Electron Configuration
Atomic Models and Energy Levels
Models of the Atom-History in Models
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Democritus (450BC, GR): smallest bit that
is still that element.
Joseph Proust (1794 FR): proposed that
compounds form in definite proportions.
Proust’s experiments and law influenced
Dalton.
Dalton (1808 UK): the atom is a solid ball
of the element. Know the five parts of his
theory.
Faraday (1830 UK): the atom has charged
pieces-magnetic field affected the stream
of particles in a cathode ray tube.
JJ Thomson (1903, UK): the atom is
positively charged with negative bits
embedded in it (plum pudding). Cathode
rays are charged particles.
Science Museum, UK
Univ. of Oregon
Atomic Models 20th Century
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Rutherford (1911 NZ): the gold
experiment demonstrated that
most of the atom is empty space.
It is a massive tiny nucleus with
electrons in a cloud.
Niels Bohr (1913 DK): applied
quantum theory of Einstein and
others to atoms.
James Chadwick (1932 UK):
proved the existence of the
neutron with its slightly greater
mass. Nucleus has both protons
and neutrons.
This includes 1932
addition of neutrons
Updated Model
Bohr
Model
How to Know the Number of Electrons
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Subatomic particle
Atoms are composed of protons, neutrons and
electrons.
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The atomic number (Z) is equal to the number of
protons.
The number of protons = number of electrons for
neutral atom.
The charge of the atom or ion is equal to the
number of protons – the number of electrons.
The number of electrons can be determined by
using the atomic number and charge:
number of electrons (negative charges) + ion charge
Energy levels, subshells, and orbitals,
and spin: Quantum Numbers
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Electrons are associated with energy, more specifically
quantum energy, and exemplify wave-like and particle-like
characteristics.
Quantum: the magnitude of a physical phenomena takes on
only discrete values
The electrons of a ground state electron are in the lowest
principal energy levels possible.
Electrons in an orbital that is higher than ground state are
excited electrons.
The electrons exist in specific energy levels, not a continuous
range.
The principal energy level, n, which is numbered 1, 2, 3, 4, …
Excited Electrons
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When excited electrons drop back down to ground state, a photon
of light is emitted.
The drop between two orbits has a specific wavelength or color.
Each element has a unique pattern, a signature.
This is used in mass spectrometry and in astronomy to determine
the elements present in a sample.
The visible spectrum of light and emission line spectra of hydrogen, neon, and iron.
Note that the heavier an element is, the more spectral bands it has.
Subshells
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Electrons are placed in energy levels.
These energy levels are sub-divided into
subshells (labeled s, p, d or f).
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The s subshells is the lowest energy and begins in
level 1.
The p subshells is higher energy and therefore
doesn’t begin until level 2.
The d is higher energy and begins in level 3 and the
f is even higher energy and begins in level 4.
The subshells are further sub-divided into orbitals
(s has 1 orbital, p has 3 orbitals, d has 5 orbitals
and f has 7 orbitals).
Each orbital can hold 2 electrons.
Valence Electrons: Element properties
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The location of valence electrons that are available to
participate in reactions can be predictive of properties
and behavior of an element.
S orbital electrons—group 1 and group 2, the alkali
metals and alkaline earth metals, have valence electrons in
the s orbital.
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The principal quantum number “n” fills the s orbital.
The maximum is two electrons for any one s orbital.
What are the ionic charges for elements
of these groups?
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P orbitals
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The p block on the periodic table includes groups 13
through 18, excepting Helium which does not have
occupied p orbitals.
There is a maximum of six electrons filling the three
orbital shapes with a maximum of two electrons each.
Think of the shape of the orbital as a map of where an
electron is most likely to be, a map of probability. It is not
a route that the particle flies along.
d Orbitals
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These are in groups 3 through 12, transition elements.
The Romans. This diverse group of metals have variable
oxidation states with the exception of Ag, Zn, and usually
Cd.
Five orbital shapes
10 electrons
f Orbitals
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No group number.
The Actinides. These larger atoms exhibit some unusual properties
(Neodymium and Praseodymium)
The Lantanides. These are the heavy elements which are unstable and
radioactive.
