Chapt10 11 fall 07

Chapt10 11 fall 07 - Chemistry 1035 - General Chemistry...

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Chemistry 1035 - General Chemistry Chapter 10 & 11 Fall 2007 Professor John G. Dillard 406A Davidson Hall 231-6926 john.dillard@vt.edu Review Sessions: Tuesday, 6:30-8:00 pm Wednesday, 8:30-10:00 pm Davidson Hall - Room 3
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Chapter 10 & 11(only 11.1-11.2) Lewis structures Hybridization sp 3 – tetrahedral sp 2 – triangular planar sp – linear dsp 3 – trigonal bipyramid d 2 sp 3 – octahedral Valence Shell Electron-Pair Repulsion (VSEPR) Molecular shapes Resonance
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Lewis structures – localized electron bonding model Bonding in compounds: atoms are “attached” or “connected” to one another by sharing pairs of electrons via atomic orbitals of the atoms. The particular electrons of interest are the valence electrons. The approach for preparing a Lewis structure is 1. Write the valence electron configuration (dot structure) for the atoms to be bonded – it is convenient to write the configuration for the central atoms and then “add” other atoms to the central atom. 2. Connect the electrons that “bond” the atoms so that each bond shares at least two electrons 3. Arrange the electrons (bonded and non-bonded) electrons about the central atom in a manner so that 8 electrons are localized about each atom.
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Block e - configuration/notation can illustrate bonding H 2 1s H + 1s H x x H H 1s H + x 3s 3p Cl x HCl 2s 2p x NH 3 3 N + 3 H x x x NH 3 HCl The N 2s electrons are not involved in bonding and are called non-bonding e -
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Writing Lewis electron dot structures is illustrated In the Lewis approach one places "dots" corresponding to the valence electrons around the atom . Following this, the pendent atoms are arranged to "share" electrons - covalent bonds. Consider the bonding in H 2 or HCl or NH 3 H 2 : H • + • H H •• H or H—H where "••" (2 dots) equals "—" a bar or a bond Lewis structures – localized electron bonding model Cl Cl H or H—Cl H + HCl: NH 3 : H N + 3 H N H H N H H H or
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The orbital overlap representation can also be used to illustrate bonding H 2 : Two 1s spherical orbitals overlap to make a sigma ( σ ) bond H H + HCl: The H 1s spherical orbital overlaps (comes in contact; shares the same space) with one of the 3p orbitals of Cl (select 3p x on Cl) + H Cl P x H H HCl: x
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The orbital overlap representation for NH 3 Each 1s orbital on H overlaps with a 2p orbital on nitrogen N P x P z P y H H H N H H H This H 1s – N 2p overlap results in the formation of 3 sigma ( σ ) bonds
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The sp 3 hybrid orbital representation for NH 3 For ammonia the arrangement of the orbitals is tetrahedral, sp 3 , and the molecular structure , the arrangement of the atoms, forms a triangular pyramid.
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The structure of molecules – arrangement of electron pairs
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Types of atomic orbital overlap – sigma and pi bonds For the N 2 molecule: the Lewis “dot” structure for N 2 would be: N N Consider the bonds "between" the nitrogens. The p orbitals are mutually perpendicular , so overlap of a p x on one N with a p x of the 2 nd N occurs "head on" and produces a σ , sigma bond.
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Chapt10 11 fall 07 - Chemistry 1035 - General Chemistry...

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