CHAPTER 5
ELECTRONS IN
ATOMS
P. 126
ERNEST RUTHERFORD’S MODEL
Discovered dense +
nucleus
•e-s move like planets
around sun
•Mostly empty space
Didn’t explain chemical
properties of elements
THE BOHR MODEL
•Danish physicist
•Student of
Rutherford
Why don’t
electrons
fall into
nucleus?
“Why don’t
e-s fall into
nucleus”?
Neils Bohr
(1885-1962)
Niels Bohr
THE BOHR MODEL
I pictured the
electrons found in
specific circular
paths around the
nucleus, and can
jump from one
level to another.
Niels Bohr
Furthermore, each
level has a fixed
amount of energy
different from other
levels
BOHR’S MODEL
fixed energy e- have called
Energy levels
 Like rungs of ladder
e- can’t exist btwn energy
levels
energy levels not evenly
spaced
• High levels closer (less energy
needed to jump)
Bohr’s model of the atom 5:17
THE QUANTUM MECHANICAL
MODEL
•
e-’s don’t move like big objects
•
• Rutherford & Bohr model
Energy - “quantized” (in chunks)
•
•
exact energy needed to move e- 1
energy level called a quantum
energy never “in btwn”
• quantum leap in energy must exist
Schrodinger
• Erwin Schrodinger (1926)
mathematically described energy &
position of e- in atom
Quantum Leap TV intro
THE QUANTUM MECHANICAL
MODEL
 energy levels for e-
 Orbits not circular
 Based on probability of
finding e- certain
distance from nucleus
 electron cloud
ATOMIC ORBITALS
• Principal Quantum # (n) - energy
level of e- (1, 2, 3 etc.)
• atomic orbitals - regions of space w/ high
probability of finding e- (not a true “orbit”)
• within each energy level
• Sublevels like rooms in a hotel
• s, p, d, and f
• Different shapes
Max # of e- that fit
in energy level is:
2n
2
How many ein level 2?
level 3?
s and p orbitals 1:20
d orbitals 3:40
atomic orbitals
review (14:28)
ATOMIC ORBITALS
# of
orbitals
(regions of
space)
s
spherical
p
dumbell
d
1
2
1st
3
2nd
5
6
10
7
14
4th
clover leaf
f
complicated
Maximum
electrons
First
possible
energy level
3rd
Summary of Principal Energy Levels, Sublevels, and
Orbitals
Principal Number
energy
of
level
sublevels
n=1
1
n=2
2
n=3
3
n=4
4
Type of
sublevel
1s (1 orbital)
2s (1 orbital
2p (3 orbitals)
3s (1 orbital)
3p (3 orbitals)
3d (5 orbitals)
4s (1 orbital)
4p (3 orbitals)
4d (5 orbitals)
4f (7 orbitals)
Max # of
Electron
electrons configuration
2
8
18
32
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
ORDER OF ELECTRON SUBSHELL
FILLING:
NOT “IN ORDER”
 Lowest
energy fill
first
1s2
2s2 2p6
3s2 3p6 3d10
4s2 4p6 4d10 4f14
5s2 5p6 5d10 5f14
6s2 6p6 6d10
7s2 7p6
Increasing energy
 energy
levels
overlap
1s2 2s2 2p6 3s2 3p64s2 3d10 4p65s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
ELECTRON
CONFIGURATION
1
1s
Principal energy level
# valence es: 1 or 2
p: 1-6
row #
d: 1-10
1-7
7 rows
group #
f: 1-14
sublevel
s, p, d, or f
4 sublevels
Total e- = Atomic #
period # = # e- energy levels
SUBLEVELS D AND F ARE
“SPECIAL”
1A
1
2
3
4
5
6
7
2A
group # = # valence e-
3A 4A 5A6A 7A
8B
3B 4B5B 6B 7B
3d
4d
1B 2B
d
5d
6d
6
7
4f
5f
f
8A
Increasing energy
SECTION 5.2 ELECTRON ARRANGEMENT
IN ATOMS P. 133
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
6d
5d
4d
3d
aufbau diagram page 133
2s
Aufbau - German for “building up”
1s
5f
4f
ELECTRON CONFIGURATIONS…
….3 rules explain how e-’s fill their
orbitals:
1) Aufbau principle – e-’s enter lowest
energy level first.
