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Ch 11 - Electron Configuration
Radiant Energy
Waves
• Light travels as both
•
Waves and Packets of
energy.
• These packets are called
photons.
Light is a form of
Electromagnetic
Radiation.
– EM Radiation
has waves in the
electric and
magnetic fields
Electromagnetic waves
• Electromagnetic waves have two basic parts.
1. electric field
2. magnetic field
• The fields are
perpendicular to each
other.
Waves
• All waves (Water or
Electromagnetic) have 4 key
characteristics:
– Amplitude
– Wavelength
– Period
– Frequency
Wave Characteristics
• Amplitude.
– Height of a wave from origin
to a peak/crest.
– Affects brightness and
intensity.
• Wavelength.
– Distance from crest to crest.
Distance for one full cycle.
– Visible light: 400-750nm.
Wave Characteristics
• Period
– time that it takes to complete
a full cycle.
– Measured in seconds
• Frequency.
– number of cycles per second.
– Measured in hertz(Hz)
– High frequency = high
energy
Wave Characteristics
• Speed of light
– Speed of light a
constant:
3.00 X 108 m/s.
• Frequency and Wavelength
related by the equation: 
=c/
Wavelength and frequency
You can also find the frequency by
rearranging the equation:  = c / 
• First, multiply by  (frequency):
   = (c/)  
• Now, divide by  (wavelength):
() /  = c / 
• Leaving:  = c/
Moving on….
Try this…
Remember:  = c / 
1) If the frequency of a wave is 93.1 x 106 ,
what is the wavelength?
2) If the wavelength of a wave is 1.54 m, what
is it’s frequency?
Try this…
1. Split up into groups of 3.
2. Draw one wave with wavelength 4 cm
3. Draw second wave with wavelength twice
of the first wave.
4. Draw a third wave with 3 times the
wavelength the first wave.
5. Draw 3 more waves with the same
wavelengths as the first set but with an
amplitude of 6 cm.
Try this…
1. Order the waves from lowest to highest
frequency.
2. Order the waves from lowest to highest
energy.
3. Order the waves from lowest to highest
amplitude.
Electromagnetic Spectrum
• Many parts including:
–
–
–
–
–
–
–
Gamma Rays (10-11 m)
X-Rays (10-9 m)
Ultra-violet (10-8 m)
Visible (10-7 m)
Infared (10-6 m)
Microwave (10-2 m)
TV/Radio (10-1 m)
Electromagnetic Spectrum
• Visible Spectrum:
ROY G BIV
– Red
– Orange
– Yellow
– Green
– Blue
– Indigo
– Violet
Electromagnetic Spectrum
(once more)
Wavelength practice…
Remember:  = c / 
For the following questions assume that the speed of light is =
3000m/s
1) If the frequency of a wave is 847 Hz, what is the
wavelength?
2) If the frequency of a wave is 4,985 Hz, what is the
wavelength?
3) If the frequency of a wave is 290 Hz, what is the
wavelength?
4) If the frequency of a wave is 38,759 Hz, what is the
wavelength?
Wavelength practice…
Remember:  = c / 
For the following questions assume that the speed of light is =
3000m/s
1) If the wavelength of a wave is 1.54 m, what is it’s
frequency?
2) If the wavelength of a wave is .875 m, what is it’s
frequency?
3) If the wavelength of a wave is 3.39 m, what is it’s
frequency?
4) If the wavelength of a wave is .657 m, what is it’s
frequency?
Electron Configuration
Quantum Theory
Early Puzzlements
• Wave model for light
was originally
accepted by scientific
community.
• This couldn’t explain
why metals heating
first emitted invisible
radiation and then
visible radiation.
• Other questions
included why elements
only emitted certain
characteristic colors of
light.
Terminology
• Ground State – when an atom is at the
lowest possible energy state.
• Excited State – when an atom has excess
energy
Line Spectra
• Def: A spectrum that
contains only certain
colors/wavelengths.
• AKA: The Atomic
Emission Spectrum
• Each element has it’s
own “fingerprint”
emission spectrum.
Line Spectra
• Assume you “energize”
some H atoms.
• There are only certain types
of photons emitted.
• We see only selected colors
that correspond to these
photons energy levels.
Line Spectra
• Each photon has a
frequency that is
proportional to the change
in energy of the electron.
