Atomic Structure
Electron Configuration
Scandium 3-D video (2:31)
3-D Graphic Examples of Atomic Orbitals
Jumping Electrons


normally electrons exist in the ground state,
meaning they are as close to the nucleus as
possible
when an electron is excited by adding energy to
an atom, the electron will absorb energy and
"jump" to a higher energy level
 heating
a chemical with a
Bunsen burner is enough
energy to do this


after a short time, this electron will
spontaneously "fall" back to a lower energy
level, giving off a quantum of light energy called
a photon
the key to Bohr's theory was the fact that the
electron could only "jump" and "fall" to precise
energy levels, thus emitting a limited spectrum of
light.
 quantum is the amount of energy required to
move an electron from one energy level to
another
Quantum Numbers (however, actual numbers
are often not used)

each electron in an atom is described by four
different quantum numbers
 think
of the 4 quantum numbers as the address of an
electron… country > state > city > street



electrons fill low energy orbitals before they fill
higher energy ones
the first three of these quantum numbers (n, l,
and m) represent the three dimensions in which
an electron could be found
the fourth quantum number (s) refers to a certain
magnetic quality called spin

Principle quantum number (n)
Quick intro,
more later.
 describes
the SIZE of the orbital or ENERGY LEVEL
(shell) of the atom.

Angular quantum number (l)
a
SUB-LEVEL (shell) that describes the type or
SHAPE of the orbital

Magnetic quantum number (m)
 the
NUMBER of orbitals
 describes an orbital's ORIENTATION in space

Spin quantum number (s)
 describes
the SPIN or direction (clockwise or counterclockwise) in which an electron spins
Principle
Quantum # (n)
LEVEL/SIZE
Angular
Quantum # (l)
ORBITAL
SHAPE or
SUBLEVEL
Magnetic
Quantum #
(m)
AXIS/
ORIENTATION
or ORBITALS
1
2
s
s p
s
p
d
s
p
d
f
1
1
1
3
5
1
3
5
7
3
3
4
1
orbital
4 total
orbitals
9 total orbitals
2 e-
8 e-
18 e-
16 total orbitals
Spin
Quantum # (s)
DIRECTION
OF
ELECTRON
SPIN
32 e-
4f
= level and sub-level
= max. # of electrons
= # of electrons
= number of orbitals
14 (7)
4d
10 (5)
4p
6 (3)
4s
2 (1)
32
3d
10 (5)
3p
6 (3)
3s
2 (1)
18
2p
6 (3)
2s
2 (1)
8
1s
2 (1)
2
Principle Quantum Number (n) or
Energy Level


values 1-7 used to specify the level the electron
is in
describes how far away from the nucleus the
electron level is
 the
lower the number, the closer the level is to the
atom's nucleus and less energy*

maximum # of electrons that can fit in an energy
level is given by formula 2n2
Angular Quantum Number (l) or
SUB-LEVELS
determines the shape of the sub-level
 number of sub-levels equal the level number

 ex.

the second level has two sub-levels
they are numbered but are also given letters
referring to the sub-level type
l=0 refers to the s sub-level
l=1 refers to the p sub-level just know this
l=2 refers to the d sub-level
l=3 refers to the f sub-level
Magnetic quantum number (m)
or ORBITALS
Electron Orbitals YouTube 1:37
 the third of a set of quantum numbers
 tells us how many sub-levels there are of
a particular type and their orientation in
space of a particular sub-level
 only two electrons can fit in an orbital

= electron

S sub-level
has only 1 orbital
only holds two electrons
P sub-level
has 3 orbitals
holds up to six electrons
D sub-level
has 5 orbitals
holds up to 10 electrons
F sub-level
has 7 orbitals
holds up to 14 electrons
Spin quantum number (s)
the fourth of a set of quantum numbers
 number specifying the direction of the spin
of an electron around its own axis.

