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Titanium metal is light and corrosion resistant. The Guggenheim Museum in Bilbao, Spain, is the largest titanium- clad building in the world. The exterior is covered with 344, ft^2 of 0.016-in.-thick titanium pieces, each with a unique shape.
Chapter 23
The d -Block Elements
Chapter 21 "Periodic Trends and the " and Chapter 22 "The " described the chemistry of the s -block and p -block elements. In this chapter, we survey the chemistry of the d -block elements, which are also called the transition metals. We again use valence electron configurations and periodic trends, as well as the principles of chemical bonding, thermodynamics, and kinetics, as tools to describe the properties and reactivity of these elements. Because all the d -block elements are metals, they do not have the extreme variability in chemistry that we saw among the elements of the p block. Instead, these elements exhibit significant horizontal and vertical similarities in chemistry, and all have a common set of characteristic properties due to partially filled d subshells.
Alloys and compounds of the d -block elements are important components of the materials the modern world depends on for its continuing technological development, while most of the first-row transition metals are essential for life. This chapter introduces some of the key industrial and biological roles of these elements. You will learn, for example, why copper, silver, and gold have been used for coins and jewelry since ancient times, how Cr3+^ impurities can be responsible for the characteristic colors of both rubies and emeralds, why an iron oxide was used in primitive compasses, why insects have greenish-blue blood, and why cobalt is an essential component of vitamin B 12.
half-filled 5 s subshell, with 5 s^14 d^4 and 5 s^14 d^6 valence electron configurations, respectively. Further complications occur among the third-row transition metals, in which the 4 f , 5 d , and 6 s orbitals are extremely close in energy. Although La has a 6 s^25 d^1 valence electron configuration, the valence electron configuration of the next element—Ce—is 6 s^25 d^04 f^2. From this point through element 71, added electrons enter the 4 f subshell, giving rise to the 14 elements known as the lanthanides. After the 4 f subshell is filled, the 5 d subshell is populated, producing the third row of the transition metals. Next comes the seventh period, where the actinides have three subshells (7 s , 6 d , and 5 f ) that are so similar in energy that their electron configurations are even more unpredictable.
As we saw in the s -block and p -block elements, the size of neutral atoms of the d - block elements gradually decreases from left to right across a row, due to an increase in the effective nuclear charge ( Z eff) with increasing atomic number. In addition, the atomic radius increases down a group, just as it does in the s and p blocks. Because of the lanthanide contraction , however, the increase in size between the 3 d and 4 d metals is much greater than between the 4 d and 5 d metals (Figure 23. "The Metallic Radii of the First-, Second-, and Third-Row Transition Metals"). (For more information on the lanthanides, see Chapter 7 "The Periodic Table and Periodic Trends", Section 7.3 "Energetics of Ion Formation".) The effects of the lanthanide contraction are also observed in ionic radii, which explains why, for example, there is only a slight increase in radius from Mo3+^ to W3+.
23.1 General Trends among the Transition Metals 2738
Figure 23.1 The Metallic Radii of the First-, Second-, and Third-Row Transition Metals
Because of the lanthanide contraction, the second- and third-row transition metals are very similar in size.
As you learned in Chapter 7 "The Periodic Table and Periodic Trends", electrons in ( n − 1) d and ( n − 2) f subshells are only moderately effective at shielding the nuclear charge; as a result, the effective nuclear charge experienced by valence electrons in the d -block and f -block elements does not change greatly as the nuclear charge increases across a row. Consequently, the ionization energies of these elements increase very slowly across a given row (Figure 7.10 "A Plot of Periodic Variation of First Ionization Energy with Atomic Number for the First Six Rows of the Periodic Table"). In addition, as we go from the top left to the bottom right corner of the d block, electronegativities generally increase, densities and electrical and thermal conductivities increase, and enthalpies of hydration of the metal cations decrease in magnitude, as summarized in Figure 23.2 "Some Trends in Properties of the Transition Metals". Consistent with this trend, the transition metals become steadily less reactive and more “noble” in character from left to right across a row. The relatively high ionization energies and electronegativities and relatively low enthalpies of hydration are all major factors in the noble character of metals such as Pt and Au.
23.1 General Trends among the Transition Metals 2739
Note the Pattern
Due to a small increase in successive ionization energies, most of the transition metals have multiple oxidation states separated by a single electron.
Note the Pattern:
Most compounds of transition metals are paramagnetic, whereas virtually all compounds of the p -block elements are diamagnetic.
