CHEM 2301 1st Edition Lecture 2Outline of Last Lecture I. Course introductionII. Drawing Lewis structuresIII. Lewis structures of common moleculesOutline of Current Lecture IV. Shorthand line drawingsV. Bond polarity and electronegativityVI. Basic principles of resonance theoryVII. Assessing the relative importance of resonance structuresCurrent Lecture- Shorthand line drawingso Line drawings are not the same things as Lewis structureso Vertex of lines indicates a carbon with all four bonds Assume all other bonds are to hydrogenso Hydrogen atoms bonded to hetero (non-carbon) atoms must be drawn Show lone pairs and formal charges to be more specifico Must show double bondso A line ending implies a methyl group (CH3)o Ring structure formulas can be written as “cyclo[formula]” or with a bracket over the molecular formula as shown below- Bond polarity and electronegativityo Direction and magnitude of polarity in bonds is critical for defining reactivity Some atoms in the molecule have partial charges because of differences in electronegativity- Molecules containing uneven electronegative atoms are polarThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute. More electronegative atoms will have a greater attraction for the electrons shared in the covalent bond, giving it a partial negative charge and the other atom in the bond a partial positive charge.  Partial positive/negative charges labeled as: δ+/δ-- Basic principles of resonance theoryo Resonance structures are not real.  An individual resonance structure does not accurately represent the structureof a molecule or ion. Only the hybrid does.o Resonance structures are not in equilibrium with each other.  There is no actual movement of electrons from one form to another (one moves electrons by arrows to show how two forms are related to each other).o Resonance structures are not isomers.  Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons. o Resonance structures are not in equilibrium with each othero Nuclei don’t differ in position between resonance structures (only electrons move).o Typically electrons are shifted in lone pairs and multiple bonds. o Typically, shift electron pairs. o Resonance stabilization important when it can delocalize charge (negative charge to electronegative atoms; positive charge to multiple C atoms, O with three bonds, N with 4 bonds).o The double arrow drawn between resonance structures show that the actual structure is a hybrid of two resonance structures. Examples of resonance structuresBecause the Cl atom has a higher electronegativity than the C atom, the electrons in the bond between them spend more time by the Cl nucleus, giving it a partial negative charge. Atoms like CH4 that are nonpolar have no partial charges.- In a resonance structure, all bond lengths are equal regardless of how many bonds there are- Assessing relative importance of resonance structureso Minimize charges (Typically less than two with nitro exception)o Electronegative atoms can bear positive charges IF they possess an octeto Avoid structures with two C atoms of opposite chargesA couple more examples of resonance