Transition metals have incompletely filled d subshells or readily give rise to ions with incompletely filled d subshells (Figure 15.1) (The Group 2B metals—Zn, Cd, and Hg— are sometimes treated as transition metals, but they do not have this characteristic electron configuration, so they really do not belong in this category.) This attribute is responsible for several notable properties, including distinctive coloring, formation of paramagnetic compounds, catalytic activity, and especially a strong tendency to form complex ions. We focus on the first-row elements from scandium to copper, the most common transition metals. Table 15.1 lists some of their properties.
As we read across any period from left to right, atomic numbers increase, electrons are added to the outer shell, and the nuclear charge increases by the addition of protons. In the third-period elements—sodium to argon—the outer electrons weakly shield one another from the extra nuclear charge. Consequently, atomic radii decrease rapidly from sodium to argon, and ionization energies and electronegativities increase steadily (see Figures 2.17, 2.22, and 3.39).
For the transition metals, the trends are different. According to Table 15.1, the nuclear charge increases from scandium to copper, but electrons are being added to the inner 3d subshell. These 3d electrons shield the 4s electrons from the increasing nuclear charge somewhat more effectively than outer-shell electrons can shield one another, so the atomic radii decrease less rapidly. For the same reason, electronegativities and ionization energies increase only slightly from scandium across to copper compared with the increases from sodium to argon.
Although the transition metals are less electropositive (more electronegative) than the alkali and alkaline earth metals, their standard reduction potentials suggest that all of them except copper should react with strong acids such as hydrochloric acid to produce hydrogen gas. However, most transition metals are inert toward acids or react slowly with them because of a protective layer of oxide. A case in point is chromium: Despite a rather negative standard reduction potential, it is quite inert chemically because of the formation on its surface of chromium(III) oxide (Cr2O3). Consequently, chromium is commonly used as a protective and noncorrosive plating on other metals. On automobile bumpers and trim, chromium plating serves a decorative as well as a functional purpose.
General Physical Properties
Most of the transition metals have relatively small atomic radii and a close-packed structure (see Figure 6.28) in which each atom has a coordination number of 12. The combined effect of small atomic size and close packing result in strong metallic bonds. Therefore, transition metals have higher densities, higher melting and boiling points, and higher heats of fusion and vaporization than the Group 1A, 2A, and 2B metals (Table 15.2).
The electron configurations of the first-row transition metals were discussed. Calcium has the electron configuration [Ar]4s2. From scandium across to copper, electrons are added to the 3d orbitals. Thus, the outer electron configuration of scandium is 4s23d1, that of titanium is 4s23d2, and so on. The two exceptions are chromium and copper, whose outer electron configurations are 4s13d5 and 4s13d10When the first-row transition metals form cations, electrons are removed first from the 4s orbitals and then from the 3d orbitals. (This is the opposite of the order in which orbitals are filled in atoms.) For example, the outer electron configuration of Fe2+ is 3d6, not 4s23d4.
Transition metals exhibit variable oxidation states in their compounds. Figure 15.2 shows that the common oxidation states for each element from scandium to copper include +2, +3, or both. The +3 oxidation states are more stable at the beginning of the series, whereas the +2 oxidation states are more stable toward the end. To understand this trend, you must examine the ionization energy plots in Figure 15.3. In general, the ionization energies increase gradually from left to right. However, the third ionization energy (when an electron is removed from the 3d orbital) increases more rapidly than the first and second ionization energies. Because it takes more energy to remove the third electron from the metals near the end of the row than from those near the beginning, the metals near the end tend to form M2+ ions rather than M3+ ions.
The highest oxidation state for a transition metal, that of manganese (4s23d5), is +7. For elements to the right of Mn (Fe to Cu), the oxidation numbers are lower. Transition metals usually exhibit their highest oxidation states in compounds with very electronegative elements such as oxygen and fl uorine. Examples include V2O5, CrO3, and Mn2O7.
Two Examples: The Chemistry of Iron and Copper
Figure 15.4 shows samples of the first-row transition metals. Here we will briefl y survey the chemistry of two of these elements—iron and copper—paying particular attention to their occurrence, preparation, uses, and important compounds.
After aluminum, iron is the most abundant metal in the Earth’s crust (6.2 percent by mass). It is found in many ores; some of the economically important ones are hematite (Fe2O3), siderite (FeCO3), and magnetite (Fe3O4) (Figure 15.5). Pure iron is a gray metal and is not particularly hard. Its ion is essential in living systems because it reversibly binds oxygen to hemoglobin, the protein in blood that carries oxygen from the lungs to the rest of the tissues of the body.
Iron reacts with hydrochloric acid to give hydrogen gas:
Concentrated sulfuric acid oxidizes the metal to Fe3+, but concentrated nitric acid renders the metal “passive” by forming a thin layer of Fe3O4 over the surface. One of the bestknown reactions of iron is rust formation. The two oxidation states of iron are +2 and +3. Iron(II) compounds include FeO (black), FeSO4 • 7H2O (green), FeCl2 (yellow), and FeS (black). In the presence of oxygen, Fe2+ ions in solution are readily oxidized to Fe3+ ions. Iron(III) oxide is reddish brown, and iron(III) chloride is brownish black.
Copper, a rarer element than iron (6.8 X 10-3 percent of Earth’s crust by mass), is found in nature in the uncombined state as well as in ores such as chalcopyrite (CuFeS2) (Figure 15.6). The reddish-brown metal is obtained by roasting the ore to give Cu2S and then metallic copper:
Impure copper can be purified by electrolysis. After silver, which is too expensive for large-scale use, copper has the highest electrical conductivity. It is also a good thermal conductor. Copper is used in alloys, electrical cables, plumbing (pipes), and coins.
Copper reacts only with hot concentrated sulfuric acid and nitric acid. Its two important oxidation states are +1 and +2. The +1 state is less stable and disproportionates in solution:
All compounds of Cu(I) are diamagnetic and colorless except for Cu2O, which is red. The Cu(II) compounds are all paramagnetic and colored. The hydrated Cu2+ ion is blue. Some important Cu(II) compounds are CuO (black), CuSO4 • 5H2O (blue), and CuS (black).