Orbital Diagrams and Electron Configuration by davebusters

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									            Orbital Diagrams and
           Electron Configuration
• Drawing orbital diagrams gives information
  not only about the orbitals that are/have
  been filled but also about the number of
  unpaired electrons.

• Orbital diagrams can be cumbersome!!



                                           1
       Electron Configuration
• A short-hand notation is commonly used in
  place of orbital diagrams to describe the
  electron configuration of an atom.

• Electron configuration:
  – a particular arrangement of electrons in the
    orbitals of an atom


                                                   2
      Electron Configuration
• The electron configuration tells the number of
  electrons found in each subshell.

• If there are three electrons in a 2p subshell, we would
  write:
                            2p3
  where the superscript (3) indicates the number of
  electrons in that subshell




                                                        3
      Electron Configuration
• The orbital diagram for an O atom:


       1s    2s      2p         3s

  The electron configuration for an O atom:

                   1s22s22p4


                                              4
    Electron Configuration
• To determine the electron configuration
  of an atom (or ion) without first writing
  the orbital diagram:
  – determine the number of electrons present
  – add electrons to each subshell in the
    correct filling order until all electrons have
    been added
     • use the “diagonal” diagram to help determine
       the filling order

                                                      5
   Electron Configuration
Example: Write the electron configuration
of a Mn atom (Z = 25).

       1s2 2s22p6 3s23p6 4s2 3d5




                                        6
    Electron Configuration
Example: Write the electron configuration
of an O2- ion (Z = 8).
An O2- ion has 8 protons and 10 electrons

              1s22s22p6




                                            7
     Electron Configuration
Write the electron configuration of a krypton
atom (Z = 36).
        1s22s22p63s23p64s23d104p6

        This is the Kr “core”  [Kr]

• The noble gas “core” can be used to write the
  electron configuration of an element using
  core notation:
      noble gas “core” + valence electrons

                                                  8
                 Core notation
To write the electron configuration using the core notation:
•   Find the Noble Gas that comes before the atom.
•   Determine how many additional electrons must be
    added beyond what that noble gas has.
    (= Atomic number of atom minus atomic number of
    noble gas)
•   Determine the period that element is in. (This
    determines the value of n of the s subshell to start with
    when adding extra electrons)
•   Add electrons starting in that “n” subshell

                                                            9
       Electron Configuration
Write the core electron configuration of Sr
(Z = 38).
   Previous noble gas: Kr (Z = 36)
   Extra electrons: 38 (e of Sr) - 36 = 2
   Period number of Sr: 5
   So: Kr core plus 2 extra e- starting in 5s
                   [Kr] 5s2


                                                10
       Electron Configuration
Write the core electron configuration of Br
(Z = 35).
   Previous noble gas: Ar (Z = 18)
   Extra electrons: 35 - 18 = 17
   Period number: 4
   So: Ar core plus 17 extra e- starting with 4s
                [Ar] 4s23d104p5


                                                   11
         Isoelectronic Series
• When atoms ionize, they form ions with the
  same number of electrons as the nearest
  (in atomic number) noble gas.
   Na = 1s22s22p63s1 = [Ne]3s1
   Na+ = 1s22s22p6     = [Ne]

  Cl = 1s22s22p63s23p5 = [Ne]3s23p5
  Cl- = 1s22s22p63s23p6 = [Ar]

                                               12
      Isoelectronic Series

• N (7 e-):      1s22s22p3
 N3- (10 e-):   1s22s22p6 = [Ne]


• O (8 e-):      1s22s22p4
 O2- (10 e-):   1s22s22p6 = [Ne]

• F (9 e-):      1s22s22p5
 F- (10 e-):    1s22s22p6 = [Ne]
                                    13
        Isoelectronic Series

• Na (11 e-):     1s22s22p63s1
 Na+ (10 e-):    1s22s22p6 = [Ne]

• Mg (12 e-):     1s22s22p63s2
 Mg2+ (10 e-):   1s22s22p6 = [Ne]

• Al (13 e-):     1s22s22p63s23p1
 Al3+ (10 e-):   1s22s22p6 = [Ne]


                                     14
1A
            Ions of the highlighted                                   8A
H    2A
                 elements are
                                               3A   4A 5A 6A 7A He
             isoelectronic with Ne.
Li Be                                           B   C   N   O    F    Ne

Na Mg                                           Al Si   P   S    Cl   Ar
          3B 4B 5B 6B 7B 8B      8B 8B 1B 2B

K    Ca Sc Ti     V   Cr Mn Fe   Co Ni    Cu Zn Ga Ge As    Se   Br Kr


Rb   Sr Y    Zr   Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb          Te I      Xe

Cs Ba La Hf       Ta W    Re Os Ir   Pt   Au Hg Tl Pb Bi    Po At Rn

Fr Ra Ac Rf       Db Sg Bh Hs Mt


                  Ce Pr   Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

                  Th Pa   U   Np Pu Am Cm Bk Cf     Es Fm Md No Lr
                                                                15
         Isoelectronic Series
• Isoelectronic: having the same number of
  electrons

• N3-, O2-, F-, Ne, Na+, Mg2+, and Al3+ form
  an isoelectronic series.
  – A group of atoms or ions that all contain the
    same number of electrons


                                                    16
         Isoelectronic Series
• Examples of isoelectronic series:
  – N3-, O2-, F-, Ne, Na+, Mg2+, Al3+

  – Se2-, Br-, Kr, Rb+, Sr2+, Y3+

  – Cr, Fe2+, and Co3+




                                        17
Periodic Properties of Elements
• Chemical and physical properties of the
  elements vary with their position in the periodic
  table.
   – Atomic size
   – Size of Atom vs. Ion
   – Size of Ions in Isoelectronic series
   – Ionization energy
   – Electron affinity
   – Metallic character
                                                      18
What Is the Size of an Atom?
The bonding atomic
radius is defined as
one-half of the
distance between
covalently bonded
nuclei.




