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Periodic  Proper.es  of  the  Elements  

How  Orbitals  Fill  The  Basis  of  the  Periodic  Table  

Trends  in  Atomic  Proper.es  of  Elements  

Mendeleev  is  given  credit  for  the  Periodic  Table  

The  following  slides  we  shall  skip  through  fast  because...  

•  ...they  were  stolen  from  the  internet  and  •  I  cannot  find  their  source.  •  But  the  informa.on  is  useful!  

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The History of the Modern

Periodic Table  

During the nineteenth century, chemists began to categorize the elements according to similarities

in their physical and chemical properties. The end result of these studies was our modern

periodic table.  

Johann Dobereiner  

1780 - 1849  Model of triads  

In 1829, he classified some elements into groups of three, which he called triads. The elements in a triad had similar chemical properties and orderly physical properties.  

(ex. Cl, Br, I and Ca, Sr, Ba)  

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John Newlands  

1838 - 1898  Law of Octaves  

In 1863, he suggested that elements be arranged in “octaves” because he noticed (after arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th element.  

John Newlands  

1838 - 1898   Law of Octaves  

Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society.

John Newlands  

1838 - 1898   Law of Octaves  

His law of octaves failed beyond the element calcium.   WHY?  

Would his law of octaves work today with the first 20 elements?  

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Dmitri Mendeleev  

1834 - 1907  

In 1869 he published a table of the elements organized by increasing atomic mass.  

Lothar Meyer  

1830 - 1895  

At the same time, he published his own table of the elements organized by increasing atomic mass.  

•  Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass.

•  Both left vacant spaces where unknown elements should fit.

So why is Mendeleev called the “father of the modern periodic table” and not Meyer, or both?  

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•  stated that if the atomic weight of an element caused it to be placed in the wrong group, then the weight must be wrong. (He corrected the atomic masses of Be, In, and U)

•  was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown.

Mendeleev...  

After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions for Sc, Ga, and Ge were amazingly close to the actual values, his table was generally accepted.  

However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights determined. By looking at our modern periodic table, can you identify what problems might have caused chemists a headache?  

Ar and K  Co and Ni  Te and I  Th and Pa  

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Henry Moseley  

1887 - 1915  

In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number.

*“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.”  

Henry Moseley  

His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.  

Glenn T. Seaborg  After co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became known as the Actinide series.

1912 - 1999  

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Glenn T. Seaborg  

He is the only person to have an element named after him while still alive.

1912 - 1999  

"This is the greatest honor ever bestowed upon me - even better, I think, than winning the Nobel Prize."  

Here  is  a  simple  periodic  table...  

Comments...  

•  This  is  an  older  style  table.  

•  The  rows  are  called  “periods.”  

•  The  columns  are  “families”  or  “groups.”  –  Type  A:    representa.ve  –  Type  B:    transi.on  –  The  guys  at  the  boSom  are  inner  transi)on  elements.  

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Elements  in  Columns  Have  Similar  Proper.es  

•  This  is  especially  true  of  type  A  (representa.ve)  elements.  

•  But,  there  are  changes  in  reac.vity,  etc.,  as  one  goes  down  a  column.  

•  This  link  gives  group  IA  as  an  example...  

hSp://www.youtube.com/watch?v=Ft4E1eCUItI  

 

•  More  about  ver.cal  trends  later!  

 

Looking  at  the  rows...  

•  The  rows  are  called  periods.  •  How  they  fill  gives  insight  into  the  atomic  structures.  

•  In  par.cular,  we  shall  examine  the  idea  of  electron  configura)ons.  

Some  Points...  

•  As  just  stated,  rows  of  the  periodic  table  are  called  periods.  

•  These  have  the  following  lengths:  o 2  o 8  o 18  o 32  

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Look  at  these  numbers  in  detail...  

•     2  =  2  x  12  •     8  =  2  x  22  •  18  =  2  x  32  •  32  =  2  x  42  

•  This  gives  some  insight  into  quantum  numbers!  

In  the  last  chapter  we  saw...  

