Chapter 8: Atomic Electron Configurations

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Chapter 8 Atomic Electron Configurations and
Chemical Periodicity
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Spin Quantum Number ms

• The electron in an atom behaves as if it has
  spin. This property is quantized, with two
  possible values, and so we need a fourth quantum
  number, ms.
• values of 1/2 (NOT dependent on n, l, or ml )
• identifies an electron in an orbital (max. 2)
• arbitrary 1/2 spin up
• - 1/2 spin down

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Magnetic Properties

• much of what we know about electron spin comes
  from the study of magnetic properties
• diamagnetic repelled by magnetic field
  electrons paired
• paramagnetic attracted to magnetic field
  unpaired electrons

4
Electron Configurations in Atoms

• in hydrogen, energy dependent solely on n
• in multielectron atom, energy depends on both n
  and l
• Aufbau Principle (Ger building up)
• the electronic configuration of lowest energy is
  the one in which each electron is placed in the
  lowest-energy hydrogen-like orbital available.

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Energy Levels for Multielectron Atoms
l 0 1
2
ml 0 -1 0 1 -2 -1 0
1 2

4fs
4d
4p
3d
4s
3p
3s
2p
2s
1s
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Electron Configurations --shorthand

• e.g. hydrogen n 1 l 0 ml 0 ms 1/2
• designated 1s1 ---- 1s
• 1s
• helium
• 1st electron n 1 l 0 ml 0 ms -1/2
• 2nd electron n 1 l 0 ml 0 ms 1/2
• designated 1s2 ---- 1s
• 1s
• This is an example of....

spectroscopic notation
box notation
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Pauli Exclusion Principle

• no two electrons in an atom can have the same set
  of four quantum numbers (n, l, ml, ms).
• a single atomic orbital may contain up to two
  electrons
• s subshell holds up to 2 electrons
• p subshell holds up to 6 electrons
• d subshell holds up to 10 electrons

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Carbon Anomaly

• with carbon (1s22s22p2) there are several
  possibilities
• Which is lowest energy?

or
or
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Hunds Rule (Empirical)

• When filling a subshell, the lowest energy state
  is obtained if the largest number of different
  orbitals in the subshell are utilized.
• When there are 2 or more electrons in half-filled
  subshells, the state of lowest energy is one
  where all the spins are parallel
• lowest state for C is
• note 1s2 2s2 2p2 Hunds Rule is implied

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Electron Configurations cont.
filled Noble Gas shell

• Ne 1s22s22p6
• so, for Na Ne 3s1
• 3s and 3p orbitals filled as might be expected
• Note 4s orbitals slightly lower in energy
  filled before 3d orbitals
• can follow electron configurations with Periodic
  Table

Noble Gas Configuration
implies above configuration
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Electron Configurations and the Periodic Table
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Electron Configurations and the Periodic Table
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Stabilities

• certain electronic configurations tend to be
  especially stable
• half filled subshell e.g. d5
• completely filled subshell e.g. s2
• generally electron configurations are as
  predicted by Aufbau, Pauli, and Hunds
  principles, but there are exceptions.
• e.g. V Ar4s23d3
• Cr Ar4s13d5
• Cu Ar4s13d10
• Note this only occurs between s and d orbitals!

as predicted
each half-filled
filled
half-filled
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Electron Configuration of Ions

• in general, to form an anion or cation from s or
  p-block elements, simply remove or add
  appropriate number of electrons from/to outer
  shell
• Ca -----gt
  Ca2
• 1s22s22p63s23p64s2 -----gt 1s22s22p63s23p6
• Cl -----gt
  Cl-
• 1s22s22p63s23p5 -----gt 1s22s22p63s23p6
• however, with transition metals, outer shell
  s-electrons removed first.
• Co Ar4s23d7 -----gt Co2 Ar3d7
• Fe Ar4s23d6 -----gt Fe3 Ar3d5

-2e-
1e-
-2e-
-3e-
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Periodic Properties Atomic Size
For the Main Group Elements, atomic radii
generally increase going down the periodic
table, due to increase in shell number
(n). decrease going across the periodic table,
due to increased effective nuclear charge
(Zeff).
values in picometers 1 x 10-12 m
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Periodic Properties Ionization Energies
Energy required to remove one electron from an
atom in the gas phase. A(g) A(g) e-
For the Main Group Elements, ionization energies
generally decrease going down the periodic
table, due to increase in shell number (n)
electron is farther from the nucleus. increase
going across the periodic table, due to
increased effective nuclear charge (Zeff).
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Periodic Properties Electron Affinity
Energy of the process when one electron is
aquired by an atom in the gas phase. A(g)
e- A-(g)
By convention these values are negative
(exothermic), with a more negative value meaning
a larger electron affinity. Exceptions in trends
(C to N) due to electronic stabilities.
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Periodic Properties Ionic Radii
Cations are always smaller than the original
neutral atom due to loss of electron density.
Anions are always larger than the original
neutral atom due to gain of electron density.
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First Ionization Energies

• Minimum energy to completely remove an e- from
  the ground state atom in the gas phase.
• H(g) -----gt H e- E 2.179 x
  10-18 J (1312 kJ/mole)
• E h? hc / ? ? hc / E 91.18nm
• i.e. corresponds to the n 1 ------gt n ?
  transition in hydrogen

