Transcript Waves - TeacherWeb

Waves
• All waves, whether they are
water waves or electromagnetic
waves, can be described in
terms of four characteristics:
•
•
•
•
amplitude
wavelength
frequency
speed
Amplitude

The amplitude of a wave is the height
of the wave measured from the origin
to its crest.
www.hyperphysics.phys-astr.gsu.edu
Wavelength

The wavelength of a wave is the
distance that the wave travels as it
completes one full cycle of upward and
downward motion.
•www.hyperphysics.phys-astr.gsu.edu
http://www.800mainstreet.com/spect/emission-flame-exp.html
Visible Light Wavelengths
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Visible light has wavelengths in the
range of 400 to 750 nm
Remember that a nanometer is 10-9
meter
You may remember the visible light
spectrum as ROYGBIV
Frequency
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The frequency of a wave tells how fast
the wave oscillates up and down.
The frequency of light is measured by
the number of times a light wave
completes a cycle of upward and
downward motion in one second.
Units can be: cycles/sec or Hertz
Speed of Light
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Light moves through space at a
constant speed of 3.00 x 108 m/s
You will use c = 3.00 x 108 m/s
Wavelength and Frequency
Calculation
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λ = c/v
If the frequency of radiation is 3 x 1015
cycles/sec, what is the wavelength?
Wavelength and Frequency
Calculation


λ = c/v
If the wavelength of radiation is
3 x 10-4 m, what is the frequency?
Electromagnetic
Spectrum

Electromagnetic radiation can be described in terms
of a stream of photons, which are massless
particles each traveling in a wave-like pattern and
moving at the speed of light. Each photon contains
a certain amount (or bundle) of energy, and all
electromagnetic radiation consists of these photons.
The only difference between the various types of
electromagnetic radiation is the amount of energy
found in the photons. Radio waves have photons
with low energies, microwaves have a little more
energy than radio waves, infrared has still more,
then visible, ultraviolet, X-rays, and ... the most
energetic of all ... gamma-rays.
Electromagnetic
Spectrum
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Radio Waves
Radio: yes, this is the same kind of energy
that radio stations emit into the air for your
boom box to capture and turn into your
favorite Mozart, Madonna, or Coolio tunes.
But radio waves are also emitted by other
things ... such as stars and gases in space.
You may not be able to dance to what these
objects emit, but you can use it to learn
what they are made of.
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Radio Waves
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Microwaves

Microwaves: they will cook your
popcorn in just a few minutes! In
space, microwaves are used by
astronomers to learn about the
structure of nearby galaxies, including
our own Milky Way!
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Infrared Radiation

Infrared: we often think of this as
being the same thing as 'heat',
because it makes our skin feel warm.
In space, IR light maps the dust
between stars.
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Visible Light

Visible: yes, this is the part that our
eyes see. Visible radiation is emitted
by everything from fireflies to light
bulbs to stars ... also by fast-moving
particles hitting other particles.
Don’t All Units Work The
Same?
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In the older "CGS" version of the metric system, the
units used were angstroms. An Angstrom is equal
to 0.0000000001 meters (10-10 m in scientific
notation)! In the newer "SI" version of the metric
system, we think of visible light in units of
nanometers or 0.000000001 meters (10-9 m). In
this system, the violet, blue, green, yellow, orange,
and red light we know so well has wavelengths
between 400 and 700 nanometers.
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Visible Spectrum
•http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
http://www.800mainstreet.com/spect/emission-flame-exp.html
Ultraviolet Radiation
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Ultraviolet: we know that the Sun is a
source of ultraviolet (or UV) radiation,
because it is the UV rays that cause
our skin to burn! Stars and other "hot"
objects in space emit UV radiation.
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
X-ray Radiation

X-rays: your doctor uses them to look
at your bones and your dentist to look
at your teeth. Hot gases in the
Universe also emit X-rays .
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Gamma Radiation
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Gamma-rays: radioactive materials (some
natural and others made by man in things
like nuclear power plants) can emit gammarays. Big particle accelerators that scientists
use to help them understand what matter is
made of can sometimes generate gammarays. But the biggest gamma-ray generator
of all is the Universe! It makes gamma
radiation in all kinds of ways.
Gamma Radiation
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Which types of radiation
reach the Earth?
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
What is Electromagnetic
Radiation?

