neutrons). For example, in a general reaction:
The following conditions are met (the total number of nucleons before and after the reaction remains unchanged):
Energy yield of a nuclear reaction
Nuclear reactions are accompanied by energy transformations. The energy yield of a nuclear reaction is the value:
where: M A and M B are the masses of initial products, M C and M D are the masses of the final reaction products. Value Δ M called mass defect. Nuclear reactions can proceed with the release ( Q> 0) or with energy absorption ( Q < 0). Во втором случае первоначальная кинетическая энергия исходных продуктов должна превышать величину |Q|, which is called reaction threshold.
In order for a nuclear reaction to have a positive energy yield, the specific binding energy of nucleons in the nuclei of the initial products must be less than the specific binding energy of nucleons in the nuclei of the final products. This means that the value Δ M
Successful, diligent and responsible implementation of these three points will allow you to show an excellent result on the CT, the maximum of what you are capable of.
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2 1. Introduction 1.1. The subject of atomic physics, its brief history of development, goals and objectives 1.2. Basic definitions. Electron, proton, neutron, atom, ion, molecule, nuclide, atomic nucleus, chemical element, isotopes 1.3. Nuclear and shell properties of the atom 1.4. Units of measurement of physical quantities in atomic physics. Electron-volt. Mole, Avogadro's constant, atomic mass unit, relative atomic mass. Scales of energies, lengths, frequencies, masses in atomic and nuclear physics 1.5. Classical, relativistic and quantum physics. Momentum and energy 1.6. Photon. Photon energy scale (electromagnetic radiation scale)
3 Physics of the atom Atomic physics (physics of the atom and atomic phenomena) is a branch of physics that studies the structure and properties of atoms, as well as elementary processes in which atoms take part. The objects of study of atomic physics are both atoms and molecules, atomic and molecular ions, exotic atoms and other microparticles In the phenomena studied within the framework of atomic physics, electromagnetic interactions play the main role. semiconductors and nanomaterials) The theoretical basis of atomic physics itself is quantum theory and quantum electrodynamics There is no clear boundary between atomic physics and other branches of physics, and in accordance with the international classification, atomic physics is included in the field of atomic, molecular physics and optics
4 A brief history of the development of atomic physics The concept of "atom" was used by ancient Greek scientists (5th - 2nd centuries BC) to refer to the smallest, indivisible particles that make up everything that exists in the world. Experimental confirmation of atomistic ideas was obtained in the 19th century in chemical and physical research The idea that an atom consists of positively and negatively charged parts was substantiated in the second half of the 19th century. In 1897, J.J. Thomson discovered the electron, and it was soon proved that it is an integral part of all atoms. The idea of an atom as a system consisting of an atomic nucleus and an electron shell was substantiated by E. physics, nuclear physics stood out and, somewhat later, elementary particle physics
5 A brief history of the development of atomic physics The foundations of modern atomic physics were laid at the beginning of the 20th century, when N. Bohr gave explanations a number of the most important properties of the atom (1913) and put forward two "quantum" postulates According to the first of them, there are special (stationary) states of the atom in which the latter does not radiate energy, although the charged particles (electrons) included in its composition make accelerated motion According to the second postulate , the radiation of an atom occurs during the transition from one stationary state to another, and the frequency ν of this radiation is determined from the condition h = E – E (Bohr's frequency rule), where h is Planck's constant, E and E are the values of the energy of the atom in the initial and final states. First postulate reflects the fact of the stability of the atom, the second discreteness of frequencies in atomic spectra
6 A brief history of the development of atomic physics Bohr's theory, which proved unable to fully explain the properties of atoms and molecules, was replaced by a consistent quantum theory created in the 1920s and 1930s (W. Heisenberg, E. Schrödinger, P. Dirac) Nevertheless, Bohr's postulates still retain their significance and are an integral part of the foundations of the physics of microscopic phenomena. Within the framework of modern quantum theory, the most complete explanation of the properties of the atom is given: the principles of formation of optical and X-ray spectra, the behavior of atoms in magnetic (Zeeman effect) and electric (Stark effect) fields, the periodic system of elements and the nature of the chemical bond were theoretically substantiated, methods were developed for calculating the electronic structure of atoms, molecules and solids (the Hartree-Fock self-consistent field method), new devices were created for studying the structure and properties of matter (electron microscope) Development of the ideas of quantum theory (gi the spin hypothesis, the Pauli principle, etc.), in turn, was based on experimental research in the field of atomic physics (line spectra of atoms, the photoelectric effect, the fine and hyperfine structure of spectral lines, the experiments of Frank and Hertz, Davisson and Germer, Stern and Gerlach, the effect Compton, the discovery of deuterium and other isotopes, the Auger effect, etc.)
