Chem 312, Au13
Problem Set 9 (slightly revised)
Due in class, Monday, December 2
1. According to Table 5.4, the reduction potentials for the alkali metals falls in the order
Cs ≅ Li < Rb ≅ K < Na (this is the beginning of the activity series,
http://en.wikipedia.org/wiki/Reactivity_series). This is an unusual order, not following any of the
periodic trends we have discussed. The thermochemical cycle for the reduction potentials (Figure
6.11) from Wulfsberg, has three terms that are specific to the metal. Describe the periodic trends
for all three of these (in general, heats of atomization decrease on going down the periodic table,
although not for the transition elements). Given these three qualitative trends, can you see how Li
and Cs could be the most reactive?
If you sum the metal-specific terms in the thermochemical cycle in Figure 6.11, do you get the
correct order of reduction potentials? Use the in the Tables in Chapter 6 (also in the Appendixes).
Revision: Don’t worry that this gives the right order for only some of the elements; I’m not sure
why it’s off for one of them in particular.
2. (a) Draw the Lewis dot structure for perchlorate, ClO4–, the way you were taught to draw it in
freshman chemistry: make Cl–O double bonds and single bonds to minimize the formal charges.
How many electrons about the chlorine are involved in bonding? What chlorine orbitals must be
used to accommodate these electrons? Are these valence orbitals in the standard sense?
(b) Now draw a Lewis dot structure for perchlorate that obeys the octet rule at Cl. This may look a little
weird, but according to all recent theoretical studies, this is a much better description of perchlorate
than the one you drew in (a). Include formal charges in your Lewis dot structure.
3. (a) Draw a Lewis dot structure of SF6, which has an octahedral structure (it’s a very stable molecule).
How many electrons are arrayed around the sulfur? Here you’re forced to exceed the octet rule, right?
(b) Even here, current calculations find that the octet rule is obeyed. This is possible because SF6 is
best considered to be the sum of resonance forms. In each resonance form, there is at least one
sulfur-fluorine interaction that is purely ionic, with no covalent bond. We talked about this for
PF5, as drawn at right. Draw three of the many
F
F
F
octet-obeying resonance structures for SF6. (I
F
F
know, they look odd.) Based on these F P F
F
P
F
P
F
F
F
resonance structures, how many total S-F
F
F
F
bonds are there in SF6? Divide that number by
six to get the average S–F bond order for SF6 based on these resonance structures. [In PF5, there is
one bond to the two axial (up and down) F’s, so the average P–F bond order is ½.]
(c) Draw the Molecular Orbital diagram for SF6. This should look very similar to the MO diagram
for an octahedral transition metal complex (like FeF63–), except that there are no d orbitals (sulfur
doesn’t have valence d orbitals). You can use p orbitals on the fluorine atoms or sp3 hybrid
orbitals, it doesn’t matter.
i. At the left put the sulfur valence s and p orbitals (which is lower in energy?) and at the right
put six fluorine orbitals. Are the F orbitals above or below the S orbitals? Why?
ii. Make a bonding and antibonding combination between the sulfur s orbital and the 6 F orbitals.
Draw the six F orbitals that interact with the sulfur s orbital, being sure to indicate their relative
phase.
iii. Consider the pair of F orbitals that lie on the ±x axis. They are trans to each other (on
opposite sides of the S, at 180° to each other). They interact with the S px orbital. The two F
orbitals can be combined into two linear combinations that are half on one F and half on the
other. One of these has its lobes that point at the S in phase, the other has the lobes out-ofphase. Which one binds to the S px? The other one is non-bonding. Place these on your MO
diagram.
Chem 312, Au13; PS 9, p. 2
iv. Because SF6 is octahedral, the x, y, and z axes are the same. So once you’ve done the bonding
on x, that along y and z is the same. Add the orbitals to your MO diagram as appropriate.
v. How many bonding orbitals do you have in the MO diagram? What is the average S–F bond
order? Does this agree with your resonance structure picture from part b above?
vi. If you were to add the S 3d orbitals, where would they go in the diagram? Would they be
close in energy to the F orbitals or not? What is the rule for how the gap in energy affects the
strength of the bonding? How does that apply in this case?
