Element of dispersion and particle experience. Rutherford's experience. Planetary model of the atom

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Rutherford's experience.

Ernst RUTHERFORD (1871-1937), English physicist, one of the founders of the doctrine of radioactivity and the structure of the atom, founder of a scientific school, foreign corresponding member of the Russian Academy of Sciences (1922) and honorary member of the USSR Academy of Sciences (1925). Director of the Cavendish Laboratory (since 1919). Discovered (1899) alpha and beta rays and established their nature. Created (1903, together with F. Soddy) the theory of radioactivity. Proposed (1911) a planetary model of the atom. Carried out (1919) the first artificial nuclear reaction. Predicted (1921) the existence of the neutron. Nobel Prize (1908).

Rutherford's experiment (1906) on the scattering of fast charged particles passing through thin layers of matter made it possible to study the internal structure of atoms. In these experiments, alpha particles were used to probe atoms - fully ionized helium atoms - resulting from the radioactive decay of radium and some other elements. Rutherford bombarded heavy metal atoms with these particles.

Rutherford knew that atoms consist of light negatively charged particles - electrons and a heavy positively charged particle. The main goal of the experiments is to find out how the positive charge is distributed inside the atom. The scattering of α - particles (that is, a change in the direction of movement) can only be caused by the positively charged part of the atom.

Experiments have shown that some of the α particles are scattered at large angles, close to 180˚, that is, they are thrown back. This is only possible if the positive charge of the atom is concentrated in a very small central part of the atom - the atomic nucleus. Almost the entire mass of the atom is also concentrated in the nucleus.

It turned out that the nuclei of various atoms have diameters of the order of 10 -14 – 10 -15 cm, while the size of the atom itself is ≈10 -8 cm, that is, 10 4 – 10 5 times the size of the nucleus.

Thus, the atom turned out to be “empty”.

Based on experiments on the scattering of α - particles on atomic nuclei, Rutherford came to to the planetary model of the atom. According to this model, an atom consists of a small positively charged nucleus and electrons orbiting around it.

From the point of view of classical physics, such an atom must be unstable, since electrons moving in orbits with acceleration must continuously emit electromagnetic energy.

Further development of ideas about the structure of atoms was made by N. Bohr (1913) on the basis of quantum concepts.

Laboratory work.

This experiment can be carried out using a special device, the drawing of which is shown in Figure 1. This device is a lead box with a complete vacuum inside it and a microscope.

Scattering (change in direction of movement) of α-particles can only be caused by the positively charged part of the atom. Thus, from the scattering of α particles, it is possible to determine the nature of the distribution of positive charge and mass inside the atom. The diagram of Rutherford's experiments is shown in Figure 1. A beam of α-particles emitted by a radioactive drug was released by a diaphragm and then fell on a thin foil of the material under study (in this case, gold). After scattering, the α-particles fell on a screen coated with zinc sulfide. The collision of each particle with the screen was accompanied by a flash of light (scintillation), which could be observed through a microscope.

With a good vacuum inside the device and in the absence of foil, a strip of light appeared on the screen, consisting of scintillations caused by a thin beam of α particles. But when foil was placed in the path of the beam, α-particles, due to scattering, were distributed over a larger area of the screen.

In our experiment, we need to examine the α-particle, which is directed at the gold core when making an angle of 180° (Fig. 2) and monitor the reaction of the α-particle, i.e. at what minimum distance will the α-particle approach the gold core (Fig. 3).

Rice. 2

Fig.3

Given:

V 0 =1.6*10 7 m/s – initial speed

d = 10 -13

r min =?

Questions:

What is the minimum distance r min between the α particle and the nucleus that can be achieved in this experiment? (Fig. 4)

Fig.4

Solution:

In our experiment, the α-particle is represented as an atom

m neutr kg

Z=2 – protons

N= Au – Z = 4 – 2 = 2 neutrons

m p = kg

Z=79 – number of protons

N= Au – Z = 196 – 79 =117 (neutrons)

Cl 2 / H ∙m 2 – electrical constant

Even in ancient times, the idea arose that the Universe consists of small indivisible particles - atoms. This idea of the structure of matter survived until the end of the 19th century, when by the beginning of the 20th century it was reliably established that each atom contains electrons. Priority in the discovery of the electron belongs to the English physicist J. Thomson. At the same time, it was known at that time that the atom is electrically neutral. Consequently, the negative charge of electrons must be compensated by the positive charge of unknown particles included in the charge of the atom.

In the 90s of the 19th century, J. Thomson’s model of the atom in the form of a homogeneous, positive spherical medium in which negatively charged electrons are dispersed, like raisins in a bun, became widespread. J. Thomson's atomic model is like a cupcake. However, soon the author of the “cupcake” model suggested a non-static position of electrons in the atom.

The most realistic seemed to be the nuclear or planetary model of the atom by E. Rutherford, proposed by the English physicist in 1911. The planetary model was the result of experiments performed by E. Rutherford and his colleagues on the scattering of alpha particles. The experiments consisted of the following. A beam of positively charged α particles is directed onto a thin foil-shaped sheet of gold. Behind the foil was a screen covered with a scintillator - a substance that emits light at the point at which the alpha particle hits. Based on J. Thomson's model, it would be expected that α particles would not be deflected through large angles, since electrons are much lighter than α particles. And, indeed, experiments showed that most alpha particles passed freely through the sheet of foil, as if it were basically empty space. And yet, some α particles were deflected at small angles, which was, as one might assume, a consequence of interaction with the positive charge of the atom. But what was unexpected and stunning was that a small number of α particles were scattered at large angles, reaching 180°. This could only happen if the positively charged α particles experienced repulsion from a massive positive charge concentrated in a small region of space.

