In 1665, Isaac Newton was a young scientist studying at Cambridge University in England. One sunny day, Newton darkened his room and made a hole in his window shutter, allowing just one beam of sunlight to enter the room. He then took a

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In 1665, Isaac Newton was a young scientist
studying at Cambridge University in England. One
sunny day, Newton darkened his room and made a
hole in his window shutter, allowing just one
beam of sunlight to enter the room. He then took a
glass prism and placed it in the sunbeam. The
result was a spectacular multicolored band of
light just like a rainbow. The multicolored band
of light is called a color spectrum.
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Newton believed that all the colors he saw were
in the sunlight shining into his room. To test
this, he placed another prism upside-down in
front of the first prism. He was right. The band
of colors combined again into white sunlight. He
was the first to show that white light is made up
of the colors that we see.
Newton also felt that light rays were made up of
particles,
but nearly 150 years later
Fresnel showed that light waves demonstrated
diffraction, a property of waves.
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A. J. Fresnel demonstrated mathematically that
the pattern of alternating light and dark lines
(diffraction pattern) produced when light travels
through an aperture, or around an obstruction,
would only occur if light moved as waves.
Nearly 20 years later, in the early 1830s
Michael Faraday demonstrated that a magnetic
field could cause an electric current to flow in
a wire. This is known as magnetic induction.
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A little more than 30 years after that James
Clerk Maxwell developed a set of equations that
confirmed Faradays idea that electricity and
magnetism are simply two parts of a single
phenomenon,
electromagnetism.
Maxwell showed that this phenomenon would produce
waves which travel at the speed of light. He
also suspected that there were light waves other
than those that produced the light that we could
see. We now refer to this collection of
different waves of electromagnetic radiation
(light) as the
electromagnetic spectrum (EMS)
Before we look at the EMS, lets talk about waves.
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There are two main types of waves,
transverse,
and
longitudinal.
http//paws.kettering.edu/drussell/Demos/waves/wa
vemotion.html
Light moves through space as two interacting
transverse wave disturbances,
one caused by the electrical field and the other
by the magnetic.
Parts of a wave you should be familiar with
wave
Another important feature of wave motion is the
inverse relationship between wavelength and
frequency.
That is, as one increases the other decreases.
This relationship is expressed in the following
equation

c
(the speed of light)
?
(wavelength)
(frequency)
?
Well
What about units, you ask?
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What is the highest point on a wave called?

1. Trough
2. Crest
3. Amplitude
4. Colgate

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the speed of light has a constant value of 3.00 x
108 m/s
Wavelength is in meters.
Frequency is in
cycles/s
or 1/s
These are also known as 1 hertz (Hz)
or s-1.
Sample problem 1
Calculate the
frequency of light with a wavelength of 5.22 x
10-10 m.
c ??
Manipulating the variables to solve for ? gives,
v c/?
/
/

3.00 x 108 m/s

5.22 x 10-10 m
/
5.75 x 1017/s

5.75 x 1017 Hz
Sample problem 2 (You try)

Calculate the wavelength of a radio station
signal with a frequency of 99.7 MHz.
c ??

/
?
c
/
/
?

3.00 x 108 m/s
/
99.7 x 106/s
3.01 m

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OK, now lets discuss the Electromagnetic
Spectrum.
microwaves
Radio waves
The spectrum goes from radio waves (long ?) on
the right

The visible part of the spectrum goes
from red (7 x 10-7m) to violet (4
x 10-7m),
to gamma rays (short ?) on the left.
,
in this sequence v increases.
ROYGBIV
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At the end of the 19th century, there were two
things about which science was certain
Matter was particles

everything was messed up, by a German
scientists named
and light was waves.
Then,
in 1900,
Max Planck.
Planck was trying to determine how the color of
light radiated by a body was related to its
temperature.
Two separate mathematical explanations already
existed.
One failed to work for light at high frequencies,
the other at low frequencies.
Planck demonstrated that the
problem could be solved by treating light as
being given off in discrete units he called
them quanta - rather than being given off
continuously, as previously assumed.
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Planck found that light energy was proportional
to its frequency.
The relationship is given by
E

h
?
E is energy.
h is Plancks constant,
6.63 x 10-34 J . s
If light is only emitted or absorbed in a
discrete quantum of energy,
If its not continuous,
then what is it?
A particle?
Evidence to support Plancks theory came in 1905
when a Swiss patent clerk explained the
photo-electric effect.
When UV light is shined on a
piece of metal connected to a circuit, electrons
are ejected into the circuit and a current is
produced.
What no one had yet explained was why the
intensity of light had no effect on the energy of
the electrons.
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The patent clerk proposed that if, like Planck
said, light consisted of discrete quanta -
photons
-
they would interact with the electrons like
particles.
He even showed the electrons had energy
related to hv !!
Who was this clerk?
Albert Einstein
Sample Problem 1
Calculate the
energy of a photon of light with a frequency
of 5.45 x 1014 Hz.
E hv
(6.63 x 10-34 J . s)