The electronegativities of the first-row transition metals increase smoothly from Sc (χ = 1.4) to Cu (χ = 1.9). Thus Sc is a rather active metal, whereas Cu is much less reactive. The steady increase in electronegativity is also reflected in the standard reduction potentials: thus E ° for the reaction M2+(aq) + 2e−^ → M^0 (s) becomes progressively less negative from Ti ( E ° = −1.63 V) to Cu ( E ° = +0.34 V). Exceptions to the overall trends are rather common, however, and in many cases, they are attributable to the stability associated with filled and half-filled subshells. For example, the 4 s^23 d^10 electron configuration of zinc results in its strong tendency to form the stable Zn2+^ ion, with a 3 d^10 electron configuration, whereas Cu+, which also has a 3 d^10 electron configuration, is the only stable monocation formed by a first-row transition metal. Similarly, with a half-filled subshell, Mn2+^ (3 d^5 ) is much more difficult to oxidize than Fe2+^ (3 d^6 ). The chemistry of manganese is therefore primarily that of the Mn2+^ ion, whereas both the Fe2+^ and Fe3+^ ions are important in the chemistry of iron.
The transition metals form cations by the initial loss of the ns electrons of the metal, even though the ns orbital is lower in energy than the ( n − 1) d subshell in the neutral atoms. This apparent contradiction is due to the small difference in energy between the ns and ( n − 1) d orbitals, together with screening effects. The loss of one or more electrons reverses the relative energies of the ns and ( n − 1) d subshells, making the latter lower in energy. Consequently, all transition-metal cations possess dn valence electron configurations , as shown in Table 23.2 for the 2+ ions of the first-row transition metals.
23.1 General Trends among the Transition Metals 2741
Note the Pattern
All transition-metal cations have dn^ electron configurations; the ns electrons are always lost before the ( n − 1) d electrons.
Table 23.2 d -Electron Configurations of the Dications of the First-Row Transition Metals
Ti2+^ V2+^ Cr2+^ Mn2+^ Fe2+^ Co2+^ Ni2+^ Cu2+^ Zn2+ d^2 d^3 d^4 d^5 d^6 d^7 d^8 d^9 d^10
The most common oxidation states of the first-row transition metals are shown in Table 23.3 "Common Oxidation States of the First-Row Transition Metals*". The second- and third-row transition metals behave similarly but with three important differences:
- The maximum oxidation states observed for the second- and third-row transition metals in groups 3–8 increase from +3 for Y and La to +8 for Ru and Os, corresponding to the formal loss of all ns and ( n − 1) d valence electrons. As we go farther to the right, the maximum oxidation state decreases steadily, reaching +2 for the elements of group 12 (Zn, Cd, and Hg), which corresponds to a filled ( n − 1) d subshell.
- Within a group, higher oxidation states become more stable down the group. For example, the chromate ion ([CrO 4 ]2−) is a powerful oxidant, whereas the tungstate ion ([WO 4 ]2−) is extremely stable and has essentially no tendency to act as an oxidant.
- Cations of the second- and third-row transition metals in lower oxidation states (+2 and +3) are much more easily oxidized than the corresponding ions of the first-row transition metals. For example, the most stable compounds of chromium are those of Cr(III), but the corresponding Mo(III) and W(III) compounds are highly reactive. In fact, they are often pyrophoric , bursting into flames on contact with atmospheric oxygen. As we shall see, the heavier elements in each group form stable compounds in higher oxidation states that have no analogues with the lightest member of the group.
23.1 General Trends among the Transition Metals 2742
The acid–base character of transition-metal oxides depends strongly on the oxidation state of the metal and its ionic radius. Oxides of metals in lower oxidation states (less than or equal to +3) have significant ionic character and tend to be basic. Conversely, oxides of metals in higher oxidation states are more covalent and tend to be acidic, often dissolving in strong base to form oxoanions.
23.1 General Trends among the Transition Metals 2744
E X A M P L E 1
Two of the group 8 metals (Fe, Ru, and Os) form stable oxides in the + oxidation state. Identify these metals; predict the stoichiometry of the oxides; describe the general physical and chemical properties, type of bonding, and physical state of the oxides; and decide whether they are acidic or basic oxides.
