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Chapter 5
Chapter 5
Quantum Theory and Electron Configurations
It’s all about color…
• In terms of atomic models, so far:
 Dalton (1803) = Tiny, solid particle
 Thomson (1897) = “Plum Pudding” model – Electrons
stuck on the outside of a big positive charge
 Rutherford (1911) = Positively-charged nucleus with
electrons moving around it
• Rutherford’s model of the atom not quite right
 Could not explain chemical properties of elements
 Could not explain color changes when metal is heated
Bohr Model of the Atom
• Niels Bohr’s model of the atom
 Electron found only on specific, circular paths around
nucleus
 Each orbit has fixed energy level
• Hypothesis: When electrons are excited (added
energy), jump into higher energy levels. When
they moved back into lower energy levels - gave
off light.
• Electrons do not exist between levels (think of
rungs on a ladder)
 Electrons absorb and emit only certain quanta
(amounts) of energy
 Quantum of energy = fixed amount of energy
required to move from one energy level to another
energy level
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Chapter 5
Bohr’s Planetary Model of the Atom
• Electrons must have enough energy
to keep moving around the nucleus
• Electrons orbit nucleus in defined
energy levels, just like planets orbit
the sun
• Each energy level assigned a
principal quantum number n.
• Lowest energy level called ground
state (n=1)
• Higher energy levels (n=2, 3, 4...)
excited states
• Model worked OK for hydrogen but
not so good for other elements
Nucleus
n=1
n=2
Bohr’s Model
• Further away from the
nucleus means more
energy.
• There is no “in
between” energy
• Energy Levels
Increasing energy
Fifth
Fourth
Third
Second
First
Nucleus
Chapter 5
Electron starts on
lowest energy level
(ground state)
Lowest energy level = ground state
Higher energy levels = excited states
Add energy
to electron –
moves to
excited state
Energy
levels are
not evenly
spaced
Energy
Level 3
Energy
Level 2
Energy
Level 1
Nucleus
Electron starts
on lowest
energy level
(ground state)
Lowest energy level = ground state
Higher energy levels = excited states
Add energy
to electron –
moves to
excited state
Energy
Energy
levels are
not evenly
spaced
Electron
returns to
lower state –
emits/gives off
quantum of
energy
Energy
Level 3
Energy
Level 2
Energy
Level 1
Nucleus
Bohr used this
theory to explain
the lines in the
atomic emission
spectra for
hydrogen
Chapter 5
Each of these lines
corresponds to different
energy changes
434 nm
410 nm
656 nm
486 nm
Chapter 5
Chapter 5
Chem I - Mon, 9/22/15
Do Now
 Get to work on the PEN worksheet
from last class



Homework
 MEAL paragraph if not finished
Agenda
History
Intro to quantum
Electron Config
Quantum-Mechanical Model of the
Atom
• Since the Bohr model had a very limited
use, a new and very different model of
the atom exists
• The Quantum Mechanical Model
(1926) contains:
 Quantum energy levels
 Dual wave/particle nature of electrons
 Electron clouds
• In the new model, don’t know exactly
where electrons are - only know
probabilities of where they could be
Heisenberg Uncertainty Principle
• Heisenberg Uncertainty Principle =
impossible to know both the velocity (or
momentum) and position of an electron at the
same time
Quantum-Mechanical Model of the
Atom
• Orbital = region around nucleus where an
electron with a given energy level will
probably (90%) be found
• Four kinds of orbitals
s - spherical in shape, lowest orbital for every
energy level
p - dumbbell shaped, second orbital
d - complex “flower” shape, third orbital
f - very complex shape, highest orbital
s-orbitals
• All s-orbitals are spherical.
• As n increases, the s-orbitals get larger.
p- orbitals
• Three p-orbitals: px, py, and pz
 Lie along the x-, y- and z- axes of a Cartesian
system.
 Dumbbell shaped, gets larger as n increases
d and f - orbitals
• There are five d and seven
f-orbitals.
Quantum Mechanical Model
• Principle Energy Levels (n)
 Labeled from 1-7
 First energy level is n=1
 Contains sublevels (s, p, d and f)
• Each energy level contains the number of sublevels equal to its value
for n
– If n=3, there are three sublevels
Chapter 5
Quantum Mechanical Model
• In each sublevel there are atomic orbitals
• Atomic orbitals – describe a space where an electron is
likely to be found
Type of
subshell
Shape of
orbitals
Number of
orbitals
Orbital
‘names’
s
Spherical
1
s
p
Dumbbell
3
px, py, pz
d
Cloverleaf
(and one
donut)
5
f
Multi-lobed
7
Quantum Mechanical Model
• Each orbital can contain two electrons.