Plank’s Theory
• Every object can only
absorb or emit a
fundamental amount
of energy.
• This amount is called
a quantum.
• The amount is like
moving up or down
steps.
Plank’s Theory
• Plank’s Theory is based
on the relationship
between frequency and
the energy of the particle.
• Energy = h x frequency
• E = h
• Plank’s Constant:
– h = 6.6262 X 10-34 J-s
Dual Nature of Radiant Energy
• Proven in 1923 by
Arthur Compton
– Showed photon could
collide with an electron
like tiny balls.
• Summary:
– Light behaves as a
wave ( = c/)
– Light behaves as a
particle (E = h)
Electron Configuration
Another Look at the Atom
The Bohr Model
• Bohr drew the connection
between Rutherford's
model of the atom and
Planks idea of
quantization.
• Energy levels labeled with
Quantum Numbers (n)
• Ground state, or lowest
energy level – n=1
• Excited State – level of
higher energy
Matter Waves
• If energy has dual nature, why not matter?
• De Broglie thought so.
– Matter Waves – the wavelike behavior of waves.
– Didn’t stand without experimental proof
• Davison and Germer proved this with experiments
in 1927.
• Why don’t we see these matter waves? Mass must
be very small to observe wavelength.
Heisenberg Uncertainty
• Uncertainty Principle
– The position and momentum of a moving
object cannot simultaneously be measured and
known exactly.
• Translation:
– Cannot know exactly where and how fast an
electron is moving at the same time.
Electron Configurations
A New Approach to the Atom
Quantum-mechanics Model
• Includes all the ideas of the atom we have
covered:
– Energy of electrons is quantized
– Electrons exhibit wavelike behavior
– Electrons position and momentum cannot be
simultaneously known
– Model does describe the probable location of
electrons around the nucleus
Probability and Orbitals
• Electron Density:
– The density of an
electron cloud.
• Atomic Orbitals:
– A region around the
nucleus of an atom
where an electron with
a given energy is likely
to be found.
• Kinds of orbitals:
– Each kind has own
different basic shape.
– Given letter
designations of s, p, d
and f.
– s-orbitals are spherical
– p-orbitals are dumbbell
– d- and f-orbitals more
complex.
Probability and Orbitals
• Electron Density:
– The density of an electron cloud.
• Atomic Orbitals:
– A region around the nucleus of an atom where an
electron with a given energy is likely to be found.
– Orbitals are nothing like orbits of a planet!
Probability and Orbitals
• Lets look at the cool animation…
Orbitlas and Energy
• Principle energy levels
(n) can be divided into
sublevels.
• Number of sublevels
is equal to the number
of the principle energy
level.
Orbitals and Energy
• Each sublevel has one
or more orbitals
–
–
–
–
s – one
P – three
d – five
f – seven
• Summary provided on
pg 372-374
Electron Spin
• Electrons have two spins:
– Up or clockwise
– Down or counterclockwise
• Only two electrons (one of each spin) can
occupy an orbital. These electrons are said
to be “paired”.
Electron Configurations
Electron Configurations
Electron Configurations
• Notation which shows how the electrons are
distributed among the various atomic orbital
and energy levels.
1s2
How it works
2
1s
• “1” refers to the principle quantum number
“n=1”
• This “n=1” stands for the energy level
• The electrons occupy the first energy level
of the atom
How it works
2
1s
• “s” refers to the angular momentum.
• Tells us electrons occupy an “s” or spherical
orbital
How it works
2
1s
• “2” refers to the total number of electrons in
that orbital (s)
How it works
2
1s
• Summary
• There are two electrons(2) in the spherical
orbital (s) at the fist energy level (1)
Reminder/Review
• Principle energy levels
(n) can be divided into
sublevels.
• Number of sublevels
is equal to the number
of the principle energy
level.
Orbitals
• An orbital is a space that can be occupied
by up to two electrons.
• Each sublevel holds different number of
orbitals.
Sublevel
# of orbitals Maximum #
s
p
d
f
1
3
5
7
of electrons
2
6
10
14
Orbitals
• To calculate # of orbitals:
– # of orbitals = n2
• ex. Thrid energy level (principle level) (n=3)
– # of orbitals = 32
– # of orbitals = 9
– 3 sublevels (s+p+d) => (1+3+5), see above chart
Orbitals
• To calculate # of electrons:
– # of electrons = 2n2
• ex. fourth energy level (principle level) (n=4)
– # of orbitals = 2(4)2
– # of orbitals = 32
Filling Sublevels with Electrons
• There is a specific
order that energy
sublevels fill up.
