The Chemistry of Transition Metals

By Laird, B.B.

Edited by Paul Ducham

TRANSITION METALS HAVE ELECTRONIC CONFIGURATIONS WITH INCOMPLETE D OR F SHELLS

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).

Electron Configurations
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.

Oxidation States
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.

Iron
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:

Eq 15.1

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
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:

Eq 15.2

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:

Eq 15.3

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).

Figure 15.1

Table 15.1

Table 15.2

Figure 15.2

Figure 15.3

Figure 15.4

Figure 15.5 and 15.6

TRANSITION METALS CAN FORM A VARIETY OF COORDINATION COMPOUNDS

Transition metals have a stong tendency to form complex ions (see p. 648). A coordination compound typically consists of a complex ion and counter ion. (Some coordination compounds such as Fe(CO)5 do not contain complex ions.) Much of what we now know about coordination compounds stems from the classic work of Alfred Werner,1 who prepared and characterized many coordination compounds. In 1893, at the age of 26, Werner proposed what is now commonly referred to as Werner’s coordination theory.
   Nineteenth-century chemists were puzzled by a certain class of reactions that seemed to violate valence theory. For example, the valences of the elements in cobalt(III) chloride and in ammonia seem to be completely satisfied, and yet these two substances react to form a stable compound having the formula CoCl3 • 6NH3. To explain this behavior, Werner postulated that most elements exhibit two types of valence: primary valence and secondary valence. In modern terminology, primary valence corresponds to the oxidation number and secondary valence to the coordination number of the element. In CoCl3 • 6NH3, according to Werner, cobalt has a primary valence of +3 and a secondary valence of +6.
   Today we use the formula [Co(NH3)6]Cl3 to indicate that the ammonia molecules and the cobalt atom form a complex ion; the chloride ions are not part of the complex but are held to it by ionic forces. Most, but not all, of the metals in coordination compounds are transition metals.
   The molecules or ions that surround the metal in a complex ion are called ligands (Table 15.3). Every ligand has at least one unshared pair of valence electrons, as these following examples demonstrate:

Eq 15.4

Ligands, therefore, act as Lewis bases, donating one or more electron pairs to the metal. On the other hand, the transition metal atom (in either its neutral or positively charged state) acts as a Lewis acid, accepting (and sharing) pairs of electrons from the ligands. As a result, the metal-ligand bonds are usually coordinate covalent bonds.
   The atom in a ligand that is bound directly to the metal atom is known as the donor atom. For example, nitrogen is the donor atom of the NH3 ligand in the [Cu(NH3)4]2 complex ion. The coordination number in coordination compounds is the number of donor atoms surrounding the central metal atom in a complex ion. For example, the coordination number of Ag in [Ag(NH3)2]+ is 2, of Cu2+ in [Cu(NH3)4]2+ is 4, and of Fe3+ in [Fe(CN)6]3- is 6. The most common coordination numbers are 4 and 6, but other coordination numbers, such as 2 or 5, are also known.
   Depending on the number of donor atoms present, ligands are classified as monodentate, bidentate, or polydentate (see Table 15.3). H2O and NH3 are monodentate ligands because they have only one donor atom each. Ethylenediamine (sometimes abbreviated “en”) is a bidentate ligand:

Eq 15.5

The two nitrogen atoms can coordinate with a metal atom as shown in Figure 15.7.
   Bidentate and polydentate ligands are also called chelating agents because they can hold the metal atom like a claw (from the Greek chele, meaning “claw”). Ethylenediaminetetraacetate ion (EDTA) is a polydentate ligand used to treat metal poisoning (Figure 15.8). Six donor atoms enable EDTA to form a very stable, water-soluble complex ion with lead. In this form, lead is removed from the blood and tissues and excreted from the body. EDTA is also used to clean up spills of radioactive metals.

Table 15.5

Fifure 15.7

OXIDATION NUMBERS OF METALS IN COORDINATION COMPOUNDS

Another important property of coordination compounds is the oxidation number of the central metal atom. The net charge of a complex ion is the sum of the charges on the central metal atom and its surrounding ligands. In the [PtCl6 ]2- ion, for example, each chloride ion has an oxidation number of -1, so the oxidation number of Pt must be +4. If the ligands do not bear net charges, the oxidation number of the metal is equal to the charge of the complex ion. Thus, in [Cu(NH3 )4]2+ each NH3 is neutral, so the oxidation number of Cu is +2.
   Example 15.1 shows how to determine the oxidation numbers of metals in coordination compounds.