Electromagnetic radiation from space is unable to reach the surface
of the Earth except at a very few wavelengths, such as the visible
spectrum, radio frequencies, and some ultraviolet wavelengths.
Astronomers can get above enough of the Earth's atmosphere to
observe at some infrared wavelengths from mountain tops or by
flying their telescopes in an aircraft. Experiments can also be taken
up to altitudes as high as 35 km by balloons which can operate for
months. Rocket flights can take instruments all the way above the
Earth's atmosphere for just a few minutes before they fall back to
Earth, but a great many important first results in astronomy and
astrophysics came from just those few minutes of observations. For
long-term observations, however, it is best to have your detector on
an orbiting satellite ... and get above it all!
http://imagine.gsfc.nasa.gov/docs/science/know_l1/emspectrum.html
Planck’s Theory
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Max Planck proposed that there is a
fundamental restriction on the
amounts of energy that an object
emits or absorbs, and he called each
of these pieces of energy a quanta.
E = hv
Planck’s constant is
h= 6.626 x 10-34 Js
Planck’s Equation
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E = hv
E = amount of energy emitted or
absorbed
h = Planck’s constant 6.626x10-34 Js
Photoelectric Effect
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Einstein used Planck’s equation E=hv to
explain the photoelectric effect.
In the photoelectric effect, electrons are
ejected from the surface of a metal when
light shines on the metal
Einstein proposed that light consists of
quanta of energy that behave like tiny
particles of light.
Energy quanta are photons.
Arthur Compton
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Compton demonstrated that a photon
could collide with an electron,
therefore a photon behaves like a
particle.
DeBroglie
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Louis de Broglie determined that
particles exhibit wavelike behavior.
Dual Nature of Light
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Light acts like particle and behaves like
a wave.
Types of Spectra
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Continuous Spectrum – a blend of
colors one into the other.
An example of a continuous spectrum
is a rainbow.
Types of Spectra
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Emission Spectrum (Bright-line spectrum)
- a spectrum that contains only certain
colors, or wavelengths
Energy is added to an element sample. The
electrons absorb the energy and jump to a
higher energy level. They only stay there for
an instant and then fall back to a lower
energy level. As the electrons fall back down
they emit photons of light. Each photon has
a specific wavelength and frequency.
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www.cms.k12.nc.us/.../notes/ch04electrons.html

