|Chemistry 8th Edition / Chang|
|Student Study Guide
ELECTRON CONFIGURATIONS AND THE AUFBAU PRINCIPLE (7.8–7.9)
The Energies of Orbitals. The way electrons are distributed among the orbitals of an atom is its electronic structure or electron configuration. Each of the principal energy levels (called shells) of an atom can be divided into subshells, which are sets of orbitals. The s subshell has one s orbital, the p subshell has three p orbitals, the d subshell has five d orbitals, and so on. In addition each subshell has an electron capacity based on a maximum of two electrons per orbital. The orbitals available to a shell and the subshell capacities are summarized in Table 7.3.
Table 7.3 Allowed Numbers of Orbitals per Energy Level
1 0 1s 1 2 2 0 2s 1 2 2 1 2p 3 6 3 0 3s 1 2 3 1 3p 3 6 3 2 3d 5 10 4 0 4s 1 2 4 1 4p 3 6 4 2 4d 5 10 4 3 4f 7 14
In a hydrogen atom all of the orbitals in a particular energy level (same n value) have the same energy. In a hydrogen atom the electron resides in the 1s orbital because it has the lowest possible energy in that orbital. The ground state electron configuration of H is 1s1, as shown below.
Atoms of elements other than hyrogen have more than one electron and are called many-electron atoms. In these atoms each principal energy level is split into several energies. These energies correspond to the orbitals of one type having a slightly different energy than orbitals of another type. This splitting is shown in Figure 7.1. The energy of an orbital in a many-electron atom depends on both n and l. For a given value of n, the energy of orbitals increases as l increases. Therefore, for atoms other than hydrogen atoms, the 2p orbital has a higher energy than the 2s orbital.
Figure 7.1 The orbital energies in a many-electron atom. Within a shell the orbital energies increase as the l quantum number increases. The energy of an orbital depends on its n and l values. Here the 4s subshell is shown to be lower than the 3d subshell.
The order of increasing energy for orbitals within a given principal energy level is
s < p < d < f < …
This trend results from the fact that s orbitals have a greater electron density near the nucleus than p orbitals, and p orbitals have greater electron density near the nucleus than d orbitals. In a many-electron atom, the energies of the orbitals increase as follows (Figure 7.1):
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s …
Note also that the higher principal energy levels (n values) have more subshells, and as these split in energy some of these subshells can actually be lower in energy than subshells from a lower principal energy level. Thus the energy of a 4s orbital is very close to the energy of a 3d orbital in a many-electron atom. In potassium and calcium atoms the energy of the 4s orbital is less than the energy of the 3d orbital.
Electron Configurations. To write the electron configuration for the ground state of a many-electron atom, recall that the order in which subshells are occupied is from lowest energy to highest energy. The order shown in Figure 7.1 can easily be remembered by drawing Figure 7.24 of the textbook. The general approach of building up the atoms of the elements by adding a proton to the nucleus and an electron to an orbital is called the Aufbau or building up principle.
Helium has two electrons both of which can occupy the 1s orbital. The electron configuration is: He 1s2. A lithium atom has three electrons. According to the Pauli principle, we can place only two electrons in the 1s orbital; the third electron must enter the next higher orbital, the 2s orbital. Its electron configuration is: Li 1s22s1. The electron configurations of the second period elements are:
With Be the 1s and 2s orbitals are filled, and so additional electrons must enter the 2p orbitals. The 2p subshell is being filled over the next six elements.
B 1s22s22p1 C 1s22s22p2 N 1s22s22p3 O 1s22s22p4 F 1s22s22p5 Ne 1s22s22p6
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