According to E. Rutherford's model, an atom consists of a massive, positively charged nucleus, in which 99.94% of the mass of the atom is concentrated. The magnitude of the positive charge is estimated by the product ze, where z is the atomic number of the chemical element in D. Mendeleev’s table; e - elementary charge. Around the nucleus inside a sphere with an outer diameter of ~10 -10 m, z electrons rotate in closed elliptical orbits, forming the electron shell of the atom. Electrons cannot be at rest in an atom, since in this case they would fall onto the nucleus under the influence of Coulomb attraction. According to E. Rutherford's estimates, the dimensions of the nucleus should be of the order of 10 -15 -10 -14 m. Comparing the sizes of the nucleus and the atom, we come to the conclusion that the electrons should be located from the nucleus at a distance of (10-100) 103 more than kernel size. And hence the second conclusion: the bulk of the atom is empty space.

The disadvantage of E. Rutherford's model is the inability to explain the fact of the exceptional stability of the atom: firstly, during collisions with other atoms; secondly, according to the laws of classical physics, the rotation of electrons around the nucleus cannot be stable, since it must be accompanied by electromagnetic radiation, like any accelerated movement of charged particles. And according to the laws of classical physics, electrons moving in a circle have centripetal acceleration. The centripetal force that holds an electron in an orbit of radius r is the Coulomb force of attraction of the electron to the nucleus:

where ε o = 8.85 10 -12 F/m - electrical constant; m e - electron mass, kg; v is the speed of the electron in orbit, m/s. Creating an electromagnetic field requires energy. The energy of the electron should gradually decrease, and with it the speed of rotation of the electron around the nucleus. The electron must eventually fall into the nucleus. However, atoms are quite stable formations and can exist for billions of years. Thirdly, according to E. Rutherford’s model, the emission spectrum of an atom should be continuous. Experiments have shown that the emission spectrum of a particular atom is discrete.

Rutherford's experience.

Thus, the atom turned out to be “empty”.

From the point of view of classical physics, such an atom must be unstable, since electrons moving in orbits with acceleration must continuously emit electromagnetic energy.

Further development of ideas about the structure of atoms was made by N. Bohr (1913) on the basis of quantum concepts.

Laboratory work.

This experiment can be carried out using a special device, the drawing of which is shown in Figure 1. This device is a lead box with a complete vacuum inside it and a microscope.

Rice. 2

Fig.3

Given:

V 0 =1.6*10 7 m/s – initial speed

d = 10 -13

= 180°

r min =?

Questions:

What is the minimum distance r min between the α particle and the nucleus that can be achieved in this experiment? (Fig. 4)

Fig.4

Solution:

In our experiment, the α-particle is represented as an atom

m neutr kg

Z=2 – protons

N=Au –Z = 4 – 2 = 2 neutrons

m p =kg

Z=79 – number of protons

N=Au –Z = 196 – 79 =117 (neutrons)

Cl 2 /H ∙m 2 – electrical constant

m 2 =6.6∙10 -27 kg

- the charge of an α-particle is equal to 2 elementary.

Answer: r min =4.3·10 -14 m

Conclusion: During this experiment, it was possible to find out that the a-particle was able to approach the atomic nucleus to a minimum distance, which was r min =4.3·10 -14 m and return back along the same trajectory along which it began to move.

When Rutherford performed the same experiment for the first time, with such an a-particle positioned relative to an angle of 180°, he said in surprise: “This is almost as incredible as if you fired a 15-inch projectile at a piece of tissue paper, and the projectile returned would come to you and strike you.”

And in truth, this is not probable, the fact is that when carrying out this experiment at smaller angles, the a-particle will certainly jump to the side, just as a pebble of several tens of grams when colliding with a car is not able to noticeably change its speed (Fig. 5). Since their mass is approximately 8000 times greater than the mass of the electron, and the positive charge is equal in magnitude to twice the charge of the electron. These are nothing more than fully ionized helium atoms. The speed of α particles is very high: it is 1/15 the speed of light. Consequently, electrons, due to their low mass, cannot noticeably change the trajectory of the α particle.

The computer program simulates Rutherford's classical experiment on probing the atom with alpha particles, based on the results of which it was proposed planetary model of atomic structure .

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 107 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. 1.

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°.

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 all positive charge and at least 99.95% of its mass. The charge of the nucleus must be equal to the total charge of all the electrons that make up the atom.

Based on classical ideas about the movement of microparticles, Rutherford proposed 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. Electrons cannot be at rest, since they would fall onto the nucleus.

Rutherford's planetary model of the atom was a major step forward in the development of knowledge about the structure of the atom. It was absolutely necessary to explain experiments on the scattering of α-particles, but it turned out to be unable to explain the very fact of the long existence of an atom, i.e., its stability. According to the laws of classical electrodynamics, a charge moving with acceleration should emit electromagnetic waves that carry away energy. In a short time (about 10–8 s), all the electrons in the Rutherford atom must waste all their energy and fall onto the nucleus. The fact that this does not happen in stable states of the atom shows that the internal processes in the atom do not obey classical laws.

The user has the opportunity:

observe the scattering of particles on a stationary gold core;
change the impact distance and initial velocity of the particle;
measure the particle scattering angle;
study the scattering curve when bombarding a gold nucleus with a stream of particles with a given energy in automatic mode.

Now I know what an atom looks like!

Ernest Rutherford, 1911

One day, in the farming hinterland of what the Maori call Aotearoa, the Land of the Long White Cloud, a young settler was digging potatoes. With enviable tenacity, the guy dug into the ground with a shovel, extracting a crop that would help his family survive difficult times. It is unlikely that he hoped to find gold nuggets there - unlike other parts of New Zealand, his area was not famous for its mines - but he was destined for a golden future.

Ernest Rutherford, who was destined to be the first to look into the depths of the atom, was born into a family of early settlers in New Zealand. His grandfather George Rutherford, a wheelwright from Dundee, Scotland, came to the Nelson colony on the tip of the South Island to help build a sawmill. When it was ready, Rutherford Sr. moved the family to the village of Brightwater (now Spring Grove) south of Nelson, in the Wairoa River valley. There, George's son James, who grew flax and earned his living by doing so, married the English emigrant Martha, who gave birth to Ernest on August 30, 1871.