 only
two electrons of opposite spin may
occupy an orbit
 the only possible values of a spin quantum
number are +1/2 or -1/2.
Principle
Quantum # (n)
SHELL/SIZE
Angular
Quantum # (l)
ORBITAL
SHAPE or
SUBSHELL
Magnetic
Quantum #
(m)
AXIS/
ORIENTATION
or ORBITALS
1
2
s
s p
s
p
d
s
p
d
f
1
1
1
3
5
1
3
5
7
3
3
4
1
orbital
4 total
orbitals
9 total orbitals
2 e-
8 e-
18 e-
16 total orbitals
Spin
Quantum # (s)
DIRECTION
OF
ELECTRON
SPIN
32 e-
Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy
Levels
Principle
energy
level (n)
1
2
3
4
Number of
orbitals
per type
Number of
orbitals
per
level(n2)
s
1
1
2
s
1
p
3
4
8
s
1
p
3
9
18
d
5
s
1
p
3
d
5
16
32
f
7
Type of
sublevel
Maximum
number of
electrons
(2n2)
“Rules” for Writing Electron
Configurations

a method of writing where electrons are
found in various orbitals around the nuclei
of atoms.
 three
1.
2.
3.
rules in order to determine this:
Aufbau principle
Pauli exclusion principle
Hund’s rule
Aufbau Principle
electrons occupy the orbitals of the
lowest energy first
 each
written represents an atomic
orbital (such as
or
or
or
….)
 electrons in the same sublevel/shell have
equal energy ( same energy as
)
 principle energy levels/shells (1,2,3,4..)
can overlap one another

 ex:
4s orbital has less energy than a 3d orbital
Pauli Exclusion Principle

Hamster
video 1:00
only two electrons in an orbital

must have opposite spins
 represents one electron

represents two electrons in an orbital
actually incorrect as well, see next slide
Hund’s Rules
every orbital in a subshell must have one
electron before any one orbital has two
electrons
 all electrons in singly occupied orbitals
have the same spin.

Writing Orbital Diagrams
Energy

Orbitals grouped in s, p, d, and f orbitals (sharp,
proximal, diffuse, and fundamental)
s orbitals
d orbitals
p orbitals
f orbitals
Boron
Atomic # 5
http://colossus.chem.umass.edu/genchem/whelan/class_ima
ges/Orbital_Energies.jpg
Boron ion (3+)
Atomic # 5
http://colossus.chem.umass.edu/genchem/whelan/class_ima
ges/Orbital_Energies.jpg
Neon
Atomic # 10
http://colossus.chem.umass.edu/genchem/whelan/class_ima
ges/Orbital_Energies.jpg
Bromine
Atomic # 35
http://colossus.chem.umass.edu/genchem/whelan/class_ima
ges/Orbital_Energies.jpg
Bromine ion (1-)
Atomic # 35
http://colossus.chem.umass.edu/genchem/whelan/class_ima
ges/Orbital_Energies.jpg
1s
2s
2p
3s
3p
Na
1s2 2s2 2p6 3s1
Mg
1s2 2s2 2p6 3s2
Al
1s2 2s2 2p6 3s2 3p1
Si
1s2 2s2 2p6 3s2 3p2
P
1s2 2s2 2p6 3s2 3p3
S
1s2 2s2 2p6 3s2 3p4
Cl
1s2 2s2 2p6 3s2 3p5
Ar
1s2 2s2 2p6 3s2 3p6
Orbital diagrams
Electron Configurations
Writing Electron Configurations

To write out the electron configuration of an atom:
 use the principal quantum number/energy
level (1,2,3, or 4…)
 use the letter term for each sub-level (s,p,d, or f);

don’t worry about orientation such as x,y,z axis but you
do have to be able to draw these for IB
 use
a superscript number indicates how many
electrons are present in each sub-level
 hydrogen =1s1.
 Lithium =1s22s1.
 don’t write anything for spin
Electron Configurations
Energy Level
4
2p
Number of electrons in
the sublevel
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
6s2 4f14… etc.
Order of Electrons
Sometimes levels are switched in order to keep the level together.
I hate when they do that! 4s requires less energy and I think it
should be before 3d.
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
Weird electron configuration video (3:24)
exceptions (don’t need to know this, just
be aware that there are exceptions)
 orbitals “like” to be empty, half filled, or full


therefore, an electron leaves the 4s (leaving it half
full) and goes to the 3d in order to make it full
Cr
we would predict:
1s2 2s2 2p6 3s2 3p6 4s2 3d4
but it is actually:
1s2 2s2 2p6 3s2 3p6 4s13d5
Cu
we would predict:
1s2 2s2 2p6 3s2 3p6 4s2 3d9
but it is actually:
1s2 2s2 2p6 3s2 3p6 4s1 3d10
Noble Gas Shortcut
same