Example 15.1

NAMING COORDINATION COMPOUNDS

Having discussed the various types of ligands and the oxidation numbers of metals, it is now time to learn how to name coordination compounds. The rules for naming coordination compounds are as follows:

1. The cation is named before the anion, as in other ionic compounds. The rule holds regardless of whether the complex ion bears a net positive or a net negative charge. In K3[Fe(CN)6] and [Co(NH3)4Cl2]Cl, for example, we name the K+ and [Co(NH3)4Cl2]+ cations first, respectively.
2. Within a complex ion, the ligands are named first, in alphabetical order, and the metal ion is named last.
3. The names of anionic ligands end with the letter o, whereas a neutral ligand is usually called by the name of the molecule. The exceptions are H2O (aquo), CO (carbonyl), and NH3 (ammine). Table 15.4 lists the names of some common ligands.
4. When several ligands of a particular kind are present, we use the Greek prefixes di-, tri-, tetra-, penta-, and hexa- to name them. Thus, the ligands in the cation [Co(NH3)4Cl2]+ are “tetraamminedichloro.” (Note that prefixes are ignored when alphabetizing the ligands.) If the ligand itself contains a Greek prefix, we use the prefixes bis (2), tris (3), and tetrakis (4) to indicate the number of ligands present. For example, the ethylenediamine ligand already contains the prefix di, so if a complex contains two such ligands then the name is bis(ethylenediamine).
5. The oxidation number of the metal is written in Roman numerals following the name of the metal. For example, the Roman numeral III is used to indicate the +3 oxidation state of chromium in [Cr(NH3)4Cl2]+, which is called tetraamminedichlorochromium( III) ion.
6. If the complex is an anion, its name ends in -ate. For example, in K4[Fe(CN)6] the anion [Fe(CN)6]4- is called hexacyanoferrate(II) ion. The Roman numeral II indicates the oxidation state of iron. Table 15.5 gives the names of anions containing metal atoms.

Examples 15.2 and 15.3 show how to apply these rules to the naming of coordination compounds.

Example 15.2

Example 15.3

Table 15.4

Table 15.5

STRUCTURE OF COORDINATION COMPOUNDS

In studying the geometry of coordination compounds, there is often more than one way to arrange the ligands around the central atom. Compounds rearranged in this fashion have distinctly different physical and chemical properties. Figure 15.9 shows four different geometric arrangements for metal atoms with monodentate ligands.
   In these diagrams the structure and coordination number of the metal atom relate to each other as follows:

Eq 15.6

   The three-dimensional spatial arrangements in coordination compounds can lead to stereoisomerism. Stereoisomers are compounds that are made up of the same types and numbers of atoms bonded together in the same sequence, but with different three-dimensional structures. The two types of stereoisomers are geometric isomers (isomers that cannot be interconverted without breaking a chemical bond) and optical isomers (isomers that are nonsuperimposable mirror images). Coordination compounds may exhibit one or both types of stereoisomerism. Many coordination compounds, however, do not have stereoisomers.
   We use the terms cis and trans to distinguish one geometric isomer of a compound from the other. Cis means that two particular atoms (or groups of atoms) are adjacent to each other, and trans means that the atoms (or groups of atoms) are on opposite sides in the structural formula. The cis and trans isomers of coordination compounds generally have quite different colors, melting points, dipole moments, and chemical reactivities. Figure 15.10 shows the cis and trans isomers of diamminedichloroplatinum(II). Although the types of bonds are the same in both isomers (two Pt—N and two Pt—Cl bonds), the spatial arrangements are different. Another example is tetraamminedichlorocobalt(III) ion, shown in Figure 15.11.
   As an example of optical isomerism in coordination compounds, Figure 15.12 shows the cis and trans isomers of dichlorobis(ethylenediamine)cobalt (III) ion and their mirror images. The trans isomer and its mirror image are superimposable, but the cis isomer and its mirror image are not. The cis isomer and its mirror image are, therefore, optical isomers. Unlike geometric isomers, optical isomers have identical physical and chemical properties, except for the way in which they interact with polarized light and in the way they react with other chiral molecules.

