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Dr. Streit Week 8 notes

by: Rachel Ferrell

Dr. Streit Week 8 notes CHEM 1030 - 003

Marketplace > Auburn University > Chemistry > CHEM 1030 - 003 > Dr Streit Week 8 notes
Rachel Ferrell
GPA 4.0

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covers beginning of chapter 7
Fundamentals Chemistry I
John D Gorden
Class Notes
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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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