7 A brief history of the development of atomic physics In the second third of the 20th century, within the framework of atomic physics and based on the ideas of quantum theory, new experimental methods of physical research were developed: electron paramagnetic resonance (EPR), photoelectron spectroscopy (PES), electron impact spectroscopy (ESI) , devices for their implementation (maser, laser, etc.) have been created. Fundamental principles of quantum theory (interference of quantum states, Lamb shift of levels, etc.) have received direct experimental confirmation, new methods for calculating the electronic structure of matter (density functional theory), and predicted new physical phenomena (superradiance) Methods have been developed for experimental studies of processes occurring with single atoms, ions and electrons held by electric and magnetic fields of a special configuration (atomic and ion "traps")
8 A brief history of the development of atomic physics New results in the field of atomic physics in the last third of the 20th century and the beginning of the 21st century are mainly associated with the use of lasers measurements with single atoms and molecules, determine the characteristics of highly excited states of atoms, study the dynamics of intraatomic and intramolecular processes lasting up to several femtoseconds (10–15 s) ), as well as the cooling of individual atoms to ultralow temperatures. Theoretical studies of recent decades in the field of atomic physics are associated with the rapid progress of computer technology and are aimed at developing efficient methods and means th structure and properties of multielectron atomic systems, taking into account the electron correlation energy, relativistic quantum mechanical and quantum electrodynamic corrections
9 Atomic physics Research in the field of atomic physics has found many scientific and practical applications For industrial purposes, to determine the elemental composition of a substance, methods of atomic spectral analysis are used, including EPR, FES and SEA To solve geological, biological and medical problems, methods of remote and local laser spectral atomic analysis, laser isotope separation is carried out for industrial and technical purposes Experimental and theoretical methods of atomic physics are used in astrophysics (determination of the composition and physical characteristics of the matter of stars and the interstellar medium, the study of Rydberg atoms), metrology (atomic clocks) and other areas of science and technology
10 Goals and objectives of the course of atomic physics The main goal of the discipline "Physics of the atom and atomic phenomena", as part of the course of general physics, is to form basic knowledge of the physics of microscopic phenomena at the atomic-molecular level and the ability to apply them to solve applied problems To achieve this goal the following tasks are solved: – analysis of the development of atomistic and formation of quantum concepts; – study of the most important experimental facts of atomic physics and their interrelation; - revealing the specifics of micro-phenomena and the failure of the classical theory to explain them; – study of the fundamentals of quantum mechanics and methods for solving quantum mechanical problems; – systematic study and explanation based on the quantum theory of the structure and properties of atoms and molecules, their behavior in external fields and in interaction with each other
12 Electron Electron is a stable elementary particle with a negative electric charge The absolute value of the electron charge is equal to the elementary charge q e = –e –1.610 –19 C The mass of the electron m e = m –31 kg The spin of the electron is ½ The magnetic moment of the electron is approximately equal to the Bohr magneton μ e – μ B - -4 eV / T The symbol e or e is used to designate an electron - Electrons form the electron shells of all atoms and ions The electron has an antiparticle positron (e +)
15 Proton Proton is a stable elementary particle with a positive electric charge The charge of the proton is equal to the elementary charge q p = e –19 C The mass of the proton m p 1836m e –27 kg The spin of the proton is ½ The magnetic moment of the proton μ p –8 eV/T The proton has an antiparticle antiproton (p-)