Figure 15.9

Figure 15.10

Figure 15.11

Figure 15.12

CRYSTAL FIELD SPLITTING IN OCTAHEDRAL COMPLEXES

Crystal field theory explains the bonding in complex ions purely in terms of electrostatic forces. In a complex ion, two types of electrostatic interaction come into play. One is the attraction between the positive metal ion and the negatively charged ligand or the partially negatively charged end of a polar ligand. This is the force that binds the ligands to the metal. The second type of interaction is electrostatic repulsion between the lone pairs on the ligands and the electrons in the d orbitals of the metals.
   The d orbitals have different orientations, but in the absence of external disturbance, they all have the same energy. In an octahedral complex, a central metal atom is surrounded by six lone pairs of electrons (on the six ligands), so all five d orbitals experience electrostatic repulsion. The magnitude of this repulsion depends on the orientation of the particular d orbital. Figure 15.13 shows, for example, that the lobes of the dx2-y2 orbital point toward the corners of the octahedron along the x and y axes, where the lone-pair electrons are positioned. Thus, an electron residing in this orbital would experience a greater repulsion from the ligands than an electron would in, say, the dxy orbital. For this reason, the energy of the dx2-y2 orbital is increased relative to the dxy, dyz, and dxz orbitals. The energy of the dz2 orbital is also greater, because its lobes are pointed at the ligands along the z axis.
   As a result of these metal-ligand interactions, the five d orbitals in an octahedral complex are split between two sets of energy levels: a higher level with two orbitals (dx2-y2 and dz2) having the same energy and a lower level with three equal energy orbitals (dxy, dyz, and dxz), as shown in Figure 15.14. The crystal field splitting (∆) is the energy difference between the two sets of d orbitals in a metal atom when ligands are present. The magnitude of ∆, which depends on the metal and the nature of the ligands, has a direct effect on the color and magnetic properties of complex ions.

Figure 15.13

Figure 15.14

COLOR

We learned that white light, such as sunlight, is a combination of all colors. A substance appears black if it absorbs all the visible light that strikes it. It appears white or colorless if it absorbs no visible light. An object appears green if it absorbs all light but refl ects the green component. An object also looks green if it refl ects all colors except red, the complementary color of green (Figure 15.15).
   What is true for refl ected light also applies to transmitted light (that is, the light that passes through a medium, such as a solution). Consider the hydrated cupric ion ([Cu(H2O)6]2+), for example. It absorbs light in the orange region of the spectrum, so an aqueous solution of CuSO4 appears blue to us. When the energy of a photon is equal to the difference between the ground state and an excited state, absorption occurs as the photon strikes the atom (or ion or compound) and an electron is promoted to a higher level. This concept makes it possible to cal- culate the energy change involved in the electron transition. The energy of a photon, given by Equation 1.3, is:

E = hv

where h represents Planck’s constant (6.626 X 10-34 J s) and v is the frequency of the radiation, which is 5.00 X 1014 s-1 for the wavelength of orange light (600 nm). Here E = ∆, the crystal field splitting, so we have

E = hv
    = (6.626 X 10-34 J s)(5.00 X 1014 s-1
    = 3.32 
X 10-19 J 

(This is the energy absorbed by one ion.) If the wavelength of the photon absorbed by an ion lies outside the visible region, then the transmitted light looks the same (to us) as the incident light—white—and the ion appears colorless.
   The best way to measure crystal field splitting is to use spectroscopy to deter- mine the wavelength at which light is absorbed. The [Ti(H2 O)6 ]3+ ion provides a straightforward example because Ti3+ has only one 3d electron (Figure 15.16). The [Ti(H2 O)6 ]3+ ion absorbs light in the visible region of the spectrum (Figure 15.17).
   The wavelength corresponding to maximum absorption in Figure 15.16(b) is 498 nm. Thus, the crystal field splitting can be calculated as follows. We start by writing

E = hv

Also

Eq 15.7

where c is the speed of light and l is the wavelength. Therefore,

Eq 15.8

This is the energy required to excite one [Ti(H2 O)6 ]3+ ion. To express this energy difference in the more convenient units of kilojoules per mole, we write

Eq 15.9

Aided by spectroscopic data for a number of complexes, all having the same metal ion but different ligands, chemists have calculated the crystal field splitting for each ligand and established a spectrochemical series, which is a list of ligands arranged in increasing order of their abilities to split the d-orbital energy levels:

Eq 15.10

These ligands are arranged in the order of increasing value of ∆. CO and CN 2, which cause a large splitting of the d-orbital energy levels, are called strong-field ligands. The halide ions and hydroxide ion, which split the d orbitals to a lesser extent, are called weak-field ligands.

