Atom

The elements that compose the ordinary matter are formed by atoms, the smallest constituent unit with well-defined and specific characteristics of each chemical element, for example, the mass and atomic number. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms.

Although it is the smallest material structure (dimensions of the order of 10-10 m) exists from the chemical point of view, the atom does not represent the ultimate constituent of matter. Globally neutral, the atom consists mainly of three types of particles:

  • electron
  • proton
  • neutron

Protons and neutrons occupy the central part of the atom, the nucleus, whose radius is of the order of 10-15 m and in which almost all of the atomic mass is therefore concentrated.

The elaboration of the atomic hypothesis

The atomic hypothesis, that is the hypothesis of the indivisibility of matter, was introduced as a philosophical speculation since the V century b.C. by Democritus and his school. Resumed in the XVIII century to interpret some results of the kinetic theory of gases, its validity was confirmed during the XIX century with the impetuous development of chemistry. The further progress of experimental physics (discovery of cathode rays and channel rays, X-rays, atomic spectra, radioactivity) put in full light, at the end of last century, the complexity of atom structure and invalidated the concept of atom as the ultimate particle of matter. The theoretical work that followed, aimed to interpret the numerous experimental data, led to the development of new areas of physics such as quantum mechanics, atomic physics and nuclear physics, thus beginning a new scientific era, the atomic era.

The introduction on scientific basis of atomic theory is due to D. Bernoulli. Starting from the consideration that matter was made by the union of elementary units of very small size and placed at a great distance from each other, Bernoulli demonstrated in his Hydrodynamics (1738) that the pressurep of a gas is inversely proportional to the volumeV occupied by it, thus giving a theoretical foundation to R. Boyle’s experimental law.

In 1803 the english physicist and chemist J. Dalton used the term atom to indicate in gas the ultimate particles of matter and then postulated that each chemical element corresponded to a specific atom, theoretical hypothesis that imposed itself to the attention of scientists for the experimental confirmations of C.L. Berthollet, J.L. Proust and L.J. Gay-Lussac his contemporaries.

The analysis of mixtures, conducted in 1808 by the French chemist Proust, highlighted the discontinuous structure of chemical substances: according to these results the matter appeared no more as a continuous and homogeneous whole, but as a union of different substances, union that implied a granular structure. Thus the molecular hypothesis was born, based on the assumption that every distinct chemical substance, or pure substance (not fractionable by freezing or distillation), was formed, in turn, by identical particles or molecules. And since pure substances were composed of simple substances in exactly defined proportions (Proust’s law) it could be deduced that molecules were formed by atoms of different chemical elements.

The classification of elements according to their chemical properties and their atomic weight, in which many chemists were interested, was elaborated in 1869 by Russian chemist D. Mendeleev who was able to order the elements according to the increasing order of their atomic weight. The success achieved by the great chemist marked the triumph of atomic hypothesis because it established unequivocally the discontinuity of matter. From a theoretical point of view the periodicity of chemical characteristics of elements remained unexplainable because it implied a relation between different atoms that was justified only after the discovery of the internal structure of atom.

The structure of the atom

While chemical laws (Dalton, Gay-Lussac, Avogadro etc..) showed that matter was not a continuous whole, infinitely divisible, but it was composed of “elementary” particles, atoms and molecules, a series of new phenomena and experiences led to discover an internal structure of atom itself. They discovered new particles constituting the atom, which appeared not indivisible, but composed of parts. The first of these particles was found by J.W. Hittorf in 1869, studying cathode rays, discovered by J. Plücker in 1858, he demonstrated that they were composed of charged particles, with charge and mass well defined.

These particles, called electrons by J. Stomey in 1891, whose mass is about 1/2000 of the mass of hydrogen (E. Wiechert, 1897), were soon recognized as universal constituents of matter and identified as the atoms of electricity postulated around 1830 to interpret Faraday’s laws of electrolysis. In 1909 C.G. Barkla, studying X ray scattering, was able to determine the number of electrons of each element, finding for light atoms a value equal to the order number of the element in Mendeleev classification.

The number of electrons composing the atom is therefore, like mass, related to each specific atom; it is called atomic number and it is indicated with the letter Z. A second type of particles universal constituent of atom, protons, was highlighted by E. Goldstein in 1886 with the discovery of channel rays; the studies of W. Wien, J.J. sir Thomson and J.B. Perrin showed that they were particles with positive electric charge, equal in absolute value to electron charge and whose mass coincided with hydrogen mass.

In these researches, Thomson was able to calculate with extreme precision the mass of ionized atoms (atoms from which were subtracted one or more electrons) and to differentiate, in 1913, in the same chemical element, atoms with different atomic masses (isotopes) represented by integers (mass numbers, A), multiples of hydrogen mass set equal to 1. The third subatomic particle was discovered by J. sir Chadwick in 1932 (discovery that earned him the Nobel prize in 1935). This new constituent of atom, with mass almost identical to proton, but electrically neutral, was called neutron.

While the first elementary components of atom were identified and the hypothesis of its indivisibility was invalidated, another discovery contributed to abandon the old notion of immutable atom: the discovery of natural radioactivity (H. Becquerel, 1896; Pierre and Marie Curie, 1898), a phenomenon in which an atom, by emitting X-rays, electrons and ionized helium atoms (α rays), can transform itself in another atom, that is to operate a transmutation.