16 Annihilation of an antiproton An antiproton (blue track) collides with a proton in a bubble chamber resulting in four positive pions (red tracks) and four negative pions (green tracks) The yellow track belongs to a muon, which is born as a result of pion decay
17 Neutron Neutron elementary particle with zero electrical charge The lifetime of a neutron in a free state is approximately 886 s The mass of a neutron m n 1839m e –27 kg The spin of a neutron is ½ Despite the absence of an electric charge, the neutron has a magnetic moment μ n – –8 eV/T Neutron denoted by the symbol n or n 0 Neutron has an antiparticle antineutron Protons and neutrons are united by the common name nucleons Atomic nuclei consist of protons and neutrons
18 Neutron Since neutrons have no electrical charge, they do not leave tracks in particle detector chambers Neutrons can nevertheless be detected by their interactions with other charged particles The colorized image shows particle tracks in a cloud chamber filled with a mixture of hydrogen gas, ethyl alcohol and water The neutron beam penetrates into the chamber from below and causes transmutations of oxygen and carbon atoms that are part of the molecules of ethyl alcohol
19 Atom An atom is a microparticle consisting of an atomic nucleus and its surrounding electrons (electron shell) A positively charged nucleus holds negatively charged electrons by the forces of electric attraction the electron charge is equal to e, then when the number of electrons in the shell is equal to the number of protons in the nucleus, the total electric charge of the atom is zero. ), however, due to the fact that the mass of the proton (as well as the neutron) is almost 2 thousand times greater than the mass of the electron, almost the entire mass of the atom () is concentrated in the nucleus
20 Gold atom Au Image of a single gold atom obtained using a transmission electron microscope Magnification times to a size of 35 mm
22 Silicon atoms Si Colorized image of silicon atoms obtained using a transmission electron microscope. The unit cell of the crystal is shown. The bonds between the atoms are also visible. Magnification times to a size of 35 mm
24 Uranium atoms U A colorized image of uranium atoms was obtained using a transmission electron microscope. Small regular dots are individual atoms, larger formations are clusters consisting of 2–20 atoms Field of view is approximately 100 Å. Magnification up to a size of 35 mm
25 Uranyl microcrystals UO 2 2+ Colorized image of uranyl microcrystals obtained using a transmission electron microscope Each speck represents a single uranium atom Magnification times to a size of 35 mm
27 Chemical element, nuclide, isotopes Atoms with a certain number of protons Z in the nucleus belong to the same chemical element. The number Z is called the atomic number of a chemical element. A set of atoms with a certain number of protons Z and neutrons N in the nucleus is called a nuclide. Nuclides are denoted by adding to the name of the element the value of the mass number A, equal to the sum of Z + N (for example, oxygen-16, uranium-235), or by placing the number A near the symbol of the element (16 O, 235 U). Nuclides of the same element are called isotopes. The mass of the lightest atom of the hydrogen atom, consisting of one proton and one electron, is equal to m H 1.67 10 –27 kg. The masses of the remaining atoms are approximately A times greater than m H. There are 90 chemical elements and more than 300 different nuclides in nature; 270 of them are stable, the rest are radioactive. About radioactive nuclides obtained artificially.
31 Ions The process of removing or attaching electrons to an atom is called ionization If the number of electrons in the shell is less than Z, a positive atomic ion is obtained, if more than Z is negative Thus, an ion is an electrically charged atom (or molecule) that is formed upon detachment or attachment one or more electrons to a neutral atom (or molecule)
32 Ions Positively charged ions are called cations, negatively charged anions. Ions are denoted by a chemical symbol with an index that indicates the multiplicity (the amount of charge in units of elementary charge) and the sign of the ion: H -, Na +, UO 2 2+ Ions can be both stable formations (usually in solutions or crystals), so and unstable (in gases under normal conditions) atomic cations can be obtained up to a charge of +(Z - 1). Thus, for example, U 90+ and U 91+ were obtained on ion accelerators. Atomic anions with a charge of 2 or more do not exist in the free state.