•  3  quantum  numbers,  n,  l,  and  ml.  •  For  each  value  of  n  we  have  n2  orbitals.  •  We  play  some  number  games  here:  

v n  =  1  →  1  orbital  v n  =  2  →  4  orbitals  (4  =  1  +  3)  v n  =  3  →  9  orbitals  (9  =  1  +  3  +  5)  v n  =  4  →  16  orbitals  (16  =  1  +  3  +  5  +  7)  

•  These  are  all  ½  the  lengths  of  the  periods!    What’s  wrong?  

Solu.on  to  this  problem...  

•  H  has  just  one  electron.  •  The  Schrödinger  equa.on  seen  earlier  needed  another  quantum  number  to  handle  mul.-­‐electron  atoms.    (Technical  detail:    This  comes  naturally  if  rela.vity  included.)  

•  This  new  quantum  number  is  ms,  the  electron  spin.    This  has  values  of  =  +1/2  and  -­‐1/2.  

•  The  physical  explana.on  is  lee  for  lecture!  

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Some  Pictures  of  Spin  

Schema'c  Picture   Language  o2en  used:  

 Spin  up,  Spin  down  

 Spin  +1/2,  Spin  -­‐1/2  

 

α        β    

|+  >          |-­‐  >  

We  observe  spin  states  with  a  magne.c  field  quite  oeen  

Spins  precess  about  the  magne.c  field...  

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Other  depic.ons  of  spin  

Electron  spin  discovered  in  the  Stern-­‐Gerlach  Experiment  

The  Guys...  

O7o  Stern   Walter  Gerlach  

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Effect  on  energy  levels...  Single  Electron  

(Schrödinger  Produces  Degeneracies)  Mul'ple  Electrons  

(Degeneracies  Li2ed!)  

Atomic  Orbitals  &  the  Pauli  Exclusion  Principle  

•  Pauli  principle:    No  two  electrons  in  the  same  atom  can  have  the  same  four  quantum  numbers.  

•  This  means,  only  two  atoms/atomic  orbital.  

•  Orbitals  fill  in  order  of  increasing  energy  (part  of  the  AuEau  principle—more  later!).  

•  E(s  orbital)  <  E(p)  <  E(d)  <  E(f),  for  a  given  value  of  n.  

Some  more  guys...  

Wolfgang  Pauli   Pauli  &  Bohr  Studying  Spin  

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Ordering  Changes  Caused  by  Shielding/Penetra.on  

Shielding     Penentra'on  

An  alternate  view  via  radial  distribu.on  func.ons...  

Orbital  Filling:    The  AuEau  Principle  

Niels  Bohr   Friedrich  Hund  

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The  AuEau  Principle  

•  Electrons  fill  atoms  in  order  of  increasing  orbital  energy.    q 1s  2s  2p  3s  3p  4s  3d  4p  5s  4d  5p  6s,  etc.  

•  Degenerate  orbitals  fill  so  as  to  maximize  the  number  of  parallel  electron  spins.  q This  is  Hund’s  rule!  

Order  of  filling  some.mes  visualized  as  follows:  

Filling  Atomic  Orbitals  Fairly  Simple  

•  We  look  at  the  first  few  elements.  

•  Note  that  we  are  comparing  a)  Electron  configura.ons  

&  

b)  Orbital  diagrams.  

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Elements  3  –  10...  

Nota.onal  Shortcut  

•  We  oeen  write  in  CORE  electrons  as  their  INERT  GAS  SHELLS.  

•  Examples:  –  Na  as  [Ne]3s1  instead  of  1s22s22p63s1  

–  Rb  as  [Kr]5s1  instead  of  a  really  long  mess!  

•  We  see  the  first  18  elements  to  the  right...  

The  periodic  table  shows  how  we  can  think  in  terms  of  s,  p,  d,  and  f  blocks...  

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One  thing  should  be  obvious...    

You  can  construct  an  electron  configura.on  

directly  from  the  periodic  table  (usually)!  