Follow the chart.
• Or….
• Read the PT like a
book!
Filling Sublevels with Electrons
• Read the PT like a book!
• Remember:
– d elements move up 1 principle #
– f elements move up 2 principle #
Practice
Let’s practice!!!
• Oxygen
• Vandium
• Europium
Example
• Oxygen
– 1s2 2s2 2p4
• Vandium
– 1s2 2s2 2p6 3s2 3p6 4s2 3d3
• Europium
– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f6
Short-hand Notaion
•
•
•
•
•
Very similar to electron configuration
Start at the selected element
Move backward till you are at a noble gas.
Write the noble gas in brackets
Continue as normal.
Example
• Oxygen
– 1s2 2s2 2p4
• Vandium
– 1s2 2s2 2p6 3s2 3p6 4s2 3d3
– [Ar] 4s2 3d3
• Europium
– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f6
– [Rn] 6s2 4f6
Let’s practice!!!
• Oxygen
• Vandium
• Europium
Try these on your own!!!
Write the long-hand and short-hand electron
notation for the following elements.
1.
2.
3.
4.
5.
Beryllium
Fluorine
Silicon
Manganese
Gallium
6. Silver
7. Samarium
8. Gold
9. Bismuth
10.Uranium
Orbital Notaion
•
•
•
•
•
Very similar to electron configuration
Just use a box instead of a superscript
Boxes represent orbitals
Arrows represent electrons
Put electrons in “up” first, then “down”.
DON’T FORGET!!!
• An orbital is a space that can be occupied
by up to two electrons.
• Each sublevel holds different number of
orbitals.
Sublevel
# of orbitals Maximum #
s
p
d
f
1
3
5
7
of electrons
2
6
10
14
Orbital Notaion
• Remember only two electrons per box
• Ex. F
• Ex. P
Practice
•
•
•
•
Li
C
Al
S
Lewis Dot Diagram
• consists of an elemental symbol and…
• the valence electrons
– outer-most electrons of the atom
– valence electrons are the available electrons
that can be involved in bonding.
Lewis Dot Diagram
• If you don’t know how many electrons are
in the valence shell then, write the electron
configuration!
• Look at the last principal number.
• Li: 1s2 2s1
• S: 1s2 2s2 2p6 3s2 3p4
• Kr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
Lewis Dot Diagram
• How to make them
– there are multiple ways
– we are going to use a specific order
• Ex. O
Practice
•
•
•
•
Li
C
Al
S
Electron Notation Poster
• Once you have your group come to me to
get your Element.
Extra practice!!!!
If an atom has 16 protons and 30 neutrons and a
charge of -2
• What is the atomic number?_____
• What is the atomic mass?_______
• What is the mass number?______
• How many electrons does it have?_____
• What is the element symbol?______
• What kind of ion is it?________
If an atom has 19 protons and 42 neutrons and a
charge of +1
• What is the atomic number?________
• What is the atomic mass?_______
• What is the mass number?_________
• How many electrons does it have?______
• What is the element symbol?______
• What kind of ion is it?_______
Write two symbols on the board.
•
•
•
•
•
•
•
How many protons are there?______
How many neutrons are there?______
How many electrons are there?_____
What is the atomic number?______
What is the atomic mass?______
What is the mass number?_______
What kind of ion is it? ________
Label the Periodic
Table!!!
Make scientist map!!!
Remember:  = c / 
E = h
=c/
1)If the wavelength of a wave is 3.39 m, what
is it’s frequency?________
2)If the wavelength of a wave is .657 m, what
is it’s frequency?________
3)If the frequency of a wave is 4,985 Hz, what
is the wavelength?_______
4)If the frequency of a wave is 290 Hz, what
is the wavelength?_______
QUIZ
• Clear your desk of everything except a pen/pencil
and paper
1)Write the abbriviated electron notation of Ru (atomic # 44)
2)Write the unabbriviated electron notation of O (atomic # 8)
3)Write the orbital notation of C (atomic # 6)
4)Write the Lewis Dot Diagram of N (atomic # 7)