34 Molecule A molecule is the smallest stable particle of a substance, consisting of more than one atom. A molecule is characterized by a certain composition of atomic nuclei, the number of electrons and a spatial structure. Chemical formulas are used to indicate the quantitative and qualitative composition of molecules: O 2 (oxygen molecule), H 2 O (molecule water), CH 4 (methane molecule), C 6 H 6 (benzene molecule), C 60 (fullerene molecule)
39 DNA molecule A colorized image of a DNA molecule was obtained using a transmission electron microscope In a high vacuum chamber, a DNA sample is coated with a thin layer of platinum Metallic coating gives a contrast image in an electron microscope
40 Nuclear and shell properties of the atom Nuclear propertiesShell properties Determined by the composition of the nucleus: radioactivity, ability to participate in nuclear reactions, etc. Determined by the structure of the electron shell: chemical, physical (electrical, magnetic, optical, etc.) 42 Energy The unit of energy in The SI is the joule (J), however, for the energy values of objects and phenomena of atomic physics, such a unit is rarely used. More commonly used is an off-system unit of energy called the electron volt (eV, eV) passing through an accelerating potential difference of 1 volt: 1 eV = J –6 eV) units of electron-volt, as well as some others: rydberg (Rydberg, Ry), hartree (hartree, Ha, or atomic unit, a. e.) Rydberg is numerically equal to the ionization energy of a hydrogen atom from the ground state in the approximation of an infinite mass of the nucleus: 1 Ry eV Hartree is equal to the absolute value of the potential energy of an electron in the ground state of the hydrogen atom in the approximation of an infinite mass of the nucleus: 1 Ha = 2 Ry eV The energies of states of atomic systems, as well as transitions between states can be measured in other units
43 Mass The unit of mass in SI is the kilogram (kg), however, to measure the masses of objects of atomic physics, an off-system unit of measurement is used, called the atomic mass unit (amu). The atomic mass unit is equal to 1/12 of the mass of an unbound, unexcited carbon-12 atom (12 C): 1 a. e. m kg 1 a. e. m. is approximately equal to the mass of one proton or neutron Relative atomic mass is the mass of an atom, expressed in a. e.m. Avogadro's constant N A is a physical constant numerically equal to the number of atoms in 12 g of pure carbon-12 isotope: N A mol –1 Mole (a unit of the amount of a substance in SI) by definition contains N A structural elements (atoms, molecules, ions).
44 Length The SI unit of length is the meter (m). 1 meter is equal to the distance traveled by light in vacuum in a time interval equal to 1/second. With the exception of measurements of wavelengths of electromagnetic radiation in the radio range, such a unit of length is rarely used in atomic physics, and instead, to measure linear dimensions, as well as wavelengths, submultiple units of a meter are used: centimeter (cm, 1 cm \u003d 10 -2 m), millimeter ( mm, 1 mm = 10–3 m), micrometer (μm, μm, 1 μm = 10–6 m), nanometer (nm, 1 nm = 10–9 m), picometer (pm, 1 pm = 10–12 m ) and others, as well as off-system units: angstrom (Å, 1 Å = 0.1 nm = 10–10 m), boron (or Bohr radius) (1 boron Å)
45 Time The SI unit of time is the second (s). atomic time standard: one second (or atomic second) is equal to the periods of electromagnetic radiation corresponding to the energy transition between two levels of the hyperfine structure of the ground state of the isotope 133 Cs (cesium-133) The duration of fast processes in atomic physics is usually measured in fractional units of a second: nano, pico- or femtoseconds (ns, ps, fs, 1 fs = 10 -15 s)
46 Scales of physical quantities in atomic and nuclear physics Phenomena of atomic physics are characterized by dimensions from 10–12 m (inner subshells of heavy atoms) to tenths of a nanometer (sizes of atoms and small molecules), energies from 10–6 eV (hyperfine structure of levels) to 10 5 eV (binding energies of electrons of inner subshells), times from tens of femtoseconds (duration of ultrashort laser pulses) to thousands of seconds (lifetimes of metastable states of atoms) Typical sizes of molecules are 0.1–1 nm. The internuclear distance of the smallest molecule (H 2) is nm. DNA macromolecules and many polymers can have macroscopic dimensions. Thus, the length of an unfolded DNA helix can reach several centimeters with a width of about 2 nm.
47 Photon A photon, or a quantum of electromagnetic radiation (field), is a massless elementary particle that does not have an electric charge In a vacuum, a photon moves at a speed c A photon has a spin equal to 1 The projections of the spin onto directions perpendicular to the direction of photon propagation determine the state of its polarization γ
· X-ray spectrum analysis · Radiospectroscopy ·
Atomic physics- a branch of physics that studies the structure and properties of atoms. Atomic physics arose in the late 19th - early 20th century as a result of experiments that established that the atom is a system of a positively charged nucleus and negatively charged electrons, and was developed in connection with the creation of quantum mechanics, which explained the structure of the atom. The structure of the atomic nucleus is studied in nuclear physics.