Given: group 8 metals
Asked for: identity of metals and expected properties of oxides in + oxidation state
Strategy:
Refer to the trends outlined in Figure 23.1 "The Metallic Radii of the First-, Second-, and Third-Row Transition Metals", Figure 23.2 "Some Trends in Properties of the Transition Metals", Table 23.1 "Valence Electron Configurations of the First-Row Transition Metals", Table 23.2, and Table 23.3 "Common Oxidation States of the First-Row Transition Metals*" to identify the metals. Decide whether their oxides are covalent or ionic in character, and, based on this, predict the general physical and chemical properties of the oxides.
Solution:
The +8 oxidation state corresponds to a stoichiometry of MO 4. Because the heavier transition metals tend to be stable in higher oxidation states, we expect Ru and Os to form the most stable tetroxides. Because oxides of metals in high oxidation states are generally covalent compounds, RuO 4 and OsO 4 should be volatile solids or liquids that consist of discrete MO 4 molecules, which the valence-shell electron-pair repulsion (VSEPR) model predicts to be tetrahedral. Finally, because oxides of transition metals in high oxidation states are usually acidic, RuO 4 and OsO 4 should dissolve in strong aqueous base to form oxoanions.
Exercise
Predict the identity and stoichiometry of the stable group 9 bromide in which the metal has the lowest oxidation state and describe its chemical and physical properties.
23.1 General Trends among the Transition Metals 2745
C O N C E P T U A L P R O B L E M S
- The transition metals show significant horizontal similarities in chemistry in addition to their vertical similarities, whereas the same cannot be said of the s - block and p -block elements. Explain why this is so.
- The energy of the d subshell does not change appreciably in a given period. Why? What effect does this have on the ionization potentials of the transition metals? on their electronegativities?
- Standard reduction potentials vary across the first-row transition metals. What effect does this have on the chemical reactivity of the first-row transition metals? Which two elements in this period are more active than would be expected? Why?
- Many transition metals are paramagnetic (have unpaired electrons). How does this affect electrical and thermal conductivities across the rows?
- What is the lanthanide contraction? What effect does it have on the radii of the transition metals of a given group? What effect does it have on the chemistry of the elements in a group?
- Why are the atomic volumes of the transition elements low compared with the elements of groups 1 and 2? Ir has the highest density of any element in the periodic table (22.65 g/cm^3 ). Why?
- Of the elements Ti, Ni, Cu, and Cd, which do you predict has the highest electrical conductivity? Why?
- The chemistry of As is most similar to the chemistry of which transition metal? Where in the periodic table do you find elements with chemistry similar to that of Ge? Explain your answers.
- The coinage metals (group 11) have significant noble character. In fact, they are less reactive than the elements of group 12. Explain why this is so, referring specifically to their reactivity with mineral acids, electronegativity, and ionization energies. Why are the group 12 elements more reactive?
23.1 General Trends among the Transition Metals 2747
S T R U C T U R E A N D R E A C T I V I T Y
- Give the valence electron configurations of the 2+ ion for each first-row transition element. Which two ions do you expect to have the most negative E ° value? Why?
- Arrange Ru3+, Cu2+, Zn, Ti4+, Cr3+, and Ni2+ in order of increasing radius.
- Arrange Pt4+, Hg2+, Fe2+, Zr4+, and Fe3+ in order of decreasing radius.
- Of Ti2+, V2+, Mn2+, Fe2+, Co2+, Ni2+, and Zn2+, which divalent ion has the smallest ionic radius? Explain your reasoning.
A N S W E R S
- Ti2+, 3 d^2 ; V2+, 3 d^3 ; Cr2+, 3 d^4 ; Mn2+, 3 d^5 ; Fe2+, 3 d^6 ; Co2+, 3 d^7 ; Ni2+, 3 d^8 ; Cu2+, 3 d^9 ; Zn2+, 3 d^10. Because Z eff increases from left to right, Ti2+ and V2+ will have the most negative reduction potentials (be most difficult to reduce).
- Hg2+ > Fe2+ > Zr4+ > Fe3+ > Pt4+
23.1 General Trends among the Transition Metals 2748
Moreover, all dissolve readily in aqueous acid to produce hydrogen gas and a solution of the hydrated metal ion: M3+(aq).
Table 23.4 Some Properties of the Elements of Groups 3, 4, and 5
Group Element Z
Valence Electron Configuration
Electronegativity
Metallic Radius (pm)
Melting Point (°C)
Density (g/cm^3 )
Sc 21 4 s^23 d^1 1.36 162 1541 2. Y 39 5 s^24 d^1 1.22 180 1522 4. La 57 6 s^25 d^1 1.10 183 918 6.