At Nelson School and later at Canterbury College in Christchurch, the largest and most English city in the South Island, Rutherford proved himself a diligent and capable student. One of the future scientist’s classmates remembered him as “a spontaneous, sincere, simple and very pleasant young man who, although he was not a child prodigy, if he saw a goal, he immediately grasped the main thing” 11 .

Ernest Rutherford (1871-1937), father of nuclear physics.

Rutherford's deft hands worked wonders with any mechanical device. The experimenter's youthful hobbies prepared him well for subtle manipulations with atoms and atomic nuclei. With a skill worthy of a surgeon, he disassembled clocks, created working models of water mills, and even made an amateur camera to take photographs. In Canterbury, having learned about electromagnetic phenomena discovered in Europe, he set out to build his own installation. Following Hertz, he assembled a radio transmitter and receiver that anticipated Marconi's invention of the wireless telegraph. Rutherford demonstrated that radio waves could travel long distances, pass through walls, and magnetize iron. His original experiments gave him the opportunity to apply for a place in Cambridge, England.

Coincidentally, in the year that Rutherford was born, a new physics laboratory was organized in Cambridge, of which Maxwell became its first director. Cavendish Laboratory, named. so in honor of the brilliant physicist Henry Cavendish (by the way, among other things, he was the first to isolate hydrogen as a chemical element), it turned into a world center of atomic physics. It is located on Free School Lane, near the center of the famous university town. Maxwell himself supervised the construction and selected equipment for the world's first physical research laboratory. After Maxwell's death in 1879, the director's chair was taken by another famous physicist, Lord Rayleigh. And in 1884, the inimitable J. J. (Joseph John) Thomson took the reins of government.

This energetic and versatile man with long dark hair, a bushy mustache and wire-rimmed glasses became one of the driving forces behind a revolution in science education that opened up enormous research opportunities for students. Previously, experimental work for physics students was done only as a dessert at the end of a long banquet where mathematical subjects were served. However, even this treat the teachers shared rather reluctantly. After the student passed all the exams in mechanics, thermal phenomena, optics and other theoretical subjects, he was sometimes allowed to touch some instruments for a short time. At Cavendish, with its top-of-the-line equipment, these short tastings turned into a full meal. Thomson enthusiastically welcomed the new system, which allowed a student from another university to come to Cambridge and conduct research under the supervision of a local scientist. Based on their results, the invitee wrote a dissertation and received a higher degree. Today we take PhD holders for granted because they are the ones who join the academic world. But at the end of the 19th century. such a system was innovative, and a revolution in physics was not long in coming.

The innovations came into full force in 1895, and Rutherford was among the first students invited. He received the "1851 Scholarship", awarded to young talented people from the countries of the British Dominion (now a Commonwealth country). Having exchanged provincial New Zealand for the university of Cambridge, Rutherford worked to benefit not only his own career, but also the entire atomic physics.

There is a legend about how Rutherford accepted this gift of fate. They say his mother received a telegram with good news and went to the field where he was digging potatoes. When she read to her son what an honor he had received, he at first did not believe his ears, but, barely realizing his happiness, he threw away the shovel and exclaimed: “Today I dug potatoes for the last time!” 12

Taking his homemade radio, Rutherford sailed to London. There he promptly slipped on a banana peel and injured his knee, but, fortunately, the entire subsequent journey through the labyrinths of the foggy city passed without a hitch. As he moved north, the mists gave way to fresh air, and the city was replaced by English landscapes and the sacred outlines of various colleges on the River Cam. Here Rutherford settled at Trinity College. The great gates of the college, founded in 1546 by King Henry VIII, and the legends of Newton's glorious deeds still dominate the reverent steps of students as they enter here. (The University of Cambridge is divided into many colleges where students study and live, and Trinity College is the largest of them.) After leaving Trinity College and enjoying a short walk, you almost immediately find yourself in the Cavendish Laboratory.

Rutherford was not the only one of the stream of students who poured from all over the world into Cambridge research laboratories. Thomson cherished the spirit of unity of difference that reigned here and every day after lunch he invited young employees to tea. He later recalled: “We talked about everything in the world, but not about physics. I didn't encourage talking about physics because we were meeting to relax... and because it's easy to learn to speak your bird language, but hard to unlearn. And if you don’t get used to it, then the ability to maintain a conversation on general topics will atrophy as unnecessary” 13 .

Despite Thomson's attempts to encourage young researchers, the pressures at Cambridge were apparently taking their toll. “When I return from the laboratory, I find myself restless and usually in a rather nervous state,” Rutherford once wrote. To relax a little, he began smoking a pipe, maintaining this habit for the rest of his life. “Sometimes I took a puff,” continues Rutherford, “and I managed to concentrate a little... Any man of science should smoke a pipe, otherwise where can he have patience? Scientists should have it like ten Jobs combined” 14.

Local students also added fuel to the fire, treating the visiting people as strangers. Rutherford's classmates from the golden youth, teasing him as a hillbilly from Antipodia, did little to lift his spirits. About one such bully, Rutherford said: “There is one laboratory assistant on whose chest I would not mind, like a true Maori, performing a war dance” 15.

Thomson was a pedantic experimenter and at one time enthusiastically studied the properties of electricity. Having assembled the original installation, he studied the combined influence of electric and magnetic fields on the so-called cathode rays - negatively charged beams of electricity coming from a negatively to a positively charged electrode (a contact connected to the corresponding pole of the battery). A negatively charged electrode generates cathode rays, and a positively charged electrode attracts them.

Charges behave differently in electric and magnetic fields. The force with which the electric field acts on a negative charge is directed opposite to the direction of the field. As for the magnetic field, the force in it acts at right angles to the field. In addition, unlike electric force, magnetic force depends on the speed of the charge. Thomson figured out how to compensate the electric and magnetic fields to determine this speed. And thanks to it, he could determine the ratio of the charge of the rays to their mass. By putting the charge of the particles in the beams equal to the charge of ionized hydrogen, Thomson discovered that their mass was several thousand times less than that of hydrogen. Simply put, cathode rays are made up of elementary particles that are much lighter than atoms. Changing the conditions and repeating the experiment over and over again, Thomson always got the same result. He called these negatively charged particles corpuscles, but they were later given a different name: since then it has become the same - electrons. They were the first to open a small window into the rich world of the atom.