General information [ | ]
Modern atomic physics is based on quantum mechanical theory, which describes physical phenomena at the atomic-molecular level. Atomic physics considers an atom as a system of positively charged nucleus and negatively charged electrons. The properties of this system and the elementary processes occurring in it are determined by the electromagnetic interaction, in contrast to nuclear physics and elementary particle physics, where strong interaction and weak interaction play a fundamental role.
Story [ | ]
Planetary model of the atom
The idea of the existence of the smallest indivisible particles - atoms, was first formulated by the ancient Greek philosophers Leucippus, Democritus and Epicurus. In the 17th century, this idea was continued in the works of the French philosophers P. Gassendi and R. Descartes, and the English chemist R. Boyle. The atomistics of this period was rather speculative, ideas about atoms were like constant, indivisible particles, of various sizes and shapes, devoid of chemical and physical properties, the combination of which all material bodies consist of. In the works of I. Newton and M. V. Lomonosov, assumptions were made about the possibility of combining atoms into more complex structures - corpuscles.
The most important milestones in the history of atomic physics were the discovery of the electron in 1897 by the English physicist J. J. Thomson and radioactive decay by the French scientists M. Sklodowska-Curie and P. Curie, they changed the idea of the atom as a system of interacting charged particles, according to the theory of the Dutch physicist X . Lorenz . On the basis of these studies, Thomson proposed in 1903 a model of the atom in the form of a sphere with a positive charge, interspersed with small particles with a negative charge - electrons, held in the atom due to the equality of the force of attraction of the positive charge to the forces of mutual repulsion of electrons. Further studies of radioactivity by F. Soddy led to the discovery of isotopes, thereby destroying scientific ideas about the absolute identity of all atoms of one chemical element. An important role was also played by the study of the photoelectric effect by A. G. Stoletov and the further explanation of this phenomenon by A. Einstein.
The planetary model of the atom had a number of shortcomings, of which the most significant was associated with the theoretically correct loss of electron energy: since the electron rotates around the atom, it is affected by centripetal acceleration, and according to the Larmor formula, any charged particle moving with acceleration radiates energy. If the electron loses energy, then eventually it must fall into the nucleus, which does not happen in reality. Refinement of the atom model became possible only from the standpoint of completely new ideas about the atom, discovered by the German physicist
From chemistry and previous sections of physics, we know that all bodies are built from individual, very small particles - atoms and molecules. Atoms are the smallest particles of a chemical element. A molecule is a more complex particle consisting of several atoms ...
§ 195. Avogadro's constant. Dimensions and masses of atoms
One of the important constants of atomic physics is the Avogadro constant (see Volume I, § 242) - the number of structural elements (atoms, molecules, ions, etc.) in a mole of a substance. Knowing the Avogadro constant, one can find quantities that characterize an individual atom: mass ...
§ 196. Elementary electric charge
The laws of electrolysis discovered by Faraday testify in favor of the existence of the smallest, indivisible quantities of electricity. During electrolysis, one mole of any - valence element transfers the charge of coulombs (- Faraday's constant). For one atom (more precisely, io ...
§ 197. Units of charge, mass and energy in atomic physics
So, the charge of any particle always contains an integer number of elementary charges. For a particle of atomic size, this integer will also be small. In view of this, in atomic physics it is convenient to take the elementary charge as the unit of electric charge. For one...
§ 198. Measurement of the mass of charged particles. mass spectrograph
From the course of electricity, we know that a charged particle moving in a magnetic field is affected by a force called the Lorentz force. The Lorentz force is perpendicular to the magnetic field and to the velocity of the particle, and its direction is determined by the left hand rule (Fig....
§ 199. The mass of the electron. Mass versus speed
In an experiment to measure the mass of an electron using a mass spectrograph, only one strip is found on a photographic plate. Since the charge of each electron is equal to one elementary charge, we conclude that all electrons have the same mass...
§ 200. Einstein's law
In the previous paragraph, we established a relationship between the kinetic energy of a body and its mass: if the body is given kinetic energy, then its mass increases by an amount. This connection is of a general nature: it applies to any bodies - large and small, dawn ...
