Solutions to Exam 4 for CHM 1311
July 26, 1999


Open Periodic Tables and Closed Everything Else.
Answer any 10 of the 11 questions


  1. As complex as nitrogen-containing molecules can become, which of the following molecular shapes about a central nitrogen atom will never be observed? Why?

    1. bent
    2. octahedral
    3. trigonal planar
    4. trigonal pyramidal
    5. trigonal bipyramidal

  2. Give the most likely Lewis structure(s) and describe the shape of the thiosulfate ion, S2O32-, around each sulfur if its skeletal structure is:
                      2-
    O   S   O   S   O

  3. The thiocyanide anion is shown in three possibilities below, each with a different choice for the central atom. They are all the best structures that can be made for that central atom choice. Only one of these structures is correct. Which is the real thiocyanide ion and why?
          ..             -        ..             -        ..         ..  -
    (a) [ :S : N ::: C: ]   (b) [ :S : C ::: N: ]   (c) [ :C :: S :: N: ]
           ¨               ¨
         (-1) (+1) (-1)          (-1) (0)   (0)          (-2) (+2)  (-1)
    

  4. Vitamin C, ascorbic acid
  5. At right is the skeletal structure of vitamin C. The multiple bonds are not shown. Determine the hybridization of the non-hydrogen atoms in its ring and from that, show where the multiple bonds, if any, must be. There are no multiple bonds beyond any that show up in or on the ring.

    (Red is oxygen, gray is carbon, and white is hydrogen. The "ring" is that cycle of atoms at the right end of the molecule comprised of one oxygen and four carbons with other stuff dangling off. So we don't get confused, let's number the ring atoms starting with the oxygen as number 1 and proceeding clockwise around the ring to number 5, the carbon just below #1.)

  6. Which of the following molecules are polar?

    1. carbon oxysulfide, OCS

      • No choice but to be polar. No geometry could cancel out unequal bond polarities!
    2. boron tribromide, BBr3

      • Nonpolar since B, having no lone pairs, makes a trigonal planar shape, and the B-Br bond polarities all cancel vectorially.
    3. nitrogen tribromide, NBr3

      • Ah, but N has a lone pair, so the Br's are pushed down in a trigonal pyramid structure. The nonplanar component of the N-Br polarity leaves this molecule polar.
    4. chloroform, HCCl3

      • Here "H" takes the place of a lone pair forcing the Cl's down to near tetrahedral bonding. So since C-H polarity is radically less than C-Cl, chloroform is a polar molecule
    5. arsenic pentachloride, AsCl5

      • Like PCl5, it's Periodic Table brother, AsCl5 must be trigonal bipyramidal with 3 of the chlorines equatorial (with vector cancelled polarity) and 2 chlorines axial (whose As-Cl polarity also cancels). Like PCl5, it is nonpolar.

  7. It may seem to you that our obsession with second row diatomic molecules is ludicrous, considering the enormous variety of all other molecules. We do it because the molecular orbitals are particularly simple in these diatomic molecules; so they make good introductory examples. But we can broaden even this simple use of MO theory to a few more examples by taking advantage of the fact that molecules with the same number of electrons (isoelectronic) have similar electronic structures!

    Take CO, carbon monoxide, for example. With what second row homonuclear diatomic, E2, is CO isoelectronic? And what does that tell you about the relative magnetic properties of CO and CO+ ?

  8. Toward which side (reactants or products) does equilibrium lie in the following and why?

    NH3 + PH4+  arrow double NH4+ + PH3

  9. With which pairs below (if any) would you expect to see some sort of chemical reaction? Balance any appropriate reaction(s).

    1. Au(s) and Cu2+(aq)

    2. Au3+(aq) and Cu(s)

  10. (Warning: This is one of Dr. Parr's "creative" questions.)

    Molecular fluorine is made by the electrolysis of the molten salt, KHF2. The existence of that salt suggests a bizarre HF2- ion! Lewis would absolutely choke on such a species, but MO makes it simple to understand. That ion would hold together only if there was a bond between hydrogen's 1s orbital and the highest occupied molecular orbital of F2-. Since F2- has an odd number of electrons, that highest MO orbital (s2px* ) has only one electron, and it stands to reason that it can bond with the single electron in hydrogen atom.

    Remember the origin and hence the shape of that s2px* orbital to prove that HF2- must be linear. (The key is orbital overlap. The x axis runs between the fluorine atoms.)

  11. What is the shape of the ion and bond order of the N O bonds in the nitrite ion, NO2- ?

  12. It is a fact that the more (legitimate) Lewis resonance structures a chemist can imagine for a molecule, the more stable that molecule turns out to be. Clearly the molecule isn't reading the chemist's mind. So what accounts for that greater stability?


Last modified 26 July 1999