Figure 15.15

Figure 15.16

Figure 15.17

MAGNETIC PROPERTIES

The magnitude of the crystal field splitting also determines the magnetic properties of a complex ion. The [Ti(H2 O)6]3+ ion, having only one d electron, is always paramagnetic. For an ion with several d electrons, however, the situation is not as clear cut. Consider, for example, the octahedral complexes [FeF6]3- and [Fe(CN)6]3- (Figure 15.18).
   Both contain Fe+3, the electron configuration of which is [Ar]3d5, but there are two possible ways to distribute the five d electrons among the d orbitals. According to Hund’s rule, maximum stability is reached when the electrons are placed in five separate orbitals with parallel spins. This arrangement can be achieved only if two of the five electrons are promoted to the higher-energy dx2-y2 and dz2 orbit- als. No such energy input is needed if all five electrons enter the dxy, dyz, and dxz orbitals. According to the Pauli exclusion principle , four of these elec- trons must pair up, leaving only one unpaired electron in this arrangement.
   Figure 15.19 shows the distribution of electrons among the d orbitals that results in low-spin and high-spin complexes. The actual arrangement of the electrons is deter- mined by the amount of stability gained by having maximum parallel spins versus the investment in energy required to promote electrons to higher d orbitals. Because F- is a weak-field ligand, the five d electrons enter five separate d orbitals with parallel spins to create a high-spin complex (see Figure 15.19). The cyanide ion is a strong-field ligand, on the other hand, so the energy needed to promote two of the five d electrons to higher d orbitals is too much, and therefore a low-spin complex is formed.
   The actual number of unpaired electrons (or spins) in a complex ion can be found by using a technique called electron spin resonance spectroscopy (ESR), and in general, experimental findings support predictions based on crystal field splitting. However, a distinction between low- and high-spin complexes can be made only if the metal ion contains more than three and fewer than eight d electrons, as shown in Figure 15.19.

Figure 15.18

Figure 15.19

TETRAHEDRAL AND SQUARE-PLANAR COMPLEXES

So far we have concentrated on octahedral complexes. The splitting of the d orbital energy levels in tetrahedral and square-planar complexes can also be accounted for satisfactorily by the crystal field theory. In fact, the splitting pattern for a tetrahedral ion is just the reverse of that for octahedral complexes. In this case, the dxy, dyz, and dxz orbitals are more closely directed at the ligands and therefore have more energy than the dx2-y2 and dz2 and dz2 orbitals (Figure 15.20). Most tetrahedral complexes are high spin. Presumably, the tetrahedral arrangement reduces the magnitude of the metal-ligand interactions, resulting in a smaller  value. This is a reasonable assumption because the number of ligands is smaller in a tetrahedral complex. 

   As Figure 15.21 shows, the splitting pattern for square-planar complexes is the most complicated. The dx2-y2 and dz2 orbital possesses the highest energy (as in the octahedral case), and the dxy orbital the next highest. However, the relative placement of the dz2 and the dxz and dyz orbitals cannot be determined simply by inspection and must be calculated.

Figure 15.20

Figure 15.21

LIGAND FIELD THEORY

Crystal field theory, although quite successful at explaining the spectral and magnetic properties of many coordination compounds, ignores the covalent character of metal- ligand bonds because it is based solely on electrostatic interactions. Thus, it would be unable to predict the fact that a neutral ligand, such as CO or ethylenediamine, can have a larger crystal field splitting than an ionic ligand, such as Cl- or F-. In order to take into account the covalent character of metal-ligand bonds, a molecular orbital approach (see Sections 3.5 and 4.5) to the electronic structure of coordination compounds, known as ligand field theory, has been developed.
  Ligand field theory is based on the idea that atomic orbitals that are close in energy will mix more effectively in molecular orbitals than those that are far apart. Consider, for example, the octahedral complex ion Eq 15.12. From Figure 15.13, we saw that the dx2-y2 and dz2 atomic orbitals are oriented toward the ligands. As such, the 3dx2-y2 and 3dz2 orbitals on the metal center will mix with the ligand lone-pair orbit- als to form two bonding and two antibonding s molecular orbitals. The remaining 3d orbitals—3dxy, 3dyz, and 3dxz —are oriented in between the ligands and, because of minimal overlap, they will not mix with the ligand lone-pair orbitals, that is, they will be nonbonding orbitals. The 4s orbital on the metal center is spherical in shape and will overlap with all of the lone-pair ligand orbitals to form a pair of bonding and antibonding σ molecular orbitals. The remaining three 4p orbitals on the metal center will overlap individually with the lone-pair ligand orbitals along the x, y, and z axes. So we have the six 4s, 4p, 3dx2-yand 3dz2 orbitals of the metal center mixing with the ligand lone-pair orbitals to form six bonding and six antibonding σ molecular orbitals oriented along the six vertices of an octahedron, with three nonbonding orbit- als corresponding to the 3dxy, 3dyz, and 3dxz orbitals. 