Thomson's stunning discovery was initially met with skepticism by the scientific community. “At first, few people believed that there were such objects - smaller than an atom,” he recalled. - Many years later, even one outstanding physicist, who was present at my lecture at a meeting of the Royal Society, told me that he was completely sure that I was “fooling everyone’s head.” His words did not surprise me. I myself resisted this explanation, and only when experiments left me no other choice did I publicly declare the existence of bodies smaller than atoms” 16.

Meanwhile, on the other side of the English Channel, the discovery of radioactive decay cast doubt on the prevailing ideas about the indivisibility of the atom. In 1896, the Parisian physicist Henri Becquerel sprinkled uranium salts on a photographic plate wrapped in black paper and was quite surprised when he saw that the plate darkened over time, which meant that some mysterious rays were coming from the salts. Unlike X-rays, Becquerel’s ones appeared on their own without any electrical devices. The scientist discovered that the radiation came from any uranium-containing compounds. Moreover, the more uranium in the compound, the more it radiated. It was logical to assume that it was the uranium atoms themselves that emitted this radiation.

Marie Skłodowska-Curie, a physicist of Polish origin, who worked in Paris, repeated Becquerel’s experiments, and also, together with her husband Pierre, found mysterious radiation in two elements they discovered: radium and polonium. The latter emitted even more intensely than uranium, and their quantity decreased over time. Maria coined the term “radioactivity,” which she used to describe the phenomenon of spontaneous decay of atoms, releasing special radiation. For their landmark discovery of the fragility of atoms in radioactive processes, Becquerel and the Curies were awarded the Nobel Prize in 1903. Dalton's timeless elements, which had reigned supreme in science for a century, were in motion.

Rutherford followed these events with great interest. While his teacher Thomson was busy discovering the electron, Rutherford turned his close attention to the fact that gases could be ionized with radioactive materials. For some reason, the rays coming from uranium and other radioactive compounds removed the gas from a state of electrical inertness and turned it into an electrically active conductor. The radioactive radiation behaved like two sticks being rubbed against each other to create a spark.

But most importantly, radioactivity struck a spark of interest in Rutherford and forced him to engage in a methodical study of its properties, which was destined to revolutionize our ideas about physics. And the beginner, who began by assembling radios and other electromagnetic devices, had to gain experience and turn into an experimenter of the highest class, capable of traveling into the world of the atom with the help of radioactive radiation. Knowing that a magnetic field deflects unlike charges in different directions, Rutherford realized that radioactive rays have positive and negative components. He gave them names, respectively, alpha and beta radiation. (Beta particles turned out to be simply electrons, and soon Rutherford’s classification was continued by Villard, who discovered a third, electrically neutral component - gamma rays.) In a magnetic field, alpha particles twist in one direction, and beta particles in the other, like horses running around the circus arena in different directions. Rutherford looked at how much each type of radiation was blocked by an obstacle and proved that beta rays penetrate deeper than alpha rays. Therefore, alpha particles are larger than beta particles.

In 1898, in the midst of his research into radioactivity, Rutherford decided to take a break to settle matters of the heart. He went to New Zealand for a short time, where he married his high school sweetheart, Mary Newton. However, they did not return to England. A married man should have a good income, Rutherford concluded, and accepted a position as a professor at McGill University in Montreal, Canada, with a salary of £500 a year - decent money in those days, about $50,000 in today's equivalent. The happy couple sailed to the cold region, where the scientist soon continued his research.

At McGill, Rutherford was more eager than ever to unmask alpha particles and reveal their true colors. Repeating Thomson's experiments to determine the charge-to-mass ratio with alpha rays instead of electrons, he suddenly saw that the charge of alpha particles was the same as that of helium ions. The suspicion was creeping in that the heaviest product of radioactive decay was actually helium traveling incognito.

Just when Rutherford could use some help in solving atomic mysteries, another tracker appeared in town. In 1900, Frederick Soddy (1877-1956), a chemist from Sussex, England, received a position at McGill University. Having learned about Rutherford's experiments, he wanted to make his contribution, and together they began to study the phenomenon of radioactivity. They hypothesized that radioactive atoms such as uranium, radium and thorium decay into simpler atoms of other chemical elements, releasing alpha particles in the process. Soddy, who was fascinated by the history of the Middle Ages, guessed that radioactive transformations were, in a sense, the embodiment of the cherished dream of alchemists who were trying to obtain gold from base metals.

In 1903, shortly after Rutherford published their joint theory of radioactive transformations, Soddy decided to join forces with William Ramsay of University College London, a recognized expert on helium and noble gases in general (neon and others). Ramsay and Soddy conducted a series of careful experiments in which alpha particles from radioactive radium were collected in a special glass tube. Then the scientists examined the spectral lines of the resulting sufficiently dense gas, which turned out to be exactly the same as those of helium. Spectral lines are narrow stripes in the vicinity of certain frequencies (in the visible part of the spectrum these are certain colors). Each element, emitting or absorbing light, produces its own set of lines. In the emission spectrum of helium, some violet, yellow, green, blue-green and red lines are always visible, as well as two characteristic bluish stripes. These “fingerprints” served as irrefutable proof in the experiments of Ramsay and Soddy that alpha particles and ionized helium are one and the same.

Soddy also coined the term “isotope,” which he used to describe varieties of the same chemical element that had different atomic weights. For example, deuterium, or “heavy” hydrogen, is chemically no different from ordinary hydrogen, but its atomic weight is approximately twice as large. The radioactive isotope of hydrogen, tritium, is generally about three times heavier than ordinary hydrogen. When it decays, it produces helium-3, a light isotope of the familiar helium. Soddy developed what he called the law of radioactive displacement: as a result of alpha decay, an element in the periodic table moves back two spaces, as if it had a bad move in a board game. Beta decay, on the contrary, gives the right to one move forward, and one of the isotopes of the element sitting in the next cell is obtained. A living example is the decay of tritium, which, turning into helium-3, jumps one cell further.

