CHEM 227 1st Edition Lecture 2 Outline of Last Lecture I. 1.1 Atomic StructureII. 1.2 Atomic Structure: orbitalsIII. 1.3 Atomic Structure: Electron configurationsIV. 1.4 Development of Chemical Bonding TheoryV. 1.5 The Nature of Chemical Bonds: Valence bond theoryVI. 1.6 sp3 Hybrid Orbitals and the structure of MethaneVII. 1.7 sp3 hybrid orbitals and the structure of ethaneVIII.1.8 sp2 Hybrid orbitals and structure of EthyleneIX. 1.9 sp Hybrid Orbitals and the structure of AcetyleneX. 1.10 hybridization of Nitrogen, Oxygen, Phosphorus, and SulfurXI. 1.11 Describing Chemical Bonds: Molecular orbital theoryOutline of Current Lecture I. 1.12 Drawing Chemical StructuresII. 2.1 Bond Polarity and ElectronegativityIII. 2.2 Polar Covalent bonds: dipole momentsIV. 2.3 Formal ChargesV. 2.4 ResonanceVI. 2.5 Rules for Resonance FormsVII. 2.6 Drawing Resonance FormsCurrent Lecture1.12 Drawing Chemical Structures- Condensed strucutres: carbon-carbon and carbon-hydrogen bonds aren’t shown, instead they are just understoodThese 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.- Skeletal Structures: even simpler than condensed structures… following these ruleso 1. Carbon atoms aren’t usually shown, instead, a carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line. Occasionally, a carbon might be indicated for emphasis or clarityo 2. Hydrogen atoms bonded to carbon aren’t shown. Because carbon always has avalence of 4, we mentally supply the correct number of hydrogen atoms for each carbono 3. Atoms other than carbon and hydrogen are shown2.1 Bond Polarity and Electronegativity- Electronegativity (EN): ability of an atom to attract electrons in a covalent bond  createsbond polarities- Nonpolar covalent bonds: atoms with similar EN- Polar covalent bonds: difference in EN of atoms <2- Ionic bonds: difference in EN of atoms >2- Inductive Effect: shifting of electrons in a bond in response to EN of neary atoms- Partial positive charge : δ+ - Partial negative charge : δ –- Whole molecules: vector summation of individual bond polarities and lone-pair contributions2.2 Polar Covalent Bonds: Dipole Moments- Dipole Moment: Net molecular polarity, due to difference in summed charges- µ - magnitude of charge Q at end of molecular dipole times distance r between charges - µ = Q × r, in debyes (D)- Large dipole moments- EN of O and N>H- Both O and N have lone-pair electrons oriented away from all nuclei- In symmetrical molecules the effects of the local dipoles may cancel each other2.3 Formal Charges- The charge of specific atoms within a molecule- A formalism and don’t imply the presence of actual ionic charges in a molecule. Instead, they are a device for electron “bookkeeping”2.4 Resonance- Some molecules can be represented by more than one Lewis structure; in these cases we draw all the contributors which differ in the position of the π bonds or lone pairs.- The electrons are delocalized; each structure is a resonance form - The resonance forms are connected by a double-headed arrow- Thinking about resonance forms:o Realize that a substance like the acetate ion is the same as any other. Acetate doesn’t jump back and forth between two resonance forms, spending part of thetime looking like one and part of the time looking like the other. Rather, acetate has a single unchanging structure that we say is a resonance hybrid of the two individual forms and has characteristics of botho Both structures are contributors to the resonance hybrid This is NOT an equilibrium2.5 Rules for Resonance forms1. Individual resonance forms are imaginary, not real. The real structure is a composite, or resonance hybrid of the different forms.2. Resonance forms differ only in the placement of their pi or nonbonding electrons. Neither the position nor the hybridization of any atom changes from one resonance form to another. 3. Different resonance forms of a substance don’t have to be equivalent. Even though two resonance forms may not be equivalent, both contribute to the overall resonance hybrid. When 2 resonance forms are nonequivalent, the actual structure of the resonance hybrid resembles the more stable form more than it resembles the less stable form. 4. Resonance forms obey normal rules of valency. A resonance form is like any other structure: the octed rule still applies to second-row, main-group atoms. Make sure these more basic rules aren’t broken.5. The resonance hybrid is more stable than any individual resonance form. In other words, resonance leads to stability. Generally speaking, the larger the number of resonance forms, the more stable a substance is because its electrons are spread out over a larger part of the molecule and are closer to more nuclei.2.6 Drawing Resonance forms- Curved Arrows: the electrons from the atom or bond at the tail of the arrow go to the atom or bond at the head of the