Chapter 24

(a) Draw the structure for \(\mathrm{Pt}(\mathrm{en}) \mathrm{Cl}_{2}\). (b) What is the coordination number for platinum in this complex, and what is the coordination geometry? (c) What is the oxidation state of the platinum? [Section 24.1]

(a) What is the difference between Werner's concepts of primary valence and secondary valence? What terms do we now use for these concepts? (b) Why can the \(\mathrm{NH}_{3}\) molecule serve as a ligand but the \(\mathrm{BH}_{3}\) molecule cannot?

(a) What is themeaning of the term coordination number as it applies to metal complexes? (b) Generally speaking, what structural feature characterizes substances that can serve as ligands in metal complexes? Give an example of a ligand that is neutral and one that is negatively charged. (c) Would you expect ligands that are positively charged to be common? Explain. (d) What type of chemical bonding is characteristic of coordination compounds? Illustrate with the compound \(\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{6} \mathrm{Cl}_{3}\)

A certain complex of metal \(\mathrm{M}\) is formulated as \(\mathrm{MCl}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\). The coordination number of the complex is not known but is expected to be 4 or 6 . (a) Would conductivity measurements provide information about the coordination number? (b) In using conductivity measurements to test which ligands are bound to the metal ion, what assumption is made about the rate at which ligands enter or leave the coordination sphere of the metal?

Indicate the coordination number of the metal and the oxidation number of the metal in each of the following complexes: (a) \(\mathrm{Na}_{2}\left[\mathrm{CdCl}_{4}\right]\) (b) \(\mathrm{K}_{2}\left[\mathrm{MoOCl}_{4}\right]\) (c) \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{Cl}_{2}\right] \mathrm{Cl}\) (d) \(\left[\mathrm{Ni}(\mathrm{CN})_{5}\right]^{3-}\) (e) \(\mathrm{K}_{3}\left[\mathrm{~V}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)_{3}\right]\) (f) \(\left[\mathrm{Zn}(\mathrm{en})_{2}\right] \mathrm{Br}_{2}\)

Indicate the coordination number of the metal and the oxidation number of the metal in each of the following complexes: (a) \(\mathrm{K}_{3}\left[\mathrm{Co}(\mathrm{CN})_{6}\right]\) (b) \(\mathrm{Na}_{2}\left[\mathrm{CdBr}_{4}\right]\) (c) \(\left[\mathrm{Pt}(\mathrm{en})_{3}\right]\left(\mathrm{ClO}_{4}\right)_{4}\) (d) \(\left[\mathrm{Co}(\mathrm{en})_{2}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)\right]^{+}\) (e) \(\mathrm{NH}_{4}\left[\mathrm{Cr}\left(\mathrm{NH}_{3}\right)_{2}(\mathrm{NCS})_{4}\right]\) (f) \(\left[\mathrm{Cu}(\mathrm{bipy})_{2} \mathrm{III}\right.\)

(a) What is the difference between a monodentate ligand and a bidentate ligand? (b) How many bidentate ligands are necessary to fill the coordination sphere of a six-coordinate complex? (c) You are told that a certain molecule can serve as a tridentate ligand. Based on this statement, what do you know about the molecule?

For each of the following polydentate ligands, determine (i) the maximum number of coordination sites that the ligand can occupy on a single metal ion and (ii) the number and type of donor atoms in the ligand: (a) ethylenediamine (en), (b) bipyridine (bipy), (c) the oxalate anion \(\left(\mathrm{C}_{2} \mathrm{O}_{4}{\underline{\phantom{xx}}}^{2-}\right)\), (d) the \(2-\) ion of the porphine molecule (Figure 24.8); (e) [EDTA]^{4- } .

Polydentate ligands can vary in the number of coordination positions they occupy. In each of the following, identify the polydentate ligand present and indicate the probable number of coordination positions it occupies: (a) \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{4}(0-\mathrm{phen})\right] \mathrm{Cl}_{3}\) (b) \(\left[\mathrm{Cr}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right] \mathrm{Br}\) (c) \(\left[\mathrm{Cr}(\mathrm{EDTA})\left(\mathrm{H}_{2} \mathrm{O}\right)\right]^{-}\) (d) \(\left[\mathrm{Zn}(\mathrm{en})_{2}\right]\left(\mathrm{ClO}_{4}\right)_{2}\)

