Electron shells and subshells

Increasing the nuclear charge of an atom (together with its number of electrons) leads to the consecutive occupation by electrons of the electronic shells and subshells of higher and higher eneigy. This produces a quasi-periodicity (sometimes called periodicity in chemistry) of the valence shells, and as a consequence, a quasi-periodicity of all chemical and physical properties of the elements (reflected in the Mendeleev periodic table). 

Energy-level diagram of electron shells and subshells of the elements. 

Energy Level Diagram of Electron Shells and Subshells of the Elements... 

There is no single best form of the periodic table since the choice depends on the purpose for which the table is used. Some forms emphasize chemical relations and valence, whereas others stress the electronic configuration of the elements or the dependence of the periods on the shells and subshells of the atomic structure. The most convenient form for our purpose is the so-called long form with separate panels for the lanthanide and actinide elements (see inside front cover). There has been a lively debate during the past decade as to the best numbering system to be used for the individual... 

These are shown in Fig. 2.3 and illustrate most convincingly the various quantum shells and subshells described in the preceding section. The energy required to remove the I electron from an atom of hydrogen is 13.606 eV (i.e. 1312 kJ per mole of H atoms). This rises to 2372 kJ mol for He (Is-) since the positive charge on the helium nucleus is twice that of the... 

The location of an electron in an atom is described by a wavefunction known as an atomic orbital atomic orbitals are designated by the quantum numbers , l, and mi and fall into shells and subshells as summarized in Fig. 1.30. 

Electrons having the same value of n in an atom are said to be in the same shell. Electrons having the same value of n and the same value of / in an atom are said to be in the same subshell. (Electrons having the same values of n, /, and m in an atom are said to be in the same orbital.) Thus, the first two electrons of magnesium (Table 17-3) are in the first shell and also in the same subshell. The third and fourth electrons are in the same shell and subshell with each other. They are also in the same shell with the next six electrons (all have n = 2) but a different subshell (/ = 0 rather than 1). With the letter designations of Sec. 17.3, the first two electrons of magnesium are in the Is subshell, the next two electrons arc in the 2s subshell, and the next six electrons are in the 2p subshell. The last two electrons occupy the 3s subshell. 

Bohr clearly distinguished chemical properties due to outer electrons from radioactivity due to the interior of the atom, substituting the notion of "shells" and "subshells" of electrons for the earlier idea of "orbits" or circles and relating the filling of these shells to properties of groups within the periodic table. 135 The full import of this paper was not appreciated until after the war, by which time Bohr had changed his mind about some aspects of the theory. 

For example, lithium has an electron arrangement 2,1, but its electronic configuration is Is 2s. The characters in red indicate the shell and subshell. The numbers in blue indicate the number of electrons in that subshell. So the two electrons in the first shell of lithium atoms are located in the Is subshell or Is orbital. The one electron in lithium s second shell is in the 2s subshell or 2s orbital. Now consider carbon. It has the electron arrangement 2, 4. The two electrons in the first shell go into the Is orbital. The next subshell to be filled is the 2s orbital, which holds a maximum of two electrons. The remaining two electrons go into the next available subshell, which is 2p. So carbon has an electronic configuration Is 2s 2pl... 

Electron configuration of an atom indicates its extranuclear structure that is, arrangement of electrons in shells and subshells. Chemical properties of elements (their valence states and reactivity) can be predicted from electron configuration. 

Orbital A region around the nucleus of an atom where an electron is mostly likely to reside (compare with principal quantum number, shell, and subshell). 

The chapter will assume some prior knowledge of the principles of atomic structure, including quantum numbers, wavefunctions, principle shells and subshells and the radial distribution function, which indicates the probability of finding an electron as a function of distance from the nucleus. 

We have seen that the term symbol for the ground state of the hydrogen atom is -5i/i. For a helium atom t = 0, 5 = 0. J = 0, and the term symbol for the ground state is 5 . For an atom such as boron, we can make use of the fact that all closed shells and subshells (such as the He example just given) contribute nothing to the term symbol. Hence both the Is- and 2jZ electrons give L = S = J = 0. The 2p electron has L = I, 5 = J. and J = 1 i, yielding -Pyz carbon there are two p electrons. 

Table 2. FUled shell and subshell 7r-electron configurations for [42]...

Based on the analysis we have used to assign peaks in photoelectron spectra to shells and subshells in atoms, why is the peak at 0.42 MJ/mole in the K spectrum assigned to the = 4 shell (as opposed to being another subshell of = 3) Refer to the data in Table 1 of ChemActivity 10 Electron Configurations. 

We have noted that the valence electrons play an important role in determining the properties of an element, both physical and chemical. In fact, it is not only the number of valence electrons present in an atom, but also their energies which are important. The ionization energies of the electrons in an atom provide a measure of how tightly the electrons in each shell (and subshell) are held by the atom. 

The letters s, p, d, and/, corresponding to values of Z = 0,1,2, and 3, respectively, are conventionally used to designate sublevels. The number of electrons present in a given sublevel is limited to 2,6,10, and 14 for s, p, d, and/, sublevels, respectively. It follows that, since each orbital can be occupied by a maximum of 2 electrons, Aere is only 1 orbital in the 5 sublevel, 3 orbitals in the p sublevel, 5 in the d and 7 in the /. The value of the principal quantum number and the letter designating the azimuthal quantum number are written in sequence to designate both the shell and subshell. For example. Ad represents the d subshell of the fourth shell. 

The atomic subshells host a number of (2/ + 1) pairs of electrons. So the p subshell hosts three pairs of electrons designated by the quantum numbers ot = +1,0, — 1. These states differ in their orientation within the electron shell and possess normally identical energy. Physicists, for some obscure reason, call such levels degenerate. When, however, an external electric field or magnetic field is applied, the orientation of the magnetic moment of the p-, d-electrons, etc., makes a difference, and the three levels of the p term mentioned above (ot = +1, 0, —1) will differ in energy (Stark effect, Zeeman effect). 

---