Ch 4

1
Ch 4

• Waves and Spectra

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Objectives

• SWBAT
• Label the parts of a wave.
• Solve wavelength to frequency calculations.

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Waves

• All waves, whether they are water waves or
  electromagnetic waves, can be described in terms
  of four characteristics
• amplitude frequency
• wavelength speed

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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
5
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
6
Visible Light Wavelengths

• 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

7
Frequency

• 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

8
Speed of Light

• 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

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Wavelength and Frequency Calculation

• ? c/v
• If the frequency of radiation is 3 x 1015
  cycles/sec,
• what is the wavelength?

10
Wavelength and Frequency Calculation

• ? c/v
• If the wavelength of radiation
• is 3 x 10-4 m, what is the frequency?

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Practice Problems

• 1. If the wavelength of radiation is 9.6 x 10 - 6
  m, what is the frequency?
• 2. If the frequency of radiation is 0.33 x 109
  cycles/sec, what is the wavelength?
• If the wavelength of radiation is 6.22 x 10 -12
  m, what is the frequency?
• 4. If the frequency of radiation is 7.8 x 1010
  cycles/sec, what is the wavelength?

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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

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• 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.

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Electromagnetic Spectrum
Retrieved from http//imagine.gsfc.nasa.gov/docs/
science/know_l1/emspectrum.html
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Radio Waves

• Radio
• this is the same kind of energy that radio
  stations emit into the air for your iPod to
  capture and turn into your favorite music.
• But radio waves are also emitted by other
  things ... such as stars and gases in space.
  You can use the information gathered from stars
  to learn what they are made of.
• Retrieved from http//imagine.gsfc.nasa.gov/docs/
  science/know_l1/emspectrum.html

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Radio Waves

• Long, Slow Waves

http//imagine.gsfc.nasa.gov/docs/science/know_l1/
emspectrum.html
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How Cell Phones Work

• A cell phone is actually a radio.
• To see how a cell phone works, check out
• http//www.howstuffworks.com/cell-phone.htm

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Microwaves

• 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
19
Infrared Radiation

• Infrared radiation we often think of this as
  being the same thing as 'heat', because it makes
  our skin feel warm.
• Have you seen IR heat lamps at fast food
  restaurants?
• In space, IR light maps the dust between stars.
• http//imagine.gsfc.nasa.gov/docs/science/know_l1/
  emspectrum.html

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Visible Light

• This is the part of the electromagnetic spectrum
  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.

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DONT ALL UNITS WORK THE SAME?

• 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 has wavelengths between 400
  and 700 nanometers.
• http//imagine.gsfc.nasa.gov/docs/science/know_l1/
  emspectrum.html

22
Visible Spectrum

• http//imagine.gsfc.nasa.gov/docs/science/know_l1/
  emspectrum.html

http//www.800mainstreet.com/spect/emission-flame-
exp.html
23
Ultraviolet Radiation

• 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

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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
25
Gamma Radiation

• Gamma-rays radioactive materials (some natural
  and others made by man in things like nuclear
  power plants) can emit gamma-rays.
• Big particle accelerators that scientists use to
  help them understand what matter is made of can
  sometimes generate gamma-rays.
• But the biggest gamma-ray generator of all is the
  Universe! It creates gamma radiation in all kinds
  of ways.

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Gamma Radiation

• Short, fast radiation

http//imagine.gsfc.nasa.gov/docs/science/know_l1/
emspectrum.html
27
Which Radiation Reaches the Earth?

• Electromagnetic radiation from space is unable
  to reach the surface of the Earth except at a few
  wavelengths, such as the visible spectrum, radio
  frequencies, and some ultraviolet wavelengths.
• Do you know what keeps the radiation from
  reaching the Earth?

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Which types of radiation reach the Earth?
?
?
?
http//imagine.gsfc.nasa.gov/docs/science/know_l1/
emspectrum.html
29
What is Electromagnetic Radiation?

• 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
30
PLANCKS THEORY

• 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
• Plancks constant is
• h 6.626 x 10-34 Js

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PLANCKS EQUATION

• E hv
• E amount of energy emitted or absorbed
• h Plancks constant 6.626x10-34 Js

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Photoelectric Effect

• Einstein used Plancks equation Ehv 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.

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Photoelectric Effect Experiment
http//zebu.uoregon.edu/2000/ph101/photoexp.gif
34
Photoelectric Effect
http//www.uwsp.edu/physastr/kmenning/images/photo
elec.gif
35
Arthur Compton

• Compton demonstrated that a photon could collide
  with an electron, therefore a photon behaves like
  a particle.

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DeBroglie

• Louis de Broglie determined that particles
  exhibit wavelike behavior.

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Dual Nature of Light

• Light acts like a particle and behaves like a
  wave.

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Types of Spectra

• Continuous Spectrum a blend of colors one into
  the other.
• An example of a continuous spectrum is a rainbow.