Indicate the likely coordination number of the metal in each of the following complexes: (a) \(\left[\mathrm{Rh}(\mathrm{bipy})_{3}\right]\left(\mathrm{NO}_{3}\right)_{3}\) (b) \(\mathrm{Na}_{3}\left[\mathrm{Co}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)_{2} \mathrm{Cl}_{2}\right]\) (c) \(\left[\mathrm{Cr}(\mathrm{o} \text { -phen })_{3}\right]\left(\mathrm{CH}_{3} \mathrm{COO}\right)_{3}\) (d) \(\mathrm{Na}_{2}[\mathrm{Co}(\mathrm{EDTA}) \mathrm{Br}]\)

(a) What is meant by the term chelate effect? (b) What thermodynamic factor is generally responsible for the chelate effect? (c) Why are polydentate ligands often called sequestering agents?

Pyridine \(\left(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{~N}\right)\), abbreviated py, is the following molecule: (a) Why is pyridine referred to as a monodentate ligand? (b) Consider the following equilibrium reaction: \(\left[\mathrm{Ru}(\mathrm{py})_{4}(\mathrm{bipy})\right]^{2+}+2 \mathrm{py} \rightleftharpoons\left[\mathrm{Ru}(\mathrm{py})_{6}\right]^{2+}+\) bipy What would you predict for themagnitude of the equilibrium constant for this equilibrium? Explain the basis for your answer.

Write the formula for each of the following compounds, being sure to use brackets to indicate the coordination sphere: (a) hexaamminechromium(III) nitrate (b) tetraamminecarbonatocobalt(III) sulfate (c) dichlorobis(ethylenediamine)platinum(IV) bromide (d) potassium diaquatetrabromovanadate(III) (e) bis(ethylenediamine) zinc(II) tetraiodomercurate(II)

Write the formula for each of the following compounds, being sure to use brackets to indicate the coordination sphere: (a) tetraaquadibromomanganese(III) perchlorate (b) bis(bipyridyl)cadmium(II) chloride (c) potassium tetrabromo(ortho-phenanthroline)cobaltate (III) (d) cesium diamminetetracyanochromate(III) (e) tris(ethylenediammine)rhodium(III) tris(oxalato)cobaltate(III)

Write the names of the following compounds, using the standard nomenclature rules for coordination complexes: (a) \(\left[\mathrm{Rh}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{Cl}_{2}\right] \mathrm{Cl}\) (b) \(\mathrm{K}_{2}\left[\mathrm{TiCl}_{6}\right]\) (c) \(\mathrm{MoOCl}_{4}\) (d) \(\left[\mathrm{Pt}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)\right] \mathrm{Br}_{2}\)

Write names for the following coordination compounds: (a) \(\left[\mathrm{Cd}(\mathrm{en}) \mathrm{Cl}_{2}\right]\) (b) \(\mathrm{K}_{4}\left[\mathrm{Mn}(\mathrm{CN})_{6}\right]\) (c) \(\left[\mathrm{Cr}\left(\mathrm{NH}_{3}\right)_{5} \mathrm{CO}_{3}\right] \mathrm{Cl}\) (d) \(\left[\mathrm{Ir}\left(\mathrm{NH}_{3}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]\left(\mathrm{NO}_{3}\right)_{3}\)

(a) Draw the two linkage isomers of \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{5} \mathrm{SCN}\right]^{2+}\). (b) Draw the two geometric isomers of \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{3} \mathrm{Cl}_{3}\right]^{2+}\). (c) Two compounds with the formula \(\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{5} \mathrm{ClBr}\) can be prepared. Use structural formulas to show how they differ. What kind of isomerism does this illustrate?

A four-coordinate complex \(\mathrm{MA}_{2} \mathrm{~B}_{2}\) is prepared and found to have two different isomers. Is it possible to determine from this information whether the complex is square planar or tetrahedral? If so, which is it?

Consider an octahedral complex \(\mathrm{MA}_{3} \mathrm{~B}_{3}\). How many geometric isomers are expected for this compound? Will any of the isomers be optically active? If so, which ones?

