CH101 Class Notes
CH101 Class Notes CH 101
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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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