Electron Configuration
The Electromagnetic Spectrum
Know what
is in the red
boxes
High frequency
lower frequency
Short wavelength
longer wavelength
High energy
lower energy
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
• Emission line spectrum
– energy is applied to a specific element
• this “excites” the element and the light is viewed
through a spectroscope
– a continuous spectrum is NOT observed, but a series of
very bright lines of specific colors with black spaces inbetween instead
– unique for every element and are used to identify atoms
(much like fingerprints are used to identify people)
More on emission line spectrum
Give off energy
when falls back
down to
normal energy
level
– the process
• electrons surround the nucleus in specific orbitals
or energy levels
• when electrons are excited (heat/electricity) they
can move to a higher energy level
• when they move back down they emit energy in
the form of electromagnetic radiation
• because electrons can only exist in certain energy
levels, only certain transitions can occur
• the color of the light emitted depends on the
frequency of the emitted photon
• http://www.youtube.com/watch?v=QI50GBUJ48s
this is a repetitive slide- just couldn’t bear to delete it
1. an electron in the
atom gains (absorbs)
energy from heating
2. electron jumps up an
energy level.
3. electron is now
unstable (unwelcome)
in this level and is
“kicked out”
4. when the electron
loses the energy and
come back to the
original level, light is
emitted
The Atomic Emission Spectrum of
Hydrogen
• the emission spectrum of hydrogen is the simplest
emission spectrum because there is only one electron
– it is not uniform, but concentrated into bright lines,
indicating the existence of only certain allowed electron
energy levels
– visible light, infrared, and UV are emitted when electrons fall
back down level
– McGraw Hill animation link
convergence
up here
(levels are
close
together)
• 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
Scandium 3-D video (2:31)
3-D Graphic Examples of Atomic Orbitals
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
s
1
2
3
s p
s p
1
1
3
1
orbital
4 total
orbitals
2 e-
8 e-
4
d
s
p
d
5
1
3
5
3
9 total orbitals
16 total orbitals
Spin Quantum #
(s)
DIRECTION OF
ELECTRON SPIN
18 e-
32 e-
f
7
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
s
1
2
3
s p
s p
1
1
3
1
orbital
4 total
orbitals
2 e-
8 e-
4
d
s
p
d
5
1
3
5
3
9 total orbitals
16 total orbitals
Spin Quantum #
(s)
DIRECTION OF
ELECTRON SPIN
18 e-
32 e-
f
7
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 rules in order to determine this:
1. Aufbau principle
2. Pauli exclusion principle
3. 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
• 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
Hamster
video 1:00
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
Electron Configurations
4
2p
Number of electrons in
the sublevel
Energy Level
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14… etc.
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
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