Dr. Streit Week 8 notes
Dr. Streit Week 8 notes CHEM 1030 - 003
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This 3 page Class Notes was uploaded by Rachel Ferrell on Tuesday March 29, 2016. The Class Notes belongs to CHEM 1030 - 003 at Auburn University taught by John D Gorden in Fall 2015. Since its upload, it has received 16 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.
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Date Created: 03/29/16
Rachel Ferrell CHEM 1030 3/21-‐3/24 Chapter 7: Molecular Geometry: • Valence Shell Electron Pair Repulsion Model (VSEPR) Model o AB x o A= central atom o X= can be any integer 2-‐6 • Basic idea of model→electrons repel each other, therefore they will arrange themselves to be as far apart as possible from each other • Creates Electron Domains, which can be… o Lone pairs o Single bonds o Double bonds o Triple bonds • Electron domain geometry= arrangement of electron domains around a central atom o Assuming that the electron domains are all bonds and not lone pairs… o 2 electron domains→ linear o 3 electron domains→ trigonal planar o 4 electron domains→ tetrahedral o 5 electron domains→ trigonal bipyramidal o 6 electron domains→ octahedral • Molecular geometry= arrangement of bonded atoms o Is not always the same as electron domain geometry o Electron domain geometry is used to find molecular geometry o When an atom has lone pairs as an electron domain→changes shape o o How to predict geometries: § 1) Draw lewis structure § 2) Count number of electron domains on the central atom § 3) Count number of lone pairs o Bond Angle= angle between 2 adjacent AB bonds (AB ) x § linear→ 180 § trigonal planar→120 § tetrahedral→109.5 § trigonal bipyramidal→90/120 § octahedral→90 Deviations from Ideal Bond Angles: • some electron domains are better than others at repelling neighboring domains o Lone pairs→ take up more space than bonded pairs of electrons § Makes bond angle slightly smaller o Double/triple bonds→ repel more strongly than single bonds § Makes bond angle slightly smaller Geometry of Molecules with more than 1 Central Atom: • Treat them as separate central atoms • The whole molecule does not have to be one shape; parts of it can be different Molecular Geometry and Polarity: • Molecular polarity= an important consequence of molecular geometry • The polarity of a molecule with 3 or more atoms depends on… o 1) The polarity of the individual bonds o 2) The molecules molecular geometry • ex. C2 →bonds are polar but the molecule is nonpolar (because dipole moments cancel each other out) H 2→ bonds= polar; molecule= polar (dipole moments do not cancel out because of bent shape) • Dipole moments→ can be used to distinguish between structural isomers o Ex. cis vs. trans molecules o Have same molecular formula, but different polarities/structure/shape Intermolecular Forces: • Intermolecular forces= attractive forces between neighboring molecules o Collectively known as van-‐der-‐waals interactions o Strong intermolecular forces→solid at room temp o Weak intermolecular forces→ liquid or gas at room temp • Dipole-‐Dipole Interactions o Attractive forces between polar molecules o Partial positive charge attracted to partial negative on other molecule o The magnitude of the attractive forces depends on the magnitude of the dipole (how electronegative atom is) • Hydrogen Bonding o Special type of dipole-‐dipole interaction o –H bonded to highly electronegative N,O, F o results in an especially strong dipole-‐dipole interaction • Dispersion Forces o Also called London dispersion forces o Result from Coulombic attraction between instantaneous dipoles of non-‐polar molecules o All molecules have dispersion forces • Ex. problem: Name the types of intermolecular forces: o A.) CCl →4 dispersion o B.) CH C3OH→ dispersion, dipole, hydrogen bonding o C.) CH C2CH → 3dispersion, dipole • Ion-‐Dipole Interactions o Coulombic attractions between ions (+ or -‐) and polar molecules o Bigger positive charge=stronger interaction (attracted to the partially negative side of the polar molecule) o More intermolecular forces in polar molecule→ stronger ion interaction Valence Bond Theory: • =atoms share electrons when atomic orbitals overlap • 1) a bond forms when single occupied atomic orbitals on 2 atoms overlap • 2) the 2 electrons shared in a region of orbital overlap must have opposite spin • 3) Formation of a bond results in a lower potential energy for the system • ex. 2H bond forms when 1s orbitals overlap HF bond forms when 1s orbital overlaps with 2p
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