Ms. Cleary
Chem 11
A model
 A representation or explanation of a reality that is so
accurate and complete that it allows the model builder
to predict events.
 Scientific Method leads to model building
 Gather data, develop a model, formulate a hypothesis,
test and modify the model.
Bohr’s Model
 Why don’t the electrons fall into the
nucleus?
 Move like planets around the sun.
 In circular orbits at different levels.
 Amounts of energy separate one level
from another.
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr postulated that:




Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away
from the nucleus
An atom with maximum number of electrons in the
outermost orbital energy level is stable (unreactive)
How did he develop his theory?
 He used mathematics to explain the
visible spectrum of hydrogen gas
 http://www.mhhe.com/physsci/chemistr
y/essentialchemistry/flash/linesp16.swf
Energy and Visible Light
High
Low
energy
energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Wavelength
Visible Light
(700 nm)
(400 nm)
The line spectrum
 electricity passed
through a gaseous
element emits light
at a certain
wavelength
 Can be seen when
passed through a
prism
 Every gas has a
unique pattern
(color)
Line spectrum of various elements
Bohr’s Triumph
 His theory helped to explain periodic law
 Halogens are so reactive because it has one e- less than
a full outer orbital
 Alkali metals are also reactive because they have only
one e- in outer orbital
Drawback
 Bohr’s theory did
not explain or show
the shape or the
path traveled by the
electrons.
 His theory could
only explain
hydrogen and not
the more complex
atoms
Increasing energy
Fifth
Fourth
Third
Second
First
 Further away
from the
nucleus means
more energy.
 There is no “in
between”
energy
 Energy Levels
Complete Bohr Diagrams for
the Following:
 Mg
 Li
 Ne
F
The Quantum Mechanical Model
 Energy is quantized. It comes in chunks.
 A quanta is the amount of energy needed to move
from one energy level to another.
 Since the energy of an atom is never “in between”
there must be a quantum leap in energy.
 Schrödinger derived an equation that described
the energy and position of the electrons in an atom
Atomic Orbitals
 Principal Quantum Number (n) = the energy
level of the electron.
 Within each energy level the complex math of
Schrödinger's equation describes several
shapes.
 These are called atomic orbitals
 Regions where there is a high probability of
finding an electron
Orbitals
 Electrons spin around the nucleus creating an electron
cloud.
 The electron clouds come in 4 different shapes, called
orbitals.
 The four orbitals are called s, p, d, and f.
 Each orbital is capable of holding different
numbers of electrons:
Orbital
# of Electrons
s
2
p
6
d
10
f
14
S orbitals
 1 s orbital for
every energy level
1s 2s
 Spherical shaped
 Each s orbital can hold 2 electrons
 Called the 1s, 2s, 3s, etc.. orbitals
3s
P orbitals
 Start at the second energy level
 3 different directions
 3 different shapes
 Each orbital can hold 2 electrons
 The p Sublevel has 3
p orbitals
The D sublevel contains 5 D rdorbitals
 The D sublevel starts in the 3 energy level
 5 different shapes (orbitals)
 Each orbital can hold 2 electrons
The F sublevel has 7 F orbitals
 The F sublevel starts in the fourth energy level
 The F sublevel has seven different shapes (orbitals)
 2 electrons per orbital
Summary
Starts at
energy
level
Sublevel
# of shapes
(orbitals)
Max # of
electrons
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
Electron Configurations
 The way electrons are arranged in atoms.
 Aufbau principle- electrons enter the lowest energy
first.
 This causes difficulties because of the overlap of
orbitals of different energies.
 Pauli Exclusion Principle- at most 2 electrons per
orbital - different spins
Electron Configurations
First Energy Level
 only s sublevel (1 s orbital)
 only 2 electrons
 1s2
Second Energy Level
 s and p sublevels (s and p orbitals are available)
 2 in s, 6 in p
 2s22p6
 8 total electrons
Levels
Third energy level
• s, p, and d orbitals
• 2 in s, 6 in p, and 10 in d
• 3s23p63d10
• 18 total electrons
Fourth energy level
• s,p,d, and f orbitals
• 2 in s, 6 in p, 10 in d, and 14 in f
• 4s24p64d104f14
• 32 total electrons
Electron Configurations
 Electron configurations are a shorthand for writing
exactly what was in the energy level diagrams.
 Electron configuration for O is:
1s22s22p4
# of electrons
period
orbital
Electron configuration for Ar is:
1s22s22p63s23p6
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
Electron Configuration
 Hund’s Rule- When electrons occupy orbitals of equal
energy they don’t pair up until they have to .
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p  The first to electrons go
into the 1s orbital
2p
 Notice the opposite
spins
 only 13 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p  The next electrons go
into the 2s orbital
2p
 only 11 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
Orbitals fill in order
 Lowest energy to higher energy.
 Adding electrons can change the energy of the orbital.
 Half filled orbitals have a lower energy.
 Makes them more stable.
 Changes the filling order
Write these electron
configurations
 Titanium - 22 electrons
 1s22s22p63s23p64s23d2
 Vanadium - 23 electrons 1s22s22p63s23p64s23d3
 Chromium - 24 electrons
 1s22s22p63s23p64s23d4 is expected
 But this is wrong!!
Chromium is actually
 1s22s22p63s23p64s13d5
 Why?
 This gives us two half filled orbitals.
 Slightly lower in energy.
 The same principal applies to copper.
Copper’s electron
configuration






Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.
Remember these exceptions
Electron Configuration and the
Periodic Table
 Groups 1 and 2 represent the s orbital
 Groups 13-18 represent the p orbital
 Groups 3-12 represent the d orbital
 Lanthanides and Actinides represent f orbital
Practice
Time to practice: Draw the following energy level
diagrams on your own filling up electron
configurations: H, He, Be, N, Na, Ni, Br,
2. Do electron configurations for the elements listed in
#1.
1.