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Chapter 2- Polar Covalent Bonds; Acids and Bases

Ashley  Piekarski,  Ph.D.    

Why do I care, Dr. P?

•  In  Chapter  1,  we  studied  valence  bond  theory  which  uses  hybrid  orbitals  to  account  for  the  observed  shapes  of  organic  molecules  

 •  In  Chapter  2,  we  will  study  how  the  electrons  are  distributed  in  covalent  bonds  and  how  that  distribuDon  affects  chemical  reacDvity  

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Electronegativity

•  What  is  electronegaDvity?  

Electronegativity

The electronegativity value !  indicates the attraction of an atom for shared

electrons !  increases from left to right going across a period

on the periodic table !  decreases going down a group on the periodic

table !  is high for the nonmetals, with fluorine as the

highest !  is low for the metals

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Electronegativity values

Low values

High values

Ionic bond

An ionic bond !  occurs between metal and nonmetal ions !  is a result of electron transfer !  has a large electronegativity difference (1.8 or more).

Examples: Atoms Electronegativity Type of Bond

Difference _____ ________ Cl–K 3.0 – 0.8 = 2.2 Ionic N–Na 3.0 – 0.9 = 2.1 Ionic S–Cs 2.5 – 0.7 = 1.8 Ionic

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Nonpolar covalent bond

A nonpolar covalent bond !  occurs between nonmetals !  has an equal or almost equal sharing of electrons !  has almost no electronegativity difference (0.0 to

0.4)

Examples: Atoms Electronegativity Type of Bond

Difference _______________ N–N 3.0 – 3.0 = 0.0 Nonpolar covalent Cl–Br 3.0 – 2.8 = 0.2 Nonpolar covalent H–Si 2.1 – 1.8 = 0.3 Nonpolar covalent

Polar covalent bond

A polar covalent bond !  occurs between nonmetal atoms !  has an unequal sharing of electrons !  has a moderate electronegativity difference

(0.5 to 1.7) Examples: Atoms Electronegativity Type of Bond

Difference __________ _ O–Cl 3.5 – 3.0 = 0.5 Polar covalent Cl–C 3.0 – 2.5 = 0.5 Polar covalent O–S 3.5 – 2.5 = 1.0 Polar covalent

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Electronegativity and bond types

Methanol

•  Draw  the  structure  for  methanol.  •  What  is  the  electronegaDvity  value  for  oxygen?  

•  What  is  the  electronegaDvity  value  for  carbon?  

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Methanol- electrostatic map

•  ElectrostaDc  potenDal  maps  show  calculated  charge  distribuDons  

•  Colors  indicate  electron-­‐rich  (red)  and  electron-­‐poor  (blue)  regions  

•  Arrow  indicate  direcDon  of  bond  polarity  

Bond polarity

•  This  is  the  basis  of  organic  chemistry.    Understanding  the  polarity  of  a  molecule  helps  to  know  the  reacDvity!  

Note:  electrostaDc  potenDal  maps  in  textbook  give  you  a  clue  to  the  electron-­‐rich  and  electron-­‐poor  atoms  in  molecules  

Red:  negaDve  Blue:  posiDve  

   

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Inductive effect

•  InducDve  effect:  the  shiXing  of  electrons  in  a  sigma  bond  in  response  to  the  electronegaDvity  of  nearby  atoms  •  Metals inductively donate electrons •  Non-metals inductively withdraw electrons

Learning check

•  Assign  δ+/δ-­‐  charges  to  show  the  direcDon  of  expected  polarity.  •  H3C-MgBr

•  H3C-SH

•  H2N-H

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Dipole Moments

•  Molecular  polarity  results  from  the  vector  summa,on  of  all  individual  bond  polariDes  and  lone-­‐pair  contribuDons.  •  The quantity measured is called a dipole

moment, µ

µ =Q × rQ = charge

r = distance

Dipole Moments

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Learning check

•  An  aromaDc  compound,  benzene,  has  a  dipole  moment  of  zero.    Why?  

Learning check

•  Chloromethane,  CH3Cl,  has  a  dipole  moment  of  1.87.    Make  a  three-­‐dimensional  drawing  of  methyl  chloride  and  show  the  direcDon  of  the  dipole  moment.  

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Formal charges

•  Formal  charges  are  only  electron  “bookkeeping”  and  do  NOT  imply  the  presence  of  actual  ionic  charges  •  gives clues to the chemical reactivity

Formal Charge = # of VE's in free atom( )- # of bonding electrons2

⎛⎝⎜

⎞⎠⎟ − # of nonbonding electrons( )

Learning check

Calculate  the  formal  charges  for  all  the  atoms  in  the  acetate  ion:  

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Resonance

•  What  is  resonance?  •  it is the way we describe electron

delocalization in a compound that has pi bonding

Resonance Hybrid

•  A  structure  with  resonance  forms  does  not  alternate  between  the  forms  

•  Instead,  it  is  a  hybrid  of  the  two  resonance  forms  

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Resonance rules!

•  Rule  1:  Individual  resonance  forms  are  imaginary,  not  real.  

Resonance rules!

•  Rule  2:  Resonance  forms  differ  only  in  the  placement  of  their  pi  or  nonbonding  electrons  

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Resonance rules!