2) Pauli Exclusion Principle - 2 e-’s
max/orbital (hotel room) - different
spins
PAULI EXCLUSION PRINCIPLE
No 2 electrons in an atom
can have the same four
quantum numbers.
To show different
direction of spin, a pair
in the same orbital is
written as:
Wolfgang Pauli
QUANTUM NUMBERS
Each e- has unique set of 4 quantum
#’s describing it
1)
2)
3)
4)
Principal quantum #
Angular momentum quantum #
Magnetic quantum #
Spin quantum #
ELECTRON CONFIGURATIONS
3) Hund’s Rule- When e-’s occupy
orbitals of same energy, they
won’t pair up until they must
write e- configuration for
Phosphorus
 all 15 e-’s must be accounted
for
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p The first 2 e-’s go into
the 1s orbital
2p
Notice opposite
direction of spins
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2s
1s
2p The next e-’s go in 2s
orbital
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
• The next e-’s go in
2p orbital
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
• The next e-’s go in
3s orbital
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The last 3 e-’s go in 3p
orbitals
2p They each go into
separate shapes (Hund’s)
• 3 unpaired e-’s
Orbital
notation
= 1s22s22p63s23p3
An internet program about electron
configurations is:
Electron Configurations
I
electron config (song) 3:24
FILLING ORBITALS
Lowest  higher energy
Adding e-’s changes energy of
orbital
•Full orbitals best situation
•half filled orbitals next best
• more stable
• Changes filling order
WRITE THE ELECTRON
CONFIGURATIONS FOR THESE
ELEMENTS:
Titanium - 22 electrons
2 2 6 2 6 2 2
 1s 2s 2p 3s 3p 4s 3d
Vanadium - 23 electrons
2 2 6 2 6 2 3
 1s 2s 2p 3s 3p 4s 3d
Chromium - 24 electrons
2 2 6 2 6 2 4 (expected)
 1s 2s 2p 3s 3p 4s 3d
But this is not what happens!!
CHROMIUM IS ACTUALLY:
1s22s22p63s23p64s13d5
Why?
2 half filled orbitals
•Half full slightly lower in energy
•Same applies to copper
COPPER’S E- CONFIGURATION
• Copper has 29 e-s so expect:
1s22s22p63s23p63d94s2
• actual configuration is:
1s22s22p63s23p63d104s1
• 1 more full orbital & 1 half filled
• Exceptions
• d4
• d9
IRREGULAR CONFIGURATIONS OF
CHROMIUM AND COPPER
Chromium steals a 4s e- to make its
3d sublevel HALF FULL
Copper steals a 4s electron
to FILL its 3d sublevel
SECTION 5.3 PHYSICS AND THE QUANTUM
Light
MECHANICAL MODEL
P. 138
• Study of light led to quantum mechanical
model
• Light is electromagnetic radiation
• EM radiation: gamma rays, x-rays, radio
waves, microwaves
• Speed of light = 2.998 x 108 m/s
• “c” - celeritas (Latin for speed)
• All EM radiation travels same in vacuum
- Page 139
“R O Y
Frequency Increases
Wavelength Longer
G
B
I V”
PARTS OF A WAVE
Crest
Wavelength
Amplitude
Trough
ELECTROMAGNETIC RADIATION
PROPAGATES THROUGH SPACE AS A
WAVE MOVING AT THE SPEED OF LIGHT.
Equation:
c =
c = is a constant (2.998 x 108 m/s)
 (lambda) = wavelength, in meters
 (nu) = frequency, in units of hertz (hz or sec-1)
WAVELENGTH AND FREQUENCY
• inversely related
• one gets bigger, other smaller
• Different frequencies = different
colors
• wide range of frequencies
(spectrum)
- Page 140
Use Equation: c =
Low
Energy
High
Energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Visible Light Wavelength
Long  =
Low Frequency
=
Low ENERGY
Short  =
High Frequency
=
High ENERGY
ATOMIC SPECTRA
White light all
colors of visible
spectrum
• prism
separates it
according to
λ
IF THE LIGHT IS NOT WHITE
heating gas with
electricity will emit
colors
• this light thru prism
is different
ATOMIC SPECTRUM
elements emit
own characteristic
colors
• composition of
stars determined
thru spectral
analysis
• atomic emission
spectrum
• Unique to each
element, like
fingerprints!