Let’s imagine that you accidentally come across an incomprehensible machine with balls and you don’t see its contents. Sometimes blue balls jump out of it, and the machine blinks once, and sometimes red balls, whose appearance is accompanied by two flashes. How can we understand from here what is happening inside? You can probably assume that the machine contains a homogeneous mixture of red and blue balls, scattered here and there, like raisins in a pudding.

By 1904, physicists knew that in radioactive processes atoms transform into each other, emitting particles with different charges and masses, but no one had a clue about the big picture. Thomson ventured to put forward the idea that positive and negative charges are evenly mixed, and the latter, since they are lighter, have greater freedom of movement. When the experimenters tasted this pudding, he hoped, they would see how good it was. But, alas, the first pudding came out lumpy. And fate decreed that this verdict would be given by Thomson’s New Zealand favorite.

The next period in Rutherford's life was perhaps the most fruitful. In 1907, the University of Manchester - Dalton's scientific path once ran through these northern English places - invited the scientist to head the department of physics. What Manchester gained was McGill's loss. By that time, Rutherford had “ridden his luck,” as he himself, not without boasting, remarked to his biographer (and student) Arthur Eve 17, and was already a noticeable figure in science. Like a true helmsman, he steered his ship with a firm hand: he hired the best young researchers, assigned them interesting tasks and fired those who did not live up to expectations. Loud, sometimes quick-tempered and sometimes cursing instruments at all costs, the professor with his invariable pipe and mustache really instilled fear in his subordinates. But the outbursts of anger quickly passed, the bright sun appeared from behind the drying clouds, and then there was no one in the world friendlier, more good-natured and more supportive than Rutherford.

At that time, the Manchester biochemist and future first president of Israel, Chaim Weizmann, became close to him. He described Rutherford as “lively, energetic and boisterous. He cared about everything, not just science. He willingly and energetically discussed everything in the world, even if he didn’t have the slightest idea about something. Going down to the dining room for dinner, I already heard the boom of his friendly voice in the corridor... He was good-natured, but he could not tolerate fools” 18.

Weizmann recalled, comparing Rutherford with Einstein, whom he also knew well: “As scientists, these two men were the opposite of each other: Einstein was all calculation, Rutherford was all experiment. But in life they were little alike. Einstein seemed unattainable, and Rutherford looked like the big, boisterous New Zealander he was. In the field of experiment, Rutherford, of course, was a genius, one of the best. He had a special flair, and whatever he touched, everything turned to gold” 19.

In Manchester, Rutherford made ambitious plans: to split an atom with alpha particles and see what was inside. He guessed that relatively large alpha particles were an ideal device For research into the deep structure of the atom. First of all, he wanted to test the strength of Thomson's pudding model and understand whether it was true that the atom was a pie of impressive positively charged pieces and small negative charges. Determined to win, Rutherford managed to steal two valuable prizes from under the noses of his competitors: the much-coveted supply of radium (they fought for it with Ramsay) and the bright head of the German physicist Hans Geiger, who had previously worked under the former head of the department. Rutherford set Geiger the task of developing a reliable way to detect alpha particles.

The method proposed by Geiger - counting sparks that jump between the electrodes of a metal tube when alpha particles, ionizing the gas sealed inside, turn it into a conductor - formed the basis of the famous counter, named after the author of the invention, the Geiger counter. This meter works due to the fact that electrical currents circulate in closed circuits. Each time the sample emits an alpha particle, a circuit through the electrodes and the conducting gas is completed and an audible click is heard. Despite Geiger's useful finding, Rutherford usually used a different registration method. He took a screen coated with zinc sulfide, a material in which bombarding alpha particles cause tiny flashes of light, called scintillations.

In 1908, Rutherford interrupted his research to go receive the Nobel Prize in Chemistry, awarded to him for his study of alpha particles. But the laboratory was empty for long. Armed with reliable detection methods, he moved on to new experiments, in which Geiger and the talented, although not yet university graduate, Ernest Marsden also took part.

The fate of 20-year-old (1909) Marsden was strikingly similar to the fate of Rutherford himself. Marsden also came from a simple background. His father worked at a provincial textile factory in Lancashire, England, making cotton fabrics. Rutherford moved from his native New Zealand to England - for Marsden, everything subsequently turned out exactly the opposite. They both started doing interesting experiments while they were still at university. As for Marsden, before he had finished his studies, he was already invited to test his talents.

Rutherford later recalled that warm-up question, which resulted in a fruitful collaboration between Geiger and Marsden. “One day Geiger came up to me and said: “Perhaps it’s time for young Marsden to do a little research?” I was already thinking about this, so I replied: “Then let him see if any alpha particles are scattered at large angles.”20

Rutherford, famous for his ability to ask the right questions at the right time, felt that if alpha particles suddenly appeared, flying back from metal, this would give a clue about the structure of matter. He was, of course, interested to see what would happen, but he did not have high hopes for a positive outcome of the experiment. But this option could not be completely ruled out. You never know, suddenly something is hidden inside from which particles will bounce. It would have been a sin not to try my luck.

In some particularly sensitive measurements, particle physicists must be like nocturnal animals prowling for prey. You can catch her slightest movement if you can only see well in the dark. In this activity, young researchers have an advantage. And not even because of better vision, but rather because of patience. It is not surprising that Rutherford and Geiger assigned twenty-year-old Marsden to monitor the scattering of alpha particles. He was instructed to curtain the windows as tightly as possible, then sit down and wait until the pupils dilate enough to catch the slightest glimpse of light from all sides. Only then was it possible to begin observations.

Marsden placed plates of different thicknesses and from different metals (lead, platinum and others) next to a glass ampoule with radium compounds and waited for the alpha particles escaping from the ampoule to hit the plate and pass through it or bounce off. A screen with zinc sulfide served as a scintillator. It showed how many particles were reflected and at what angles. Having finished with the next metal, Marsden showed the data with all the sparkles that his keen eyes noticed to Geiger. Together they found that thin sheets of gold produced the most reflections. But they also mostly let alpha particles through, as if the foil were from the other world. And when reflections occurred occasionally, the particles tended to bounce at very large angles (90 degrees or more). Consequently, they apparently dissipated on some solid condensations in the depths of gold.

