1) (12 points) The formulas for three chemical species are given below. For each

Develop a structure.

Indicate the stereochemistry at each atom that is bonded to 2 or more other atoms

Indicate the hybridization adopted by each such atom

State whether the substance should function as a Lewis acid or a Lewis base, and why

acetonitrile, CH₃CN

Structure has 3 C-H single bonds, a C-C single bond, and a CN triple bond. There is a lone pair on the N atom.

Stereochemistry at methyl carbon is tetrahedral, sp3. Stereochemistry at nitrile carbon is linear, sp.

Can function as a Lewis base at the N atom (lone pair to donate) and as a Lewis acid at the nitrile C atom (underused p orbitals).

dioxonitronium cation, NO₂⁺

Structure is identical to that of CO2, with the + charge on the central N atom. There are 2 lone pairs on each O.

Stereochemistry at central N is linear, sp hybrids.

Can function as a Lewis acid at the N atom (underused p orbitals). It is conceivable that it could be a Lewis base at the O atoms.

boron hydroxide, H₃BO₃

Structure has each O single bonded to an H and to B. It is tempting to draw a double bond between B and one O, but this leads to unreasonable formal charges. The structure is best thought of as involving 3 single bonds to B.

Stereochemistry at B is trigonal planar, sp2. Stereochemistry at oxygen is bent, sp3.

B has an unused valence p orbital. The molecules is acidic at the B atom.

2) (8 points) A chemist claims that s/he has discovered a previously unknown allotrope of sulfur, in which sulfur atoms are arranged in 2-dimensional sheets. Suppose that you are asked to comment on the validity of the claim. What are your comments?

At first glance the structure looks odd because each S atom makes 3 bonds, contrary to the stable allotropes in which each S atom makes 2 bonds. 2 bonds is what we would expect based on the valence electron configuration (2 unpaired electrons). The structure drawn is skeletal, in that it does not exlicitly show all sulfur electrons. Only 3 of 6 electrons are shown for each S. So begin by putting a lone pair and a single electron on each atom. This shows that a structure with all single bonds is not reasonable because it leaves electrons unpaired. However, if we use the odd electrons to form double bonds, we get a structure that has some similarity to the structure of benzene, with delocalized pi bonds. The lone pairs on the S atoms prevent the structure from being identical to that of benzene of course. Looked at this way, the structure seems not unreasonable!