Sketch all the possible stereoisomers of (a) tetrahedral \(\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2} \mathrm{Cl}_{2}\right]\), (b) square-planar \(\left[\mathrm{IrCl}_{2}\left(\mathrm{PH}_{3}\right)_{2}\right]^{-}\), (c) octa- hedral \(\left[\mathrm{Fe}(0 \text { -phen })_{2} \mathrm{Cl}_{2}\right]^{+}\)

(a) A complex absorbs light with wavelength of \(530 \mathrm{~nm}\). Do you expect it to have color? (b) A solution of a compound appears green. Does this observation necessarily mean that all colors of visible light other than green are absorbed by the solution? Explain. (c) What information is usually presented in a visible absorption spectrum of a compound? (d) What energy is associated with the absorption at \(530 \mathrm{~nm}\) in \(\mathrm{kJ} / \mathrm{mol}\) ?

In crystal-field theory, ligands are modeled as if they are point negative charges. What is the basis of this assumption, and how does it relate to the nature of metalligand bonds?

Explain why the \(d_{x y}, d_{x z}\), and \(d_{y z}\) orbitals lie lower in energy than the \(d_{z^{2}}\) and \(d_{x^{2}-y^{2}}\) orbitals in the presence of an octahedral arrangement of ligands about the central metalion.

(a) Sketch a diagram that shows the definition of the crystal-field splitting energy \((\Delta)\) for an octahedral crystal field. (b) What is the relationship between the magnitude of \(\Delta\) and the energy of the \(d-d\) transition for a \(d^{1}\) complex? (c) Calculate \(\Delta\) in \(\mathrm{kJ} / \mathrm{mol}\) if a \(d^{1}\) complex has an absorption maximum at \(590 \mathrm{~nm}\).

The \(\left[\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}\) ion has an absorption maximum at about \(725 \mathrm{~nm}\), whereas the \(\left[\mathrm{Ni}\left(\mathrm{NH}_{3}\right)_{6}\right]^{2+}\) ion absorbs at about \(570 \mathrm{~nm} .\) Predict the color of each ion. (b) The \(\left[\mathrm{Ni}(\mathrm{en})_{3}\right]^{2+}\) ion absorption maximum occurs at about \(545 \mathrm{~nm}\), and that of the [Ni(bipy) \(\left._{3}\right]^{2+}\) ion occurs at about \(520 \mathrm{~nm}\). From these data, indicate the relative strengths of the ligand fields created by the four ligands involved.

Give the number of \(d\) electrons associated with the central metal ion in each of the following complexes: (a) \(\mathrm{K}_{3}\left[\mathrm{TiCl}_{6}\right]\), (b) \(\mathrm{Na}_{3}\left[\mathrm{Co}\left(\mathrm{NO}_{2}\right)_{6}\right]\), (c) \(\left[\operatorname{Ru}(\mathrm{en})_{3}\right] \mathrm{Br}_{3}\), (d) \([\mathrm{Mo}(\mathrm{EDTA})] \mathrm{ClO}_{4},(\mathrm{e}) \mathrm{K}_{3}\left[\mathrm{ReCl}_{6}\right]\).

Give the number of \(d\) electrons associated with the central metal ion in each of the following complexes: (a) \(\mathrm{K}_{3}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]\), (b) \(\left[\mathrm{Mn}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]\left(\mathrm{NO}_{3}\right)_{2}\), (c) \(\mathrm{Na}\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]\) (d) \(\left[\mathrm{Cr}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{Br}_{2}\right] \mathrm{ClO}_{4}\), (e) \([\mathrm{Sr}(\mathrm{EDTA})]^{2-}\)

For each of the following metals, write the electronic configuration of the atom and its \(2+\) ion: (a) \(\mathrm{Mn},(\mathrm{b}) \mathrm{Ru}\), (c) Rh. Draw the crystal-field energy-level diagram for the \(d\) orbitals of an octahedral complex, and show the placement of the \(d\) electrons for each \(2+\) ion, assuming a strong-field complex. How many unpaired electrons are there in each case?

For each of the following metals, write the electronic configuration of the atom and its \(3+\) ion: (a) \(\mathrm{Ru}\), (b) Mo, (c) Co. Draw the crystal-field energy-level diagram for the \(d\) orbitals of an octahedral complex, and show the placement of the \(d\) electrons for each \(3+\) ion, assuming a weak-field complex. How many unpaired electrons are there in each case?