   The MO energy-level diagram for Eq 15.12 is shown in Figure 15.22. The arrangement of the highest occupied molecular orbitals (3dxy, 3dyz, 3dxz, and the two Eq 15.13 orbitals) is identical to that predicted by crystal field theory with the 3dx2-y2 and dzorbitals replaced with the two Eq 15.13 molecular orbitals. However, the MO-based ligand field theory is more complete as it provides an understanding of the dependence of the crystal field splitting on the ligand type. In this octahedral complex, the 3dxy, 3dyz, and 3dxz orbitals are not oriented along the metal-ligand bond axes, but instead are oriented in between the ligands. As such, they will interact most strongly with the orbitals of the ligand (p or π) that are perpendicular to the metal-ligand bond. In the case of monatomic ionic ligands, such as OH-, F-, Cl-, Br-, and I-, the electrostatic repulsion from the full p orbitals on these ligands raises the energy of the 3dxy, 3dyz, and 3dxz orbitals resulting in a reduction of the crystal field splitting ()—thus, these species are weak-field ligands. This repulsion is stronger if the p electrons are less tightly bound to the ligand, that is, if the ligand is less electronegative, thus explaining why I- has a lower crystal field splitting than F- in the spectrochemical series. On the other hand, ligands with unoc- cupied antibonding π* orbitals, such as CO or CN-, tend to lower the energy of the 3dxy, 3dyz, and 3dxz orbitals, leading to an increase in the crystal field splitting. Such species are strong-field ligands. The lowering of the 3dxy, 3dyz, and 3dxz orbital energies in this fashion is called back bonding and is due to mixing with the unoccupied π* orbitals on the ligand, which are oriented away from the ligand double (or triple bond) and toward the 3dxy, 3dyz, and 3dxz orbitals of the metal center (see Figure 3.18). Neutral ligands, such as NH3 and H2O, neither have a strong repulsion from p orbitals nor back bonding of the π orbitals and so are generally intermediate between weak- and strong- field ligands in the spectrochemical series.

Figure 15.22

APPLICATIONS OF COORDINATION COMPOUNDS

Coordination compounds are found in living systems and have many uses in the home, in industry, and in medicine.

Metallurgy
Coordination compounds can be used to extract and purify metals, such as gold and silver. Although gold and silver are usually found in the uncombined state in nature, in other metal ores they may be present in relatively small concentrations and are more difficult to extract. In a typical process, the crushed ore is treated with an aqueous cyanide solution in the presence of air to dissolve the gold by forming the soluble complex ion [Au(CN)2 ]- :

Eq 15.14

The complex ion [Au(CN)2 ]- (along with some cation, such as Na+ ) is separated from other insoluble materials by filtration and treated with an electropositive metal such as zinc to recover the gold:

Eq 15.15

Figure 15.23 shows an aerial view of a “cyanide pond” used for the extraction of gold.

Therapeutic Chelating Agents
The chelating agent EDTA is used in the treatment of lead poisoning. Certain platinum-containing compounds can effectively inhibit the growth of cancerous cells.

Chemical Analysis
Although EDTA has a great affinity for a large number of metal ions (especially +2 and +3 ions), other chelates are more selective in their binding. For example, dimeth- ylglyoxime,

Eq 15.16

forms an insoluble brick-red solid with Ni2+ and an insoluble bright-yellow solid with Pd2+. These characteristic colors are used in qualitative analysis to identify nickel and palladium. Furthermore, the quantities of ions present can be determined by gravimetric analysis as follows: To a solution containing Ni2+ ions, we add an excess of dimethyl- glyoxime reagent, and a brick-red precipitate forms. The precipitate is then filtered, dried, and weighed. Knowing the formula of the complex (Figure 15.24), we can read- ily calculate the amount of nickel present in the original solution.

Detergents
The cleansing action of soap in hard water is hampered by the reaction of the Ca2+ ions in the water with the soap molecules to form insoluble salts or curds. In the late 1940s, the detergent industry introduced a “builder,” usually sodium tripolyphosphate, to cir- cumvent this problem. The tripolyphosphate ion is an effective chelating agent that forms stable, soluble complexes with Ca2+ ions. Sodium tripolyphosphate revolutionized the detergent industry. However, because phosphates are plant nutrients, phosphate- containing wastewater discharged into rivers and lakes caused algae to grow, depleting the waters of oxygen. As a result, most or all aquatic life eventually died. This process, called eutrophication, led many states to ban phosphate detergents beginning in the1970s, thus forcing manufacturers to reformulate their products to eliminate phosphates.

Figure 15.23

Figure 15.24