Valence subshells

The loss of these r-subshell valence electrons explains the common +1 and +2 charges on ions of these elements, except for helium, which is chemically inert. 

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... 

This problem clearly did not worry Stoner, who just went ahead and assumed that three quantum numbers could be specified in many-electron atoms. In any case, Stoner s scheme solved certain problems present in Bohr s configurations. For example, Bohr had assigned phosphorus the configuration 2,4,4,41, but this failed to explain the fact that phosphorus shows valencies of three and five. Stoner s configuration for phosphorus was 2,2,2,4,2,2,1, which easily explains the valencies, since it becomes plausible that either the two or the three outermost subshells of electrons form bonds. 

Boron (Z = 5) has five electrons. Two enter the ls-orbital and complete the n = 1 shell. Two enter the 2s-orbital. The fifth electron occupies an orbital of the next available subshell, which Fig. 1.41 shows is a 2p-orbital. This arrangement of electrons is reported as the configuration 1 s22s22p1 or He 2s22/ 1 (5), showing that boron has three valence electrons. 

FIGURE 2.2 When a main-group metal atom forms a cation, it loses its valence s-and p-electrons and acquires the electron configuration of the preceding noble-gas atom. The heavier atoms in Croups 1 S/lll and 14/IV retain their complete subshells of d-electrons. 

The valences of the rare-earth metals are calculated from their magnetic properties, as reported by Klemm and Bommer.14 It is from the fine work of these investigators that the lattice constants of the rare-earth metals have in the main been taken. The metals lutecium and ytterbium have only a very small paramagnetism, indicating a completed 4/ subshell and hence the valences 3 and 2, respectively (with not over 3% of trivalent ytterbium present in the metal). The observed paramagnetism of cerium at room temperature corresponds to about 20% Ce4+ and 80% Ce3+, that of praseodymium and that of neodymium to about 10% of the quadripositive ion in each case, and that of samarium to about 20% of the bipositive ion in equilibrium with the tripositive ion. 

The elements may be divided into types (Fig. 17-10), according to the position of the last electron added to those present in the preceding element. In the first type, the last electron added enters the valence shell. These elements are called the main group elements. In the second type, the last electron enters a d subshell in the next to last shell. These elements are the transition elements. The third type... 

Transition-metal and rare-earth atoms that contain partially occupied d or f valence subshells also give rise to spectral tine structure, often with very complicated multiplet splitting [2,27,28]. The spin-unpaired valence d or f electrons can undergo spin-orbit coupling with the unpaired core electron (remaining in the orbital from which the photoelectron was removed), producing multiple non-degenerate final states manifested by broad photoelectron peaks [2,27]. 

It is noteworthy that Rydberg orbital occupancies on the central atom (rY, final column of Table 3.29) are relatively negligible (0.01-0.03e), showing that d-orbital participation or other expansion of the valence shell is a relatively insignificant feature of hyperbonded species. However, the case of HLiH- is somewhat paradoxical in this respect. The cationic central Li is found to use conventional sp linear hybrids to form the hydride bonds, and thus seems to represent a genuine case of expansion of the valence shell (i.e., to the 2p subshell) to form two bonding hybrids. However, the two hydride bonds are both so strongly polarized toward H (93%) as to have practically no contribution from Li orbitals, so the actual occupancy of extra-valent 2pu orbitals ( 0.03< ) remains quite small in this case. 

For brevity, many chemists record the electron configuration of an atom by giving only its outermost subshell, like As for potassium or 4/ for calcium. These electrons are most distant from the positive nucleus and, therefore, are most easily transferred between atoms in chemical reactions. These are the valence electrons. 

In Table 4-3, the common oxidation numbers in the last column are interpreted as the result of either losing the valence electrons (leaving a positive ion) or gaining enough electrons to fill that valence subshell. Table 4-4 compares three ions and a neutral atom. 

The three long rows of metallic elements in the middle of the periodic table, constituting the rectangle from scandium (21) to mercury (80), are the transition metals. Each of these three rows reflects the filling of a -type subshell that holds up to 10 electrons. Figure 4-5 shows the valence subshell of the first series of transition metals. Notice the general increase in the number of electrons occupying the id subshell. 

You may recall from the discussion of electron transfer (see Table 4-4) that a stable configuration precisely filled an r-type subshell and a />-type subshell. Only five elements have atoms with their valence y>-subshells filled these are the inert gases in the far right column of the periodic table. Their lack of chemical reactivity is explained by their stable electron configurations. 

All of the elements in the left hand column of the periodic table, the alkali metals, have what number of valence electrons in what type of subshell ... 

Symbol Nd atomic number 60 atomic weight 144.24 a rare earth lanthanide element a hght rare earth metal of cerium group an inner transition metal characterized by partially filled 4/ subshell electron configuration [Xe]4/35di6s2 most common valence state -i-3 other oxidation state +2 standard electrode potential, Nd + -i- 3e -2.323 V atomic radius 1.821 A (for CN 12) ionic radius, Nd + 0.995A atomic volume 20.60 cc/mol ionization potential 6.31 eV seven stable isotopes Nd-142 (27.13%), Nd-143 (12.20%), Nd-144 (23.87%), Nd-145 (8.29%), Nd-146 (17.18%), Nd-148 (5.72%), Nd-150 (5.60%) twenty-three radioisotopes are known in the mass range 127-141, 147, 149, 151-156. 

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. 

The electronic configuration of an atom describes the number of electrons that an atom possesses, and the orbitals in which these electrons are placed. The arrangements of electrons in orbitals, subshells and shells are called electronic configurations. Electronic configurations can be represented by using noble gas symbols to show some of the inner electrons, or by using Lewis structures in which the valence electrons are represented by dots. 

The single-bond radii for the lanthanons (La to Lu, Fig. 11-10) show some interesting features. The magnetic properties26 require that most of the lanthanons have metallic valence 3, but that two of these metals, europium and ytterbium, have metallic valence 2. The interatomic distances reflect these valences dearly, as shown in Figure 11-10. The stability of metallic valence 2, rather than 3, for these two metals may be attributed to the special stability of a half-filled or completely filled 4/ subshell. 

Its importance lies in the fact that it may be used, at least for simple molecules, to place the measured photoabsorption spectrum on an absolute scale. However, great care should be exercised in the use of (11.31) and (11.32) since the summation includes transitions from valence orbitals to inner orbitals already occupied. Such transitions cannot be seen, of course, and thus calibration of subshell photoabsorption spectra by this method will give results that are only approximate.123... 

Aluminum loses three valence electrons to form the Al3+ ion, which has the same electron configuration as neon. When the atoms of metals on the left of the p block in Periods 4 and higher lose their s- and p-elec-trons, they leave a noble-gas core surrounded by an additional, complete subshell of d-electrons. For instance, gallium forms the ion Ga3+ with the configuration [Ar]3d10. The d-electrons of the p-block atoms are gripped tightly by the nucleus and, in most cases, cannot be lost. We saw in Section 1.18 that the inert-pair effect implies that the elements listed in Fig. 1.44 can lose either their valence p-electrons alone or all their valence p- and s-electrons. 

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