Draw the crystal-field energy-level diagrams and show the placement of \(d\) electrons for each of the following: (a) \(\left[\mathrm{Cr}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}\) (four unpaired electrons), (b) \(\left[\mathrm{Mn}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}\) (high spin), (c) \(\left[\mathrm{Ru}\left(\mathrm{NH}_{3}\right)_{5} \mathrm{H}_{2} \mathrm{O}\right]^{2+}\) (low spin), (d) \(\left[\mathrm{IrCl}_{6}\right]^{2-}\) (low spin), (e) \(\left[\mathrm{Cr}(\mathrm{en})_{3}\right]^{3+}\), (f) \(\left[\mathrm{NiF}_{6}\right]^{4-}\)

Draw the crystal-field energy-level diagrams and show the placement of electrons for the following complexes: (a) \(\left[\mathrm{VCl}_{6}\right]^{3-}\), (b) \(\left[\mathrm{FeF}_{6}\right]^{3-}\) (a high-spin complex), (c) \(\left[\mathrm{Ru}(\mathrm{bipy})_{3}\right]^{3+}\) (a low-spin complex), (d) \(\left[\mathrm{NiCl}_{4}\right]^{2-}\) (tetrahedral), (e) \(\left[\mathrm{PtBr}_{6}\right]^{2-}\), (f) [Ti(en) \(\left._{3}\right]^{2+}\).

The complex \(\left[\mathrm{Mn}\left(\mathrm{NH}_{3}\right)_{6}\right]^{2+}\) contains five unpaired electrons. Sketch the energy-level diagram for the \(d\) orbitals, and indicate the placement of electrons for this complex ion. Is the ion a high-spin or a low-spin complex?

The ion \(\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]^{3-}\) has one unpaired electron, whereas \(\left[\mathrm{Fe}(\mathrm{NCS})_{6}\right]^{3-}\) has five unpaired electrons. From these results, what can you conclude about whether each complex is high spin or low spin? What can you say about the placement of \(\mathrm{NCS}^{-}\) in the spectrochemical series?

(a) A compound with formula \(\mathrm{RuCl}_{3} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) is dissolved in water, forming a solution that is approximately the same color as the solid. Immediately after forming the solution, the addition of excess \(\mathrm{AgNO}_{3}(a q)\) forms \(2 \mathrm{~mol}\) of solid \(\mathrm{AgCl}\) per mole of complex. Write the formula for the compound, showing which ligands are likely to be present in the coordination sphere. (b) After a solution of \(\mathrm{RuCl}_{3} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) has stood for about a year, addition of \(\mathrm{AgNO}_{3}(a q)\) precipitates \(3 \mathrm{~mol}\) of \(\mathrm{AgCl}\) per mole of complex. What has happened in the ensuing time?

Sketch the structure of the complex in each of the following compounds: (a) \(c i s-\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]\left(\mathrm{NO}_{3}\right)_{2}\) (b) \(\mathrm{Na}_{2}\left[\mathrm{Ru}\left(\mathrm{H}_{2} \mathrm{O}\right) \mathrm{Cl}_{5}\right]\) (c) \(\operatorname{trans}-\mathrm{NH}_{4}\left[\mathrm{Co}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]\) (d) \(c i s-\left[\operatorname{Ru}(e n)_{2} C l_{2}\right]\)

The molecule dimethylphosphinoethane \(\left[\left(\mathrm{CH}_{3}\right)_{2} \mathrm{PCH}_{2}-\mathrm{CH}_{2} \mathrm{P}\left(\mathrm{CH}_{3}\right)_{2}\right.\), which is abbreviated \(\left.\mathrm{dmpe}\right]\) is used as a ligand for some complexes that serve as catalysts. A complex that contains this ligand is \(\mathrm{Mo}(\mathrm{CO})_{4}(\) dmpe \()\). (a) Draw the Lewis structure for dmpe, and compare it with ethylenediammine as a coordinating ligand. (b) What is the oxidation state of Mo in \(\mathrm{Na}_{2}\left[\mathrm{Mo}(\mathrm{CN})_{2}(\mathrm{CO})_{2}(\mathrm{dmpe})\right] ?\) (c) Sketch the structure of the \(\left[\mathrm{Mo}(\mathrm{CN})_{2}(\mathrm{CO})_{2}(\mathrm{dmpe})\right]^{2-}\) ion, including all the possible isomers.

Although the cis configuration is known for \(\left[\mathrm{Pt}(\mathrm{en}) \mathrm{Cl}_{2}\right]\) no trans form is known. (a) Explain why the trans compound is not possible. (b) Suggest what type of ligand would be required to form a trans-bidentate coordination to a metal atom.

