Electrons in Atoms
The modern view of the atom was
Ernest
Rutherford (1871-1937).
developed by
Screen 2.9
Ernest
Rutherford
of New Zealand
(1871-1937).
Neils Bohr and the Atom
•Proposed electrons are found in concentric orbits
around the nucleus call energy levels
•Originated the idea of quantum energy
Atomic Line Spectra and
Niels Bohr
Niels Bohr
(1885-1962)
Bohr’s theory was a great
accomplishment.
Rec’d Nobel Prize, 1922
Problems with theory —
 theory only successful for
H.
 introduced quantum idea
artificially.
 So, we go on to QUANTUM
or WAVE MECHANICS
ATOMIC ELECTRON
CONFIGURATIONS AND
PERIODICITY
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
Principle Energy Levels (n)
SUBLEVELS (l)
ORBITALS (ml)
Principle Energy Levels
n=1
n=2
n=3
n=4
Assigning Electrons to Atoms
Electrons
generally assigned to
orbitals of successively higher
energy. Aufbau Principle
Sublevel Filling of Electrons
Electrons fill orbitals from the bottom up: Aufbau Principle
Orbitals
Pauli Exclusion Principle
 No more than 2 e- assigned to an orbital
 Orbitals grouped in s, p, d (and f) sublevels

s sublevels
d sublevels
p sublevels
s orbitals
p orbitals
d orbitals
s orbitals
p orbitals
d orbitals
No.
orbs.
1
3
5
No.
e-
2
6
10
s Orbitals
All s orbitals are
spherical in
shape.
1s Orbital
2s Orbital
3s Orbital
p Orbitals
 The
three p orbitals lie 90o
apart in space
Types of Atomic
Orbitals
Figure 7.15, page 275
Writing Atomic Electron
Configurations
Two ways of
writing configs.
One is called
the spdf
notation.
spdf notation
for H, atomic number = 1
1
1s
value of n
no. of
electrons
value of l
Writing Atomic Electron
Configurations
Two ways of
writing
configs.
Other is
called the
orbital box
notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
2
1s
1s
Arrows
depict
electron
spin
Electro
n
Filling
Order
Figure 8.5
See “Toolbox” for Electron Configuration tool.
Electron Configurations
and the Periodic Table
Figure 8.7
Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total
electrons
3p
3s
2p
2s
1s
Beryllium
3p
3s
2p
2s
1s
Group 2A
Atomic number = 4
1s22s2 ---> 4 total
electrons
Boron
3p
3s
2p
2s
1s
Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons
Carbon
3p
3s
2p
2s
1s
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
Here we see for the first time
HUND’S RULE. When
placing electrons in a set of
orbitals having the same
energy, we place them singly
as long as possible.
Nitrogen
3p
3s
2p
2s
1s
Group 5A
Atomic number = 7
1s2 2s2 2p3 --->
7 total electrons
Oxygen
3p
3s
2p
2s
1s
Group 6A
Atomic number = 8
1s2 2s2 2p4 --->
8 total electrons
Fluorine
3p
3s
2p
2s
1s
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
9 total
electrons
Neon
3p
3s
2p
2s
1s
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
Note that we have
reached the end of the
2nd period, and the 2nd
shell is full!
Sodium
Group 1A
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Note that we have begun a new
period.
All Group 1A elements have
[core]ns1 configurations.
Electron Configurations
of p-Block Elements
Aluminum
Group 3A
Atomic number = 13
1s2 2s2 2p6 3s2 3p1
[Ne] 3s2 3p1
All Group 3A elements
have [core] ns2 np1
configurations where n
is the period number.
3p
3s
2p
2s
1s
Phosphorus
Group 5A
Atomic number = 15
1s2 2s2 2p6 3s2 3p3
[Ne] 3s2 3p3
All Group 5A elements
have [core ] ns2 np3
configurations where n
is the period number.
3p
3s
2p
2s
1s
Calcium
Group 2A
Atomic number = 20
1s2 2s2 2p6 3s2 3p6 4s2
[Ar] 4s2
All Group 2A elements have
[core]ns2 configurations where n
is the period number.
Electron Configurations
and the Periodic Table
Transition Metals
Table 8.4
All 4th period elements have
the configuration [argon] nsx
(n - 1)dy and so are “d-block”
elements.
Chromium
Iron
Copper
Transition Element
Configurations
3d orbitals used
for Sc-Zn (Table
8.4)
Lanthanides and Actinides
All these elements have the
configuration [core] nsx (n - 1)dy (n 2)fz and so are “f-block” elements.
Cerium
[Xe] 6s2 5d1 4f1
Uranium
[Rn] 7s2 6d1 5f3
Lanthanide Element
Configurations
4f orbitals used for
Ce - Lu and 5f for
Th - Lr (Table 8.2)
Arrangement of
Electrons in Atoms
Each orbital can be assigned no more than
2 electrons! Pauli Exclusion Principle
This is tied to the existence of a 4th
electron
spin quantum number, ms.
quantum number, the
Electron
Spin
Quantum
Number,
ms
Can be proved experimentally that electron
has a spin. Hunds Rule
Electron Spin Quantum
Number
Diamagnetic: NOT attracted to a magnetic
field
Paramagnetic: substance is attracted to a
magnetic field. Substance has unpaired
electrons.
Assigning Electrons to
Subshells
 In H atom all
subshells of same n
have same energy.
 In many-electron
atom:
a) subshells increase in
energy as value of n
+ l increases.
b) for subshells of
same n + l, subshell
with lower n is lower
in energy.
2px
Orbital
3px
Orbital
d Orbitals
s orbitals have no planar
node (l = 0) and so are
spherical.
p orbitals have l = 1, and
have 1 planar node,
and so are “dumbbell”
shaped.
This means d orbitals
(with l = 2) have
2 planar nodes
See Figure 7.16
3dxy Orbital
3dxz Orbital
3dyz Orbital
3dx
2
-y
2
Orbital
3dz Orbital
2
Orbital Filling: The Aufbau
Principle & Hund’s Rule
Aufbau Principle: Lower energy orbitals fill first.
Hund’s Rule:
Degenerate orbitals (those of the same energy) are filled with
electrons until all are half filled before pairing up of electrons
can occur.
Pauli exclusion principle:
Individual orbitals only hold two electrons, and each should
have different spin.
“s” orbitals can hold 2 electrons
“p” orbitals hold up to 6 electrons
“d” orbitals can hold up to 10
electrons
Orbital Filling: The Aufbau
Principle & Hund’s Rule
Aufbau Principle: Lower energy orbitals fill first.
Hund’s Rule:
Degenerate orbitals (those of the same energy) are filled with
electrons until all are half filled before pairing up of electrons
can occur.
Pauli exclusion principle:
Individual orbitals only hold two electrons, and each should
have different spin.
“s” orbitals can hold 2 electrons
“p” orbitals hold up to 6 electrons
“d” orbitals can hold up to 10
electrons
Pauli Exclusion Principle
No two electrons in the
same atom can have
the same set of 4
quantum numbers.
That is, each electron has a
unique address.