                               19
  Periodic Properties--Atomic Size
• The relative size (radius) of an atom of an
  element can be predicted by its position in the
  periodic table.

• Trends
   – Within a group (column), the atomic radius
     tends to increase from top to bottom

   – Within a period (row), the atomic radius tends
     to decrease as we move from left to right



                                                      20
Sizes of Atoms
        Bonding atomic
        radius tends to…
        …decrease from left to
         right across a row

        …increase from top to
         bottom of a column



                            21
Periodic Properties--Atomic Size
• Example: Which element would have the
  larger atomic radius, Ar or Br?

   Br should have the larger radius
      more towards the bottom
      more towards the left side




                                          22
Periodic Properties – Atom vs. Ion Size


  • Trends to know:
    – Cations (+) are smaller than their
      parent atoms.
       • Electrons are removed from the outer
         shell.
    – Anions (-) are larger than their
      parent atoms.
       • Electron-electron repulsion causes the
         electrons to spread out more in space.
                                                  23
         Sizes of Ions - Trends
• Ions increase in size
  as you go down a
  column.




                                  24
        Sizes of Ions - Trends




• In an isoelectronic series, ions have the same
  number of electrons.
• Ionic size decreases with an increasing nuclear
  charge.

                                                    25
         Ionization Energy

• The ionization energy is the amount of
  energy required to remove an electron
  from the ground state of a gaseous
  atom or ion to form a cation or more
  positively charged cation.
  – The first ionization energy is that energy
    required to remove first electron.
  – The second ionization energy is that
    energy required to remove second
    electron, etc.
                                                 26
Ionization of Gaseous Sodium:

          Na (g)  Na+ (g) + e-

• As the ionization energy increases, it
  becomes harder to remove an electron.


                                           27
           Ionization Energy
• It requires more energy to remove each
  successive electron.
• When all valence electrons have been removed,
  the ionization energy takes a quantum leap.




                                              28
Trends in First Ionization
       Energies
              • As one goes down a
                column, less energy
                is required to remove
                the first electron.
                – valence electrons are
                  farther from the
                  nucleus.
                Within each row, the
                  ionization energy
                  increases from left to
                  right
                                       29
            Ionization Energy
Which element has the higher ionization energy,
Br or Ca? Which one will lose an electron easier?

 Br has the higher ionization energy
    further to the right

 Ca will lose an electron easier because its ionization
 energy is lower.




                                                      30
             Periodic Properties
              Electron Affinity
• The energy change that occurs when an electron is
  added to a gaseous atom is called the electron affinity.

                    Cl (g) + e-  Cl- (g)

• The electron affinity becomes increasingly negative as
  the attraction between an atom and an electron
  increases
   – more negative electron affinity = more likely to gain an
     electron and form an anion

                                                             31
                Electron Affinity
• Trends:
   – Halogens have the most negative electron affinities.

   – Electron affinities become increasing negative moving
     from the left toward the halogens.

   – Electron affinities do not change significantly within a
     group.
   – Noble gases will not accept another electron.
      • To do so would require adding an electron to a
        new electron shell (significantly higher in energy)

                                                              32
             Metallic Character
• Metals:
  – shiny luster
  – malleable and ductile
  – good conductors of heat and electricity
  – form cations

• Metallic character
  – increases from top to bottom
  – Increases from right to left


                                              33
Properties of Metal, Nonmetals,
        and Metalloids




                              34
   Metals versus Nonmetals




Differences between metals and nonmetals tend
to revolve around these properties.



                                            35
     Metals versus Nonmetals
• Metals tend to form cations.
• Nonmetals tend to form anions.




                                   36
Metals
    They tend to be
    lustrous, malleable,
    ductile, and good
    conductors of heat
    and electricity.




                           37
                   Metals
• Compounds formed
  between metals and
  nonmetals tend to be
  ionic.
• Metal oxides tend to
  be basic.




                            38
Nonmetals
     • These are dull, brittle
       substances that are
       poor conductors of
       heat and electricity.
     • They tend to gain
       electrons in reactions
       with metals to acquire
       a noble gas
       configuration.

                             39
               Nonmetals
• Substances
  containing only
  nonmetals are
  molecular
  compounds.
• Most nonmetal oxides
  are acidic.




                           40
Metalloids
     • These have some
       characteristics of
       metals and some of
       nonmetals.
     • For instance, silicon
       looks shiny, but is
       brittle and fairly poor
       conductor.


                                 41

								
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