Give brief statements about the relevance of the following complexes in living systems: (a) hemoglobin, (b) chlorophylls, (c) siderophores.

Write balanced chemical equations to represent the following observations. (In some instances the complex involved has been discussed previously in the text.) (a) Solid silver chloride dissolves in an excess of aqueous ammonia. (b) The green complex \(\left[\mathrm{Cr}(\mathrm{en})_{2} \mathrm{Cl}_{2}\right] \mathrm{Cl}\), on treatment with water over a long time, converts to a brown- orange complex. Reaction of \(\mathrm{AgNO}_{3}\) with a solution of the product precipitates 3 mol of \(\mathrm{AgCl}\) per mole of Cr present. (Write two chemical equations.) (c) When an \(\mathrm{NaOH}\) solution is added to a solution of \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}, \mathrm{a}\) precipitate forms. Addition of excess \(\mathrm{NaOH}\) solution causes the precipitate to dissolve. (Write two chemical equations.) (d) A pink solution of \(\mathrm{Co}\left(\mathrm{NO}_{3}\right)_{2}\) turns deep blue on addition of concentrated hydrochloric acid.

Some metal complexes have a coordination number of \(5 .\) One such complex is \(\mathrm{Fe}(\mathrm{CO})_{5}\), which adopts a trigonal bipyramidal geometry (see Figure 9.8). (a) Write the name for \(\mathrm{Fe}(\mathrm{CO})_{5}\), using the nomenclature rules for coordination compounds. (b) What is the oxidation state of Fe in this compound? (c) Suppose one of the CO ligands is replaced with a \(\mathrm{CN}^{-}\) ligand, forming \(\left[\mathrm{Fe}(\mathrm{CO})_{4}(\mathrm{CN})\right]^{-}\) How many geometric isomers would you predict this complex could have?

Which of the following objects is chiral? (a) a left shoe, (b) a slice of bread, (c) a wood screw, (d) a molecular model of \(\mathrm{Zn}(\mathrm{en}) \mathrm{Cl}_{2}\), (e) a typical golf club.

The complexes \(\left[\mathrm{V}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\) and \(\left[\mathrm{VF}_{6}\right]^{3-}\) are both known. (a) Draw the \(d\) -orbital energy-level diagram for \(\mathrm{V}(\mathrm{III})\) octahedral complexes. (b) What gives rise to the colors of these complexes? (c) Which of the two complexes would you expect to absorb light of higher energy? Explain.

Oxyhemoglobin, with an \(\mathrm{O}_{2}\) bound to iron, is a low-spin \(\mathrm{Fe}(\mathrm{II})\) complex; deoxyhemoglobin, without the \(\mathrm{O}_{2}\) molecule, is a high-spin complex. (a) Assuming that the coordination environment about the metal is octahedral, how many unpaired electrons are centered on the metal ion in each case? (b) What ligand is coordinated to the iron in place of \(\mathrm{O}_{2}\) in deoxyhemoglobin? (c) Explain in a general way why the two forms of hemoglobin have different colors (hemoglobin is red, whereas deoxyhemoglobin has a bluish cast). (d) A 15-minute exposure to air containing 400 ppm of CO causes about \(10 \%\) of the hemoglobin in the blood to be converted into the carbon monoxide complex, called carboxyhemoglobin. What does this suggest about the relative equilibrium constants for binding of carbon monoxide and \(\mathrm{O}_{2}\) to hemoglobin?

Consider the tetrahedral anions \(\mathrm{VO}_{4}^{3-}\) (orthovanadate ion), \(\mathrm{CrO}_{4}^{2-}\) (chromate ion), and \(\mathrm{MnO}_{4}^{-}\) (permanganate ion). (a) These anions are isoelectronic. What does this statement mean? (b) Would you expect these anions to exhibit \(d-d\) transitions? Explain. (c) As mentioned in "A Closer Look" on charge-transfer color, the violet color of \(\mathrm{MnO}_{4}^{-}\), is due to a ligand-to-metal charge transfer (LMCT) transition. What is meant by this term? (d) The LMCT transition in \(\mathrm{MnO}_{4}^{-}\) occurs at a wavelength of \(565 \mathrm{~nm}\). The \(\mathrm{Cr} \mathrm{O}_{4}^{2-}\) ion is yellow. Is the wavelength of the LMCT transition for chromate larger or smaller than that for \(\mathrm{MnO}_{4}^{-} ?\) Explain. (e) The \(\mathrm{VO}_{4}{\underline{\phantom{xx}}}^{3-}\) ion is colorless. Is this observation consistent with the wavelengths of the LMCT transitions in \(\mathrm{MnO}_{4}^{-}\) and \(\mathrm{CrO}_{4}^{2-}\) ?

