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Organic Chemistry; Chapter 1 Part B and Chapter 2 Notes

by: Amanda Biddlecome

Organic Chemistry; Chapter 1 Part B and Chapter 2 Notes Chemistry 2230

Marketplace > Clemson University > Chemistry 2230 > Organic Chemistry Chapter 1 Part B and Chapter 2 Notes
Amanda Biddlecome
GPA 4.0

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About this Document

These notes cover what we did in the second week of class and include the end of Chapter 1: Remembering General Chemistry, and all of Chapter 2: Acids and Bases.
Organic Chemistry 1
Dr. Schroeder
Class Notes
Organic Chemistry
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This 4 page Class Notes was uploaded by Amanda Biddlecome on Wednesday January 6, 2016. The Class Notes belongs to Chemistry 2230 at Clemson University taught by Dr. Schroeder in Fall 2016. Since its upload, it has received 51 views.


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Date Created: 01/06/16
Organic  Chemistry  2230   Chapter  1  Part  B   January  11,  2016   Amanda  Biddlecome     1)  Geometries     -­‐Linear:  draw  Lewis  Structure;  then  look  at  bond  angles-­‐180  degrees;  2     regions  with  electrons     -­‐Trigonal  Planar:  3  regions  with  electrons;  bond  angles-­‐120  degrees;  draw     Lewis  Structures  with  formal  charges       *Formal  Charge=Group  number-­‐(lone  electron  pairs  +  ½  bonding         electrons)     -­‐Tetrahedral:  most  common;  4  regions  of  electron  density;  Lewis  Structure     first;  bond  angle-­‐109.5  degrees       *when  you  have  lone  pair  electrons,  bond  angles  change  slightly  and         you  have  a  different  molecular  geometry       *trigonal  pyramidal:  4  regions  of  electrons;  1  lone  pair;  bond  angle-­‐       107  degrees       *bent:  4  regions  of  electrons;  2  lone  pairs;  bond  angle-­‐104.5  degrees     2)  Orbitals=where  the  electrons  are  located     -­‐S  Orbital       *spherical;  groups  1  and  2     -­‐P  Orbital       *dumbbell  shaped;  groups  13-­‐18       *3  possible  orientations:  2px,  2py,  or  2pz     -­‐Bonding       *start  as  atomic  orbitals  and  when  two  of  those  combine,  you  get  a         molecular  orbital       *2  types:  bonds  with  s  orbitals=sigma  bonds;  bonds  with  p  orbitals=pi       bonds     3)  Hybridization     -­‐Carbon  forms  4  single  bonds  by  causing  an  s  orbital  electron  to  become     excited  and  jump  to  the  p  orbital;  now  there  is  1  electron  in  the  s  orbital  and     3  electrons  in  the  p  orbital  all  unpaired;  they  go  through  hybridization  and     create  a  new  orbital:  sp3;  tetrahedral;  all  bonds=sigma  (4  sigma  bonds)     -­‐Carbon  forms  a  double  bond  by  creating  a  new,  excited  orbital:  sp2;  3     regions  of  electron  density;  trigonal  planar;  bond  angle=120  degrees;  double     bond=pi  bond  and  others=sigma  bonds  (1  pi  bond  and  1  sigma  bond)     -­‐Carbon  forms  a  triple  bond  by  forming  a  new  hybridized  orbital:  sp;  triple     bond  has  1  sigma  bond  and  2  pi  bonds;  2  regions  of  electron  density;  linear;     180  degree  bond  angles       4)  Types  of  Bonds     -­‐sigma  is  an  overlap  of  s  orbitals     -­‐pi  is  an  overlap  of  p  orbitals               Organic  Chemistry  2230   Chapter  2:  Acids  and  Bases   January  13,2016   Amanda  Biddlecome     1)  Acids  and  Bases     -­‐common  acids:  HClO4,  H2SO4,  HBr,  HCl,  HNO3,  H3PO4     -­‐common  bases:  NaOH,  KOH,  Ba(OH)2     -­‐Bronsted-­‐Lowry  Acids  and  Bases       *Acid=proton  donor:  H^+=H3O^+=All  7  strong  acids  because  they         dissociate  completely;  always  positively  charged       *Base=proton  acceptor;  always  negatively  charged         *in  reactions,  negative  always  attacks  positive  (general  rule)       -­‐an  acid  +  a  base    =  conjugate  acid  +  conjugate  base       *acid-­‐base  reaction     -­‐Keq=concentration  of  products/concentration  of  reactants     -­‐Ka=acid-­‐dissociation  constant       *every  acid  has  one       *if  it’s  large,  then  the  acid  is  strong       *Ka  values  are  very  small  values  and  hard  to  quantify,  so  we  use  the         pKa  to  describe  acid  strength       -­‐pKa=  -­‐logKa       *the  smaller  the  pKa,  the  stronger  the  acid       *used  mostly  in  organic  chemistry     2)  Acid  Strengths     -­‐the  smaller  the  pKa,  the  stronger  the  acid     -­‐the  larger  the  pKa,  the  stronger  the  base     -­‐we  don’t  find  the  pKb  in  