If you need additional help with this, you can Google (or Bing or …) “SF6 Molecular Orbital
Diagram” and you’ll get lots of hits. Try to avoid the ones that use Group Theory (you’ll see a1g
and t1u and maybe odd-looking tables with numbers) but it wasn’t too hard to find some sites
with useful text and pictures. I found starting with the images was helpful. This question is hard
so don’t hesitate to get help this way.
(d) Do you think that SH6 or S(CH3)6 can be made?
4. Read the paper “Addition of Ammonia to AlH3 and BH3. Why Does Only Aluminum Form 2:1
Adducts?” M. Czerw, A. S. Goldman, K. Krogh-Jespersen Inorg. Chem. 2000, 39, 363-369.
(a) What experimental fact does this paper try to explain? Do the authors’ calculations agree with the
experiments?
(b) Do the authors believe that d-orbitals are involved in bonding at aluminum?
(c) The “Steric Factors” [size effects] discussed by the authors are pretty similar to the “radius ratio
rules” discussed in class? Do the authors believe that such size effects can explain the difference
between boron and aluminum?
(d) What is the authors’ preferred explanation? Do you think this is similar to what is discussed in
questions 3c(iii) above?
(e) Main group elements such as Al, P, S, and Cl don’t exceed the octet rule and don’t utilize d orbitals
in bonding. So why do you think that this is what’s taught in high school and freshman chemistry?
5. Read the paper “A Polar Copper−Boron One-Electron σ‑Bond” M.-E. Moret, L. Zhang, J. C. Peters
J. Am. Chem. Soc. 2013, 135, 3792.
(a) In one or two sentences, what is this paper about? Do you think it’s interesting, or a peculiar odd
molecule that is not relevant to anything?
(b) The unusual TBP ligand that binds to the copper (drawn three times in Figure 1) is uncharged
(neutral) when it is not bonded to copper. With this assumption, what are the electronic
configurations of the copper center in the cation, neutral, and anion? (Your answer should be like
[Ar]4sx… ). Does this violate one of the strong statements I made repeatedly in class a month ago
(“the … is always empty”)?
(c) An abbreviated MO diagram is shown at the right of Figure 3. It’s a little odd because it shows the
non-valence (core) copper 1 s orbital. That’s because it is a diagram to explain the X-ray spectroscopy results at the left. X-rays are photons just like visible light, but have much higher energies.
The scale of the spectrum is thousands of electron volts (eVs), 103 higher than a photon of green
light (wavelength 532 nm, energy 2.33 eV), so the X-rays can promote a core electron. This picture
shows a copper 4p orbital [because of the symmetry selection rules, the 1s à 4p transition is
allowed but the 1s à 4s transition is forbidden (g à g, if you recall)]. For the purposes of this
question, assume that the copper 4s orbital is not too much lower than 4p, and is still above the B
orbital. In the neutral TBP ligand, and in (TBP)Cu+, the boron orbital is empty. As electrons are
added, do you think they will go in an orbital that is mostly Cu or mostly B in character, based on
this picture? [Hint: The dotted arrows in the figure show where the electrons go.] Is this consistent
with the results of the DFT calculations for the neutral complex described at the top of the second
column of the second page?
Chem 312, Au13; PS 9, p. 3
(d) On the basis of your answer in part (c), let’s revisit your answer to part (b). Are the electronic
configurations you wrote really correct? Is this really a violation of my “always” statement?
6. Watch the interview with Harry Gray towards the bottom of
http://pubs.acs.org/page/inocaj/multimedia/voices.html
Professor Gray is one of the most influential and articulate inorganic chemists, and a winner of the
Presidential Medal of Science (that he got from President Ronald Reagan in 1986). He spoke at the
meeting I went to in Cincinnati at the end of October. The interviewer, Prof. Richard Eisenberg, was a
graduate student working in Gray’s laboratory. Give me your impression of this interview. I’m a little
old fashioned, but I think Harry tells some pretty good stories, and gives a good sense of the human
side of chemistry. I particularly like his answer to the question of what were the greatest moments of
his career, about half way through the interview, and the comment soon after that about why his
students have been so successful.