chemical reaction kinetics<\/em> (101\/102 – reaction rates, activation energies (Ea, \u0394H\u2021), rate constants (k)).<\/li><\/ul>\n\n\n\nAll 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.<\/p>\n\n\n\n
Here is a detailed<\/em> 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.)<\/p>\n\n\n\n(added 8 Aug, 2019)<\/em> 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”<\/em>. 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.<\/p>\n\n\n\nChem 101\/102 Topics that Apply to Organic Chemistry<\/strong><\/p>\n\n\n\nLewis Structures<\/strong><\/p>\n\n\n\n- Drawing
- Know the Chem 201 Atoms of Interest (201 AoI)<\/strong>: C\/Si, N\/P, O\/S, F\/Cl\/Br\/I, H\/Li\/Na\/K, Mg, B\/Al<\/li>
- Know where the 201 AoI<\/strong> appear on the Periodic Table and know how many valence electrons are held by each type of atom<\/li>
- Know what symbols to use for atoms, bonds, nonbonding electrons, formal charges<\/li>
- 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:
- covalent bond = 2 shared electrons<\/li>
- Lewis octet = ideal electron pattern for each atom<\/li>
- formal charge computation<\/li><\/ol><\/li><\/ol><\/li>
- Interpretation
- Identify bond orders (single, double, triple)<\/li>
- Identify polar bonds (between atoms of different electronegativity)
- Know relative atom electronegativities<\/li>
- Rank bonds from least to more polar<\/li><\/ol><\/li>
- Identify non-zero formal charges (are they plausible?)<\/li>
- Identify atoms that might be associated with strong intermolecular forces<\/li>
- Characterize molecule’s energy as “low” (chemically stable) or “high” (chemically unstable)<\/li><\/ol><\/li><\/ol>\n\n\n\n
Molecular Geometry<\/strong><\/p>\n\n\n\n- Bond distances
- Know relative atom sizes for 201 AoI<\/strong><\/li>
- Predict effect of atom size and bond order (from Lewis structure) on relative bond distances<\/li><\/ol><\/li>
- Bond angles
- Know VSEPR (= electron domain theory) for predicting bond angles around a central atom based on its Lewis structure
- Know standard angles<\/li>
- Know how electron pair type (electron domain size) might perturb angles (= deviate from standard values)<\/li><\/ol><\/li>
- Know atom geometry labels: linear, bent, trigonal, tetrahedral, pyramidal, trigonal bipyramidal, square planar, octahedral<\/li><\/ol><\/li>
- 3-D Molecular models
- Build (plastic, computer) models that accurately reflect VSEPR predictions<\/li>
- Interpret (plastic, computer) models
- Label atom geometry (linear, bent, etc.)<\/li>
- Infer steric number (or number of electron domains) that supports the model’s depiction of atom geometry<\/li><\/ol><\/li><\/ol><\/li><\/ol>\n\n\n\n
Resonance Structures (= resonance contributors = resonance forms)<\/strong><\/p>\n\n\n\n- Drawing
- Use double-headed arrows to “connect” resonance structures<\/em><\/li>
- Use partial bonds & partial charges to show electron distribution in resonance hybrid<\/em><\/li>
- Push electrons (draw curved arrows) to show how one resonance structure<\/em> turns into another<\/li><\/ol><\/li>
- Interpretation
- Identify (and rank) “major” and “minor” resonance contributors<\/em><\/li>
- Construct resonance hybrid that reflects rankings of resonance contributors<\/em><\/li>
- Predict molecular geometry (adjustments to bond distances, bond angles, and atom geometries required by resonance)<\/li>
- Predict effect of resonance on intermolecular forces (assess location and size of partial charges)<\/li>
- Predict effect of resonance on molecular energy (assess degree of resonance stabilization; small? large?)<\/li><\/ol><\/li><\/ol>\n\n\n\n
Quantum Mechanical Models of Electronic Structure<\/strong><\/p>\n\n\n\n- Electrostatic Potential Maps
- Identify atoms (and regions near atoms) that are relatively electron-rich (carry a partial negative charge)<\/li>
- Identify atoms (and regions near atoms) that are relatively electron-poor (carry a partial positive charge)<\/li>
- Compare the polarity of two molecules<\/li>
- Compare the degree of charge delocalization (requires resonance) in two molecules<\/li>
- Use location and degree of charge build-up to predict possible intermolecular forces<\/li>
- Use location and degree of charge build-up to predict molecular energy<\/li><\/ol><\/li>
- Hybrid Orbitals
- Assign atom’s hybridization based on Lewis (or resonance) structure<\/li>
- Identify hybrid orbital (or atomic orbital) that contains nonbonding electrons (lone pairs)<\/li>
- Identify hybrid orbital (or atomic orbital) that contains electrons that participate in a specific covalent bond (sigma & pi bonds are treated differently!)<\/li>
- Use orbital type (atomic: 1s, 2s, 3s, … 2p, 3p, …) (hybrid: sp, sp2, sp3) to identify electrons that are held more tightly\/loosely<\/li><\/ol><\/li>
- Molecular Orbitals (MO)
- Drawing – for any covalent bond in a molecule, construct an orbital mixing diagram
- Identify the 2 orbitals (one from each atom) that overlap and share electrons<\/li>
- 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<\/li>
- Label all 4 orbitals<\/strong> (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)<\/li>
- 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”)<\/li>
- Draw electrons that <\/li>
- 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)<\/li>