• Since negative-negative repel, these electrons occupy
the orbital with opposite spins.
Quantum Mechanical Model
• The total number of orbitals of an energy level is n2.
 For the third principle energy level, n=3, which means there are
9 orbitals
• These orbitals are 3s, 3px, 3py, 3pz and the 5 d orbitals
• Remember, we no longer think of orbitals as concentric
circles, but we can say that n=4 extends farther from the
nucleus than n=1.
Valence Electrons
• Only those electrons in the highest principle energy
level
Electron Configuration and Orbital Notation
• Aufbau Principle – electrons fill lower energy orbitals first,
“bottom-up”
 n=1 fills before n=3
Energ
y
• Will an electron fill the 1s or the 2s orbital first?
2s
1s
2px
2py
2pz
Electron Configuration &Orbital Notation
• Hund’s Rule – electrons enter same energy orbitals so that
each orbital has one electron before doubling up
 Each of the first electrons to enter the equal energy orbitals must
have the same spin
Energ
y
 If we have 7 electrons, how will they fill in the below orbitals?
2s
1s
2px
2py
2pz
Electron Configuration and Orbital Notation
• Pauli Exclusion Principle – an orbital can contain no more
than 2 electrons. Electrons in the same orbital must have
different spins.
Energ
y
• If we have 8 electrons, how will they be arranged?
2s
1s
2px
2py
2pz
Apartment Analogy
•
•
•
•
Atom is the building
Floors are energy levels
Rooms are orbitals
Only two people per room
Orbital Diagrams
• Draw each orbital as a box.
• Each electron is represented using an arrow.
 Up arrows – clockwise spin
 Down arrows – counter-clockwise spin
• Determine the total number of electrons involved.
• Start with the lowest energy level (1s) and start filling in
the boxes according the rules we just learned.
Orbital Diagram
4p
3d
Energy
4s
3p
3s
2p
2s
1s
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3d
4s
3p
3s
2p
2s
1s
Chapter 5
5f
4f
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3d
4s
3p
3s
2p
2s
5f
• The first to electrons go into the
1s orbital
• Notice the opposite spins
• only 13 more
1s
Chapter 5
4f
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2p
4f
3d
4s
3p
5f
• The next electrons go into the 2s
orbital
• only 11 more
2s
1s
Chapter 5
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3d
4s
3s
2s
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
1s
Chapter 5
5f
4f
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3d
4s
3s
2s
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
1s
Chapter 5
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
Chapter 5
Orbital Diagrams
• Orbital diagrams are used
to show placement of
electrons in orbitals.
• Need to follow three
rules (Aufbau, Pauli,
Hund’s) to complete
diagrams
Li
Be
B
C
N
Ne
Na
Orbitals and Energy Levels
Principal
Sublevels
Energy Level
Orbitals
n=1
1s
1s (one)
n=2
2s , 2p
2s (one) + 2p (three)
n=3
3s , 3p , 3d
3s (one) + 3p (three) + 3d (five)
n=4
4s, 4p, 4d, 4f
4s (one) + 4p (three) + 4d (five)
+ 4f (seven)
Chapter 5
Summary
shapes Max
electrons
s
p
d
f
Chapter 5
Starts at
energy level
Orbitals and Energy Levels
n=4
Increasing energy
n=3
n=2
n=1
1s
2p
2s
3d
3p
3s
4f
4d
4p
4s
and so on....
Electron Configuration
• Let’s determine the electron configuration for
Phosphorus
• Need to account for 15 electrons
Chapter 5
Writing Electron Configuration
• Determine the total number of electrons.
• Write the principle energy level number as a coefficient,
the letter for the subshell, and an exponent to represent
the number of electrons in the subshell.
• He: 1s2
The Kernel (Noble Gas) Notation
• Determine the total number of electrons
• Find the previous noble gas and put its symbol in brackets
• Write the configuration from that noble gas forward as
usual
Writing electron configurations
• Examples
•
•
•
O
Ti
Br
•
Core format
•
•
•
O
Ti
Br
1s2 2s2 2p4
1s2 2s2 2p6 3s2 3p6 3d2 4s2
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5
[He]
[Ar]
[Ar]
2s2 2p4
3 d2 4s 2
3d10 4s2 4p5
Chapter 5