Atoms  take  on  electrons  in  an  orderly  manner—usually  but  irregulari.es  can  occur.    Do  you  see  any  

here?  

A  corollary  that  follows  from  Hund’s  rule  &  AuEau...  

•  Half-­‐filled  subshells  are  par.cularly  stable.  

•  Fully-­‐filled  subshells  are  par.cular  stable.  

•  What  appears  to  be  an  anomaly  really  is  not!  

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Atoms-­‐first  Viewpoint  Here...  

•  Electronic  configura.ons  build  atoms.  •  Atoms  with  similar  configura.ons  are  similar  in  proper.es.  

•  The  periodic  table  and  configura.ons  are  in.mately  related.  

•  In  the  next  few  slides,  we  shall  show  a  few  of  these  happy  families  of  elements.  

The  most  metallic  of  the  metals...  

Alkali  Metals   Alkaline  Earth  Metals  

The  most  noxious  and  snobbiest  elements...  

The  Halogens   The  Noble  Metals  

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PreSy  Pictures  of  the  Halogens  

Fluorine   Cl,  Br,  and  I  

One  last  family  

•  The  last  family  we  list  is  the  chalcogens.  

•  Some.mes,  oxygen  is  lee  out  of  the  group.  

•  The  group  portrait  is  to  the  right...  

The  black  sheep  of  the  families  

•  The  elements  in  the  2nd  row  of  the  periodic  table  oeen  have  quite  different  proper.es  than  the  ones  below  them.  

•  The  elements  to  watch  here  are  Li,  Be,  B,  C,  N,  O,  and  F.  

•  Why  this  happens  will  be  explained  later.    But,  for  now:    Beware  of  the  2nd  row!  

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Notable  things  about  the  excep.ons...  

•  Li  reacts  directly  with  N2.    No  other  element  does.  

•  Be  is  a  deadly  poison.    Mg  and  Ca  are  nontoxic  and  the  others  are  mildly  toxic  in  this  family.  

•  B,  N,  C,  and  O  can  form  double  and  triple  bonds.    Elements  below  them  cannot.  

•  F  does  not  form  any  ions  other  than  F-­‐.    (Cl,  for  instance,  forms  ClO-­‐,  ClO2

-­‐,  ClO3-­‐,  and  ClO4

-­‐.)  

A  Useful  Aside:    Elements  that  have  predictable  ion  charges...  

Enough  about  groups!  

 

Let  us  now  discuss  

PERIODS!!!  

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Trends  across  periods...  

•  We  go  let  to  right.  •  Here  are  some  things  we  shall  examine:  – Atomic  size  (radii)  –  Ion  size  –  Ioniza.on  energies  – Electron  affini.es  – Metallic  Character  

•  Trends  are  generally  smooth  but  there  are  “bumps”  to  discuss!  

Atom  Size  (Atomic  radius)  

Defini.ons  of  atomic  size:    

Nonbonding  atomic  radius  (van  der  Waals)  Bonding  atomic  radius  

Nonmetals:    One-­‐half  the  distance  for  two  iden.cal  atoms  bonded  together  

Metals:    One-­‐half  the  distance  between  nuclei  of  two  touching  atoms  in  the  metal  crystal    

Compare  Two  of  These  

van  der  Waals  radius  (Kr)   Bromine  radius  in  Br2  

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Trends  in  Atomic  Radius  

•  A  slight  surprise:    Atoms  get  smaller  from  lee  to  right  in  a  period!  

•  But,  there  are  a  few  bumps.  

•  There  is  NO  surprise  with  columns.    Things  get  bigger  as  you  go  down!  (Example:    Li  <  Na  <  K  <  Rb  <  Cs)  

•  We  shall  discuss  these  things  verbally  as  the  next  two  slide  are  presented.  

Atom  Size  (vs.  Z)  Shown  Graphically  

Atom  Size:    Periodic  Table  View  

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Causes  of  these  trends...  

•  As  we  move  down  a  column  in  the  periodic  table,  the  principal  quantum  number  of  the  electrons  increases  resul.ng  in  larger  orbitals  and,  hence,  larger  atomic  radii.  