§ 201. Masses of atoms, isotopes
Consider the results of experiments on measuring the mass of positive ions. On fig. 352 is a neon positive ion mass spectrogram. Three bands of different intensity are clearly visible on the spectrogram. Comparing the distances from the strips to the slit, we can...
§ 202. Separation of isotopes. Heavy water
All isotopes of a given element enter into the same chemical reactions and form chemical compounds that are almost indistinguishable in solubility, volatility, and similar properties used in chemistry to separate elements. Therefore, conventional chemical methods...
§ 203. Nuclear model of the atom
In the previous paragraphs, we got acquainted with the data on the sizes and masses of atoms. Let us now turn to the question of the internal structure of the atom. The discovery of the phenomena of radioactivity contributed to the study of the structure of the atom. We will discuss these phenomena in detail in Chap. X...
§ 204. Energy levels of atoms
Experiments on scattering - particles discovered the existence in atoms of a heavy positive nucleus and an electron shell. Further information about the properties of atoms was given by the study of such atomic processes, which are accompanied by a change in the internal energy of the atom. WITH...
§ 205. Forced emission of light. quantum generators
The concept of the quantum energy levels of atoms was introduced into physics by N. Bohr in 1913. It very naturally explained the line atomic spectra as a result of the processes of spontaneous (spontaneous) emission and resonant (selective) ...
§ 206. Hydrogen atom. The peculiarity of the laws of motion of an electron in an atom
The existence of discrete energy levels is a fundamental property of atoms (as well as molecules and atomic nuclei). Let's try to apply the laws of physics known to us in order to imagine the structure of the atom, which explains the discreteness of its energy ...
§ 207. Multi-electron atoms. Origin of optical and x-ray spectra of atoms
Just as in the hydrogen atom, in more complex atoms, electrons can only move around the nucleus in certain chosen orbits. Various experimental data indicate that the possible orbits of electrons in an atom are grouped into a system of shells...
§ 208. Periodic system of elements of Mendeleev
The periodic law of change in the chemical properties of elements, discovered by D. I. Mendeleev, is a reflection of the deep laws of the structure of atoms; it is therefore of paramount importance not only for chemistry but also for physics. The correct theory of structure ...
§ 209. Quantum and wave properties of photons
As noted in § 184, the laws of the photoelectric effect were explained in 1905 by A. Einstein using the concept of light quanta (photons). According to these ideas, the energy of an electromagnetic field cannot be divided into arbitrary parts, but is radiated and absorbed...
§ 210. The concept of quantum (wave) mechanics
The study of the structure of the atom led to the conclusion that the behavior of electrons in an atom, as well as the behavior of photons, contradicts the usual laws of classical physics, i.e., the laws established in experiments with bodies of macroscopic dimensions. The existence of discrete...
§ 211. Discovery of radioactivity. radioactive elements
Uranium, thorium and some other elements have the property of continuously and without any external influences (that is, under the influence of internal causes) to emit invisible radiation, which, like X-rays, is capable of penetrating through opaque ...
§ 212. a-, b- and y-radiation. Wilson chamber.
As we have seen, radioactive radiation has an ionizing and photographic effect. Both of these actions are characteristic of both fast charged particles and X-rays, which are electromagnetic waves. To find out if it has...
§ 213. Methods for detecting charged particles
In the development of knowledge about the "microworld", in particular in the study of the phenomena of radioactivity, an exceptional role was played by devices that make it possible to register the insignificant effect of a single particle of atomic dimensions. One of these great tools is...
§ 214. The nature of radioactive radiation
1. radiation. The properties of radiation are similar to X-rays. Like x-rays, it ionizes the air, acts on a photographic plate, and is not deflected by a magnetic field. When passing through crystals, radiation, like X-rays, ...
§ 215. Radioactive decay and radioactive transformations
The study of radioactivity convinces us that radioactive radiation is emitted by the atomic nuclei of radioactive elements. This is obvious in relation to particles, since they simply do not exist in the electron shell. The Nuclear Origin of Particles Is Proved by Chem...
§ 216. Applications of radioactivity
1. Biological actions. Radioactive radiation has a disastrous effect on living cells. The mechanism of this action is associated with the ionization of atoms and the decomposition of molecules inside cells during the passage of fast charged particles. Particularly sensitive to and...
