Summer Suggestion #1. Review key topics from Chem 101/102

The first 5 or 6 weeks of Chem 201 races through a lot of material you covered in Chem 101/102. In particular, we discuss,

  • how molecules are held together (101/102 – atomic structure, electrons, bonds, orbitals)
  • how to draw and interpret electron patterns (101/102 – Lewis structures, resonance structures, orbitals)
  • how to connect molecular structure to other molecular properties (101/102 – charge distribution, dipole moment, intermolecular forces, molecular energy)
  • chemical reactions, beginning with acid-base equilibria (101/102 – pH, pKa, acid/base strength, Keq, reaction thermodynamics (ΔG, ΔH, ΔS))
  • chemical reaction kinetics (101/102 – reaction rates, activation energies (Ea, ΔH‡), rate constants (k)).

All of these topics should look at least a little familiar, but how well will you recall and (even more important) be able to use them next fall? One simple way to make sure you get off to a good start in Chem 201 is to review these key topics from 101/102.

Here is a detailed list of topics that you can use to guide your 101/102 review. It includes a large number of pointers regarding key facts, skills, and explanations. To put it another way, it’s a long list so don’t let it overwhelm. Pick-and-choose a few topics for summer study, and we will get the rest of it in 201. (Note: my other two summer review suggestions also cover some of the material on this list.)

(added 8 Aug, 2019) An alternative to my detailed list exists and that is to dive into a smallish (250 page) book, “Preparation for Organic Chemistry” by I. David Reingold. The book is available in paperback and also in a $9.99 Kindle-friendly version. The author describes the book in this way, “Reviews material from general chemistry that is relevant to organic, set in an organic context”. I haven’t seen more of the book than the Table of Contents, but even that limited topic list overlaps well with the topic list that I give you below.

Chem 101/102 Topics that Apply to Organic Chemistry

Lewis Structures

  1. Drawing
    1. Know the Chem 201 Atoms of Interest (201 AoI): C/Si, N/P, O/S, F/Cl/Br/I, H/Li/Na/K, Mg, B/Al
    2. Know where the 201 AoI appear on the Periodic Table and know how many valence electrons are held by each type of atom
    3. Know what symbols to use for atoms, bonds, nonbonding electrons, formal charges
    4. Know the rules for constructing plausible Lewis structures (“plausible” means the electron pattern in the Lewis structure is a plausible guide to the electron distribution in the molecule). The three most important rules are:
      1. covalent bond = 2 shared electrons
      2. Lewis octet = ideal electron pattern for each atom
      3. formal charge computation
  2. Interpretation
    1. Identify bond orders (single, double, triple)
    2. Identify polar bonds (between atoms of different electronegativity)
      1. Know relative atom electronegativities
      2. Rank bonds from least to more polar
    3. Identify non-zero formal charges (are they plausible?)
    4. Identify atoms that might be associated with strong intermolecular forces
    5. Characterize molecule’s energy as “low” (chemically stable) or “high” (chemically unstable)

Molecular Geometry

  1. Bond distances
    1. Know relative atom sizes for 201 AoI
    2. Predict effect of atom size and bond order (from Lewis structure) on relative bond distances
  2. Bond angles
    1. Know VSEPR (= electron domain theory) for predicting bond angles around a central atom based on its Lewis structure
      1. Know standard angles
      2. Know how electron pair type (electron domain size) might perturb angles (= deviate from standard values)
    2. Know atom geometry labels: linear, bent, trigonal, tetrahedral, pyramidal, trigonal bipyramidal, square planar, octahedral
  3. 3-D Molecular models
    1. Build (plastic, computer) models that accurately reflect VSEPR predictions
    2. Interpret (plastic, computer) models
      1. Label atom geometry (linear, bent, etc.)
      2. Infer steric number (or number of electron domains) that supports the model’s depiction of atom geometry

Resonance Structures (= resonance contributors = resonance forms)

  1. Drawing
    1. Use double-headed arrows to “connect” resonance structures
    2. Use partial bonds & partial charges to show electron distribution in resonance hybrid
    3. Push electrons (draw curved arrows) to show how one resonance structure turns into another
  2. Interpretation
    1. Identify (and rank) “major” and “minor” resonance contributors
    2. Construct resonance hybrid that reflects rankings of resonance contributors
    3. Predict molecular geometry (adjustments to bond distances, bond angles, and atom geometries required by resonance)
    4. Predict effect of resonance on intermolecular forces (assess location and size of partial charges)
    5. Predict effect of resonance on molecular energy (assess degree of resonance stabilization; small? large?)

