Читать книгу Foundations of Chemistry - Philippa B. Cranwell - Страница 70

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 2p_x orbital, although it could equally occupy 2p_y or 2p_z. 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
Li	3	2	1
Be	4	2	2
B	5	2	2	1
C	6	2	2	2
N	7	2	2	3
O	8	2	2	4
F	9	2	2	5
Ne	10	2	2	6
Na	11	2	2	6	1
Mg	12	2	2	6	2
Al	13	2	2	6	2	1
Si	14	2	2	6	2	2
P	15	2	2	6	2	3
S	16	2	2	6	2	4
Cl	17	2	2	6	2	5
Ar	18	2	2	6	2	6
K	19	2	2	6	2	6		1
Ca	20	2	2	6	2	6		2
Sc	21	2	2	6	2	6	1	2
Ti	22	2	2	6	2	6	2	2
V	23	2	2	6	2	6	3	2
Cr	24	2	2	6	2	6	5	1
Mn	25	2	2	6	2	6	5	2
Fe	26	2	2	6	2	6	6	2
Co	27	2	2	6	2	6	7	2
Ni	28	2	2	6	2	6	8	2
Cu	29	2	2	6	2	6	10	1
Zn	30	2	2	6	2	6	10	2
Ga	31	2	2	6	2	6	10	2	1
Ge	32	2	2	6	2	6	10	2	2
As	33	2	2	6	2	6	10	2	3
Se	34	2	2	6	2	6	10	2	4
Br	35	2	2	6	2	6	10	2	5
Kr	36	2	2	6	2	6	10	2	6

Within Table 1.4, there are various points to note. Firstly, electron filling appears to follow the rules just listed up to element 18, argon, when the 3p orbitals are full. The next electron might be expected to occupy the first 3d orbital; however, it can be seen that the electron in potassium actually occupies the 4s orbital. The reason for this is that the energy of the electrons in a 4s orbital is slightly lower than the energy of the electrons in the 3d orbitals, because the 4s orbital is larger and more diffuse. This means the electron in potassium occupies the 4s orbital preferentially, giving potassium a single outer s electron and placing the element in the s block. The electron configuration of potassium is 1s²2s²2p⁶3s²3p⁶4s¹.

Secondly, the next electron at calcium also occupies the 4s orbital, giving an electron configuration of 1s²2s²2p⁶3s²3p⁶4s². However, once the 4s orbital is filled, the next electron at scandium occupies the first 3d orbital, as this is now lower in energy than the 4p orbitals.

A shorthand way to write electron configurations is to use the symbol for the noble gas element (Group 8) to represent electrons in filled shells. So the electron configuration for potassium, 1s²2s²2p⁶3s²3p⁶4s¹ can be written as [Ar]4s¹, where [Ar] represents the configuration 1s²2s²2p⁶3s²3p⁶.

Finally, you may notice another anomaly in filling atomic orbitals at Cr and Cu. These two elements have just a single 4s electron, whereas the preceding elements have a filled 4s shell. Cr has one electron in each of its 3d orbitals, and copper has two electrons in each 3d orbital, leaving just one 4s electron in each case. The usual reason given for this relates to the extra stability associated with having a filled or half‐filled set of d orbitals.