Periodic Trends
Increases rapidly across periodic table. Dramatic
drop from end of one row to beginning of
next. Gradual decrease down a column on
periodic table. Some minor exceptions across
(e.g. Be --gtB N --gt O).
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Explanation For Trends
What if the electrons are at similar distances
from the nucleus? Compare H and He the
electron being removed from He feels a 2
nuclear charge, compared to a 1 charge in H.
Therefore, should take more energy to remove the
electron from He. Observed!
e- 2 e- He
e- 1 H
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Explanation For Trends
What if the electrons are at similar distances
from the nucleus?
e- 2 e- He
e- 1 H
e- 3 e- Li
e-
For Li would expect 1st ionization energy to be
greater than for He. (NOT observed). What if 3
electrons are not equidistant from nucleus?
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Explanation For Trends

e- 2 e- He
e-
e- 1 H
e- 3 e- Li
For Li would expect 1st ionization energy to be
greater than for He. (NOT observed). What if 3
electrons are not equidistant from nucleus? The
nuclear charge felt by the outer electron would
be shielded by inner two electrons and be equal
to 1. In this case it would be easier to remove
than 1st electron from either H or He.
(Observed!)
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Explanation For Trends
e-
e-
e-
e-
e-
e-
e- 11 e- Na
e- 10 e- Ne
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
For Li through Ne there is an almost linear
increase in 1st ionization energy because of the
linear increase in effective nuclear charge felt
by the outermost electrons (no additional
shielding). But, dramatic drop for Na. Next
electron at farther distance from nucleus, so
only feels 1 core effective nuclear charge.
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Explanation For Trends
n 1
n 2
n 3
Electrons found in groups called Shells, which
we will designate 1, 2, 3, etc. (can hold 2e-,
8e-, 8e-, respectively). Within the shells, the
electrons are similar in distance from the
nucleus. Increasing nuclear charge across
Periodic Table results in more energy being
required to remove electron. Subsequent shells
are relatively further from nucleus and are
shielded.
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Photoelectron Spectroscopy (PES)
First ionization energy removes electron from
outermost shell. PES measures the energy to
remove one electron from any shell of a neutral
atom. Energy of entering photon (h?) is larger
than ionization energy (Ie), so the electron
leaves the atom with excess kinetic energy (KE),
which PES measures. h?
Ie KE known (scanned)
measured determined
The difference between the energy of the photon
and the kinetic energy of the ejected photon
yields the ionization energy.
KE Ie
n ?
h?
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Photoelectron Spectroscopy (PES)

H peak at 1.31 MJ/mole (1312kJ) He peak at
2.37 MJ/mole i.e energy to remove He e- gt H
e- note peak He 2X peak H Li peak at 6.26
MJ/mole 2e- peak at 0.52 MJ/mole
1e- Be peak at 11.5 MJ/mole
2e- peak at 0.90 MJ/mole 2e- B
peak at 19.3 MJ/mole 2e- peak at
1.36 MJ/mole 2e- peak at 0.80
MJ/mole 1e- BUT! Saw with ionization
energies that the three outer electrons for B
belong to the same shell (n 2).




1st ionization energy
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Photoelectron Spectroscopy (PES)
1s1
BUT! Saw with ionization energies that the three
outer electrons for B belong to the same shell (n
2). Must, therefore, refine the model to
include SUBSHELLS---members of the same shell,
but with different energies. Subshells given
designations s, p, d or f, in order of
increasing PES energy. Exponent indicates number
of electrons in that subshell.
1s2
1s2
2s1
1s2
2s2
1s2
2s2
2p1
Same shell, but different subshells.
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Photoelectron Spectroscopy (PES)

• PES Data show only three peaks for B through Ne.
• Peak height of third peak increases in size from
  1 to 6.
• Conclusions
• p subshell can hold up to 6 electrons.
• s subshell can hold only up to 2 electrons

p
s
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Photoelectron Spectroscopy (PES)

• Both Na and Mg show 4 peaks.
• New peak due to 3s subshell
• Al through Ar 5 peaks.
• New peak due to 3p subshell (holds up to 6
  electrons).
• K and Ca show 6 peaks
• New peak due to 4s subshell

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Photoelectron Spectroscopy (PES)

• PES RESULTS WITH TRANSITION METALS
• Would expect next electron from Sc (Z 21) to
  begin new subshell.
• Previous elements new subshell at lower
  ionization energy but, for Sc through Zn, this
  new peak is at greater energy than 4s shell.
• 0.77MJ/mole (peak height 1) vs 0.63MJ/mole (peak
  height 2)
• this new subshell, 3d, is seen to hold up to
  10electrons (Sc ---gt Zn)
• Sc 1s22s22p63s23p64s23d1
• electron configurations listed by the order
  subshells are filled.
• But, because 4s orbital is easier to ionize
  electron from, these 2 electrons are removed
  first when forming ions
• e.g. V2 1s22s22p63s23p63d3

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Energy Levels for Multielectron Atoms
l 0 1
2
ml 0 -1 0 1 -2 -1 0
1 2

4fs
4d
PES Spectrum for Calcium
4p
?
3d
4s
3p
?
3s
PES Energy
2p
?
2s
?
1s