In 2001, chemists at SUNY-Stony Brook succeeded in synthesizing the complex trans-[Fe(CN) \(\left._{4}(\mathrm{CO})_{2}\right]^{2-}\), which could be a model of complexes that may have played a role in the origin of life. (a) Sketch the structure of the complex. (b) The complex is isolated as a sodium salt. Write the complete name of this salt. (c) What is the oxidation state of Fe in this complex? How many d electrons are associated with the \(\mathrm{Fe}\) in this complex? (d) Would you expect this complex to be high spin or low spin? Explain.

When Alfred Werner was developing the field of coordination chemistry, it was argued by some that the optical activity he observed in the chiral complexes he had prepared was because of the presence of carbon atoms in the molecule. To disprove this argument, Werner synthesized a chiral complex of cobalt that had no carbon atoms in it, and he was able to resolve it into its enantiomers. Design a cobalt(III) complex that would be chiral if it could be synthesized and that contains no carbon atoms. (It may not be possible to synthesize the complex you design, but we won't worry about that for now.)

Generally speaking, for a given metal and ligand, the stability of a coordination compound is greater for the metal in the \(3+\) rather than in the \(2+\) oxidation state. Furthermore, for a given ligand the complexes of the bivalent metal ions of the first transition series tend to increase in stability in the order \(\mathrm{Mn}(\mathrm{II})<\mathrm{Fe}(\mathrm{II})<\mathrm{Co}(\mathrm{II})<\) \(\mathrm{Ni}(\mathrm{II})<\mathrm{Cu}(\mathrm{II})\). Explain how these two observations are consistent with one another and also consistent with a crystal-field picture of coordination compounds.

Suppose that a transition-metal ion was in a lattice in which it was in contact with just two nearby anions, located on opposite sides of the metal. Diagram the splitting of the metal \(d\) orbitals that would result from such a crystal field. Assuming a strong field, how many unpaired electrons would you expect for a metal ion with six \(d\) electrons? (Hint: Consider the linear axis to be the z-axis).

Metallic elements are essential components of many important enzymes operating within our bodies. Carbonic anhydrase, which contains \(\mathrm{Zn}^{2+}\), is responsible for rapidly interconverting dissolved \(\mathrm{CO}_{2}\) and bicarbonate ion, \(\mathrm{HCO}_{3}^{-} .\) The zinc in carbonic anhydrase is coordinated by three nitrogen-containing groups and a water molecule. The enzyme's action depends on the fact that the coordinated water molecule is more acidic than the bulk solvent molecules. Explain this fact in terms of Lewis acid-base theory (Section 16.11).

Two different compounds have the formulation \(\mathrm{CoBr}\left(\mathrm{SO}_{4}\right) \cdot 5 \mathrm{NH}_{3} .\) Compound \(\mathrm{A}\) is dark violet, and compound \(\mathrm{B}\) is red-violet. When compound \(\mathrm{A}\) is treated with \(\mathrm{AgNO}_{3}(a q)\), no reaction occurs, whereas compound \(\mathrm{B}\) reacts with \(\mathrm{AgNO}_{3}(a q)\) to form a white precipitate. When compound \(\mathrm{A}\) is treated with \(\mathrm{BaCl}_{2}(a q)\), a white precipitate is formed, whereas compound \(B\) has no reaction with \(\mathrm{BaCl}_{2}(a q)\). (a) Is Co in the same oxidation state in these complexes? (b) Explain the reactivity of compounds \(\mathrm{A}\) and \(\mathrm{B}\) with \(\mathrm{AgNO}_{3}(a q)\) and \(\mathrm{BaCl}_{2}(a q)\). (c) Are compounds \(A\) and \(B\) isomers of one another? If so, which category from Figure \(24.17\) best describes the isomerism observed for these complexes? (d) Would compounds \(A\) and \(B\) be expected to be strong electrolytes, weak electrolytes, or nonelectrolytes?