Beaming with joy, Geiger ran to Rutherford and, to the latter’s utter delight, announced: “We have finally found bouncing alpha particles!”

“It was the most incredible event of my life,” Rutherford recalled. “It’s almost as incredible as if you threw a 15-inch grenade at a tissue paper screen and it bounced back at you.”21

If the atom, as Thomson thought, really is like raisin pudding, then the amorphous mixture of charges inside the gold atoms should not strongly deflect the alpha particles flying into the foil, and then they would scatter more often at small angles. But Geiger and Marsden did something different. It’s as if a good baseball player is sitting inside an atom: when the projectile is in the strike zone, the hitter hits it with all his might, and if the projectile goes beyond this zone, it flies freely further.

In 1911, Rutherford decided to propose his own instead of Thomson's model. “I think I came up with a much better atom than Jayjay,” he shared with a colleague 22 . In the article, he outlined the revolutionary idea that every atom has a tiny, positively charged nucleus at its center, which contains the bulk of the atom's mass. When alpha particles collided with gold atoms, it was this bat that knocked them back, and even then only the most accurate ones who managed to hit the bull's eye.

It turns out that the atom consists mainly of emptiness. The core occupies a pitiful fraction of its volume, everything else is bottomless nothingness. If you enlarge an atom to the size of the Earth, the nucleus would be about the diameter of a football stadium. Rutherford colorfully compared firing at a target cannonball to trying to locate a mosquito in the Royal Albert Hall, a huge concert hall in London.

Despite its small size, the nucleus plays an important role in determining the properties of the atom. Rutherford conjectured that the position of an atom in the periodic table, or, in other words, the atomic number, depends on the magnitude of the positive charge of the nucleus. In hydrogen, the nuclear charge is equal in absolute value to the electron charge, and for other elements this amount of charge must be multiplied by the atomic number. For example, gold, the 79th element, has a positive nuclear charge equal to seventy-nine electron charges. The positive central charge is balanced by a corresponding number of negatively charged electrons. As a result, the atom, if not ionized, is electrically neutral. As Rutherford argued, these electrons are evenly distributed over a sphere centered at the nucleus.

Although Rutherford's model took physics to a new level, some questions remained open. It perfectly explained the Geiger-Marsden scattering experiments, but many of the experimental properties of the atom known at that time remained a mystery. Take, for example, spectral lines - within the framework of the model, it was not clear why in hydrogen, helium and other elements they form a characteristic pattern. If the electrons in an atom are uniformly mixed, why are atomic spectra so heterogeneous? And where in the overall picture can we find a place for Planck’s quantum idea and Einstein’s photoelectric effect, in which the electron receives and releases energy in finite portions?

By a happy coincidence, in the spring of 1912, a guest from Denmark arrived at Rutherford’s laboratory, whose knowledge came in handy. Niels Bohr, a strongly built young man with large features, had recently defended his dissertation in Copenhagen and, after spending six months with Thomson in Cambridge, left for Manchester. He wrote a letter to Rutherford in advance, saying that he would not mind working in radioactivity. From Thomson he knew about Rutherford's idea of the nucleus, and he wanted to study its consequences in more detail. Once, when Bohr was calculating the process of collisions of alpha particles with atoms, a hypothesis came to his mind: what if the energy of an electron oscillating near the nucleus takes on strictly defined values, multiples of Planck’s constant? Thus, in one fell swoop, Bohr plunged atoms into the kaleidoscope of quantum theory.

Returning to Copenhagen in the summer of that year, Bohr continued to think about the structure of the atom. He was interested in the question of why atoms do not spontaneously collapse. Something must hold the negative electrons so that they do not crash into the positively charged nucleus, like a meteorite into the Earth. In Newtonian physics there is a special conserved quantity, angular momentum (angular momentum). Simply put, when a body rotates, both the number of revolutions and the direction of the axis tend to remain unchanged. Namely, the product of mass, speed and orbital radius is often a constant value. It is not for nothing that a skater begins to spin faster when he presses his orbits only with a certain energy. That is, electrons can only be located at certain distances from the atomic nucleus, or, in other words, occupy discrete levels - quantum states.

Bohr's guess immediately made it possible to make great progress in the question of why the sets of spectral lines in atoms are exactly the way they are and not others. In the Bohr model of the atom, electrons, if they are in some specific quantum state, do not receive or release energy - as if they, like a planet, fly in an absolutely stable, ideal orbit. According to Bohr, electrons are, roughly speaking, something like small Mercurys and Venuses revolving around the sun-core. But instead of the force of gravity, they are pulled towards the center by an electrostatic force acting from the positively charged nucleus. This, however, is where the analogy with the solar system ends, and Bohr's theory then takes on a completely different turn. Unlike planets, electrons sometimes jump from one quantum state to another, to the nucleus or, conversely, from the nucleus. Jumps are unpredictable and instantaneous, and the electron gains or loses energy depending on whether it jumps to a higher or lower level. As in the photoelectric effect, the frequency of the resulting radiation can be calculated by dividing the transferred energy by Planck's constant. The portions of energy themselves were later called photons, or light particles. So, the characteristic color lines in the emission spectra of hydrogen and other elements are explained by the fact that the electron, throwing off the light ballast, makes a kind of dive. The deeper he goes, the higher the frequency. Bohr's model was a triumph. Her predictions coincided surprisingly accurately with the known formulas that give the distance between the spectral lines of hydrogen.

In the winter of 1913, Bohr reported the results to Rutherford and, to his disappointment, received a rather mixed response from him. Practically thinking Rutherford found what seemed to him to be a big flaw in the model. He wrote to Bohr: “I have discovered a serious difficulty in connection with your hypothesis, of which you are, no doubt, fully aware; it is this: how can an electron know at what frequency it should oscillate when it passes from one stationary state to another? It seems to me that you are forced to assume that the electron knows in advance where it is going to stop." 23

With this apt remark, Rutherford identified one of the main inconsistencies in Bohr's atomic model. How do you know when exactly an electron will abandon the tranquility of its current state and go looking for adventure? How do you know which state he will choose? Borov's model was powerless here. This is exactly what Rutherford didn't like.

