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Bohr’s single quantum number (n) was
expanded to a total of four quantum numbers
n, l, ml, and ms
These four quantized values describe an
electron in an atom (quantized values are
restricted to certain discrete values)
These values add order to our description of
the electron in the atom
Table 1: Summary of Quantum Numbers
Table 2:
Values and Letters for the Secondary Quantum Numbers
Orbits
Orbitals
2-D path
3-D region in space
Fixed distance from
nucleus
Circular or elliptical
path
2n2 electrons per orbit
Variable distance from
nucleus
No path; varied shape
or region
2 electrons per orbital
Table 4: Energy Levels, Orbitals, and Shells
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The first two quantum numbers (n and l)
describe electrons with different energies
under NORMAL circumstances
The last two quantum numbers (ml and ms)
describe electrons with different energies
under SPECIAL conditions (e.g. magnetic field)
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Moving forward, we will be focusing on the
electrons position in space (not energy), the
language will change
◦ Main (principal) energy level = shell
◦ Energy sublevel = subshell
◦
◦
◦
◦
WHY?
Its easier!
1s orbital can be communicated as n=1, l=0
2p orbital can be communicated as n=2, l=1
Table 5:
Classification of Energy Sublevels (subshells)
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Figure 1: Energylevel diagrams
show the relative
energies of
electrons in
various orbitals
under normal
conditions (each
orbital can hold a
maximum of 2 e-)
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The energy of an electron increases with an
increasing value of principal quantum
number, n
For a given number of n, the subshells
increase in energy, in order, s<p<d<f
When creating energy-level diagrams, an
electron in an orbital is represented by
drawing an arrow, pointed up or down in a
specific circle, but two arrows in a circle
MUST be in opposite directions
Figure 2: Energy-level diagrams
for (a) hydrogen (b) helium
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Pauli Exclusion Principle – no two electrons in
an atom can have the same four quantum
numbers; no two electrons in the same
atomic orbital can have the same spin, only
two electrons with opposite spins can occupy
any one orbital
What order do we fill the orbitals?
Aufbau Principle – each electron is added to
the lowest energy orbital available in an atom
or ion
An energy sublevel must be filled before
moving onto the next higher sublevel
Figure 3:
In this aufbau
diagram, start at
the bottom (1s) and
add electrons in the
order shown by the
diagonal arrows.
You work your way
from the bottom
left corner to the
top right corner.
Figure 4: Classification of elements by
the sublevels that are being filled

Hund’s Rule – one electron occupies each of the
several orbitals at the same energy before a
second electron can occupy the same orbital
SEATWORK
 Read pp. 189 – 199 - Drawing energy-level
diagrams for atoms, anions, cations
 Practice p. 191 UC # 3, 4
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If you’re thinking this is too easy to be true,
you’re right!
There are a few complications as the atoms
get larger
As the energy level gets farther from the
nucleus, the distance between energy levels
decreases
As a matter of fact, it is believed that the
energy levels actually overlap
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Therefore, some energy levels start filling
orbitals before the previous energy level is
finished filling its subshell
The first time this is encountered is with
potassium, in which the 4s starts to fill before
the 3d
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The second complication has to do with a
variation of Hund’s Rule that takes into
account the minimizing of the electronelectron repulsion
It states, the most stable arrangement of
electrons is the arrangement with the
maximum number of unpaired electrons.
So, when the transition metals’ orbitals are
filling with electrons, at d4 and d9, an electron
from the s JUMPS up into the d5 and d10
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Overall energy state of the atom is lower after
the promotion of the electrons
Half-filled and filled subshells are more
stable (lower energy) than unfilled subshells
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A method of communicating the location and
number of electrons in electron energy levels
(presents same information as energy-level
diagrams BUT much more concise)
Figure 5:
Example of electron
configuration
Writing Electron Configurations
•The electron configuration below represents
a boron atom in its ground state.
•The superscripts indicate the number of
electrons occupying each sublevel.
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Writing out electron configurations can become
awkward as the atoms increase in the number of
electrons
The shorthand involves using the abbreviation of
the last noble gas (placed in square brackets) to
indicate that all the orbitals to that point are full.
Then the configuration is continued as usual.
Nitrogen
1s22s22p3
[He] 2s22p3
Chromium
1s22s22p63s23p64s23d4
[Ar] 4s23d4
Strontium
1s22s22p63s23p64s23d104p65s2
[Kr] 5s2
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Learning Checkpoint
Read pp. 192 – 193
◦ Understand FULL electron configuration and
Shorthand (NOBLE GAS CORE) electron
configurations
◦ Add the summary for “Procedure for Writing an
Electron Configuration” on p. 193

Complete “Electron Configuration” worksheet

Practice Questions p. 194 UC # 6, 8, 9, 10

Section 3.6 Questions p. 197
UC # 2, 3, 4, 5, 6, 7, 8, 9 10, 11, 12, 13, 14
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