http://www.800mainstreet.com/spect/emission-flame-exp.html
Creating Emission
Spectrum Lines
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http://www.800mainstreet.com/spect/emission-flame-exp.html
Helium Spectrum
http://hyperphysics.phy-astr.gsu.edu/hbase/quantum/atspect.html
Colorful Chemicals
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Try this web site to see the colorful spectra that
different metals can create.
http://webmineral.com/help/FlameTest.shtml
Bohr’s Model of the Atom
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Bohr listened to a lecture by Rutherford (about his model of
the atom). He realized how Planck’s idea of quantization could
be applied to this model to explain line spectra.
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He decided that electrons can be found only in specific energy
levels with specific amounts of energy.
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Each energy level was assigned a quantum number
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The ground state is the lowest energy level, n=1, this
energy level is closest to the nucleus
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Electrons absorb a specific quanta of energy and jump to an
excited state, n=2 or above
Bohr’s Explanation of
Hydrogen’s Spectral Lines
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Bohr proposed that when radiation is
absorbed, an electron jumps from the
ground state to an excited state.
Radiation is emitted when the electron
falls back from the higher energy level
to a lower one. The energy of the
absorbed or emitted radiation equals
the difference between the two energy
levels involved.
Bohr’s Explanation of
Hydrogen’s Spectral Lines
Continued
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Bohr used his model and Planck’s equation,
E=hv, to calculate the frequencies observed
in the line spectrum of hydrogen.
This model worked well for hydrogen with
one electron, but not for elements with
larger numbers of electrons.
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Planck, Einstein and Bohr described
light as consisting of photons – quanta
of energy that have some of the
characteristics of particles.
Heisenberg’s Uncertainty
Principle
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Heisenberg stated that the position
and the momentum of a moving
object cannot simultaneously be
measured and known exactly.
Probability of Locating an
Electron in an Atom
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think of the electrons as residing in a
cloud
more dense areas have a higher
probability of finding an electron
Draw a diagram of an atom with a
surrounding electron cloud
Quantum Mechanical Model
of the Atom
Atomic Orbitals
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An atomic orbital is a region around
the nucleus of an atom where an
electron with a given energy is likely
to be found.
The amount of energy an electron has
determines the kind of orbital it
occupies.
s sublevel
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s orbitals are spherical in shape
http://www.chemsoc.org/exemplarchem/entries/2004/dublin_fowler/sorbitals.html
p sublevels
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p orbitals are dumbbell shaped
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Px
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Py
Pz
http://www.chemsoc.org/exemplarchem/entries/2004/dublin_fowler/sorbitals.html
d sublevels
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d orbitals can be several shapes
dz2
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dx2-y2
http://www.chemsoc.org/exemplarchem/entries/2004/dublin_fowler/sorbitals.html
dxy
d sublevels
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dxz
dyz
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http://www.chemsoc.org/exemplarchem/entries/2004/dublin_fowler/sorbitals.html
f sublevels
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f orbitals are complicated 3D shapes
that need to be computer
generated
see
http://nobel.scas.bcit.ca/chem0010/un
it3/3.3.3_QM_econfig.htm#here
“To Do” Activity
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Color the s, p, d, f blocks on the
periodic table.
Principal Energy Levels
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Principal energy levels in an atom are
designated by the quantum number, n.
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“n” must be an integer
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look at the left hand margin of the periodic
table to find the principal quantum numbers
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As “n” increases (i.e. from 1 to 2), the
electron energy increases
Sublevels
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Each principal energy level is divided
into one or more sublevels.
The number of sublevels in each
principal energy level equals the
quantum number, n , for that energy
level.
How do you tell the
difference between
sublevels?
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Sublevels can be distinguished by
their:
– shapes
– sizes
– energies
Sublevels
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If “n” = 1
one sublevel
“s”
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If “n” = 2
2 sublevels “s” and “p”
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If “n” = 3
3 sublevels “s”, “p”, “d”
Orbitals
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Each sublevel consists of one or more
orbitals.
There can never be more than 2
electrons in each orbital.
Electrons in Orbitals
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Electrons behave as if they are
spinning on their own axis.
A spinning charge creates an electric
and magnetic field.
Pauli’s Exclusion Principle
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1. Each orbital in an atom can hold at
most 2 electrons
2. Each of these electrons must have
opposite spins.
Electron Pairing
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Two electrons with opposite spins (in
the same orbital) are paired.
Sublevel
Sublevel
Sublevel
Sublevel
“s” holds
“p” holds
“d” holds
“f” holds
2e
6e
10 e
14 e
Orbital Diagram “To Do”
Activity
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Acquire a clean orbital diagram.
Only use pencil on your diagram.
Use an orbital to build an electron
configuration.
Aufbau Principle
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Electrons are added one at a time to
the lowest energy orbitals available
until all of the electrons of the atom
have been accounted for.
Pauli Exclusion Principle
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An orbital can hold up to 2 electrons
Electrons in the same orbital must
have opposite spins
Hund’s Rule
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Electrons occupy equal energy orbitals
so that a maximum number of
unpaired electrons results.
This is commonly known as the “Seat
on the Bus” Rule