- 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<\/li><\/ol><\/li>
- Predict the overall bond order between two atoms by considering the number of electrons assigned to all anti\/bonding orbitals involving both atoms<\/li>
- Predict the effect of MO occupancy (the number of electrons in each MO) on molecular energy<\/li><\/ol><\/li><\/ol>\n\n\n\n
Intermolecular Forces<\/strong><\/p>\n\n\n\n- Identify & characterize possible forces
- Electrostatic forces
- Can be attractive or repulsive<\/li>
- Involves atoms carrying full or partial charges<\/li>
- Hydrogen bonds (attractive, one atom is positively charged H)<\/li><\/ol><\/li>
- London dispersion forces (= van der Waals forces)
- Can be attractive or repulsive<\/li>
- Involves “touching” or “overlapping” atoms<\/li><\/ol><\/li><\/ol><\/li>
- Express effect of forces on molecular energy
- Attractive forces lower molecular energy (stabilize molecules)<\/li>
- Repulsive forces raise molecular energy (destabilize molecules)<\/li><\/ol><\/li><\/ol>\n\n\n\n
Structure-Energy Correlations<\/strong><\/p>\n\n\n\n- Attraction = stabilization = low energy<\/li>
- Repulsion = destabilization = high energy<\/li>
- Stabilizing factors (partial list)
- Sharing electrons (strong covalent bonds)<\/li>
- Charges supported by atom electronegativities (surplus electrons located on MORE electronegative atoms; surplus positive charges located on LESS electronegative atoms)<\/li>
- Resonance<\/li>
- Surplus electrons located on more electronegative atoms (related: Surplus positive charges located on less electronegative atoms)<\/li>
- Attractive intermolecular forces (may involve interactions with solvent molecules or ions of opposite charge)<\/li><\/ol><\/li>
- Destabilizing factors
- Ineffective or insufficient electron sharing (weak covalent bonds, atoms without octets)<\/li>
- Charges NOT supported by atom electronegativities (surplus electrons located on LESS electronegative atoms; surplus positive charges located on MORE electronegative atoms)<\/li>
- Repulsive intermolecular forces<\/li><\/ol><\/li><\/ol>\n\n\n\n
Chemical Equilibria & Concentration<\/strong><\/p>\n\n\n\n- Mass-Action Law (equation connecting Keq to reactant\/product concentrations)<\/li>
- Thermodynamics (equation connecting Keq to \u0394G) (equation connecting \u0394G to \u0394H and \u0394S)
- Predicting how molecular structure affects molecular energy (see above) and Keq (reaction favorability)<\/li><\/ol><\/li>
- Le Chatelier’s Principle (effect of experimental conditions on equilibrium concentrations)<\/li>
- Acid-Base equilibria
- Convert [H3O+] to pH and to [HO-]<\/li>
- Ka and pKa
- relationship between Ka and pKa<\/li>
- relationship between Ka and acid strength<\/li>
- Henderson-Hasselbach equation (connects Ka and [H3O+])<\/li><\/ol><\/li><\/ol><\/li><\/ol>\n\n\n\n
Chemical Reaction Rates (Kinetics)<\/strong><\/p>\n\n\n\n- Rate Laws (equations connecting reaction rate – d[Product]\/dt – to reactant concentrations)<\/li>
- 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<\/em>
- Arrhenius equation ? (equation connecting k to energy barrier, Ea)<\/li>
- Transition state equation ? (equation connecting k to \u0394G\u2021)<\/li><\/ol><\/li>
- Reaction Mechanisms
- Reaction intermediates
- electron-pushing (drawing curved arrows to show how electron patterns & molecular geometries change during a chemical reaction)<\/li><\/ol><\/li>
- Transition states<\/li><\/ol><\/li>
- Energy changes & Reaction (or Potential) Energy Diagram
- Interpretation
- Identify energy minima & maxima<\/li>
- Identify reaction barriers<\/li>
- Identify reaction energies (overall, step-by-step)<\/li>
- Label reactant, product, intermediate(s), transition state(s)<\/li><\/ol><\/li>
- Prediction
- Is reaction (overall, step-by-step) favorable or unfavorable?<\/li>
- Given diagrams for two competing (and possibly hypothetical reactions),
- Which reaction is more (less) favorable?<\/li>
- Which reaction is faster (slower)? <\/li><\/ol><\/li><\/ol><\/li><\/ol><\/li><\/ol>\n","protected":false},"excerpt":{"rendered":"
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 […]<\/p>\n","protected":false},"author":55,"featured_media":0,"comment_status":"closed","ping_status":"closed","sticky":false,"template":"","format":"standard","meta":{"footnotes":""},"categories":[1],"tags":[],"class_list":["post-6361","post","type-post","status-publish","format-standard","hentry","category-uncategorized"],"_links":{"self":[{"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/posts\/6361","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/users\/55"}],"replies":[{"embeddable":true,"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/comments?post=6361"}],"version-history":[{"count":4,"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/posts\/6361\/revisions"}],"predecessor-version":[{"id":6373,"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/posts\/6361\/revisions\/6373"}],"wp:attachment":[{"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/media?parent=6361"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/categories?post=6361"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/blogs.reed.edu\/chem201202\/wp-json\/wp\/v2\/tags?post=6361"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}