Foundations of Chemistry. Philippa B. Cranwell
orbital.
Just like s orbitals, p orbitals can only hold a maximum of two electrons. However as there are three p orbitals per energy level, there can be a maximum of six p electrons in total in a p sub‐shell. Elements that have filled or partially filled p orbitals in their outer shells are found in the p block of the periodic table, which consists of Groups 3–8 (also known as Groups 13–18).
d orbitals
The third principal quantum shell in an atom contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d). The shapes of the d orbitals are shown in Figure 1.8. These are slightly more complex than p orbitals and again point in different directions along the x, y, and z axes. Four of the d orbitals have a similar shape (although three have lobes that point between the axes and the fourth has lobes that point along the axes). The fifth d orbital (dz2) has a dumb‐bell shape with a donut ring around the middle. Again, each d orbital can hold a maximum of 2 electrons, so the maximum number of d electrons in any one energy level is 10. Elements that have outer electrons in d orbitals are found in the d block of the periodic table.
Figure 1.8 (a) The dxy orbital; (b) the dxz orbital; (c) the dyz orbital; (d) the dx2–y2orbital; (e) the dz2 orbital.
1.2.5 Describing electronic configurations
Now that you are familiar with the concept of principal quantum levels and atomic orbitals within these levels, the arrangement of electrons in atoms can be seen to follow a logical procedure.
Atomic orbitals are regions in space where there is a high probability of finding an electron. The n = 1 shell has one s orbital. The n = 2 shell has one s orbital and three p orbitals. The n = 3 shell has one s orbital, three p orbitals, and five d orbitals. Each atomic orbital can hold a maximum of two electrons.
As we move further away from the nucleus, the electrons and energy levels have higher energies. Within a principal quantum shell, the electrons in different types of orbitals have different energies. Electrons in s orbitals have a lower energy than electrons in p orbitals, which have a lower energy than electrons in d orbitals. So the order of energies of electrons in atomic orbitals is s < p < d. This is shown in Figure 1.9.
Figure 1.9 The relative energies of the orbitals in an atom. Relative energies not to scale.
Electrons fill atomic orbitals according to the following rules:
1 Electrons always go into the lowest energy orbital possible. This is called the Aufbau principle, sometimes known as the building‐up principle.
2 If there is more than one orbital of the same energy available, electrons always fill an unoccupied orbital first: Hund's rule.
3 If two electrons occupy the same orbital, they will have opposite spin: the Pauli exclusion principle.
We have already seen that the electrons for hydrogen and helium obey these rules. The lowest energy orbital is the 1s orbital, and this is filled at helium. The electrons are forced to have opposite spins in helium. The electron configuration for the atom is a shorthand that describes the occupancy of the orbitals. For hydrogen, the electron configuration is 1s1. For helium, the electron configuration is 1s2. The superscripted number represents the number of electrons in the 1s orbital for each element.
Once the 1s orbital is full, the next electron at lithium (Z = 3) must enter the second energy level, as this is the next lowest in energy. Within this energy level, the 2s orbital is of lower energy than the 2p orbitals. Thus the electron occupies the 2s orbital. The electron configuration of lithium is 1s22s1.
The next element, beryllium (Z = 4), has four electrons, and the fourth electron must pair up with the electron in the 2s orbital, as this is of lower energy than the 2p orbitals. The electron configuration is 1s22s2. The two electrons in the 2s orbital have opposite spins.
At boron (Z = 5), the 2s orbital is full, and so the next electron must occupy a 2p orbital. All are empty and of the same energy, and so we arbitrarily place the electron in the 2px orbital, although it could equally occupy 2py or 2pz. The electron configuration for boron is 1s22s22p1.
The next two electrons at carbon and nitrogen enter the other empty 2p orbitals. Nitrogen has the electron configuration 1s22s22p3. As all 2p orbitals are half‐filled at nitrogen, the next electron of oxygen must pair with another p electron to give one fully occupied and two half‐occupied 2p orbitals: 1s22s22p4. This can be visualised more easily by using the representation with electrons in boxes, as in Figure 1.10 for oxygen.
Figure 1.10 Electron arrangements in lithium, oxygen, and chlorine. Outer shell electrons are shown in red boxes.
The next electrons complete the remaining two 2p orbitals so that at neon (Z = 10), we have a filled second shell of electrons: 1s22s22p6. We will see that this is a very stable arrangement of electrons and has significant consequences for the chemical reactivity of the element.
Once the second shell is full at neon, the next electron at sodium (Z = 11) enters the third shell and occupies the 3s orbital. The filling process is then repeated as for the second row of elements. The electron arrangement for chlorine (Z = 17) is shown in Figure 1.10. Table 1.4 shows how electrons fill atomic orbitals for the first 36 elements – i.e. up to the end of the fourth row of the periodic table.
Table 1.4 Electron configurations for the first 36 elements.
| Element symbol | Atomic number | n = 1 shell | n = 2 shell | n = 3 shell | n = 4 shell | ||||
|---|---|---|---|---|---|---|---|---|---|
| 1sH | 2s1 | 2p1 | 3s | 3p | 3d | 4s | 4p | ||
| He | 2 | 2 |