Absolutely Small - Michael D. Fayer [72]
Now things change again in a big way. The energy level diagram for many electron atoms (Figure 11.1) shows that the 3d orbitals are above the 4s orbitals in energy, but they are below the 4p orbitals. As mentioned earlier in this chapter, the interposition of the 3d orbitals between the 4s and the 4p orbitals gives rise to the first transition series in the Periodic Table. There are five 3d orbitals. The Pauli Principle states that there can be at most two electrons in an orbital. Then there can be 10 electrons in the five 3d orbitals, resulting in the 10 transition metals, scandium through zinc (see the Periodic Table). So after Ca come 10 elements that arise from filling the 3d orbitals. All of these are metals. Many of them are very common metals that we are familiar with in everyday life, such as iron (Fe), copper (Cu), nickel (Ni), zinc (Zn), and chromium (Cr). They can readily form ions. The first two elements in a row, such as K and Ca or Na and Mg, always form cations with a particular charge, +1 for the first column (Na+1 and K+1) and +2 for the second column (Mg+2 and Ca+2). However, the transition metals can form a variety of cations. They are said to have various oxidation states. When a metal loses an electron it is said to be oxidized. The oxidation state is the number of electrons it loses.
Consider iron. It can form the oxidation states +2 and +3, that is, it forms the cations Fe+2 and Fe+3. Fe+2 is readily understandable. Fe can lose the two 4s electrons just like Ca to make the +2 oxidation state. In addition, Fe has six 3d electrons. Hund’s Rule says that electrons will stay unpaired if possible. Five electrons can go into one of each of the five 3d orbitals. This half-filled configuration is particularly stable. Iron is one 3d electron past the half-filled 3d orbitals, so it will readily lose a 3d electron in addition to the two 4s electrons to give an oxidation state of +3. So Fe can form salts like FeCl2 and FeCl3.
In addition to giving rise to the first transition series (first group of transition metals), the 3d electrons are involved in another important molecular phenomenon. We discussed that oxygen will form two covalent bonds (share two electrons with other atoms) to obtain the Ne configuration. An example is water, H2O. Sulfur, which is directly below oxygen, forms H2S, analogous to H2O. However, it can also form SF6 through involvement of the 3d orbitals, which are close in energy to the 3p orbitals. There is no equivalent for oxygen because the first set of d orbitals, the 3d’s, are much higher in energy than the 2s and 2p orbitals that are involved in bonding in the second row of the Periodic Table.
After the first series of transition metals are completed by filling the 3d orbitals, the next element is gallium (Ga). Ga is a metal, and like aluminum, it will form +3 ions. The configuration in which the 3d orbitals are completely filled is very stable, so Ga only forms +3 cations. The stability of the filled 3d orbitals can also be seen in zinc. Zn only forms +2 ions by losing the two 4s electrons. Following Ga are germanium (Ge), arsenic (As), and selenium (Se), which tend to form four, three, and two covalent bonds, respectively, to obtain the krypton (Kr) closed shell configuration. Like the elements immediately above Ge, As, and Se, additional bonds can be formed by involving the 4d electrons, which are very close in energy to the 4p orbitals. The next element is bromine, which is a halogen, and forms a -1 anion to obtain the Kr closed shell configuration. Finally, the row ends with krypton, which has a closed shell.
Larger Atoms and the Lanthanides and Actinides
The elements in the fifth row of the Periodic Table follow the same trends as those in the fourth row. The fifth row has the second series of transition metals. The elements in the sixth and seventh rows behave like those