Atomic models

The identification of atomic constituents opened a new problem: to theoretically explain how these particles could form a stable building and to study its structure. Many atomic phenomena (emission of optical spectra, photoelectric effect, X-ray emission, Zeeman effect etc.) derive from the internal complexity of the atom and to interpret them it is necessary to build models able to account coherently of all experimental results.

Thomson and Rutherford models

The first attempt to give a concrete image of the atom is due to Thomson, who described it as a sphere of positive electricity, homogeneous and indivisible, within which electrons are immersed in conditions of electrostatic equilibrium.

To this scheme J. Perrin, in 1901, opposes a planetary model: the atom would be composed by a central nucleus, positively charged, around which, at relatively immense distances, electrons rotate, thus maintaining the balance between coulombian attraction force and centrifugal force.

Ten years later, in 1911, E. Rutherford resumed and developed this model to interpret his experiments of diffusion of alpha particles in matter (in particular it was not explained with Thomson full atom the great path that these particles were able to do in matter). The size of nucleus deduced from Rutherford diffusion experiments is of the order of 10-15÷10-14 m, while the radius of electron orbit is about 10-10 m.

Diffusion studies also allowed to determine the number of electrons of many atoms (atomic number Z) and the fact that this number was the same for isotopes of the same element allowed to confirm that chemical properties of various elements are determined only by orbital electrons.

Bohr-Sommerfeld model

Rutherford’s model, although it introduced the fundamental concept of nucleus and drew attention to the role of peripheral electrons, exposed itself to two important objections: the intrinsic instability of such an atom and the impossibility to interpret the existence of discontinuous atomic spectra (see spectroscopy). In fact electrons, during their trajectory, according to electrodynamic laws should emit radiation, gradually losing their kinetic energy until they are attracted in the nucleus, thus giving rise to a continuous emission spectrum, phenomena in contradiction with experience.

Danish physicist N. Bohr, who in 1922 was awarded the Nobel prize for his studies on atomic structure, even if he always accepted Rutherford’s model of planetary atom, he abandoned any attempt to interpret experimental data in classical electrodynamics and he extended also to atom the new quantum hypothesis elaborated by M. Planck at the beginning of XX century. The postulates of Bohr’s atomic theory can be summarized as follows:

  • mechanical postulate, according to which the electron can describe around the nucleus only a well determined series of trajectories in which it does not cause emission of electromagnetic radiation so that the energy of the atom remains constant; the latter can assume only a series of discrete values that constitute the energy levels of the atom (quantization of orbits);
  • optical postulate, for which an electron can pass from an orbit to which corresponds an energy Ei to another one to which corresponds an energy Ef by absorption (or emission, if Ef is less than Ei) of a quantum (minimum quantity of a physical quantity that depends on two values of this one) of electromagnetic energy to which is associated frequency:

\[\nu =\dfrac{E_f-E_i}{h}\]

The constant h = 6,625⋅10-34 J⋅s is called Planck constant. Bohr’s model applied to hydrogen atom allows, in excellent agreement with experimental results, precise calculations of its size (characterized by minimum radius or Bohr radius: RB = 0,529⋅10-10 m), energy levels and atomic emission spectra studied by J.J. Balmer (1885), T. Lyman (1906), L.C.H. Paschen (1908) and others.

Among the different possible energy levels, marked by the symbol n (n = 1, 2, 3, …, main quantum number) and corresponding to the successive electronic layers enveloping the nucleus, indicated with the letters K, L, M, …, the atom tends to stay, in the absence of external stress, in the stable state of lowest energy (n = 1, electron in the K layer), called fundamental state. Following the absorption of radiation by electron, this passes to a higher energy level (energy transition) bringing the atom in an excited state. This state is however unstable and after some time the atom returns to its fundamental configuration releasing the excess energy by emitting a photon.

Electron can also move away from nucleus if atom itself absorb an energy higher than electronic bond energy (13,6 eV for hydrogen). The atom is not electrically neutral anymore, but it has a positive charge; it is reduced to ionic state. A.J.W. Sommerfeld generalized Bohr’s model postulating the existence of elliptical electronic orbits (instead of circular) with the nucleus in one of the foci. The result is that for each of the allowed values of energy, for example for the nth, there are n stable elliptical orbits, each characterized by an integer number (azimuthal quantum number l = 0, 1, 2, 3, …, n-1) that is interpreted as the momentum of the electron orbital with respect to the nucleus.

In this model results quantized also the spatial orientation of the orbits and precisely the orientations allowed for an electron in a state which competes the azimuthal quantum number “l” are in number equal to 2l+1. Each orientation is marked by an integer, positive or negative, m (|m|≤l) called magnetic quantum number.

The imperfections of Bohr-Sommerfeld theory did not take long to appear: its validity was limited to hydrogen atom and even in this case it did not give a satisfactory interpretation of the intensity of spectral lines that make up the atomic emission spectrum. To solve these problems, distinguished physicists such as L. de Broglie, E. Schrödinger, P.A.M. Dirac, M. Born, W. Pauli and W. Heisenberg (all awarded the Nobel Prize for their studies) impressed to the researches, between 1920 and 1930, even bolder orientations that had to revolutionize physics.

  • Electron affinity

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