3
Ac 89 7 s^26 d^1 1.10 188 1051 10. Ti 22 4 s^23 d^2 1.54 147 1668 4.
4 Zr 40 5 s^24 d^2 1.33 160 1855 6. Hf 72 6 s^25 d^24 f^14 1.30 159 2233 13. V 23 4 s^23 d^3 1.63 134 1910 6. 5 Nb 41 5 s^24 d^3 1.60 146 2477 8. Ta 73 6 s^25 d^34 f^14 1.50 146 3017 16.
The group 3 metals react with nonmetals to form compounds that are primarily ionic in character. For example, reacting group 3 metals with the halogens produces the corresponding trihalides: MX 3. The trifluorides are insoluble in water because of their high lattice energies, but the other trihalides are very soluble in water and behave like typical ionic metal halides. All group 3 elements react with air to form an oxide coating, and all burn in oxygen to form the so-called sesquioxides (M 2 O 3 ), which react with H2O or CO 2 to form the corresponding hydroxides or carbonates, respectively. Commercial uses of the group 3 metals are limited, but “mischmetal,” a mixture of lanthanides containing about 40% La, is used as an additive to improve the properties of steel and make flints for cigarette lighters.
Group 4 (Ti, Zr, and Hf)
Because the elements of group 4 have a high affinity for oxygen, all three metals occur naturally as oxide ores that contain the metal in the +4 oxidation state resulting from losing all four ns^2 ( n − 1) d^2 valence electrons. They are isolated by initial conversion to the tetrachlorides, as shown for Ti:
23.2 A Brief Survey of Transition-Metal Chemistry 2750
Equation 23.
2FeTiO 3 (s) + 6C(s) + 7Cl 2 (g) → 2TiCl 4 (g) + 2FeCl 3 (g) + 6CO(g)
followed by reduction of the tetrachlorides with an active metal such as Mg.
Note the Pattern
The chemistry of the group 4 metals is dominated by the +4 oxidation state. Only Ti has an extensive chemistry in lower oxidation states.
In contrast to the elements of group 3, the group 4 elements have important applications. Titanium (melting point = 1668°C) is often used as a replacement for aluminum (melting point = 660°C) in applications that require high temperatures or corrosion resistance. For example, friction with the air heats the skin of supersonic aircraft operating above Mach 2.2 to temperatures near the melting point of aluminum; consequently, titanium is used instead of aluminum in many aerospace applications. The corrosion resistance of titanium is increasingly exploited in architectural applications, as shown in the chapter-opening photo. Metallic zirconium is used in UO 2 -containing fuel rods in nuclear reactors, while hafnium is used in the control rods that modulate the output of high-power nuclear reactors, such as those in nuclear submarines.
Consistent with the periodic trends shown in Figure 23.2 "Some Trends in Properties of the Transition Metals", the group 4 metals become denser, higher melting, and more electropositive down the column (Table 23.4 "Some Properties of the Elements of Groups 3, 4, and 5"). Unexpectedly, however, the atomic radius of Hf is slightly smaller than that of Zr due to the lanthanide contraction. Because of their ns^2 ( n − 1) d^2 valence electron configurations, the +4 oxidation state is by far the most important for all three metals. Only titanium exhibits a significant chemistry in the +2 and +3 oxidation states, although compounds of Ti2+^ are usually powerful reductants. In fact, the Ti2+(aq) ion is such a strong reductant that it rapidly reduces water to form hydrogen gas.
Reaction of the group 4 metals with excess halogen forms the corresponding tetrahalides (MX 4 ), although titanium, the lightest element in the group, also forms dihalides and trihalides (X is not F). The covalent character of the titanium halides increases as the oxidation state of the metal increases because of increasing
23.2 A Brief Survey of Transition-Metal Chemistry 2751
Each titanium atom is surrounded by an octahedral arrangement of six sulfur atoms that are shared to form extended layers of atoms. Because the layers are held together by only van der Waals forces between adjacent sulfur atoms, rather than covalent bonds, the layers slide past one another relatively easily when a mechanical stress is applied.
Group 5 (V, Nb, and Ta)
Like the group 4 elements, all group 5 metals are normally found in nature as oxide ores that contain the metals in their highest oxidation state (+5). Because of the lanthanide contraction, the chemistry of Nb and Ta is so similar that these elements are usually found in the same ores.
