Transcript Slide 1

Electron Configuration
Where are the electrons around the
nucleus?
Order of sublevel energies
• S<P<D<F
Hund’s Rule
• Electrons are added from lowest to highest
energy orbitals.
• 1 electron is added to each degenerate orbital
in a subshell before adding 2 electrons are
added to any one orbital
Build up
(Aufbau)
Orbital Diagrams
• we often represent an orbital as a square and the
electrons in that orbital as arrows
– the direction of the arrow represents the spin of the
electron
unoccupied
orbital
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orbital with
1 electron
orbital with
2 electrons
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Periodic table
Periods– horizontal rows
Groups—Verticle columns with numbers at the top
1. each period begins with element in 1A column
2. 1A column marks beginning of filling a new
principle energy level
3. each consecutive element on the table has one
more electron than the last
4. Each element in a group has the identical outer
shell configuration except n changes
5. Outer shell (valence shell) configuration accounts
for the properties of the elements
Z effective
Zeff
Z effective
how much the electron “feels” the
nucleus
Penetration & Shielding
Effective Nuclear Charge
• in a multi-electron system, electrons are simultaneously
attracted to the nucleus and repelled by each other
• outer electrons are shielded from full strength of
nucleus
– screening effect
• effective nuclear charge is net positive charge that is
attracting a particular electron
• Z is nuclear charge, S is electrons in lower energy levels
– electrons in same energy level contribute to screening, but
very little
– effective nuclear charge on sublevels trend, s > p > d > f
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Zeffective = Z - S
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Screening & Effective Nuclear Charge
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Trends in Atomic Radius
Transition Metals
• increase in size down the Group
• atomic radii of transition metals roughly the
same size across the d block
– must less difference than across main group
elements
– valence shell ns2, not the d electrons
– effective nuclear charge on the ns2 electrons
approximately the same
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Atomic radius
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Which atom of the following pairs has
a larger radius?
1.
2.
3.
4.
N or F
C or Ge
N or Al
Al or Ge
Choose the
Larger Atom in Each Pair
1)
2)
3)
4)
N or FF, N is further left
C or Ge
Ge, Ge is further down
N or Al
Al, Al is further down & left
Al or Ge? opposing trends
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Electron Configuration of Cations in
their Ground State
• cations form when the atom loses electrons
from the valence shell
• for transition metals electrons, may be
removed from the sublevel closest to the
valence shell
Al atom =
Al+3 ion =
Fe atom =
Fe+2 ion =
Fe+3 ion =
Cu atom =
Cu+1 ion =
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1s22s22p63s23p1
1s22s22p6
1s22s22p63s23p64s23d6
1s22s22p63s23p63d6
1s22s22p63s23p63d5
1s22s22p63s23p64s13d10
1s22s22p63s23p63d10
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Trends in Ionic Radius
• Ions in same group have same charge
• Ion size increases down the group
– higher valence shell, larger
• Cations smaller than neutral atom; Anions bigger
than neutral atom
• Cations smaller than anions
– except Rb+1 & Cs+1 bigger or same size as F-1 and O-2
• Larger positive charge = smaller cation
– for isoelectronic species
– isoelectronic = same electron configuration
• Larger negative charge = larger anion
– for isoelectronic series
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Magnetic Properties of
Transition Metal Atoms & Ions
• electron configurations that result in unpaired electrons
mean that the atom or ion will have a net magnetic field
– this is called paramagnetism
– will be attracted to a magnetic field
• electron configurations that result in all paired electrons
mean that the atom or ion will have no magnetic field –
this is called diamagnetism
– slightly repelled by a magnetic field
• both Zn atoms and Zn2+ ions are diamagnetic, showing
that the two 4s electrons are lost before the 3d
– Zn atoms [Ar]4s23d10
– Zn2+ ions [Ar]4s03d10
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Write the Electron Configuration
and Determine whether the Fe
atom and Fe3+ ion are
Paramagnetic or Diamagnetic
Ionization Energy
• minimum energy needed to remove an electron
from an atom
– gas state
– endothermic process
– valence electron easiest to remove
– M(g) + IE1  M1+(g) + 1 e– M+1(g) + IE2  M2+(g) + 1 e• first ionization energy = energy to remove electron from
neutral atom; 2nd IE = energy to remove from +1 ion; etc.
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Ionization energy
• Ionization energy: energy required to remove
an electron from the atom or ion.
• Li Li+ + 1 e-
Li+  Li2+ + 1e-
IE = 513.3 kJ/mol
(1st ionization energy)
IE = 7298.0 kJ/mol
(2nd ionization energy)
Choose the Atom in Each Pair with the
Higher First Ionization Energy
•
•
•
•
Al or S
As or Sb
N or Si
O or Cl
Choose the Atom in Each Pair with the Higher
First Ionization Energy
1)
2)
3)
4)
Al or SS, Al is further left
As or Sb
Sb, Sb is further down
N or Si,
Si Si is further down & left
O or Cl? opposing trends
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Ionization Energy
• 1st ionization energy
E + Al  Al+ + eE + Al+  Al2+ +eE + Al2+  Al3+ + eE + Al3+  Al4+ + e-
kJ/mol
I=580
I = 1815
I=2740
I = 11,600
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• Metals
Metallic Character
–
–
–
–
–
–
malleable & ductile
shiny, lusterous, reflect light
conduct heat and electricity
most oxides basic and ionic
form cations in solution
lose electrons in reactions - oxidized
–
–
–
–
–
–
brittle in solid state
dull
electrical and thermal insulators
most oxides are acidic and molecular
form anions and polyatomic anions
gain electrons in reactions - reduced
• Nonmetals
• metallic character increases left
• metallic character increase down
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Choose the
More Metallic Element in Each Pair
•
•
•
•
Sn or Te
P or Sb
Ge or In
S or Br
Choose the
More Metallic Element in Each Pair
1)
2)
3)
4)
Sn or Te
Te, Sn is further left
P or Sb,
Sb Sb is further down
Ge or In
In, In is further down & left
S or Br? opposing trends
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