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Types of Spectra

• 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

41

• http//www.800mainstreet.com/spect/emission-flame-
  exp.html

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Helium Spectrum
http//hyperphysics.phy-astr.gsu.edu/hbase/quantum
/atspect.html
43
Colorful Chemicals

• Try this web site to see the colorful spectra
  that different metals can create.
• http//webmineral.com/help/FlameTest.shtml

44
BOHRS MODEL OF THE ATOM

• Bohr listened to a lecture by Rutherford (about
  his model of the atom). He realized how Plancks
  idea of quantization could be applied to this
  model to explain line spectra.
• He decided that electrons can be found only in
  specific energy levels with specific amounts of
  energy.
• Each energy level was assigned a quantum number
• The ground state is the lowest energy level, n1,
  this energy level is closest to the nucleus
• Electrons absorb a specific quanta of energy and
  jump to an excited state, n2 or above

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BOHRS EXPLANATION OF HYDROGENS SPECTRAL LINES

• 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.

46
BOHRS EXPLANATION OF HYDROGENS SPECTRAL LINES
CONTINUED

• Bohr used his model and Plancks equation, Ehv,
  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.

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HEISENBERGS UNCERTAINTY PRINCIPLE

• Heisenberg stated that the position and the
  momentum of a moving object cannot simultaneously
  be measured and known exactly.

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Probability of Locating an Electron in an Atom

• 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

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Quantum Mechanical Model of the Atom

• Draw a picture of the quantum mechanical model.

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Atomic Orbitals

• 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.

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Drill

• Pick up a copy of the 4-5 Prac Prob WS
• Try problems 1, 6, 7, 8 and 9

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Objectives

• iWBAT
• Color a periodic table to enable me to locate the
  s, p, d, f sublevels on the periodic table.
• Determine the elemental composition of a star by
  using various emission spectra graphical
  information.

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TO DO ACTIVITY

• Color the s, p, d, f blocks on the periodic
  table.

56
(No Transcript)
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SHAPES OF SUBLEVELS

• The next several slides will show the shape of
  the s, p, d, f sublevels.

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s sublevel

• s orbitals are spherical in shape

http//www.chemsoc.org/exemplarchem/entries/2004/d
ublin_fowler/sorbitals.html
59
p sublevels

• p orbitals are dumbbell shaped
• Px Py Pz
• http//www.chemsoc.org/exemplarchem/entries/2004/d
  ublin_fowler/sorbitals.html

60
d sublevels

• d orbitals can be several shapes
• dz2 dx2-y2 dxy
• http//www.chemsoc.org/exemplarchem/entries/2004/d
  ublin_fowler/sorbitals.html

There are 2 more on the next slide.
61
d sublevels

• dxz dyz
• http//www.chemsoc.org/exemplarchem/entries/2004/d
  ublin_fowler/sorbitals.html

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f sublevels

• f orbitals are complicated 3D shapes that need
  to be computer generated
• see http//nobel.scas.bcit.ca/chem0010/unit3/3.3.3
  _QM_econfig.htmhere

63
DRILL

• Pick up a copy of the Reading an e-
  Configuration WS drill on the front desk.
• Answer the questions
• .

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OBJECTIVES

• iWBAT
• Distinguish between principle energy levels and
  sublevels
• Use spectra data for various elements to
  determine the composition of Stars

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STAR SPECTRA

• Pick up a copy of the Star Spectra WS
• I will use the document camera to show you how to
  complete this worksheet.

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Principal Energy Levels

• Principal energy levels in an atom are designated
  by the quantum number, n.
• n must be an integer
• look at the left hand margin of the periodic
  table to find the principal quantum numbers
• As n increases (i.e. from 1 to 2), the electron
  energy increases

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Sublevels

• 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.

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How do you tell the difference between sublevels?

• Sublevels can be distinguished by their
• shapes
• sizes
• energies

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Sublevels

• If n 1 sublevel s
• If n 2 sublevels s and p
• If n 3 sublevels s, p, d

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Orbitals

• Each sublevel consists of one or more orbitals.
• There can never be more than 2 electrons in each
  orbital.

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Electrons in Orbitals

• Electrons behave as if they are spinning on their
  own axis.
• A spinning charge creates an electric and
  magnetic field.

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PAULIS EXCLUSION PRINCIPLE

• 1. Each orbital in an atom can hold at
  most 2 electrons
• 2. Each of these electrons must have
  opposite spins.

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Electron Pairing

• Two electrons with opposite spins (in the same
  orbital) are paired.
• Sublevel s holds 2 e-
• Sublevel p holds 6 e-
• Sublevel d holds 10 e-
• Sublevel f holds 14 e-

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YOU KNOW THERE ARE SCIENTISTS

• The next three slide list the scientist/proper
  name for some of the rules that we follow when we
  fill an orbital diagram.

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Aufbau Principle

• 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.

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Pauli Exclusion Principle

• An orbital can hold up to 2 electrons
• Electrons in the same orbital must have opposite
  spins

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HUNDS RULE

• 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

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ORBITAL DIAGRAM TO DO ACTIVITY

• Acquire a clean orbital diagram.
• Only use pencil on your diagram.
• Use an orbital to build an electron configuration.