• ID’s elements
LIGHT IS A PARTICLE?
Energy is quantized
Light is energy…..
light must be quantized
photons smallest pieces of light
Photoelectric effect –
• Matter emits e- when it absorbs energy
• Albert Einstein Nobel Prize in chem
Energy & frequency: directly related
ENERGY (E ) OF
ELECTROMAGNETIC RADIATION
DIRECTLY PROPORTIONAL TO
FREQUENCY () OF RADIATION.
Planck-Einstein Equation:
E = h
E = Energy, in units of Joules (kg·m2/s2)
(Joule…metric unit of energy)
h = Planck’s constant (6.626 x 10-34 J·s)
(reflecting sizes of energy quanta)
 = frequency, units of hertz (hz, sec-1)
THE MATH IN CHAPTER 5
There are 2 equations:
1) c = 
2) E = h
Put these on your 3 x 5
notecard!
EXAMPLES
1) What is the wavelength of
blue light with a frequency of
8.3 x 1015 hz?
2) What is the frequency of red
light with a wavelength of 4.2
x 10-5 m?
3) What is the energy of a
photon of each of the above?
EXPLANATION OF
ATOMIC SPECTRA
electron configurations written in
lowest energy.
energy level, and where electron
starts from, called it’s ground
state - lowest energy level.
CHANGING THE ENERGY
Let’s look at a hydrogen atom, with
only one electron, and in the first
energy level.
Changing the energy
Heat, electricity, or light can move e-’
up to different energy levels. The
electron is now said to be “excited”
Changing the energy
As electron falls back to ground state,
it gives energy back as light
Experiment #6, page 49-
Changing the energy
may fall down in specific steps
Each step has different energy
Lyman series (UV)
Balmer series
(visible)
Paschen series
(infrared)
Ultraviolet
Visible
Infrared
further they fall, more energy
released = higher frequency
orbitals also have different energies
inside energy levels
All electrons can move around.
WHAT IS LIGHT?
Light is a particle - it comes in chunks.
Light is a wave - we can measure its
wavelength and it behaves as a wave
combine E=mc2 , c=, E = 1/2 mv2 and E
= h, then we can get:
= h/mv
(from Louis de Broglie)
Calculates wavelength of a particle.
called de Broglie’s equation
• He said particles exhibit properties of waves
WAVE-PARTICLE DUALITY
J.J. Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The
electron is
a particle!
The electron
is an energy
wave!
CONFUSED? YOU’VE GOT
COMPANY!
“No familiar conceptions can be
woven around the electron;
something unknown is doing we
don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
THE PHYSICS OF THE VERY
SMALL
Quantum mechanics explains
how very small particles behave
•Quantum mechanics is an
explanation for subatomic
particles and atoms as waves
Classical mechanics describes
the motions of bodies much
larger than atoms
HEISENBERG UNCERTAINTY
PRINCIPLE
impossible to know exact location
and velocity of particle
better we know one, less we know
other
Measuring changes properties.
True in quantum mechanics, but
not classical mechanics
HEISENBERG UNCERTAINTY
PRINCIPLE
“One cannot simultaneously
determine both the position
and momentum of an
electron.”
Werner Heisenberg
You can find out where the
electron is, but not where it is
going.
OR…
You can find out where the
electron is going, but not where
it is!
IT IS MORE OBVIOUS WITH
THE VERY SMALL OBJECTS
To measure where e-, we use
light
But light energy (photon)
moves e- due to small mass
And hitting e- changes
frequency of light
After
Before
Photon
Moving
Electron
Photon
wavelength
changes
Electron
velocity changes
Fig. 5.16, p. 145