Seven orbital shapes
14 electrons
Periodic Table and Orbitals
Aufbau Principle
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Rules for writing electron configurations
The Aufbau principle states that energy levels
must be filled from the lowest to the highest and
you may not move on to the next level unless the
previous level is full. Use the periodic table as a
guide (read left to right):
Aufbau Diagram (Diagonals)
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It is important to keep in
mind that the Aufbau
principle represents an
approximate trend that holds
in most cases. There are
however exceptions to these
rules.
Why is gold, a noble metal,
that is positioned in the
middle of the transition
elements?
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Exceptions tend to be where
shells can be filled (or half-filledone electron in each orbital).
List the Levels and Sublevels in Order
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Start from the lowest energy level to the highest.
Use the periodic table as a guide.
Energy Levels
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Energy increases as electron positions move away from
the nucleus.
Hund’s Rule: or I want my own orbit.
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Hund’s Rule says that when placing electrons in
orbitals of equal energy, place one in each orbital
before doubling up in order to arrive at the lowest
energy configuration.
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The Pauli Exclusion Principle states that when
electrons do share an orbital, they must be of
different “spin.” No two electrons can have the
same quantum numbers: each one has an unique
number address.
The Elements
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Each time an electron is
added it goes to an open
orbital in the same
energy level FIRST.
In larger atoms, both the
s and p may have one
electron only to get a
configuration with one
electron in every orbital.
See Cr and Cu.
This creates a more
balanced geometry.
Periodic Table
Lewis Dot Table (s and p orbitals combined)
s1
s2
p1
p2
p3
p4
p5
p6
Noble: complete shells
Lewis Dot with Levels
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As electrons are added, they fill electron shells in an order determined by which
configuration will give the lowest possible energy. The first shell (n=1) can have only
2 electrons, so that shell is filled in helium, the first noble gas.
In the periodic table, the elements are placed in "periods" and arranged left to right
in the order of filling of electrons in the outer shell. So hydrogen and helium
complete the first period.
Quantum Numbers
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Principal quantum number (n) is the shell (energy level)
that an electron belongs to.
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If an electron is in the lowest possible energy state, then it is in its
ground state.
If an electron is in a higher energy state, then it is in an excited
state.
The total number of orbitals for n = n2. 1st level has 1*1 orbitals. 2nd
level has 2*2 orbitals or 4 total.
n may equal 1, 2, 3, 4, 5, 6, …
Angular momentum (secondary, azimuthal) quantum
number (l) specifies the shape of the orbit or subshell.
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l
0
1
2
3
4
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Letter
s
p
d
f
g
Quantum Numbers
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Magnetic quantum number (ml) specifies the orientation
in space.
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For s orbitals, l is zero. For s orbitals, ml is zero
For p orbitals, l is -1, zero, +1 (orientation to x, y, or z axis)
For d orbitals, l is -2, -1, zero, +1, +2 (five orientations)
For f orbitals, l is -3, -2, -1, zero, +1, +2, +3 (seven orientations)
ml = -l, …, 0, …, l
Spin quantum number (ms) specifies one of two possible
spin orientations within a single orbital.
This is often referred to as up or down.
Pauli Exclusion Principle states that no two electrons can
have exactly the same four values for their quantum numbers.
Two electrons in the same orbital must have opposite spins.
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Quantum Numbers (1st Three)
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n
l
(energy)
1
2
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(shape)
ml
(orientation)
Number of Orbital
orbitals
Name
0
0
1
0
0
1
1 -1, 0, +1
3
3
0
0
1
1 -1, 0, +1
3
2 -2, -1, 0, +1, +2
5
4
0
0
1
1 -1, 0, +1
3
2 -2, -1, 0, +1, +2
5
3 -3, -2, -1, 0, +1, +2, +3 7
Spin (ms)is written as -½ or +½ .
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
Number of
electrons
2
2
6
2
6
10
2
6
10
14
Resources
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http://chemwiki.ucdavis.edu/Inorganic_Chemistry
http://www.colorado.edu/physics/2000/quantumzone/bohr
.html Bohr model animation with photon interactions.