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CH101 Class Notes

by: Rebecca Sharp

CH101 Class Notes CH 101

Rebecca Sharp

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These notes cover what we went over in class.
General Chemistry
Paul Rupar
Class Notes
General Chemistry
25 ?




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This 2 page Class Notes was uploaded by Rebecca Sharp on Thursday October 13, 2016. The Class Notes belongs to CH 101 at University of Alabama - Tuscaloosa taught by Paul Rupar in Fall 2015. Since its upload, it has received 3 views. For similar materials see General Chemistry in Chemistry at University of Alabama - Tuscaloosa.

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Date Created: 10/13/16
Electron Geometry  Based on the number of electron groups o 2 electron groups = linear = ideal bond angles of 180 o 3 electron groups = trigonal planar = ideal bond angles of 120 o 4 electron groups = tetrahedral = ideal bond angles of 109.5 o 5 electron groups = trigonal bipyramidal = ideal bond angels of 90 & 120  Trigonal bipyramidal have 2 types of bonds, axial and equatorial  Equatorial are all on the same plane, they’re basically just trigonal planar (the 120 degree ones)  Axial are stuck on the top like a radio antenna (the 90 degree one) o 6 electron groups = octahedral = ideal bond angles of 90  The Lone Pair Effect o Lone pairs count as an electron group o They like to spread out around an atom. Bond pairs do not. Bond pairs stay pretty concentrated. o Lone pair/lone pair repulsion is much stronger than lone pair/bonding pair repulsion, which is all still stronger than bonding pair/bonding pair repulsion. o This means when a central atom has a lone pair, the bond angles shrink away from it. This decreases the ideal bond angles. Molecular Geometry  Based on the number of electron groups as well as the number of bonds  Example of a bent molecules; NO2-  Example of a trigonal planar molecules; BF3  Example of a tetrahedral molecule; CH4  Example of a trigonal bipyramidal molecule; PCl5  Example of an octahedral molecule; BrF5 Electron Geometry (EG) = Molecular Geometry (MG) if there are no lone pairs EG =/= MG if there are any lone pairs Non Polar v Polar Molecules  The difference between the electronegativity values for each element determines the polarity of the molecule.  Dipoles point towards the more electronegative element.  When summing dipoles to determine polarity strength and direction, sum the dipoles up two at a time. ↗+↙=0 ←+↓=↙ →+↑=↗ ↓+↑=0 o Note; she has said we will only need to be able to approximate the direction, not the exact value or angle.  A molecule can have all polar bonds and still be non-polar, if all the polar bonds cancel one another out.  Polar and non-polar molecules will not mix  Quick Hint; tetrahedral like CF4, where there’s only 1 type of bond, (in this case C-F) will almost always be non-polar because the bonds will cancel out.


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