Atomic Structure and the
Periodic Table
OR
“The world makes sense now!”
• Calculations show that
– Any s sublevel may contain one pair of
electrons
– Any p sublevel three pairs
– Any d sublevel five pairs
– Any f sublevel seven pairs
Each pair in a given sublevel has a different
place in space.
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Maximum Number of Electrons
In Each Sublevel
Maximum Number of Electrons In Each Sublevel
Sublevel
Number of Orbitals
Maximum Number
of Electrons
s
1
2
p
3
6
d
5
10
f
7
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LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 146
Energy-Level Diagrams
Refer to Pg 162
• Indicate which orbital energy levels are
occupied by electrons for a particular atom or
ion.
• Show the relative energies of electrons in
various orbitals under normal conditions.
• Notice the overlapping of the sublevels in the
fourth, fifth and subsequent orbitals – the
effect is that an electron has a lower overall
energy if the 4s sublevel is of lower energy
than the 3d sublevel.
Energy Level Diagram of a Many-Electron Atom
6s
6p
5d
5s
5p
4d
4p
3d
4f
32
18
4s
18
Arbitrary
Energy Scale
3s
3p
2s
2p
8
8
1s
2
NUCLEUS
O’Connor, Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 177
2
Rules for Energy-Level Diagrams
• An atom is assumed to be its lowest, or
ground state.
• Used as a tool to predict and understand the
concepts of reactivity and chemical bonding.
• There are 3 rules for creating energy-level
diagrams.
• Follow this procedure until the number of
electrons placed in the energy-level diagram
for the atom is equal to the atomic number
for the element.
1. Pauli Exclusion Principle
– Recall that no two electrons can have the same
quantum numbers.
– An electron in an orbital is shown by drawing an
arrow, pointed up or down to represent the
electron spin.
Two electrons occupying the same orbital
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2. Aufbau Principle
– Fill the lower energy levels first (the ones
closest to the nucleus).
– Electrons (arrows) are placed into the
orbitals by filling the lowest energy
orbitals first. An energy sublevel must be
filled before moving on to the next higher
level.
How do I know which order to list
the energy levels and orbitals?
• The energy descriptions of the first
and second quantum numbers fit
perfectly with both the arrangement of
electrons and the structure of the
periodic table.
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3. Hund’s Rule
– Place one electron into each energy level
of the same energy before pairing the
electrons.
- Spread out the electrons as much as
possible before pairing them.
• Practice…
• Draw the energy level diagram for
oxygen:
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Energy-Level Diagrams for
Anions
• Same procedure as for atoms, but add
the extra electrons corresponding to
the ion charge to the orbitals.
• Example:
Draw the energy level diagram for the
sulfide ion S2-.
Energy-Level Diagrams for
Cations
• Draw the energy-level diagram for the
corresponding neutral atom first, and then
remove the number of electrons
(corresponding to the ion charge) from the
orbitals with the highest principal quantum
number, n.
• The electrons removed might not be the
highest-energy electrons, but, in general, this
produces the correct arrangement of energy
levels based on experimental evidence.
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• Example:
Draw the energy-level diagram for the
zinc ion Zn2+
Electron Configurations
• Energy-level diagrams are a good way to
visualize the different energy levels of
electrons in atoms, but they are very
cumbersome.
• Electron configurations provide the same
information as energy-level diagrams but in a
much more concise format.
• It is a list of the number and kinds of
electrons in order of increasing energy.
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Example – electron configuration
for the oxygen atom
1s22s22p4
Energy-level diagram for the
oxygen atom
Electron configuration for the
Oxygen atom
Writing Electron
Configurations
• Follow the periodic table writing out the
filled orbitals, then carefully count the
final number of electrons in the outer
orbital and make sure it corresponds
with the atomic number.
• Example:
Identify the element whose atoms have the
following electron configuration:
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p4
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• Example:
Write the electron configuration for
the tin atom and the tin (II) ion
Shorthand Electron
Configurations
• The core electrons of an atom are
expressed by using a symbol to
represent all of the electrons of the
preceding noble gas.
• This reflects the stability of the noble
gases and the theory that only the
electrons beyond the noble gas are
chemically important for explaining
chemical properties.
• Cl: 1s22s22p63s23p5
Shorthand [Ne]: 3s23p5
• Sn:
1s22s22p63s23p64s23d104p65s24d105p2
Shorthand [Kr]: 5s24d105p2
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Anomalous Configurations
• Transition metals: An element whose
higher-energy electrons are in dorbitals.
• Draw the expected shorthand
configuration of chromium (Cr):
[Ar]: 4s23d4
• However, the observed configuration is:
[Ar]: 4s13d5
• Chromium is an exception to the Aufbau
principal. One explanation for this
anomaly is provided by experimental
evidence that indicates that unfilled
subshells are less stable than half-filed
and filled subshells, and that unfilled
subshells have higher energy,
• In the chromium atom, an s electron
moves to the d-subshell and creates two
half-filled s and d subshells.
[Ar]: 4s23d4 becomes [Ar]: 4s13d5
• This movement of electrons creates an
overall energy state that is lower and
therefore more stable.
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• Another example is Copper:
Expected shorthand configuration of
copper (Cu):
[Ar]: 4s23d9
However, the observed configuration is:
[Ar]: 4s13d10
Explaining Magnetism
• Iron, cobalt, and nickel are all strong
ferromagnetic materials.
• Presence of several unpaired electrons, is an
initial explanation for this phenomena.
• However, ruthenium, rhodium, and palladium
all have unpaired electrons yet are weakly
magnetic and consider paramagnetic.
Ferromagnetic – Strongly
Magnetic
• Ferromagnetism is based
on the properties of a
collection of atoms
rather than just one
atom.
• Effective magnets are
able to orient themselves
in the presence of a
magnetic field. Groups
of tightly packed domains
are aligned in the same
direction.
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Ferromagnetism
Paramagnetism
• Any atom with an odd number of electrons.
• Some atoms with even numbers of electrons
are paramagnetic if their electrons are
unpaired in p, d or f orbitals.
• Unpaired electrons account for some
magnetism but not for strong ferromagnetism
• Domains do not form.
• Based on the magnetism of individual atoms.
• Some examples include: magnesium,
molybdenum, lithium, and tantalum.
Paramagnetism of a Transition
Metal Salt
• Paramagnetism of oxygen
• Dancing oxygen droplets
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• Diamagnetic – electron spins are paired,
the magnetic effects cancel out and the
atom is diamagnetic (no magnetic field).
• Some examples include: copper, silver,
and gold.
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