Discoveries that led to the atomic model

5th century BC: Democritus, Greek philosopher believed that matter consists of tiny particles that cannot be broken into smaller pieces. The particles, he named “atom”, possessed weight and extent but were not perceptible to the senses, according to his belief. In addition to atomistics, several other interpretations of the nature of matter have spread. For millennia, they all remained just hypotheses. They relied solely on philosophical considerations and none of them had an experiential basis.
1803: The first experimental evidence of atomic theory, the law of multiple proportions, is associated with the name of John Dalton. According to him, chemical elements consist of unchanged and indivisible particles, possessing the properties of the elements. He stated that: „When two elements combine with each other to form two or more compounds, the ratios of the masses of one element that combines with the fixed mass of the other are simple whole numbers”.
1808: Joseph Louis Gay-Lussac‘s gas law was another step forward.
1811: According to Amedeo Avogadro’s law, gases of the same volume, temperature, and pressure contain the same number of molecules. This made it possible to interpret atomic and molecular weights as well as the formula of molecules.
1815: William Prout, starting from the fact that the weight of atoms is an integer multiple of the weight hydrogen, concluded that atoms are made up of hydrogen, atoms and therefore are divisible. It was later revealed that both the basic assumption and the conclusion are wrong.
1833: Michael Faraday’s law describing electrolysis projected the relationship between atomic structure and electric charge.
1850-1880: The atomic approach of matter became supported by kinetic gas theory, which was developed by Rudolf Clausius, James Clerk Maxwell and Ludwig Boltzmann. The model was able to interpret the pressure and specific heat of a gas and the velocity distribution of gas particles.
1865: Johann Josef Loschmidt determined the Avogadro constant (Loschmidt number): \( N_{A} = 6.02214\times10^{23} \frac{1}{mol} \) .
1869: Dmitry Ivanovich Mendeleev realized that if the elements are arranged according to their atomic weights, then a periodicity can be observed in their physical and chemical properties (periodic system).
1895: Wilhelm Conrad Röntgen discovered X-ray radiation. In addition to the fluorescence that accompanies cathode radiation, Röntgen observed a type of radiation that blackened the photographic plate. Röntgen’s results were presented by the famous mathematician , Henri Poincaré at the Paris Academy. It was then that he posed the question that led to the discovery of radioactivity: „Is there a connection between X-rays and fluorescence?”
1896: Antoine Henri Becquerel studied uranium salts: fluorescent and non-fluorescent ones. Experiments showed that there is no relationship between fluorescence and the detected radioactive radiation that blackened the photographic plate. The research was continued by Becquerel’s students, Marie Curie and Pierre Curie. They also observed that this radiation penetrates paper, glass, and thin metal foils; independent of excitation (e.g. heating); and characteristic of uranium as an element.
1897: Joseph John Thomson discovered the electron by observing the electrical conduction (cathode radiation) in sparse gases and measured the specific charge (charge to mass ratio) of the electron: \( 1.76\times10^{11} \frac{C}{kg} \) .
1898: M. Curie and P. Curie constructed a device for detecting radioactivity (ionization chamber). With this, new radioactive elements were discovered: polonium and radon.
1898: Ernest Rutherford by examining the penetration of radioactive rays discovers two types of radiation: shorter-range alpha and the more penetrating beta radiation. During the study of \(\beta\)-particles, it was observed that their specific charge is the same as that of electrons’.
1900: Paul Ulrich Villard discovered the even more penetrating gamma-radiation.
1900: Rutherford observed an exponential decrease in the intensity of radioactive radiation with respect to time, and based on this, he introduced the concept of half-life.
1902: Rutherford described the process of radioactive decay (fission). Electrons have been detached from atoms in several ways (with a strong electric field, excitation), which suggested that the electron is part of the atom.
1904: Thomson proposed the first atomic model, according to which the atom is a sphere with approximately \(10^{-10} m\) radius, uniformly filled with positively charged material in which the electrons are evenly distributed. (plum pudding model).
However, the model contradicted Philipp Lenard’s and later Rutherford’s experiments.
1905: Albert Einstein formulated an interpretation of Brownian motion.
1908: Rutherford proved that alpha particles are doubly positive helium ions (\(He^{++}\)).
1910: Robert Millikan measured the charge of the electron (charge of ions with valence of 1): \( e = 1.6022\times10^{-19} {C} \). Knowing the “elementary charge” and the specific charge, the (resting) mass of the electron can be calculated: \( m_{el} = 9.11\times10^{-31} {kg} \).
1909-11: In his experiments, Rutherford studied the scattering of \(\alpha\)-particles on a thin gold foil. He and his colleagues found that in addition to particles that scatter at a small angle, there are also those that drastically (at an angle greater than \(90^\circ\)) change the direction of their original trajectory. Some particles „bounce off” the film by nearly \(180^\circ\).
Based on the results of the measurement and calculation, in 1911, Rutherford proposed a new atomic model: the atom consists of a small positively charged nucleus of \(~10^{-14} {m}\) in its center and electrons, which are distributed around it in a relatively large space \(10^{-10} {m}\). Because the mass of the electrons is very small, virtually the entire mass of the atom is concentrated in the nucleus. According to Rutherford, the centripetal force required for the circulation of electrons around the nucleus comes from the Coulomb force between the positive nucleus and the electrons. The idea is similar to the planets orbiting the Sun, so it was named „solar system model”. However, such an atom would be flat and also cannot be stable. Due to the repulsion between the electrons, the tiny particles would have to push each other off their orbits. In addition, according to the laws of electrodynamics, every accelerating charge (including the circular electron) necessarily radiates off some of its energy. As a result, the energy of the system decreases; the electrons gradually approach the nucleus, and eventually fall into it. During this process, the orbital period of the electrons changes and the frequency of the radiation emitted by the atom must also change. Because electrons do not fall into the nucleus and thus atoms are „long-lived”, the solar system model does not accurately describe reality either.
1912: Max von Laue observed the diffraction of X-rays on crystals in his experiments.
1912-13: Most phenomena observed on atomic scale cannot be explained by classical physics. Based on Planck’s postulate of quantized oscillators and Einstein’s theory of the photoelectric effect, Niels Bohr further developed Rutherford’s model. With the model and its associated quantum conditions, many phenomena of atomic physics became interpretable. Two basic assumptions (postulate) of Bohr’s theory: – The electron can orbit the nucleus only in certain, so-called stationary orbits. As long as the electron is in such a quantum orbit, the atom does not emit photons (its energy is unchanged). – Light emission only occurs when an electron jumps from one orbit to another. If the atom changes from a state of energy \(E_m\) to \(E_n\), then the following relation holds for the frequency of the emitted photon: \(E_m – E_n = hf\) (Bohr’s frequency condition) The state with the lowest energy \(E_1\) is called the ground state. The energy levels denoted \(E_2\), \(E_3\) etc. are called the excited states. The model only gave an accurate explanation for the spectrum lines of the hydrogen atom (and one-electron ions), so Bohr’s theory needs further improvements.