subshells, and orbitals.* The Pauli exclusion principle requires that no two electrons in an atom may have identical values of all four quantum numbers. Because each orbital corresponds to a unique set of the first three quantum numbers and the spin quantum number has only two possible values, two electrons with opposite spins may occupy a given orbital. In Chapter 8 we will see that the properties of the nucleus are also dictated by quantum mechanics, and that the nucleus may also be thought of as having a shell structure.
Each shell corresponds to a different value of the principal quantum number. The periodic nature of chemical properties reflects the filling of successive shells as additional electrons (and protons) are added. Each shell corresponds to a ‘period’, or row, in the periodic table. The first shell (the K shell) has one subshell, the 1s, consisting of a single orbital (with quantum numbers n = 1, l = 0, m = 0. The 1s orbital accepts up to two electrons. Thus period 1 has two elements: H and He. If another proton and electron are added, the electron is added to the first orbital, 2s, of the next shell (the L shell). Such a configuration has the chemical properties of lithium, the first element of period 2. The second shell has 2 subshells, 2s (corresponding to l = 0) and 2p (corresponding to l = 1). The p subshell has 3 orbitals (which correspond to values for m of −1, +1, and 0), px, py, and pz, so the L shell can accept up to eight electrons. Thus, period 2 has eight elements.
There are some complexities in the filling of orbitals beyond the M shell, which corresponds to period 3. The 3d subshell is vacant in period 3 element in their ground states, and in the first two elements of period 4. Only when the 4s orbital is filled do electrons begin to fill the 3d orbitals. The five 3d orbitals are filled as one passes up the first transition series metals, Sc through Zn. This results in some interesting chemical properties, because which of the 3d orbitals are filled depends on the atom's environment, as we shall see in Chapter 7. Similarly, the second and third transition series metals correspond to filling of the 4d and 5d orbitals. The lanthanide and actinide rare earth elements correspond to the filling of the 4f and 5f shells (again resulting in some interesting properties, which we will consider subsequently). The predicted sequence in which orbitals are filled and their energy levels are shown in Figure 1.2. Figure 1.3 shows the electronic configuration of the elements.
Figure 1.2 The predicted sequence of orbital energies for electrons in atoms. S levels can hold 2 electrons, p, d, and f can hold 6, 10, and 14 respectively.
Figure 1.3 The periodic table of naturally occurring elements showing the electronic configuration of the elements. Only the last orbitals filled are shown, thus each element has electrons in the orbitals of all previous Group 18 elements (noble gases) in addition to those shown. Superscripts indicate the number of electrons in each subshell.
1.5.3 Some chemical properties of the elements
It is only the most loosely bound electrons, those in the outermost shells, that participate in chemical bonding, so elements sharing a similar outermost electronic configuration tend to behave similarly. Elements within the same column of the periodic table, or group, share outer electronic configurations and hence behave in a similar manner. Thus, the elements of Group 1, the alkalis, all have one electron in the outermost s orbital, and behave in a similar manner. The Group 18 elements, the noble, or rare, gases, all have a filled p subshell, and behave similarly.
Let's now consider several concepts that are useful in describing the behavior of atoms and elements: ionization potential, electron affinity, and electronegativity. The first ionization potential of an atom is the energy required to remove (i.e., move an infinite distance away) the least tightly bound electron. This is energy gained by the electron in reactions such as:
(1.1)
The first ionization potential of the elements is illustrated in Figure 1.4. The Second Ionization Potential is the energy required to remove a second electron, etc. The electron affinity is the energy given up in reactions such as:
(1.2)
Electronegativity is another parameter that is often used to characterize the behavior of the elements. It is a relative, unitless quantity determined from the differences in bond energy between an A–B molecule and the mean energies of A–A and B–B molecules. Electronegativity quantifies the tendency of an element to attract a shared electron when bonded to another element. For example, F has a higher electronegativity than H (the values are 3.8 and 2.5, respectively), thus the bonding electron in hydrogen fluoride, HF, is more likely to be found in the vicinity of F than of H. It is also useful in characterizing the nature of chemical bonds between elements, as we shall see in a subsequent section. Electronegativities of the elements are shown in Figure 1.5.
In general, first ionization potential, electron affinity, and electronegativities increase from left to right across the periodic table, and to a lesser degree from bottom to top. This reflects the shielding of outer electrons, particularly those in s orbitals, by inner electrons, particularly those in p orbitals, from the charge of the nucleus. Thus the outer 3s electron of neutral sodium is effectively shielded from the nucleus and is quite easily removed. On the other hand, the 2p orbitals of oxygen are not very effectively shielded, and it readily accepts two additional electrons. With the addition of these two electrons, the 2p orbital is filled and the 3s orbital effectively shielded, so there is no tendency to add a third electron. With the outer p (and s) orbitals filled, a particularly stable configuration is reached. Thus, Ne and other noble gases have little tendency to either add or give up an electron.
Figure 1.4 First ionization potential of the elements.
Figure 1.5 Electronegativities of the elements. Nonmetals are characterized by high electronegativity, metals by low electronegativity.
