Subshells relative energies

FIGURE 1.41 The relative energies of the shells, subshells, and orbitals in a many-electron atom. Each of the boxes can hold at most two electrons. Note the change in the order of energies of the 3d- and 4s-orbitals after Z = 20. 

FIGURE 9. Approximate Relative Energy Levels of Subshells... 

Before continuing the discussion of bonding theory, it is necessary to review briefly the electronic structure of the atom. Note that electrons in atoms are described as occupying orbitals which in turn constitute subshells and shells. Each orbital can hold no more than two electrons and electrons occupy the lowest energy orbitals of the atom. Figure 2.1 illustrates relative energies of atomic orbitals, where small circles represent orbitals. Note there is one s orbital in each shell, and three p orbitals, five d orbitals, and seven / orbitals in each shell. 

Figure 3.7 The relative energies and electron-filling order for shells and subshells.

FIGURE 4.5 Approximate relative energies of the subshells in the many-electron atom, reflecting the shifts due to nuclear shielding. The Is and 2s orbital energies are too low to appear in this plot. Actual relative energies will vary with number and excitation of electrons. 

FigM e 2.6 Schematic representation of the relative energies of the electrons for the various shells and subshells. 

This can be explained in terms of the relative stability of different electronic configurations and thus provides evidence for these electronic configurations. To help you understand this, you have to appreciate that there is a special stability associated with a filled subshell or a half-filled subshell - for example, the p subshell when it contains three or six electrons. Likewise, the d subshell is most stable when it contains five or ten electrons. The more stable the electronic configuration, then the more difficult it is to remove an electron and therefore the ionisation energy is higher. 

In looking for other trends in the data of Figure 6.6, the near-zero Eea s of the alkaline earth metals (Be, Mg, Ca, Sr, Ba) are particularly striking. Atoms of these elements have filled s subshells, which means that the additional electron must go into a p subshell. The higher energy of the p subshell, together with a relatively low Zeff for elements on the left side of the periodic table, means that alkaline earth atoms accept an electron reluctantly and have Eea values near zero. 

A specified electronic configuration is a statement of how many electrons are present in each energy level, described in terms of orbitals. For example, if we say that the Ti2+ ion has the d2 configuration, we are simply saying that there are two electrons in the fivefold-degenerate 3d subshell. The statement does not specify whether these electrons are in the same 3d orbital, or in two different orbitals it does not specify which orbital(s) is/are occupied, or the relative spins of the electrons if they are in different 3d orbitals. 

---