(5.45 x 1014/s)

/
/
3.61 x 10-19 J
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Now that we have some knowledge about light, we
can examine Niels Bohrs solution to the electron
orbit problem.
Bohr combined the science of
spectroscopy with Plancks quantum theory to
develop his explanation.
What is spectroscopy?
Lets find out.
If white light is separated into its
spectrum, the colors form a continuum.
If the light given off from a single element is
separated, however, distinct lines are formed.
Distinct dark lines are also formed when white
light has passed through the gas of a single
element.
This is called an
emission spectrum.
This is known as an
absorption spectrum.
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Why was the spectrum of an element composed of
separate lines, and not continuous like the
spectrum of white light?
The important question was,
Bohrs answer to that question would be his
explanation of how the electrons behaved in an
atom.
In 1885 Johann Balmer had
shown mathematically that the wavelengths of
the lines in the visible emission spectrum of
hydrogen resulted from some whole number
transition.
Bohr suggested that this transition corresponded
to an electron jumping from one possible orbit to
another and emitting a photon of light energy.
red
violet
blue
blue-green
4 visible spectral lines of H
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In Bohrs model of the atom, the electron can
only exist in these specific orbits in an atom.
Since an electron would have to possess a
specific amount of potential energy to be in one
of these orbits,
they are known as
energy levels.
Here is how it works
The electron is attracted to the positive charge
of the nucleus, so the electron has to have more
energy to be
far from the nucleus than to be
close.
Normally the electron would be in
its low- est available energy level,
this is called its ground state.

If the atom is
exposed to an energy
source the electron can absorb a quantum of
energy (photon)
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and the electron will make a quantum leap to a
higher energy level.
The electron will then
drop back down to a lower energy level.
In order for this to happen,
the
electron has to give off a quantum of light
energy (photon).

The energy of this photon would correspond
exactly to the energy difference between the two
levels.
Bohr model
Using Plancks equation, Bohr was able to
calculate the energy value for each level in the
hydrogen atom.
It is important to note that the electron can
jump from one level to another,
but it cannot go in between them.
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Wave-Particle Duality

• Louis De Broglie (1924)
• Proposed that ALL matter has wave and particle
  properties, not just electrons.
• E E ? E h? or E hc/? E mc2
• hc/? mc2 ? hc mc2? ? h mc?
• ? h/mc OR ? h/m?
• Example
• ? of baseball (mass .2 kg and ? 30 m/s)
• ? of an electron (mass 9.11 x 10-31 kg and ?
  3 x 108 m/s)

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Wave-Particle Duality

• Heisenberg (1927)
• Said that because of size and speed it is
  impossible to know both exact position and
  momentum of and electron at the same time.
• This is referred to as Heisenberg Uncertainty
  Principle
• To see an electron we strike it with something
  of similar size and observe its behavior.
• We cannot see an electron directly.
• We use photons of energy to do this.

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Quantum Mechanics

• The work of de Broglie and Heisenberg led to the
  study of quantum mechanics (motion in
  increments)
• 1. classical physics
• describes the motion of bodies much larger than
  the atoms of which they are composed.
• energy can be gained or lost in any amount
• 2. quantum physics
• describes the motion of atoms and subatomic
  particles as waves.
• particles gain or lose energy in packets called
  quanta

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Quantum Mechanical Model

• Schroedinger (1887-1961)
• Developed the quantum mechanical model of the
  atom
• He used the following equation to produce
  scatterplots that are now called electron
  clouds
• E 2?2me2/h2n2
• These electron clouds are areas in which there is
  a great probability of finding an electron (90).
• The cloud is more dense where the probability of
  finding an electron is high.
• The cloud is less dense where the probability of
  finding an electron is low.
• This is called an orbital a region in space
  in which there is a high probability of finding
  an electron.

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http//scienceworld.wolfram.com/physics/Schroeding
erEquation.html http//www.uark.edu/misc/julio/orb
itals/index.html
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Electrons and Electron Configurations (Not on
note sheet)

• Electrons have an address with four quantum
  numbers.
• We have 3 General rules for distributing these
  electrons.
• Pauli Exclusion Principal Orbitals contain no
  more than two electrons. OrEach address
  describes the location of only one electron.
• Hund Rule When filling orbitals, assign one
  electron to each orbital (of that type) before
  doubling up with two electrons per orbital.
• Aufbau Electrons fill lowest orbitals first,
  then proceed to higher energy levels.

