Download D01 Atomic Models.notebook

D01 Atomic Models.notebook
May 26, 2016
Atomic Models
1) Students will be able to describe the evolution of atomic models.
2) Students will be able to describe the role of experimental evidence in changing models of the atom.
3) Students will be able to use photon energies to describe the energy levels of atoms.
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D01 Atomic Models.notebook
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Teachers' notes
Subject
Physics 30
Topic
topic
Title
title
Grade(s)
Cross­curricular link(s)
Prior knowledge
12
curr.
know.
Intended learning outcome(s)
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Lesson notes
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Focusing questions:
1. What are atoms made of?
2. What holds atoms together?
3. How do we know any of this stuff?
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Early models of the atom:
Democritus (ancient Greece)
­coined the term atom (from atomos...indivisible)
­believed that matter consisted of small, indivisible atoms that had space between them
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Dalton's "Billiard ball" model (1800's)
­all matter is made of atoms
­all atoms of an element are identical, and different from those of a different element
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J.J. Thomson's q/m ratio of the electron and the "Plum Pudding" model
Recall: cathode rays were emitted from the cathode when a voltage was placed across two plates that were in a vacuum
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Consequences of Thomson's experiments:
­determined that the atom was divisible ­the atom could be divided into separate negative and positive charges
­calculated the q/m ratio for the electron (needed Millikan to eventually get me)
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­revised the model of the atom to consist of a homogenous distribution of negative charge embedded in a "soup" of positive charge
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Rutherford Scattering Experiment (see p. 767)
­Rutherford fired alpha particles at a thin sheaf of gold foil. A detector could be moved to determine where the scattered α2+ had been deflected.
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Expectations:
α2+ were known to be relatively large (about 8000X the size of an electron). If α2+ were directed at gold atoms they should pass straight through, as there was nothing large enough in the atom (according to Thomson's model) to cause a deflection.
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Actual results:
­the vast majority of particles went straight through the gold foil, but some underwent major deflections (up to 180 degrees)
"It was almost as incredible as if you fired a 15­inch shell at a piece of tissue paper and it came back to hit you." 15
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Consequences of the experiment: the nuclear model of the atom:
­most of the mass (and all of the + charge) of the atom is located in a central, dense nucleus
­most of the atom consists of empty space
­the electrons circle the nucleus like "planets around the Sun"
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See p. 768
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Problem with the nuclear model:
­Maxwell had predicted that accelerating charges emit EMR. e­ circling the nucleus would be accelerating and should emit EMR and lose energy, spiraling into the nucleus... yet this doesn't happen.
What are we to do with the electrons?
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More unsettling experimental evidence that needs to be reconciled with theory...
Continuous Spectra:
­produced by hot objects
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Emission Spectra (Bright line spectra):
­produced by a hot, low density gas
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Hydrogen
Iron
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­emmission spectra consist of specific wavelengths of light that are emitted by an element when it is excited
­these wavelengths act as a "fingerprint" for that element
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Absorption spectra (dark line spectra):
­produced when white light is passed through a cool, low­density gas
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The bright lines produced by an element's emission spectrum correspond to the dark lines produced by the same elements absorption spectrum (i.e. one is the "photo­negative" of the other).
Problem: What kind of model of the atom can explain the phenomenon of bright line and dark line spectra?
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Neils Bohr and the Stationary State Model
Bohr suggested that:
­ electrons can orbit the nucleus at specific distances from the nucleus. These distances are some multiple of the smallest radius possible
Orbits in an atom are quantized.
­each of these orbits has a specific energy, which is also a multiple of the energy of the smallest radius
Electron energy levels (stationary states) are quantized.
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Key concept: The energy of the photon is equal to the difference in the energy of the stationary states.
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Aurora Borealis
Recall...the sun emits a solar wind, which consists of highly energetic particles. These particles are deflected by Earth's magnetic field toward the magnetic north and south poles.
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When the particles in the solar wind interact with gases in the atmosphere (primarily oxygen and nitrogen), they cause electrons in these gases to reach an excited state. As these electrons undergo a transition to lower states, they emit characteristic wavelengths of light.
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Relevant Reading: p. 754­779 (Be aware that you are not responsible for much of the "math" in this section of the textbook)
Check and Reflect, p. 780 #5­9,11
15.4 eTest
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Lasers
Laser: Light amplification by stimulated emission of radiation
An excited electron will drop down to a less energy energy state spontaneously. As it does so, it emits a photon. This is spontaneous emission of radiation.
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Einstein predicted that an excited electron could be "encouraged" to drop to its lower energy state. This is accomplished by directing at the excited atom photons that have the same frequency as the photons that will be emitted during the electron transition. These incident photons are not absorbed.
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This process is called stimulated emission of radiation.
Some properties of laser light produced by this process:
­monochromatic (i.e. all one wavelength)
­the light is coherent (i.e. the EMR is in phase)
­coherent light tends to focus in a tight beam
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Major problem with Bohr's stationary state model:
­it doesn't really explain why energy is quantized, or why orbiting electrons don't emit EMR
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Electron Waves and the Model of the Atom
Recall...de Broglie had predicted that electrons should exhibit wave behaviours. Evidence of this behaviour was found when electrons were directed at a crystal (which acted as a diffraction grating) and interference patterns were observed.
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The principle of standing waves can be applied to electrons and the atom.
For an electron wave to be produced such that there is constructive interference, the circumference of the orbit must be equal to some whole number of wavelengths.
Circumference = nλ
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A stationary state can
exist at this circumference
because a whole # of λ's fit (i.e. we have constructive
interference).
A stationary state can't exist at
this location because
the electron wave experience
destructive interference.
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Consequences of this wave understanding:
­the orbital isn't "the path" that electrons are likely to follow. Electrons (acting as waves) do not have a defined position. This means that the orbitals show the likelihood (or probability) of an electron being in a specific location.
The quantum model of the atom is a highly statistical model because of this idea of quantum indeterminacy (we can't be sure...we can just calculate the probability).
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