Three-fourths of the vanadium produced annually is used in the production of steel alloys for springs and high-speed cutting tools. Adding a small amount of vanadium to steel results in the formation of small grains of V 4 C 3 , which greatly increase the strength and resilience of the metal, especially at high temperatures. The other major use of vanadium is as V 2 O 5 , an important catalyst for the industrial conversion of SO 2 to SO 3 in the contact process for the production of sulfuric acid. (For more information on sulfuric acid production, see Chapter 2 "Molecules, Ions, and Chemical Formulas", Section 2.6 "Industrially Important Chemicals".) In contrast, Nb and Ta have only limited applications, and they are therefore produced in relatively small amounts. Although niobium is used as an additive in certain stainless steels, its primary application is in superconducting wires such as Nb 3 Zr and Nb 3 Ge, which are used in superconducting magnets for the magnetic resonance imaging of soft tissues. Because tantalum is highly resistant to corrosion, it is used as a liner for chemical reactors, in missile parts, and as a biologically compatible material in screws and pins for repairing fractured bones.
Note the Pattern
The chemistry of the two heaviest group 5 metals (Nb and Ta) is dominated by the +5 oxidation state. The chemistry of the lightest element (V) is dominated by lower oxidation states, especially +4.
As indicated in Table 23.4 "Some Properties of the Elements of Groups 3, 4, and 5", the trends in properties of the group 5 metals are similar to those of group 4. Only vanadium, the lightest element, has any tendency to form compounds in oxidation
23.2 A Brief Survey of Transition-Metal Chemistry 2753
Figure 23.4 Aqueous Solutions of Vanadium Ions in Oxidation States of +2 to +
Because vanadium ions with different oxidation states have different numbers of d electrons, aqueous solutions of the ions have different colors: in acid V(V) forms the pale yellow [VO 2 ]+ ion; V(IV) is the blue vanadyl ion [VO]2+; and V(III) and V(II) exist as the hydrated V3+^ (blue-green) and V2+^ (violet) ions, respectively.
states lower than +5. For example, vanadium is the only element in the group that forms stable halides in the lowest oxidation state (+2). All three metals react with excess oxygen, however, to produce the corresponding oxides in the +5 oxidation state (M 2 O 5 ), in which polarization of the oxide ions by the high-oxidation-state metal is so extensive that the compounds are primarily covalent in character. Vanadium–oxygen species provide a classic example of the effect of increasing metal oxidation state on the protonation state of a coordinated water molecule: vanadium(II) in water exists as the violet hydrated ion [V(H 2 O) 6 ]2+; the blue-green [V(H 2 O) 6 ]3+^ ion is acidic, dissociating to form small amounts of the [V(H 2 O) 5 (OH)]2+ ion and a proton; and in water, vanadium(IV) forms the blue vanadyl ion [(H 2 O) 4 VO]2+, which contains a formal V=O bond (Figure 23.4 "Aqueous Solutions of Vanadium Ions in Oxidation States of +2 to +5"). Consistent with its covalent character, V 2 O 5 is acidic, dissolving in base to give the vanadate ion ([VO 4 ]3−), whereas both Nb 2 O 5 and Ta 2 O 5 are comparatively inert. Oxides of these metals in lower oxidation states tend to be nonstoichiometric.
Although group 5 metals react with the heavier chalcogens to form a complex set of binary chalcogenides, the most important are the dichalcogenides (MY 2 ), whose layered structures are similar to those of the group 4 dichalcogenides. The elements of group 5 also form binary nitrides, carbides, borides, and hydrides, whose stoichiometries and properties are similar to those of the corresponding group 4 compounds. One such compound, tantalum carbide (TiC), has the highest melting point of any compound known (3738°C); it is used for the cutting edges of high-speed machine tools.
Groups 6 and 7
Group 6 (Cr, Mo, and W)
As an illustration of the trend toward increasing polarizability as we go from left to right across the d block, in group 6 we first encounter a metal (Mo) that occurs naturally as a sulfide ore rather than as an oxide. Molybdenite (MoS 2 ) is a soft black mineral that can be used for writing, like PbS and graphite. Because of this similarity, people long assumed that these substances were all the same. In fact, the name molybdenum is derived from the Greek molybdos , meaning “lead.” More than 90% of the molybdenum produced annually is used to make steels for cutting tools, which
23.2 A Brief Survey of Transition-Metal Chemistry 2754