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Energy Components in Electrons

• Each component is given a letter a name we
  call them quantum number values
• 1. n principal
• distance from the nucleus (energy level)
• 2. l azimuthal
• indicates the type of orbital in which the
  electron moves (sublevel)
• 3. m magnetic
• indicates the orientation about the three axes in
  space of the orbital (specific orbital)
• 4. s spin
• indicates the direction of the spin of the
  electron either clockwise or counterclockwise.
• Using these we can pinpoint the exact location of
  an e-.

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Location

• n principal energy level
• n l energy sublevel, defines the type of
  orbital that the electron is in
• n l m specific orbital (axis orientation)
• n l m s spin (exact electron), identifies
  the exact electron and its location
• ANALOGY

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Orbital Types

• S-orbital spherical shape, only 1 of them
• P-orbital gumdrop or dumbell shape, 3 of them
  one on each axis (x,y,z)
• D-orbital donut shape, 5 of them
• F-orbital cigar shape, 7 of them
• Each orbital contains a max of 2 electrons
• Orbit path of an electron (according to Bohr)
• Orbital region in space where there is a high
  probability of finding an electron

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Orbital Shapes
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Orbital Sites

• http//www.colby.edu/chemistry/OChem/DEMOS/Orbital
  s.html
• http//micro.magnet.fsu.edu/electromag/java/atomic
  orbitals/index.html
• http//itl.chem.ufl.edu/ao_pict/ao_pict.html
• http//winter.group.shef.ac.uk/orbitron/AOs/1s/ind
  ex.html

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ENERGY LEVELS ORBITAL TYPES OF ORBITALS OF ELECTRONS
n 1 s 1 2
n 2 s,p 4 8
n 3 s,pd, 9 18
n 4 s,p,d,f 16 32
n 5 s,p,d,f,g 25 50
Energy level the number of orbital types Total
number of orbitals in an energy level n2 Total
number of electrons in any energy level 2n2
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Which scientists is responsible for developing
the quantum mechancial model of the atom?

• A. Einstein
• B. De Broglie
• C. Schroedinger
• D. Heisenberg

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This quantum number indicates the shape of the
orbital that the electron moves.

• A. Principal
• B. Azimuthal
• C. Magnetic
• D. Spin

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How many energy sublevels are in the 3rd energy
level?

• A. 1
• B. 2
• C. 3
• D. 4

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What are the shapes of the energy sublevels in
the 2nd energy level?

• A. Spherical
• B. Spherical and Dumbbell shaped
• C. Spherical and Donut shaped
• D. Spherical, Dumbbell, Donut, and Cigar shaped

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What is the total number of electrons in the 1st
energy level?

• A. 2
• B. 8
• C. 18
• D. 32

33
What is the maximum number of p orbitals that can
be in any one energy level?

• A. 1
• B. 3
• C. 5
• D. 7

34
How many total electrons can the d orbitals hold?

• A. 2
• B. 6
• C. 10
• D. 14

35
What is the total number of orbitals in the 4th
energy level?

• A. 1
• B. 4
• C. 9
• D. 16

36
What types of orbitals are in the 2nd energy
level?

• A. s orbitals only
• B. s and p orbitals
• C. s, p, and d orbitals
• D. s, p, d, and f orbitals

37
How many energy sublevels are in the 1st energy
level?

• A. 1
• B. 2
• C. 3
• D. 4

38
What is the total number of electrons in the 3rd
energy level?

• A. 2
• B. 8
• C. 18
• D. 32

39
Which quantum number indicates the energy level
the electron is in?

• A. Principal
• B. Azimuthal
• C. Magnetic
• D. Spin

40
Electron Configuration
A method we use to keep track of how electrons
are arranged in an atom.
It helps us to explain why atoms react the way
they do.
of electrons in sub-level
1s1
H
Sub-level
The arrangement of the electrons may be
represented in one of three ways.
Energy Level
This is how hydrogen would be shown using
spectroscopic notation
Electron configuration may also be shown with box
diagrams.
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Electron Configuration
A box is used to represent each orbital.
An arrow is used to represent each electron.
These may also be called orbital box diagrams or
orbital filling diagrams.
H 1s1
Circles may also be used instead of boxes.
Helium has 2 electrons
He
1s2
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Electron Configuration
Here are some more
Li
1s2 2s1
F
1s2 2s2 2p5
Ne
1s2 2s2 2p6
Be
1s2 2s2
1s2 2s2 2p1
B
This arrangement (filled s p in the outer
energy level) is called a stable octet.
C
1s2 2s2 2p2
N
1s2 2s2 2p3
1s2 2s2 2p4
O