•  As  we  move  to  the  right  across  a  row,  the  effec.ve  nuclear  charge  experienced  by  the  outermost  electrons  increases  (for  a  given  orbital)  and,  hence,  the  radius  shrinks.                  (Zeff  =  Z  –  s,  where  s  is  the  screening  constant.)  

Let’s  Look  at  Monatomic  Ions  

•  Ca.ons  are  posi.ve  (atoms  lose  electrons)  •  Anions  are  nega.ve  (atoms  gain  electrons)  

•  Most  ions  form  for  one  of  the  following  reasons:  – Ge|ng  an  inert  gas  core  – ASaining  a  half-­‐filled  d-­‐  or  f-­‐subshell    – ASaining  a  fully-­‐filled  d-­‐  or  f-­‐subshell    

Examples  of  Inert  Core  Ions...  

•  Each  set  has  the  same  electronic  structure  (hence,  we  call  these  isoelectronic)  

•  First  set:    N3-­‐,  O2-­‐,  F-­‐,  Ne,  Na+,  Mg2+,  Al3+  

•  Second  set:    P3-­‐,  S2-­‐,  Cl-­‐,  Ar,  K+,  Ca2+,  Sc3+,  Ti4+  

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Here  are  some  transi.on  metals...  

Ag  loses  just  1  electron   Zn  needs  to  lose  2  

Some  At  the  Board  Examples...  

Diamagne'sm  

•  Electrons  all  paired.  •  Weakly  repelled  by  magnet.  

•  Do  Al3+  and  S2-­‐  as  examples.  

Paramagne'sm  

•  One  or  more  electrons  unpaired  (either  an  odd  number  or  caused  by  Hund’s  rule).  

•  We  look  at  Fe2+  and  Fe3+  as  examples.  

Ionic  Radii  of  Ca.ons  (pm)  

•  Removing  an  electron  obviously  makes  an  ion  smaller.  

•  Each  row  of  atoms  to  the  right  consists  of  isoelectronic  sets.  

•  Ions  get  smaller  as  we  go  to  the  right.  

•  Ions  get  larger  as  we  go  down.  

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Ionic  Radii  of  Anions  (pm)  

•  Here  the  trends  are  different  in  that  ions  get  larger  as  we  go  to  the  right.  

•  They  s.ll  get  larger  as  we  go  down.  

•  All  the  pairs  are  isoelectronic.  

Ioniza.on  Energies  

•  Ca.on  forma.on  is  generally  endothermic.  •  It  takes  energy  to  remove  an  electron.  

•  Only  valence  electrons  get  removed  in  ordinary  chemical  reac.ons.  

•  Core  electrons  cannot  be  removed  in  ordinary  reac.ons!  

•  The  next  charts  show  trends  in  first  ioniza)on  energies.  

First  Ioniza.on  Energy  Trends  

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Second  and  Higher  Ioniza.on  Energies  

•  Only  valence  electrons  can  be  removed.  •  Each  successive  electron  harder  to  remove  since  it  must  now  be  separated  from  a  ca.on  as  opposed  to  a  neutral  atom.  

•  Core  electrons  stay  put!  •  We  show  these  trends  in  the  next  two  slides.  

Comparing  Na  and  Mg  

Table  for  Na  –  Ar  (kJ/mol)  

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Electron  Affini.es  

•  This  is  the  energy  (usually)  given  off  when  a  neutral  atom  receives  an  electron  to  become  an  anion.  

•  Note  that  this  is  also  (sign  change!)  the  energy  needed  to  convert  an  electron  back  into  a  neutral  atom.  

•  The  next  chart  shows  some  of  these  for  X  +  e-­‐  →  X-­‐.  

A  Short  Table  of  E.A.’s  (kJ/mol)  

Three  types  of  elements  

•  Metals  •  Nonmetals  

•  Metalloids  

•  We  discuss  all  these  while  looking  at  the  next  slide.  

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Trends  in  Metallic  Character  

A  pre|er  selec.on!  

At  last...