Chapter 23

Which transition metal atom is present in each of the following biologically important molecules: (a) hemoglobin, (b) chlorophyls, (c) siderophores, (d) hemocyanine.

Carbon monoxide, CO, is an important ligand in coordination chemistry. When \(\mathrm{CO}\) is reacted with nickel metal, the product is \(\left[\mathrm{Ni}(\mathrm{CO})_{4}\right],\) which is a toxic, pale yellow liquid. (a) What is the oxidation number for nickel in thiscompound? (b) Given that \(\left[\mathrm{Ni}(\mathrm{CO})_{4}\right]\) is a diamagnetic molecule with a tetrahedral geometry, what is the electron configuration of nickel in this compound? (c) Write the name for \(\left[\mathrm{Ni}(\mathrm{CO})_{4}\right]\) using the nomenclature rules for coordination compounds.

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?

One of the more famous species in coordination chemistry is the Creutz-Taube complex: It is named for the two scientists who discovered it and initially studied its properties. The central ligand is pyrazine, a planar six-membered ring with nitrogens at opposite sides. (a) How can you account for the fact that the complex, which has only neutral ligands, has an odd overall charge? (b) The metal is in a low-spin configuration in both cases. Assuming octahedral coordination, draw the \(d\)-orbital energy-level diagram for each metal. (c) In many experiments the two metal ions appear to be in exactly equivalent states. Can you think of a reason that this might appear to be so, recognizing that electrons move very rapidly compared to nuclei?

Solutions of \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{6}\right]^{2+},\left[\mathrm{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}(\) both octahedral \()\) and \(\left[\mathrm{CoCl}_{4}\right]^{2-}\) (tetrahedral) are colored. One is pink, one is blue, and one is yellow. Based on the spectrochemical series and remembering that the energy splitting in tetrahedral complexes is normally much less that that in octahedral ones, assign a color to each complex.

Oxyhemoglobin, with an \(\mathrm{O}_{2}\) bound to iron, is a low-spin Fe(Il) complex; deoxyhemoglobin, without the 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) \(\mathrm{A} 15\) -minute exposure to air containing 400 \(\mathrm{ppm}\) 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? (e) \(\mathrm{CO}\) is a strong-field ligand. What color might you expect carboxyhemoglobin to be?

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 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 \(\mathrm{CrO}_{4}^{2-}\) ion is yellow. Is the wavelength of the LMCT transition for chromate larger or smaller than that for MnO \(_{4}^{-}?\) Explain. (e) The VO \(_{4}^{3-}\) ion is colorless. Do you expect the light absorbed by the LMCT to fall in the UV or the IR region of the electromagnetic spectrum? Explain your reasoning.

Given the colors observed for \(\mathrm{VO}_{4}^{3-}\) (orthovanadate ion), \(\mathrm{CrO}_{4}^{2-}\) (chromate ion), and \(\mathrm{MnO}_{4}^{-}\) (permanganate ion (see Exercise \(23.84 ),\) what can you say about how the energy separation between the ligand orbitals and the empty \(d\) orbitals changes as a function of the oxidation state of the transition metal at the center of the tetrahedral anion?

In 2001 , chemists at SUNY-Stony Brook succeeded in synthesizing the complex trans-\(\left[\mathrm{Fe}(\mathrm{CN})_{4}(\mathrm{CO})_{2}\right]^{2-}\), 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 Fein this complex? How many d electrons are associated with the 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 due to 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 will not 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 (for metals that form stable \(+3\) ions in the first place). Suggest an explanation, keeping in mind the Lewis acid-base nature of the metal-ligand bond.

Many trace metal ions exist in the blood complexed with amino acids or small peptides. The anion of the amino acid glycine (gly), can act as a bidentate ligand, coordinating to the metal through nitrogen and oxygen atoms. How many isomers are possible for (a) \(\left[\mathrm{Zn}(\mathrm{gly})_{2}\right]\) (tetrahedral), \((\mathbf{b})[\mathrm{Pt}(\mathrm{g}] \mathrm{y})_{2} ]\) (square planar), \((\mathbf{c})\left[\operatorname{Cog}(\mathrm{gly})_{3}\right](\) octahedral)? Sketch all possible isomers. Use the symbol to represent the ligand.

The coordination complex \(\left[\mathrm{Cr}(\mathrm{CO})_{6}\right]\) forms colorless, diamagnetic crystals that melt at \(90^{\circ} \mathrm{C}\) . (a) What is the oxidation number of chromium in this compound? (b) Given that \(\left[\mathrm{Cr}(\mathrm{CO})_{6}\right]\) is diamagnetic, what is the electron configuration of chromium in this compound? (c) Given that \(\left[\mathrm{Cr}(\mathrm{CO})_{6}\right]\) is colorless, would you expect CO to be a weak-field or strong-field ligand? (d) Write the name for \(\left[\mathrm{Cr}(\mathrm{CO})_{6}\right]\) using the nomenclature rules for coordination compounds.

A manganese complex formed from a solution containing potassium bromide and oxalate ion is purified and analyzed. It contains \(10.0 \% \mathrm{Mn}, 28.6 \%\) potassium, \(8.8\%\) carbon, and 29.2\(\%\) bromine by mass. The remainder of the compound is oxygen. An aqueous solution of the complex has about the same electrical conductivity as an equimolar solution of \(\mathrm{K}_{4}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right] .\) Write the formula of the compound, using brackets to denote the manganese and its coordination sphere.

