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Electron  Configuration  &  Chemical   Periodicity
Wednesday,   October  21,  2 015

8:42  A M

Each  orbital  can  be   described  by  three quantum  numbers
...

• m(sub(s)): describes  spin,  either   positive  or  negative   1/2
○ Up  arrow: "spin  up"  electron
○ Down  arrow: "spin  down"  electron
• Pauli  Exclusion  Principle: two  e lectrons in  the   same atom  cannot have  the  
same  four  quantum  numbers
...
 
• Electron  configuration: describes  how  electrons  are   distributed  in  atomic  
orbitals
○ H  -­‐ 1s  ( n  =   1,  l   =  0,  ml  =   0,  m(sub(s))  =  -­‐1/2
○ He   -­‐ 1s2  ( n  =   1,  l   =  0,  ml  =   0,  m(sub(s))  =  1/2
• Sublevel  energies   and  order  for  filling   electrons:
○ 3d  is   higher  than  4s
○ Electrons  enter   orbitals  of  lower  energy   first
○ Orbitals  are   filled   before  moving  to  a   higher  energy  orbital   (Aufbau  
Principle)
• Ground  state: lowest  energy  config
...
 than  ground  state
• Hund's  Rule: when  orbitals   of  equal  energy   are   available,  the   ground  state  has  
the  max   number  of  unpaired  electrons  with  parallel   spins
• Paramagnetic: atom  with  one  or   more  unpaired  electrons,  attracted   to  
magnetic  field
Exceptions  to  Aufbau  Principle  (know  just  these  5  e xceptions)
Cr:  [Ar]  4s2  3d4  is  wrong,  really   [Ar]  4s1  3d5  ( also  true  for  Mo);  half-­‐filled  d  orbitals  
are   lower  energy  than  d4
Cu:  [Ar]  4s2  3d9  is  wrong,  really   [Ar]  4s1  3d10  ( also  Ag  and  Au);  completely   filled   d  
orbitals  are   lower  energy  than  d9
Some  definitions  
...
e
...
e
...

• Fe:  [Ar]  4s2  3d6  
○ Inner  (core)  e lectrons -­‐ in  common  w/  previous  noble  gas  and  any  filled   d  
and  f  levels;  do  not  participate   in  bonding,  i
...
 18  electrons  from  [Ar]
○ Valence  e lectrons -­‐ highest  n  level   plus  any  unfilled  d  and  f  levels,  i
...
  8  
electrons   from  4s2  3d6
○ Outer  e lectrons -­‐ highest  n  value  electrons   that  are   lost  when  making  a  
cation,  i
...
 4s2  electrons
• Ga:  [Ar]  4s2  3d10  4p1
○ Inner  -­‐ [Ar]  and  3d10
○ Valence  -­‐ 4s2  and  4p1
○ Outer  electrons   -­‐ 4s2  and  4p1
Periodic  Trends
1
...
Ga   <  Al  due  to  the   increase   in  
Z(sub(eff))  from  transition  metals
2
...

a
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As  n increases,  energy increases,  distance from  nucleus  increases
b
...
5) 2s2  (13) 2p6  (15)
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Increasing energy and  increasing distance from  the  nucleus
Closer  electrons   shield attraction  b/w  nucleus  and  farther  electrons
c
...
As  you  move  up  the  periodic   table,   electrons  are   held  more   tightly,  and  as  
you  more   from  left  to  right,  electrons   are   held  more  tightly  ( first  
ionization  energy)

c
...
As  you  move  up  the  periodic   table,   electrons  are   held  more   tightly,  and  as  
you  more   from  left  to  right,  electrons   are   held  more  tightly  ( first  
ionization  energy)
i
...
 The  nucleus  
didn't  lose  a   proton  so  remaining  e-­‐ are   held   even  tighter
ii
...
Exception:   it  is   actually  harder  to  take   away  an  e-­‐ from  an  O  than  an  
N  
...
Electron   affinity  ( energy  change  when  one   mole  of  e-­‐ is  added  to  one  
mole  of  gaseous  ions
i
...
More  tightly  =   more  negative  EA
iii
...

◊ Upper  right:  do  not  like   to  give   up  e-­‐ (high  ionization  energy),  like   to  accept   e-­‐
(large  negative   electron   affinity),  like  to  make  anions,  called   reactive   nonmetals
◊ Lower  left:  do  not  hold  onto  e-­‐ tightly  ( low  ionization  energy),  like   to  give   up  e-­‐
(small  negative  electron   affinity),  like   to  make  cations,  called   reactive   metals
◊ Noble  gases:  do  not  like   to  give   up  e-­‐ (high  ionization  energy),  do  not  like   to  
accept   e-­‐ (small  negative  electron   affinity),  don't  make  ions,  inert
What  charge  will   the  ion  have?
◊ Aiming  for  the   closest  noble  gas  config
...

A) Fe  3+
B) Fe  2+
C) Fe
D) K
E) Ca



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