A response to Rutherford's remarks was received only in 1925, but it also led to confusion among many. By that time, Bohr had acquired his own Institute of Theoretical Physics in Copenhagen (now the Niels Bohr Institute), and a whole galaxy of young scientists worked under his leadership. Among them, the German physicist Werner Heisenberg (1901-1976), who was educated in Munich and Göttingen, stood out. It was he who proposed an alternative description of how electrons behave in an atom. His model also did not explain, Why electrons jump, but made it possible to accurately calculate how likely they were to do so.

Heisenberg’s “matrix mechanics” introduced new abstract concepts into physics, which greatly confused old-school scientists and were met with hostility by some prominent physicists who understood what these concepts entailed. One of the striking examples is Einstein, who was an implacable opponent of matrix mechanics. She threw a blanket of uncertainty over the atom - and over all of nature on these and smaller scales, declaring: not all physical properties can be measured at once.

With the spirit of rebellion characteristic of youth, Heisenberg began his exposition by rejecting most of the ideas that reigned supreme among his elders. He refused to perceive the electron as an orbiting particle and replaced it with a pure abstraction: a mathematical state. To calculate position, momentum (mass times velocity), and other observable physical properties, Heisenberg multiplied this state by various quantities. His scientific supervisor, Göttingen physicist Max Born, proposed writing these quantities in the form of tables or matrices. Hence the term “matrix mechanics” (synonymous with quantum mechanics). Armed with a powerful mathematical apparatus, Heisenberg no longer saw any obstacles on his way to the depths of the atom. Then he recalled: “I had the feeling that something amazingly beautiful was being revealed to me through the surface of atomic phenomena, and I was almost dizzy at the thought that I was about to plunge into this rich world of mathematical structures that nature so generously presented to me.” spread out" 24.

In classical Newtonian physics, position and momentum can be measured simultaneously. In quantum mechanics, as Heisenberg elegantly showed, this is not the case at all. If you act on the state with the coordinate and momentum matrices, the order of these operations is of great importance. When you apply the coordinate matrix first, and then the momentum matrix, the answer will most likely be different than when you do the opposite: momentum first, and coordinates later. Operations where the order of execution matters are called non-commutative. We are all very familiar with commutative options: in arithmetic these are multiplication and addition (“from changing the places of terms…”). Due to non-commutativity, it becomes impossible to simultaneously know both physical quantities with perfect accuracy. Heisenberg formulated this fact in the form of the uncertainty principle.

For example, if you fix the position of an electron, the Heisenberg uncertainty principle in quantum mechanics ensures that the momentum is smeared as much as possible. But momentum is proportional to speed, which means that the electron cannot tell us at the same time where it is and at what speed it is flying. The electron has not only seven, but no one knows how many Fridays in a week. If the planets behaved like electrons, the ancient astrologers would have abandoned their work before they even started.

Although, according to Heisenberg, quantum mechanics by its very nature is inherent in uncertainty, it provides a recipe for how to calculate probability. That is, it does not guarantee that you will win the bet, but it does tell you what your chances are. Let's say, quantum mechanics gives the probability that an electron will jump from a given position to some other one. If this probability is zero, you know for sure that such a transition is prohibited. If not, it is resolved, and lines with the corresponding frequency can be seen in the atomic spectrum.

In 1926, physicist Erwin Schrödinger proposed an easier-to-understand version of quantum mechanics, called wave mechanics. Developing the theory built by the Frenchman Louis de Broglie, Schrödinger began to interpret electrons as “waves of matter.” Something like light waves, but represented not by electromagnetic radiation, but by material particles. How these wave functions react to physical forces is described by the Schrödinger equation. Let's say, in an atom, the wave functions of electrons under the influence of electrostatic attraction from the nucleus form “clouds” of different shapes, energies and with different average distances from the center. These clouds have no material content. They only show how likely it is for an electron to end up at a particular point in space.

These wave structures can be likened to the vibrations of a guitar string. A standing wave appears on a string fixed at both ends after plucking. Lying on the beach, we see running waves that roll onto the shore. In contrast, a standing wave is destined to move only up and down. But even with such a limitation, it can have several peaks (maxima): one, two or more - the main thing is that this number must be an integer, not a fraction. Wave mechanics establishes a correspondence between the principal quantum number of an electron and the number of maxima, which naturally explains why these particular states exist and not others.

Much to Heisenberg's chagrin, many of his colleagues preferred Schrödinger's painting. Perhaps because wave processes were somehow closer to them - there was an analogy with both sound and light... The matrices looked too abstract. However, the insightful Viennese physicist Wolfgang Pauli proved that the Heisenberg and Schrödinger models are completely equivalent. It's like digital and analogue displays - none of them is inferior to the other, and which one to choose is a matter of taste.

Pauli himself left a legacy to quantum mechanics: the idea that two electrons cannot occupy the same quantum state. The Pauli exclusion principle led two Dutch scientists, Samuel Goudsmit and Georg Uhlenbeck, to the idea that the electron can align in two directions, that is, it has spin. As the name suggests spin -"fast rotation"), spin characterizes the internal angular momentum of the electron. But, above all, the properties of spin in relation to the magnetic field are interesting. If you place an electron in a vertical magnetic field (say, inside a magnetic coil), the electron, like a mini-magnet, will face either in the direction of the field (“spin up”) or against it (“spin down”).

The electron is the servant of two masters: it usually exists in a mixed state, where the “spin up” and “spin down” positions are represented in equal shares. Wait, how can the same particle have two mutually exclusive properties? In everyday life, a compass needle cannot point both north and south at the same time, but in the quantum world there are different rules of the game. Until we measure the spin, according to the uncertainty principle, it does not have a clearly defined value. But then the experimenter turns on the external magnetic field, and then the electron turns its spin either up or down - a collapse of the wave function occurs, as they say.

