Читать книгу Geochemistry - William M. White - Страница 18

We stated above that the atomic number of an element is its most important property. This is true because the number of electrons is determined by atomic number, and it is the electronic structure of an atom that largely dictates its chemical properties. The organization of the elements in the periodic table reflects this electronic structure.

The electronic structure of atoms, and indeed the entire organization of the periodic table, is determined by quantum mechanics and the quantization of energy, angular momentum, magnetic moment, and spin of electrons. Four quantum numbers, called the principal, azimuthal, magnetic, and spin quantum numbers and conventionally labeled n, l, m, and ms, control the properties of electrons associated with atoms. The first of these, n, which may take values 1, 2, 3, ..., determines most of the electron's energy as well as its mean distance from the nucleus. The second, l, which has values 0, 1, 2, ... n−1, determines the total angular momentum and the shape of the orbit. The third, m, which may have values −l, ... 0 ... l, determines the z component of angular momentum and therefore the orientation of the orbit. The fourth, ms, may have values of –½ or +½ and determines the electron's spin. The first three quantum numbers result in the electrons surrounding the nucleus being organized into shells, 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.