Lecture Outline More Resonanc

Name ______________________________________ 1 In Ch 3 of the text, there is a series of "rules" for writing resonance structures. Rules schmules. The molecules didn't go to college, so they didn't learn any rules. I bring this up because "rule 1" says "... may not..." and "...may..." Oh yeah? How are the molecules going to know what they "may" or "may not" do? Okay, perhaps you can read that as rules for you to follow, but then it just seems like a bunch of arbitrary commands. Take a shot at rewriting "Rule 1" so that it's clear why nuclei "may not" move but electrons "may". Think about what ideas we're trying to convey by writing resonance structures and why the molecules wouldn't care about any of our "rules" for resonance structures even if they had learned them. 2 What about the other rules? Once you're comfortable with the concept of resonance, Rule 2 seems fairly obvious. Rules 3 and 4 — that we can draw a bunch of Lewis structures, then do a weighted average to get an approximation to reality — is a postulate, but one that always seems to work (wow). Rule 5 is a little over-simplified, however, so here's an example that will help clarify it. Two isomeric anions are shown below. Which is more stable? To answer this, you will need to draw all the other "important" resonance structures by moving that one charge around. (be careful with the skeletal notation... can you see those Hs? Look carefully...) One of these has 2 more res structs, the other has one more. Delocalization of charge is good. So why can't we just count 'em up like rule 5 seems to imply? O O Lecture outline More resonance — Here are some common shifts of electrons that we'll use to draw resonance structures of conjugated π-systems. These are the same moves that we'll use to illustrate electron flow in reactions. 1. Find a 3-atom unit with a lone pair next to a π-bond — lone pair in, bond pair out H C H2C CH2 now do the same thing in skeletal notation... O NH O ... but why will this type of shift NOT work for these? (Hint: find the conjugated series of – p-orbitals. Another hint: the molecules never learned about conjugation, so try shifting e s anyway — what happens? i.e., if you do this accidentally, how will you recognize that it doesn't work?) O O 2. Find a 3-atom unit with an incomplete octet next to a π-bond, and do the "windshield wiper". H C H2C CH2 now do it in skeletal notation... ... but why will the windshield wiper shift NOT work for these? (hint, hint: see above) NH3 3. Find a 2-atom unit with an incomplete octet next to a lone pair and make a bond to complete the octet 4. Find a π-bond between two atoms of very different electonegativities and give the e–s to the atom that wants them more. shift 3 H2C NH2 H2C NH2 shift 4 H2C NH2 Notice that in the process of illustrating these two shifts, we ended up drawing the first structure again. We don't generally repeat structures like that. It wastes time, and time is money. Note that these shifts also involve a pair of p-orbitals. These aren't normally called "conjugated" — "partially π-bonded" perhaps — but the same ideas about geom and hybridization and orbital overlaps apply here. Which of the two structures drawn above is better? Why? What does the actual cation + look like? (hint: Use δ s to show the charge, and use dashes for a partial bond, but don't try to draw partial lone pairs or your brain will swell.) ... same shifts are possible for these... O H O We may as well throw in a couple more that will be useful later... 5. Aromatic rings — 6-membered rings with three "double bonds" are very common among organic molecules because, as we'll see soon, this structural feature gives them special stability. Draw the second resonance structure of benzene (below)... ...and draw the actual structure. Why is the term double bonds in quotes above? Do the same kind of thing for naphthalene (used in moth balls) — you should be able to draw two more resonance structures in addition to the one given. Now, benzene's CC bonds all have the same length, but naphthalene has some shorter bonds and some longer bonds. Why??? 6. Free radicals (or just “radicals”) — these have odd numbers of electrons, so one electron unpaired. This makes radicals extremely reactive. Free radicals participate in lots of important organic reactions, including processes that exchange Hs in hydrocarbons for much more useful halogens, additions of atoms and groups to double bonds, and polymerizations. However, they've also been implicated in spoilage of food, stiffening of cell membranes in aging, and damage to DNA. Here are two resonance structures for the allyl radical, shown in skeletal notation. Draw the full structures so you can see exactly what's going on. For radicals, it's best — at least for now — not to use curved arrows, but to imagine chopping one bonding e– pair in half, then recoupling the "loose dots" in the other direction. Now try it for these — the first has 2 more res structs; the second has one. Look carefully. Guidelines for determining relative importance of resonance structures 1. Octets. In general, an all-octet structure beats one with an incomplete octet. We'll occasionally see structures with incomplete octets on B or C, but never ever on N, O, or halogens. 2. No charge separation is better than having a + and a – (but rule #1 almost always takes priority over this one) and similarly, adjacent like charges (e.g., + +) and multiple charges (e.g., 2–) should be avoided 3. Other things being equal, charge placements that follow electronegativities are better. (but again, rule #1 is almost always more important) So which structures are worth drawing? There are often many electrons that can be shifted around in lots of different ways, so this could go on forever, right? Well, no, not if it's done correctly. It's always difficult to know where to draw the line. Your text doesn't do a very good job of distinguishing between cases where there are several major resonance contributors and cases where we invoke minor contributors to explain subtle effects of bonding, charge distribution, and stabilities, so let's see if we can break this down into discrete categories... Ions. In most cases there will be a single + or – charge. Focus on moving that one charge around — delocalization is always a good thing energetically, and we'll get a pretty good idea of which atoms have some charge and about how much. These are all relatively important contributors to the real structure. Don't create more formal charges— this will just make a lot of ugly (i.e. waaaaay less important) resonance structures that contribute virtually nothing. Examples: Neutrals. Most neutral molecules are pretty well described by a single Lewis structure. So if we're just trying to convey the basics of bonding, geometry, etc, we draw it and stop. But there are a few cases where we draw resonance structures of neutrals: (1) Molecules for which one "perfect" Lewis structure doesn't exist. The best we can do is draw a set of "best possible" structures (i.e. the resonance structures) and mix 'em all together. Examples: (2) Molecules with aromatic rings. We can always shuffle e–s around the ring(s) and draw a series of nominally equivalent resonance structures. So we do this until we get tired of it. This will probably happen around Ch 17 if not sooner. (3) Molecules with one "perfect" Lewis structure (90 or 95%) and one or more minor contributors (5 or 10%) that give us necessary insight into the structure and charge distribution, or imply a bit of extra stabilization. Two cases that come to mind (there may be others, but two is enough for now) — (a) lone pair conjugated with a π-system (we've seen that, we've used it, we like it). This creates a minor all-octet structure (Yea!) but with charge separation (Boooooooo...) that often opposes electronegativities (booo.) (b) One π-bond or a conjugated series of π-bonds that can be shifted to give a + on C and a – on something electronegative like O. Yes that's a non-octet C (Boooo!), but at least the charges follow the electronegativities (rah.), so it's not a complete disaster.