Occupied electron shells

Ligand field theory mainly considers the last contribution. For this contribution the geometric distribution of the ligands is irrelevant as long as the electrons of the central atom have a spherical distribution the repulsion energy is always the same in this case. All half and fully occupied electron shells of an atom are spherical, namely d5 high-spin and dw (and naturally d°). This is not so for other d electron configurations. 

A familiar way of handling this question is offered by the notion of electronic shells. By definition, an electronic shell collects all the electrons with the same principal quantum number. The K shell, for example, consists of U electrons, the L shell collects the 2s and 2p electrons, and so on. The valence shell thus consists of the last occupied electronic shell, while the core consists of all the inner shells. This segregation into electronic shells is justified by the well-known order of the successive ionization potentials of the atoms. 

Valence shell The outermost occupied electron shell of an atom. 

The first horizontal row of the Periodic Table contains only two elements, hydrogen and helium. Each has a single occupied electron shell. 

The second horizontal row contains eight elements, and every one of the eight has two occupied electron shells. We could say that each of the eight has an outer shell with eight places and that one or more of these "places is occupied by an electron. 

Fig. 1. Self-consistent effective potential for Na2o in the spherical jellium model. The occupied electronic shells are indicated, as well as the lowest unoccupied shell.

The three electrons of the lithium atom, for instance, are arranged 2,1 among the electron shells the eleven electrons of the sodium atom are arranged 2,8,1 the nineteen electrons of the potassium atoms are arranged 2,8,8,1 and so on. Each of the alkali metals has the electrons of its atoms so arranged that the outermost occupied electron shell contains just one electron. 

The completely and partially occupied electron shells in the atom. 

Main group element Element whose atoms are characterized by the filling of s or p orbitals of the outermost shell (the occupied electronic shell with the greatest quantum number). Referred to also as a representative element. 

The ionisation energy of an atom is particularly high if the outermost shell is fully occupied. Fully occupied electron shells are energetically favourable so that atoms tend to attain configurations with completely filled outermost... 

In PMD radicals, the bond orders are the same as those in the polymethines with the closed electron shell, insofar as the single occupied MO with its modes near atoms does not contribute to the bond orders. Also, an unpaired electron leads the electron density distribution to equalize. PMD radicals are characterized by a considerable alternation of spin density, which is confirmed by epr spectroscopy data (3,19,20). 

The orbitals in an atom are organized into different layers, or electron shells, of successively larger size and energy. Different shells contain different numbers and kinds of orbitals, and each orbital within a shell can be occupied by two electrons. The first shell contains only a single s orbital, denoted Is, and thus holds only 2 electrons. The second shell contains one 2s orbital and three 2p orbitals and thus holds a total of 8 electrons. The third shell contains a 3s orbital, three 3p orbitals, and five 3d orbitals, for a total capacity of 18 electrons. These orbital groupings and their energy levels are shown in Figure 1.4. 

FIGURE 15.19 The atomic and ionic radii of the halogens increase steadily down the group as electrons occupy outer shells of the atoms. The values shown are in picometers. In all cases, the radii of the anions (represented by the green spheres) are larger than the atomic radii. 

Among many examples of -orbital interaction, only the following two are selected to illustrate the feature of HO—LU conjugation. One is the cyclooctadiene-transition metal complex ">. The figure indicates the symmetry-favourable mode of interaction in a nickel complex. The electron configuration of nickel is (3d)8 (4s)2. The HO and LU of nickel can be provided from the partly occupied 3d shell from which symmetry-allowed occupied and unoccupied d orbitals for interaction with cyclo-octadiene orbitals are picked up. 

The beautiful Bohr atomic model is, unfortunately, too simple. The electrons do not follow predetermined orbits. Only population probabilities can be given, which are categorized as shells and orbitals. The orbitals can only accommodate two electrons. Shells and orbitals can also merge ("hybridization"). In the case of carbon, the 2s orbital and the three 2p orbitals adopt a configuration in the shape of a tetrahedron. Each of these sp3 orbitals is occupied by one electron. This gives rise to the sterically directed four-bonding ability of carbon. 

The elements in group IA can add an electron with the release of energy (a small amount) because their singly occupied outer shells can hold two electrons. 

The possible states of electrons are called orbitals. These are indicated by what is known as the principal quantum number and by a letter—s, p, or d. The orbitals are filled one by one as the number of electrons increases. Each orbital can hold a maximum of two electrons, which must have oppositely directed spins. Fig. A shows the distribution of the electrons among the orbitals for each of the elements. For example, the six electrons of carbon (B1) occupy the Is orbital, the 2s orbital, and two 2p orbitals. A filled Is orbital has the same electron configuration as the noble gas helium (He). This region of the electron shell of carbon is therefore abbreviated as He in Fig. A. Below this, the numbers of electrons in each of the other filled orbitals (2s and 2p in the case of carbon) are shown on the right margin. For example, the electron shell of chlorine (B2) consists of that of neon (Ne) and seven additional electrons in 3s and 3p orbitals. In iron (B3), a transition metal of the first series, electrons occupy the 4s orbital even though the 3d orbitals are still partly empty. Many reactions of the transition metals involve empty d orbitals—e.g., redox reactions or the formation of complexes with bases. 

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