•  Rule  3:  Different  resonance  forms  of  a  substance  don’t  have  to  be  equivalent.  

Resonance rules!

•  Rule  4:  Resonance  forms  obey  normal  rules  of  valency  

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Learning check

•  Draw  the  resonance  forms  for  the  2,4-­‐pentanedione  aXer  it  has  reacted  with  a  strong  base.  

Acids and Bases

According to the Brønsted–Lowry theory, !  acids donate a proton (H+) !  bases accept a proton (H+)

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Acids and Bases

In the reaction of ammonia and water, !  NH3 is the base that accepts H+

!  H2O is the acid that donates H+

Learning check

•  Draw  the  reacDon  of  aceDc  acid  reacDng  with  sodium  hydroxide.    What  is  the  acid  and  what  is  the  base?    What  are  the  products  formed?  

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Acid and Base Strength

•  What  is  the  generic  reacDon  scheme  for  a  weak  acid  reacDng  with  water?  

•  What  is  the  Ka  of  this  reacDon?  •  Based on this expression, if your Ka is very

large is it a strong or weak acid?

pKa

•  Do  you  remember  how  to  find  pKa?  •  Would  a  stronger  acid  have  a  smaller  or  larger  pKa?  •  For convenience, acid strengths are

expressed using pKa values •  In organic chemistry it is good to start learning

the pKa of acids and their conjugate basesà gives you a clue to their reactivity

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pKa Table

Predicting acid-base reactions

•  pKa  values  are  related  as  logarithms  to  equilibrium  constants  

•  Useful  for  predicDng  whether  a  given  acid-­‐base  reacDon  will  take  place  

•  The  stronger  base  holds  on  the  proton  more  Dghtly  

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Learning check

Organic acids

•  Organic  acids  •  characterized by the presence of a positively-

charged hydrogen atom

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Organic acids

•  Those  that  lose  a  proton  from  O-­‐H,  such  as  methanol  or  aceDc  acid  

•  Those  that  lose  a  proton  from  C-­‐H,  usually  from  a  carbon  atom  next  to  a  C=O  double  bond  

Conjugate bases- electrostatic maps

•  Which  element  now  has  a  substanDal  amount  of  negaDve  charge  aXer  deprotonaDon?  

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Organic bases

•  Have  an  atom  with  a  lone  pair  of  electrons  that  can  bond  to  H+  

•  Nitrogen-­‐containing  compounds  derived  from  ammonia  are  the  most  common  organic  bases  

•  Oxygen-­‐containing  compounds  can  react  as  bases  when  a  strong  acid  or  as  acids  with  strong  bases  

Lewis Definition

•  Lewis  acids  are  electron  pair  acceptors  and  Lewis  bases  are  electron  pair  donors  

•  The  Lewis  definiDon  leads  to  a  general  descripDon  of  many  reacDon  pajerns  

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Lewis acids

•  The  Lewis  definiDon  of  acidity  includes  metal  caDons,  such  as  Mg2+    

•  They accept a pair of electrons when they form a bond to a base

•  Group  3A  elements,  such  as  BF3  and  AlCl3,  are  Lewis  acids  because  they  have  unfilled  valence  orbitals  and  can  accept  electron  pairs  from  Lewis  bases  

•  TransiDon-­‐metal  compounds,  such  as  TiCl4,  FeCl3,  ZnCl2,  and  SnCl4,  are  Lewis  acids  

•  Organic  compounds  that  undergo  addiDon  reacDons  with  Lewis  bases  (discussed  later)  are  called  electrophiles  and  therefore  Lewis  Acids  

•  The  combinaDon  of  a  Lewis  acid  and  a  Lewis  base  can  shown  with  a  curved  arrow  from  base  to  acid  

Lewis acid-base reaction

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Note: curved arrows

•  A  curved  arrow  always  means  that  a  pair  of  electrons  move  from  the  atom  at  the  tail  of  the  arrow  to  the  atom  at  the  head  of  the  arrow  

Lewis bases

•  Lewis  bases  can  accept  protons  as  well  as  Lewis  acids,  therefore  the  definiDon  encompasses  that  for  Brønsted  bases  

•  Most  oxygen-­‐  and  nitrogen-­‐containing  organic  compounds  are  Lewis  bases  because  they  have  lone  pairs  of  electrons  

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Noncovalent interactions

•  Dipole-­‐dipole  forces  •  Dispersion  forces  •  Hydrogen  bonds  

Dipole-Dipole

•  Occur  between  polar  molecules  as  a  result  of  electrostaDc  interacDons  among  dipoles  

•  Forces  can  be  ajracDve  or  repulsive  depending  on  orientaDon  of  the  molecules  

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Dispersion Forces

•  Occur  between  all  neighboring  molecules  and  arise  because  the  electron  distribuDon  within  molecules  that  are  constantly  changing  

Hydrogen Bond Forces

•  Most  important  noncovalent  interacDon  in  biological  molecules  

•  Forces  are  a  result  of  ajracDve  interacDon  between  a  hydrogen  bonded  to  an  electronegaDve  O  or  N  atom  and  an  unshared  electron  pair  on  another  O  or  N  atom  

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Applications

Hydrogen  bonding  in  DNA