Let's say two electrons are in a bundle. Then, if one has a spin that sticks up, the other immediately turns down. This flip takes place even if the electrons are far apart. In this counterintuitive phenomenon, Einstein saw the tricks of the “ghost of long-range action.” Because of these strange relationships, Einstein was convinced that quantum mechanics would one day be replaced by a deeper and clearer theory.

As for Bohr, he did not renounce paradoxes; on the contrary, he felt like a fish in water among incompatible concepts. For example, it was he who formulated the principle of complementarity, which states that an electron is both a wave and a particle. From time to time Bohr was also not averse to uttering another aphorism. He once said: “A deep truth is a truth whose opposite is also a deep truth.” It was completely in his spirit to place the Taoist symbol of the unity of opposites - yin-yang - in the very center of his coat of arms.

Despite his intransigent philosophical position, Einstein agreed with Bohr that quantum mechanics was an excellent explanation of experimental data. One of the signs of recognition of her merits was Einstein's nomination of Heisenberg and Schrödinger for the Nobel Prize in Physics. Heisenberg was awarded it in 1932, and Schrödinger shared the honor with British quantum mechanics Paul Dirac in 1933. (Einstein and Bohr were laureates in 1921 and 1922, respectively)

Rutherford, however, continued to treat quantum theory with caution and devoted his main attention to experimental studies of the atomic nucleus. In 1919, Thomson resigned his title as Cavendish professor and left the post of director of the Cavendish Laboratory, and Rutherford took over this honorary position. During his last year in Manchester and the first years after moving to Cambridge, he bombarded various nuclei with fast alpha particles. Marsden once noticed that from the place where alpha particles hit the hydrogen gas, even faster particles with higher penetrating power begin to fly. These turned out to be the nuclei of hydrogen atoms. Rutherford repeated Marsden's experiments, but replaced hydrogen with nitrogen. Imagine his surprise when hydrogen nuclei also began to fly out of nitrogen. True, scintillations from hydrogen nuclei entering the fluorescent screen were not very bright, and they could only be seen through a microscope. But they provided undeniable evidence that nitrogen atoms can emit particles from their depths. The discovery of radioactivity demonstrated that atoms could spontaneously transform into each other (undergo transmutation), and from Rutherford's bombardment experiments it was possible to change the appearance of atoms artificially.

Rutherford began to call the positively charged particles that are part of all nuclei protons. Other scientists wanted to call them “positive electrons,” but Rutherford strongly opposed it. He replied that protons are much heavier than electrons and in general they have little in common. When Dirac's prediction came true and a real positively charged electron was discovered, it was given the name "positron". Positrons became the first known representative of the so-called antimatter, which is similar in every way to ordinary matter, but has a charge of the opposite sign. Protons, in turn, are an integral part of matter that is familiar to us.

A new particle detector, a cloud chamber, came to the aid of Rutherford and his collaborators. It made it possible to observe traces of particles (for example, protons) flying from the target nucleus. While scintillation and Geiger counters only provided a stream of emitted particles, a cloud chamber could show how these particles moved through space, hence helping to better understand their properties.

It was invented by Scottish physicist Charles Wilson. While climbing Mount Ben Nevis, he noticed that in humid air, water droplets condense more readily in the presence of ions, that is, charged particles. The charges attract molecules, and they are precipitated from the air, leaving a condensation trail in the area saturated with electricity. Wilson realized that this way it was possible to register particles invisible to the eye. He took the chamber, filled it with cold, damp air and began to observe the chains of condensed steam from charged particles flying past. Jet planes leave the same mark in the sky. These tracks, captured in photographs, provide a wealth of valuable information about the progress of the experiment.

Although Wilson assembled the first prototype of his chamber in 1911, they began to be used in nuclear physics only in 1924. It was then that Patrick Blackett, a graduate student in Rutherford's group, used this device to detect protons from the radioactive decay of nitrogen. His data were in excellent agreement with Rutherford's scintillation experiments, thereby providing irrefutable evidence of artificial nuclear decay.

The nucleus is inhabited not only by protons. In 1920, with his legendary sixth sense, Rutherford guessed that in addition to protons, the nucleus also served as a refuge for some neutral particles. Twenty years later, Rutherford's student James Chadwick discovered a neutron - the same mass as a proton, but without a charge, and Heisenberg soon after wrote a historical article "On the Structure of the Atomic Nucleus", where he outlined the now accepted model of a nucleus consisting of protons and neutrons.

This picture can explain the different types of radioactivity. Alpha decay occurs when a nucleus emits two protons plus two neutrons at the same time—an exceptionally stable combination. Beta decay then takes place when a neutron produces a proton and an electron. Beta radiation consists of these very electrons. But, as Pauli showed, the story does not end there: in the decay of a neutron, a certain amount of momentum and energy disappears somewhere. Pauli decided to attribute them to an almost elusive particle, which was later discovered and named neutrino. Finally, the gamma component occurs when a nucleus transitions from a high-energy quantum state to a low-energy state. Alpha and beta decay change the number of protons and neutrons in the nucleus, and a new chemical element is formed, while gamma rays leave the composition of the nucleus unchanged.

Rutherford's brilliant discoveries and methods taught us a lesson: in order to peer into the natural world at small distances, we must turn to elementary particles. At the dawn of nuclear physics, their source was radioactive substances gushing with alpha particles. They were ideally suited for scattering experiments, from which Geiger and Marsden saw that the atom had a miniature nucleus. But Rutherford already understood: without more energetic tools there is nothing to think about in order to penetrate more seriously and deeply into the nature of the nucleus. For a nuclear fortress, you will need a particularly strong ram, or rather, rams - particles accelerated under artificial conditions to phenomenally high speeds. Rutherford, not without reason, decided that the Cavendish Laboratory would be able to build a particle accelerator, although its implementation, the scientist admitted, would require certain theoretical efforts. Fortunately, one clever young man managed to sneak out of Stalin's fortress and take a bag of quantum knowledge with him to Free School Lane.

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