THE ELECTRONIC STRUCTURES OF ATOMS

Note that in the fourth shell the orbital energies have this order As 4p Ad Af. Within any shell the order of energies is s p d f. However, also note that the 4s orbital is lower in energy than the 2d orbitals and that the 6s orbital is lower in energy than the 4f orbitals. Thus orbitals of one shell may be of higher energy than some of those in the next shell. 

The energies of orbitals are influenced by the nuclear charge of the atom and by the nature and number of the other electrons in the atom. Therefore, in the potassium atom, the 2d orbital is of higher energy than the As in scandium the 

Having briefly discussed the electronic structures of atoms, it is now 

The nohle gas elements (He, Ne, Ar, BCr, Xe, and Rn) are unreactive only recently have compounds of these elements been prepared. It has long been recognized that compounds in which each atom can, by electron sharing with other atoms, gather around itself a number of electrons equal to that foimd in a noble gas also tend to be very stable. Professor N. V. Sidgwick of Oxford University applied this observation to metal complexes. He postulated the central metal would surround itself with sufficient ligands so that the total number of electrons around the metal would be the same as that in the next noble gas. The number of electrons surrounding the coordinated metal is 

Similarly determined EAN values for other metal complexes in many cases equal the atomic numbers of noble gases. There are, however, many exceptions to this rule examples are [Ag(NH3)2] and [Ni(en)3 with EAN values of 50 and 38, respectively. This is unfortunate for if the EAN of the central metal always exactly equaled the atomic number of a noble gas, then it would be possible to estimate the coordination number of metal ions. 

In view of what has been said about electrons up to now it would be tempting to account for the fact that electrons are held to atoms by using a simple electrostatic theory. 

Accordingly, a nucleus with charge %e would attract an electron with charge e. The force F between these particles can be expressed by means of Coulomb s Law 

The picture would be similar to that of an orbiting satellite, the gravitational pull being replaced by electrostatic attraction. 

The electron then would possess potential energy due to its position in the electrostatic field of the nucleus and kinetic energy by virtue of its motion. Its total energy would be the sum of these two energy terms. Designating the electrostatic potential by V, its value can be derived from Coulomb s Law 

For a particle moving in a curved path the centripetal force is related in mechanics to the mass, the velocity and the radius of the path by the expression 

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The concept of chemical periodicity is central to the study of inorganic chemistry. No other generalization rivals the periodic table of the elements in its ability to systematize and rationalize known chemical facts or to predict new ones and suggest fruitful areas for further study. Chemical periodicity and the periodic table now find their natural interpretation in the detailed electronic structure of the atom indeed, they played a major role at the turn of the century in elucidating the mysterious phenomena of radioactivity and the quantum effects which led ultimately to Bohr s theory of the hydrogen atom. Because of this central position it is perhaps not surprising that innumerable articles and books have been written on the subject since the seminal papers by Mendeleev in 1869, and some 700 forms of the periodic table (classified into 146 different types or subtypes) have been proposed. A brief historical survey of these developments is summarized in the Panel opposite. 

As we saw in Chapter 19, chlorine represents the other extreme in chemical reactivity. Its most obvious chemical characteristic is its ability to acquire electrons to form negative chloride ions, and, in the process, to oxidize some other substance. Since the tendency to lose or gain electrons is a result of the details of the electronic structure of the atom, let us try to explain the chemistry of the third-row elements on this basis. 

Each equivalent atom (the same element, the same number of bonds and lone pairs) has the same formal charge. A check on the calculated formal charges is that their sum is equal to the overall charge of the molecule or ion. For an electrically neutral molecule, the sum of the formal charges is zero. Compare the formal charges of each possible structure. The structure with the lowest formal charges represents the least disturbance of the electronic structures of the atoms and is the most plausible (lowest energy) structure. 