The \(E^{\circ}\) values for two low-spin iron complexes in acidic solution are as follows: \(\left[\mathrm{Fe}(o-\mathrm{phen})_{3}\right]^{3+}(a q)+\mathrm{e}^{-} \Longrightarrow\) \(\quad\quad\quad\quad\quad\quad\quad\) \(\left[\mathrm{Fe}(o-\mathrm{phen})_{3}\right]^{2+}(a q) \quad E^{\circ}=1.12 \mathrm{V}\) \(\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]^{3-}(a q)+\mathrm{e}^{-} \rightleftharpoons\) \(\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\) \(\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]^{4-}(a q) \quad E^{\circ}=0.36 \mathrm{V}\) (a) Is it thermodynamically favorable to reduce both Fe(III) complexes to their Fe(II) analogs? Explain. (b) Which complex, \(\left[\mathrm{Fe}(o-\mathrm{phen})_{3}\right]^{3+}\) or \(\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]^{3-},\) is more difficult to reduce? (c) Suggest an explanation for your answer to (b).

A palladium complex formed from a solution containing bromide ion and pyridine, \(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{N}\) (a good electron-pair donor), is found on elemental analysis to contain 37.6\(\%\) bromine, 28.3\(\%\) carbon, 6.60\(\%\) nitrogen, and 2.37\(\%\) hydrogen by mass. The compound is slightly soluble in several organic solvents; its solutions in water or alcohol do not conduct electricity. It is found experimentally to have a zero dipole moment. Write the chemical formula, and indicate its probable structure.

(a) In early studies it was observed that when the complex \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{Br}_{2}\right] \mathrm{Br}\) was placed in water, the electrical conductivity of a 0.05\(M\) solution changed from an initial value of 191 \(\mathrm{ohm}^{-1}\) to a final value of 374 \(\mathrm{ohm}^{-1}\) over a period of an hour or so. Suggest an explanation for the observed results.(See Exercise 23.69 for relevant comparison data.) (b) Write a balanced chemical equation to describe the reaction. (c) \(A 500\)-mL solution is made up by dissolving 3.87g of the complex. As soon as the solution is formed, and before any change in conductivity has occurred, a 25.00-mL portion of the solution is titrated with 0.0100 \(\mathrm{M} \mathrm{AgNO}_{3}\) solution. What volume of AgNO \(_{3}\) solution do you expect to be required to precipitate the free \(\operatorname{Br}^{-}(a q) ?(\mathbf{d})\) Based on the response you gave to part (b), what volume of \(\mathrm{AgNO}_{3}\) solution would be required to titrate a fresh 25.00 -mL sample of \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{Br}_{2}\right] \mathrm{Br}\) after all conductivity changes have occurred?

The total concentration of \(\mathrm{Ca}^{2+}\) and \(\mathrm{Mg}^{2+}\) in a sample of hard water was determined by titrating a 0.100-L sample of the water with a solution of EDTA \(^{4-} .\) The EDTA \(^{4-}\) chelatesthe two cations: $$\begin{array}{c}{\mathrm{Mg}^{2+}+[\mathrm{EDTA}]^{4-} \longrightarrow[\mathrm{Mg}(\mathrm{EDTA})]^{2-}} \\\ {\mathrm{Ca}^{2+}+[\mathrm{EDTA}]^{4-} \longrightarrow[\mathrm{Ca}(\mathrm{EDTA})]^{2-}}\end{array}$$ It requires 31.5 \(\mathrm{mL}\) of 0.0104 \(\mathrm{M}[\mathrm{EDTA}]^{4-}\) solution to reach the end point in the titration. A second 0.100-L sample was then treated with sulfate ion to precipitate \(\mathrm{Ca}^{2+}\) as calcium sulfate. The \(\mathrm{Mg}^{2+}\) was then titrated with 18.7 \(\mathrm{mL}\) of 0.0104 \(M[\mathrm{EDTA}]^{4-} .\) Calculate the concentrations of \(\mathrm{Mg}^{2+}\) and \(\mathrm{Ca}^{2+}\) in the hard water in \(\mathrm{mg} / \mathrm{L} .\)

Carbon monoxide is toxic because it binds more strongly to the iron in hemoglobin (Hb) than does \(\mathrm{O}_{2}\) , as indicated by these approximate standard free-energy changes in blood: $$\begin{array}{cl}{\mathrm{Hb}+\mathrm{O}_{2} \longrightarrow \mathrm{HbO}_{2}} & {\Delta G^{\circ}=-70 \mathrm{kJ}} \\\ {\mathrm{Hb}+\mathrm{CO} \longrightarrow \mathrm{HbCO}} & {\Delta G^{\circ}=-80 \mathrm{kJ}}\end{array}$$ Using these data, estimate the equilibrium constant at 298 K for the equilibrium $$\mathrm{HbO}_{2}+\mathrm{CO} \rightleftharpoons \mathrm{HbCO}+\mathrm{O}_{2}$$

A Cu electrode is immersed in a solution that is 1.00\(M\) in \(\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\) and 1.00 \(\mathrm{M}\) in \(\mathrm{NH}_{3} .\) When the cathode is a standard hydrogen electrode, the emf of the cell is found to be \(+0.08 \mathrm{V} .\) What is the formation constant for \(\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+} ?\)