Take Home Lessons from Lecture 1
MOLECULAR AND ELECTRONIC STRUCTURE
GR1. Atoms prefer filled valence shells. This rule explains why atoms make bonds, and the type of bonds created. A corollary is that centers of electron density (bonds and lone pairs of electrons) repel each other so they stay as far apart as possible. This latter rule, the basis for the so-called VSEPR model, explains 3-dimensional molecular structure.
2. The VSEPR model assumes areas of electron density repel each other, so areas of electron density are distributed around an atom to be as far apart as possible. Remember to count both lone pairs and bonds as areas of electron density (note that a double or triple bond counts as a single area of electron density). 4 areas of electron density = tetrahedron, 3 areas of electron density = trigonal planar, 2 areas of electron density = linear.
3. Molecular orbital theory and valence bond theory (hybridization) are two different ways to think about bonding. For this class, we will concentrate on the valence bond approach for description of sigma bonding, because this provides a more intuitive description (electrons centered between nuclei). We will use molecular orbital theory to describe pi bonds and therefore bonding in resonance hybrids.
4. Possible hybridization states of first row elements: sp hybridization (one s and one 2p hybridize, two 2p left over), sp2 hybridization (one s and two 2p hybridize, one 2p left over), sp3 hybridization (one s and all three 2p hybridize)
6. According to the Huckel definition of aromaticity, monocyclic molecules will be aromatic if they are 1) flat, 2) all the ring atoms are sp2 hybridized (sp in rare cases) and 3) there are 4n + 2 pi electrons, where n = 0,1,2,3,4,5,6......
KEY examples: benzene, pyrolle vs. pyridine, cyclopentadienyl anion
RESONANCE CONTRIBUTING STRUCTURES
7. Resonance contributing structures are used when more than one structure are required to describe accurately how the electrons and charges are distributed in a molecule.
8. Resonance contributing structures do not represent equilibrating structures, rather the hybrid (blending) of them is the true molecular representation. (Think of blue and red primary colors used to create a purple color).
9. When drawing resonance forms, you should move pi bonds (one bond of a double or triple bond) and lone pair electrons, not charges, atom nuclei, or sigma bonds
10. For pi bonding and therefore pi delocalization to occur over more then two atoms, parallel 2p orbitals are needed on all of the atoms involved, explaining why these atoms must be sp2 (or sp hybridized) and why these sytems are planar.
KEY examples: enolates and amides
KEY exceptions to VSEPR: Exocyclic aromatic amines and amides
FUNCTIONAL GROUPS
11. Functional groups are characteristic groups of atoms that have unique reactivities. You MUST be able to look at complex molecules and be able to recognize functional groups such as acid chlorides, alcohols, aldehydes, alkenes, alkynes, amines, anhydrides, epoxides, esters, ethers, haloalkanes (alkyl halides), hydrazines, hydrazones, imines, ketones, lactones, lactams, nitriles, nitro groups.
KEY examples: Acetals, cyclic acetals, amides
NOMENCLATURE
12. You should learn nomenclature rules on your own. To get you started, here is a handout on alkane nomenclature. There are many problems in each chapter of the text to help you get familiar with naming the other functional groups.
ACIDS AND BASES
13. When predicting acid strength, ANALYZE THE CHARGED SPECIES (i.e. deprotonated conjugate base that is an anion). THE MORE STABLE THE CONJUGATE BASE ANION, THE STRONGER THE PARENT ACID. You should become familiar with relative acidity values. Use five rules when determining relative stabilities of anions (i.e. when you are predicting relative acid strength) These five rules are actually application of the following two principles: 1) negative charge is neutralized by nuclear positive charge and 2) delocalizing negative charge over a larger area is stabilizing.
a) Across a row of the Periodic Table, negative charge on a more electronegative atom is more stable. (Principle 1)
b) Down a column of the Periodic Table, larger anions are more stable than smaller ones. (This is more confusing than you think, so make sure you understand it). (Principle 2 dominates here)
c) Charges distributed over more atoms are better. (Nature hates isolated charges, so spreading a charge around by resonance is very stabilizing) (Pinciple 2)
d) Inductive Effect: Electronegative atoms such as F on atoms adjacent to the atom(s) with the negative charge will pull some of the charge away, thus spreading it out and leading to stabilization. (Principle 1 and 2)
e) Hybridization: The more S orbital character of the hybridization on the atom that has the negative charge, the more stable the anion (stability of anions is in the order sp>sp2>sp3) Principle 1)
KEY examples: Carboxylic acids vs. alcohols, trifluoroacetic acid, phenol vs. cyclohexanol, enolates and beta-dicarbonyls
SUMMARY OF SKILLS YOU MUST MASTER FOR THE MCAT:
1) Correctly assign hybridization statesto all atoms in a molecule
2) Describe all sigma bonds in terms of the overlap of hybrid orbitals and all pi bonds in terms of the overlap of unhybridized 2p orbitals
3) Recognize aromaticity
4) Draw appropriate resonance contributing structures to describe molecules with "pi-ways"
5) Recognize the functional groups in molecules
6) Name molecules according to IUPAC rules
7) Predict relative acidity and the position of equilibrium in acid-base reactions