organic  chemistry,  so  we  have  to  determine  base     strength  by  looking  at  the  pKa     -­‐anytime  you  have  an  equilibrium  arrow,  it  can  shift  to  either  the  reactant     side  or  the  product  side       *look  at  pKa  values  to  determine  which  sides  have  weak  and  strong         acids  and  bases;  the  formation  of  the  weaker  acid  and  base  is  always         favored  in  equilibrium,  so  the  equation  will  always  shift  to  the  side         with  the  weaker  acid  and  base     -­‐strong  acids  form  weaker  conjugate  bases     3)  Organic  Acids       -­‐R=Carbon  or  Hydrogen  or  anything  really       *place-­‐holder  in  Lewis  structures     -­‐Carboxylic  Acid:  pKa=5     -­‐Alcohol:  pKa=15     -­‐Mines  (Nitrogen  based):  pKa=10     -­‐Strong  Acids:  pKa<10     4)  Acidity  of  Hydrogen  Halides     -­‐Group  7     -­‐atomic  size  increases  going  down  the  group     -­‐electronegativity  decreases  going  down  the  group     -­‐acid  strength  increases  going  down  the  group       *deals  with  bond  length;  the  longer  the  bond,  the  easier  it  is  to  break     -­‐base  strength  decreases  going  down  the  group     -­‐the  weaker  the  base,  the  more  stable  it  is       *this  is  why  equilibrium  favors  weaker  bases     -­‐acid  strength  increase  with  electronegativity  but  competes  with  atomic  size     5)  Acidity  of  Carbon  Compounds     -­‐shorter  bonds  are  stronger     -­‐the  more  s  character,  the  shorter  and  stronger  the  bond  and  the  larger  the     bond  angle     -­‐single  bond:  pKa>60;  double  bond:  pKa=44;  triple  bond:  pka=25     -­‐electronegativity  increases  from  single  bond  to  triple  bond  because  there’s     more  electrons  involved  in  more  bonds     -­‐acid  strength  increases  with  electronegativity         *atomic  size  doesn’t  help  you  much     6)  Atomic  Radius  versus  Electronegativity     -­‐when  atoms  are  similar  in  size,  the  stronger  acid  has  its  proton  attached  to     the  more  electronegative  atom     -­‐atomic  size  wins  here     -­‐atomic  radius  and  electronegativity  compete  for  control       *increasing  electronegativity  causes  increasing  acidity       *we  use  electronegativity  to  judge  acidity  because  the  differences  in         electronegativity  between  atoms  is  much  more  significant  than  most         atomic  radii     7)  Resonance  and  Acidity       -­‐resonance  structures  distribute  charge  across  atoms  to  increase  stability  of     compounds       *a  stable  conjugate  base  means  you  had  a  strong  acid  which  is           displayed  in  the  pKa  value     8)  Substituent  Affects  and  Acidity     -­‐substituents  are  things  bonded  to  carbon  that  influence  the  molecules  ability     to  lose  electrons       *when  you  put  halides  in,  the  greatest  effect  is  noticed     -­‐Inductive  Electron  Withdrawal-­‐when  you  put  something  very     electronegative  on  a  carbon,  it  pulls  on  all  of  the  electrons  in  the  molecule,     stabilizes  it,  and  increases  its  acidity     -­‐Inductive  Effect  diminishes  as  you  move  farther  away  from  the  central  atom,     or  the  acidic  site     9)  pH  Effects  on  Organic  Compounds     -­‐if  you  put  a  molecule  into  a  solution  with  a  pH  equal  to  its  pKa  value,  it  will     be  equally  acidic  and  basic  (neutral)       *buffers     -­‐Henderson-­‐Hasselbalch  Equation     pKa=pH+log([HA]/[A])       -­‐when  pH=pKa,  the  compound  is  neutral     -­‐when  pH<pKa,  the  compound  is  acidic     -­‐when  pH>pKa,  the  compound  is  basic     -­‐the  magnitude  of  the  deviation  from  neutral  doesn’t  matter     -­‐you  can  have  more  than  one  acidic  sites  on  a  single  molecule  that  can  have     different  pKa  values       *if  placed  in  a  solution  with  a  pH  equal  to  one  of  the  pKas,  the  equal         pKa  would  remain  the  same  (neutral)  and  the  other  would  either         become  acidic  or  basic  and  would  behave  appropriately       *a  pH  value  between  the  two  pKa  values  would  result  in  both  groups         being  utilized       -­‐pH  of  human  mouth=6.8;  stomach=1.5-­‐3.5;  intestine=8.5       *these  values  determine  which  foods  get  broken  down  where  in  the         digestive  tract     10)  Lewis  Acids  and  Bases     -­‐acid=non-­‐proton-­‐donating;  will  accept  two  electrons     -­‐base=will  donate  two  electrons      


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