The Hook

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Chapter 4
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The Hook

• Burning money demo
• Draw a picture of what happened.
• What color did you see? Any idea why?
• Why didnt the money burn?

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Chapter 4

• Everything you ever wanted to know about where
  the electrons hang out!

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Building on the Atomic Theory

• What did Thompson determine?
• What did Rutherfords gold
• foil experiment prove?
• Just write the words we will talk in class!

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Section 1 Early 1900s

• Scientists started doing a lot of experiments
  looking at the absorption and emission of light
  by matter.
• Found that there is a relationship between light
  and an atoms electrons.

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Light can behave as a wave
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Draw the Wave!

• Amplitude height of the wave from the origin to
  the crest
• Wavelength (? ) the distance between the crests
  (m, cm, nm)
• Frequency (v) number of waves to pass a given
  point per unit of time (waves/second Hz)

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An Important Relationship

• The frequency and wavelength of all waves,
  including light, are inversely related.
• As the wavelength of light increases, the
  frequency decreases.

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C ?v

• Where
• C speed of light 3.00 x 108 m/sec
• ? wavelength (m, cm, nm)
• v frequency (1/sec or sec-1)

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Electromagnetic Radiation

• Includes radio waves, radar, microwaves, visible
  light, infrared light, ultraviolet light, X-rays,
  and gamma rays

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Photoelectric Effect

• Looks at the emission of electrons from a metal
  when light shines on the metal.
• Light causes electrons to be ejected from the
  metal.

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Photoelectric Effect

• Only occurs at certain frequencies!

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Wave Particle Duality

• Explained by Dr. Quantum
• Leave some space here to write a reflection on
  the video clip.

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Sometimes Light Acts Like Particles!

• What would happen if the frequency of the wave
  increased so much that you could hardly tell
  where one wave ended and another began?
• Light would start acting more like a particle
  than a wave.

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Max Plank

• Objects emit small packets of energy- Quanta
• Quantum- the minimum quantity of energy that can
  be lost or gained by an atom.
• E hv
• E Energy
• h 6.626 x 10 -34 Js (Joule x sec)
• V frequency (1/sec)

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Take a look at the WS

• Let the units be your guide!!!!!

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The Photon

• Photon- a particle of electromagnetic radiation
  having no mass, carrying a quantum of energy.

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So, what happens when photons hit an atom and
eject an electron?

• The electron goes from the ground state to an
  excited state.
• As the electron returns to the ground state, it
  gives off the energy that it gained- LIGHT

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Energy Levels

• Energy levels are
• not evenly spaced
• Energy levels
• become more closely
• spaced the greater
• the distance from the
• nucleus

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Another Cool Illustration
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The Visible Spectrum

• From about 400nm to 700nm in wavelength.
• Blue (400nm) has a shorter wavelength than red
  (700nm).

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Spectral Analysis of Emitted Light from Excited
Atoms

• When emitted light from excited atoms is passed
  through a prism a spectrum of discrete lines of
  separate colors (separate energies) is observed
  rather than a continuous spectrum of ROY G BIV.

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Each element has a unique line-emission spectra
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Interpretation of Line Spectrum of Elements

• The light atoms give off contain very specific
  wavelengths called a line spectrum
• light given off emission spectrum

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Continuous Spectrum
Atomic Line Spectrum
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Atomic Spectrum Activity
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Interpretation of Atomic Spectra

• The line spectrum is related to energy
  transitions of electrons in the atom.
• Absorption atom gaining energy
• Emission atom releasing energy

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• All samples of an element give the exact same
  pattern of lines because every atom of that
  element must have certain, identical energy
  states
• Energy of an atom is quantized limited to
  discrete values
• If the atom could have all possible energies,
  then the result would be a continuous spectrum
  instead of lines

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Bohr Model

• Electrons orbit around a nucleus
• Each orbit has a fixed energy and because of this
  cannot lose energy and fall into the nucleus
• Energy Level of an electron is the region around
  the nucleus where the electron is likely to be
  moving

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This helped explain the spectral lines

• Absorption- the electron gains energy and moves
  to a higher energy level.
• Emission- when the electron falls to a lower
  energy level.

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The Quantum ModelFinally the truth (as we know
it!)

• Electrons can behave as both waves and particles.
• Electrons can be considered waves with specific
  frequencies confined to the space around the
  nucleus.
• Electrons can also be considered negatively
  charged particles.

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Where are the electrons?

• Heisenberg Uncertainty Principle
• It is impossible to know the position and the
  velocity of an electron at the same time.

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Schrodinger Wave Equation

• Developed an equation that treated electrons as
  waves and described the location of electrons.
• Helped lay the foundation for modern quantum
  theory (atomic model).

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Quantum Theory

• Estimates the probability of finding an electron
  in a certain position
• We denote the position of the electron as a
  fuzzy cloud
• This volume of space where an electron is most
  likely to be found is called an orbital.
• The atomic orbitals have distinct shapes