Quantum Mechanical Models of Electronic Structure

  1. Electrostatic Potential Maps
    1. Identify atoms (and regions near atoms) that are relatively electron-rich (carry a partial negative charge)
    2. Identify atoms (and regions near atoms) that are relatively electron-poor (carry a partial positive charge)
    3. Compare the polarity of two molecules
    4. Compare the degree of charge delocalization (requires resonance) in two molecules
    5. Use location and degree of charge build-up to predict possible intermolecular forces
    6. Use location and degree of charge build-up to predict molecular energy
  2. Hybrid Orbitals
    1. Assign atom’s hybridization based on Lewis (or resonance) structure
    2. Identify hybrid orbital (or atomic orbital) that contains nonbonding electrons (lone pairs)
    3. Identify hybrid orbital (or atomic orbital) that contains electrons that participate in a specific covalent bond (sigma & pi bonds are treated differently!)
    4. Use orbital type (atomic: 1s, 2s, 3s, … 2p, 3p, …) (hybrid: sp, sp2, sp3) to identify electrons that are held more tightly/loosely
  3. Molecular Orbitals (MO)
    1. Drawing – for any covalent bond in a molecule, construct an orbital mixing diagram
      1. Identify the 2 orbitals (one from each atom) that overlap and share electrons
      2. Draw orbital mixing diagram that shows energy levels of the 2 orbitals that combine, and the 2 MO (bonding, antibonding) that are created by this combining
      3. Label all 4 orbitals (atomic “cartoons” of the overlapping orbitals, and “cartoons” of the MO that can be made by combining these orbitals (one MO is bonding, the other is antibonding)
      4. Draw orbital labels next to each energy level identifying orbital type (see 4.B.iv) (for MO identify as 1) bonding/BMO or antibonding ABMO, and identify as “sigma” or “pi”)
      5. Draw electrons that
      6. Know how relative magnitude of overlap between two orbitals affects energy gap between bonding and antibonding MO, and bond strength (sigma overlap/gap/strength > pi overlap/gap/strength)
      7. Draw an orbital mixing diagram that shows 1) the relative energies of the 2 uncombined and the 2 molecular orbitals, and 2) the electrons that are assigned to each of these orbitals
    2. Predict the overall bond order between two atoms by considering the number of electrons assigned to all anti/bonding orbitals involving both atoms
    3. Predict the effect of MO occupancy (the number of electrons in each MO) on molecular energy

Intermolecular Forces

  1. Identify & characterize possible forces
    1. Electrostatic forces
      1. Can be attractive or repulsive
      2. Involves atoms carrying full or partial charges
      3. Hydrogen bonds (attractive, one atom is positively charged H)
    2. London dispersion forces (= van der Waals forces)
      1. Can be attractive or repulsive
      2. Involves “touching” or “overlapping” atoms
  2. Express effect of forces on molecular energy
    1. Attractive forces lower molecular energy (stabilize molecules)
    2. Repulsive forces raise molecular energy (destabilize molecules)

Structure-Energy Correlations

  1. Attraction = stabilization = low energy
  2. Repulsion = destabilization = high energy
  3. Stabilizing factors (partial list)
    1. Sharing electrons (strong covalent bonds)
    2. Charges supported by atom electronegativities (surplus electrons located on MORE electronegative atoms; surplus positive charges located on LESS electronegative atoms)
    3. Resonance
    4. Surplus electrons located on more electronegative atoms (related: Surplus positive charges located on less electronegative atoms)
    5. Attractive intermolecular forces (may involve interactions with solvent molecules or ions of opposite charge)
  4. Destabilizing factors
    1. Ineffective or insufficient electron sharing (weak covalent bonds, atoms without octets)
    2. Charges NOT supported by atom electronegativities (surplus electrons located on LESS electronegative atoms; surplus positive charges located on MORE electronegative atoms)
    3. Repulsive intermolecular forces

Chemical Equilibria & Concentration

  1. Mass-Action Law (equation connecting Keq to reactant/product concentrations)
  2. Thermodynamics (equation connecting Keq to ΔG) (equation connecting ΔG to ΔH and ΔS)
    1. Predicting how molecular structure affects molecular energy (see above) and Keq (reaction favorability)
  3. Le Chatelier’s Principle (effect of experimental conditions on equilibrium concentrations)
  4. Acid-Base equilibria
    1. Convert [H3O+] to pH and to [HO-]
    2. Ka and pKa
      1. relationship between Ka and pKa
      2. relationship between Ka and acid strength
      3. Henderson-Hasselbach equation (connects Ka and [H3O+])

Chemical Reaction Rates (Kinetics)

  1. Rate Laws (equations connecting reaction rate – d[Product]/dt – to reactant concentrations)
  2. Rate Constants (proportionality constant ‘k’ that appears in rate law) Note: chemists use both of the following equations, but only review what is familiar to you from 101/102
    1. Arrhenius equation ? (equation connecting k to energy barrier, Ea)
    2. Transition state equation ? (equation connecting k to ΔG‡)
  3. Reaction Mechanisms
    1. Reaction intermediates
      1. electron-pushing (drawing curved arrows to show how electron patterns & molecular geometries change during a chemical reaction)
    2. Transition states
  4. Energy changes & Reaction (or Potential) Energy Diagram
    1. Interpretation
      1. Identify energy minima & maxima
      2. Identify reaction barriers
      3. Identify reaction energies (overall, step-by-step)
      4. Label reactant, product, intermediate(s), transition state(s)
    2. Prediction
      1. Is reaction (overall, step-by-step) favorable or unfavorable?
      2. Given diagrams for two competing (and possibly hypothetical reactions),
        1. Which reaction is more (less) favorable?
        2. Which reaction is faster (slower)?
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