In Chap. 3 the elementary structure of the atom was introduced. The facts that protons, neutrons, and electrons are present in the atom and that electrons are arranged in shells allowed us to explain isotopes (Chap. 3), the octet rule for main group elements (Chap. 5), ionic and covalent bonding (Chap. 5), and much more. However, we still have not been able to deduce why the transition metal groups and inner transition metal groups arise, why many of the transition metals have ions of different charges, how the shapes of molecules are determined, and much more. In this chapter we introduce a more detailed description of the electronic structure of the atom which begins to answer some of these more difficult questions. 

The modern theory of the electronic structure of the atom is based on experimental observations of the interaction of electricity with matter, studies of electron beams (cathode rays), studies of radioactivity, studies of the distribution of the energy emitted by hot solids, and studies of the wavelengths of light emitted by incandescent gases. A complete discussion of the experimental evidence for the modern theory of atomic structure is beyond the scope of this book. In this chapter only the results of the theoretical treatment will be described, These results will have to be memorized as rules of the game, but they will be used so extensively throughout the general chemistry course that the notation used will soon become familiar. 

The first plausible theory of the electronic structure of the atom was proposed in 1914 by Niels Bohr (1885-1962), a Danish physicist. In order to explain the hydrogen spectrum (Fig. 17-1), he suggested that in each hydrogen atom, the electron revolves about the nucleus in one of several possible circular orbits, each having a definite radius corresponding to a definite energy for the electron. An electron in the orbit closest to the nucleus should have the lowest energy. With the... 

For example, E. G. Mazurs (note 2, p. 105) expresses the discord as follows The periodicity of atomic structure must be accepted as a Natural Law. Therefore, scientists have to change their minds, get away from the conservatism that accepts only Mendeleev s chemical table as right, and adjust the other phenomena to this phenomenon that is, derive the chemical and physical properties of the elements from the electronic structure of the atoms. ... 

In this chapter, you learned about the electronic structure of the atom in terms of the older Bohr model and the newer quantum mechanical model. You learned about the wave properties of matter, and how to describe each individual electron in terms of its four quantum numbers. You then learned how to write the electron configuration of an atom and some exceptions to the general rules. 

The electronic structure of the atoms. The electronic structure of the atoms of the different elements and their relation to the characteristics of the Periodic Table are based on a number of experimental data and theoretical models which are fully discussed in many elementary and advanced texts of inorganic chemistry such as Cotton et al. (1999), Greenwood and Earnshaw (1997), Huheey etal. (1997), Wells (1984). 

In the process of examining the patterns outlined below, you will learn the filling order for atoms of elements in periods 5, 6, and 7. You will also see why the shape and organization of the periodic table is a direct consequence of the electronic structure of the atoms. 

Bohrium - the atomic number is 107 and the chemical symbol is Bh. The name derives from the Danish physicist Niels Bohr, who developed the theory of the electronic structure of the atom. The first synthesis of this element is eredited to the laboratory of the GSI (Center for Heavy-Ion Research) under the leadership of the German scientists Peter Armbruster and Gunther Mhnzenberg at Darmstadt, Germany in 1981, using the reaction ° Bi ( Cr, n) Bh. The longest half-life associated with this unstable element is 17 second Bh. 

The many-body perturbation theory [39] [40] [41] was used to model the electronic structure of the atomic systems studied in this work. The theory developed with respect to a Hartree-Fock reference function constructed from canonical orbitals is employed. This formulation is numerically equivalent to the M ler-Plesset theory[42] [43]. 

The explanation for the reduced energy of additional bonds is very similar to the through-the-surface mechanism of lateral interactions (see Section 2.1). The first bond changes the electronic structure of the atom. As a consequence the electronic structure is less favorable for forming another bond, so that such a bond yields a smaller energy gain. Because this mechanism is so similar to the one for lateral interactions, it is natural to use BOC or UBI EP to describe such interactions. This has been done for kMC simulations first by Lombardo and Bell and more recently by Baranov et al. and by Hansen and Neurock. 2a, 26... 

With the preceding discussion of the electronic structure of the atom and of electronic energies, the stage has been set to examine the question Why do atoms combine As is well known, the answer to this question is intimately related to the energy changes produced when two atoms approach each other. It is therefore worthwhile to study more closely what happens as two atoms come into closer proximity. 

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