Arrangement of Electrons in Atoms

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Chapter 4

• Arrangement of Electrons in Atoms

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RUTHERFORD MODEL
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CONCLUSIONS

• The atom contains a dense center called the
  nucleus.
• The nucleus contains a positive charge and most
  of the mass of the atom.  
• The nucleus is approximately 100,000 times
  smaller than the atom.

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CONCLUSIONS

• The nucleus is positively charged. The amount of
  positive charge balances the negative charge of
  the electrons.
• The electrons move around in the empty space of
  the atom surrounding the nucleus.

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Incomplete Model

• Did not explain where the electrons are located
  in the space surrounding the nucleus.
• Investigations into the absorption and emission
  of light by matter led to a new atomic model
• Studies revealed a relationship between light and
  an atoms electrons.

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Dual Nature of Light
Electromagnetic radiation exhibits 1) Wave
Properties 2) Particulate Properties
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Electromagnetic Radiation

• Radiant energy that exhibits wavelength-like
  behavior and travels through space at the speed
  of light in a vacuum.

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Classification of Electromagnetic Radiation
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Waves

• Waves have 3 primary characteristics
• 1. Wavelength distance between two peaks in
  a wave.
• 2. Frequency number of waves per second
  that pass a given point in space.
• 3. Speed speed of light is
• 3.0 ? 108 m/s.

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The Nature of Waves
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Wavelength and frequency can be interconverted.

• ? c/?
• ? frequency (s?1)
• ? wavelength (m)
• c speed of light (m s?1)

Inverse Proportion.
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Calculate the wavelength of radio waves broadcast
from an FM radio station at 98.5 MHz.

• c ? ?
• c/ ?
• 3.0 x 108 m/s ? 98.5 x 106 s-1
• 3.0 meter

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Particle Description of Light

• First arose due to the observation of the
• Photoelectric effect
• Emission of electrons from metal
• When light shines on the metal.
• Requires light with a certain minimum
• Frequency.
• See Figure 4-3 on page 93.

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Wave theory would have predicted that light of
any frequency could supply enough energy to
eject an electron.
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Max Planck (early 1900s)

• Studied emission of light by hot objects.
• Proposed objects emit energy in small specific
  amounts called Quanta.
• Quantum minimum quantity of energy
• that can be lost or gained by an atom.

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Plancks Constant
Transfer of energy is quantized, and can only
occur in discrete units, called quanta.

• ?E change in energy, in J
• h Plancks constant, 6.626 ? 10?34 J s
• ? frequency, in s?1
• ? wavelength, in m

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Surprise! Energy seems to have particle-like
properties!
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Albert Einstein

• Proposed the dual wave-particle nature of light.
• Light exhibits many wave-like properties, but
  also can be thought of as a stream of particles.
• Light is composed of a stream of Particles called
  Photons.

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Electromagnetic Radiation
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• Photon A particle of electromagnetic
• radiation having zero mass and
• carrying a quantum of energy.

Energy of a photon E h? h c/?
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Photoelectric Effect

• To emit or eject an electron from a metal, the
  electron must be struck by a photon with a
  certain mim. Energy to knock that electron loose.
• If the photon energy is too low, the electrons
  remain bound to the metal surface.

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Electrons in different metals, are bound more or
less tightly, Therefore, different metals
require different mim. Frequencies to
exhibit the photoelectric effect.
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Ground State Atom in its lowest energy state.
Excited State Atom with a higher
potential Energy.
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A Change between Two Discrete Energy Levels
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Atomic Spectrum of Hydrogen

• Continuous spectrum Contains all the
  wavelengths of light.
• Line (discrete) spectrum Contains only some of
  the wavelengths of light.

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A Continuous Spectrum (a) and A Hydrogen Line
Spectrum (b)
See also Fig. 4-6 Page 95.
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• Emission Spectrum of each element
• is Unique.
• Can be used to identify samples.
• Can help determine components of
• the stars.

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BOHR MODEL
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Niels Bohr

• Developed Quantum Model for the Hydrogen Atom.
• Model could explain the observed line emission.
• The Electron in a Hydrogen Atom moves around the
  nucleus only in certain allowed circular orbits.

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Bohr Model

• Lowest energy state
• electron is in the orbit closest to the
  nucleus.
• Higher energy states
• electron is in orbits that are successively
  farther from the nucleus.

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Electronic Transitions in the BohrModel for the
Hydrogen Atom
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• Bohr Model
• successful in explaining the spectral
• Lines of hydrogen
• did not explain spectra of atoms
• With more than one electron.
• did not explain the chemical
• behavior of atoms.