Atomic Physics

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Atomic Physics

• Early Atomic Models
• Dalton Atom
• Thompson Atom
• Rutherford Atom

• Matter Waves
• Wave-particle duality
• Physical meaning
• When wave properties are important

• The Dual Nature of Light
• Wave Properties
• Particle Properties

• Bohr Atom
• Quantization
• Atomic Spectra

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The Dalton Atom

• Dalton's atomic model was developed to explain
  chemical behavior.
• In this model, atoms had no internal structure
  and could be distinguished only by mass and
  chemical properties.

Oxygen
Carbon
Hydrogen
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Thompson Atom

• J. J. Thompson proposed a model in which
  electrons were embedded within a positively
  charged substrate.
• This became known as the "plum pudding model".

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Rutherford's Atom

• Rutherford's experiments with alpha particles
  showed that the positive part of the atom had to
  be smaller and denser than in Thompson's model.

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Solar System Model

• In this model, the electrons are assumed to orbit
  the atom in much the same way as planets orbit
  the sun.
• The classical version, however, predicts that the
  electrons will lose energy by EM radiation and
  quickly crash.

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The Dual Nature of Light

• Some properties of light, such as diffraction and
  interference, can only be explained using a wave
  model of light.
• Others, such as blackbody radiation and the
  photoelectric effect, can only be explained by
  assigning particle properties to light.

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Wave Properties of Light

• Diffraction refers to the spreading of a wave
  after it encounters a small obstacle or opening.
• Interference refers to the cancellation of two
  waves as they pass through the same region.

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Blackbody Radiation

• Blackbody radiation is the EM radiation emitted
  by any object with a temperature above absolute
  zero.
• The wavelength at which the most energy is
  emitted is inversely proportional to the object's
  temperature.

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Planck's Constant

• Max Planck found that he could explain blackbody
  behavior by assuming that vibrating atoms emit
  energy in discrete amounts.
• The constant h is now known as Planck's constant.

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Photoelectric Effect

• The photoelectric effect occurs when light ejects
  electrons from the surface of a metal.
• No electrons are ejected if the frequency of the
  light is below a critical value, known as the
  threshold frequency.

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Particle Properties of Light

• In both cases, the observed results can be
  explained by assuming that the energy of an EM
  wave comes in discrete lumps.
• In effect, light consists of particles called
  photons, with energy proportional to their
  frequency

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Wave particle Duality

• It appears that neither a particle nor a wave
  model can completely describe light (EM waves).
• Light behaves like a wave when moving through
  space, but interacts with matter like a particle.
• This dual behavior turns out to be characteristic
  of everything.

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Waves and Probability

• The "wave" associated with a "particle"
  determines the probability of finding the
  particle at a particular location at a particular
  time.

High Probability Low Probability
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DeBroglie Relations

• For material objects (such as electrons), the
  wavelength is inversely proportional to the
  product of mass and velocity.
• This wavelength becomes important when the mass
  of the object is very small.

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When Wave Properties are Important

• Because Planck's constant is so small, the wave
  properties of ordinary objects cannot be
  detected.
• Electrons and atoms are very small, however, and
  its wave properties must be considered to
  understand atoms.

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The Bohr Model

• The Bohr model of an atom is based on
  Rutherford's model, but takes the wave properties
  of the electron into account.
• It can explain not only why electrons don't crash
  into the nucleus, but also the existence of
  discrete atomic spectra.

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Orbit Constraints

• Electrons orbiting an atom are subject to two
  fundamental constraints.
• The first of these is comes from Newton's laws,
  the second from quantum physics.
• The net result of these is that the orbits are
  quantized, that is, the possible orbit sizes are
  separated by finite amounts.

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Force Condition

• For any radius, only one speed will maintain the
  electron in a circular orbit.

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Quantum Condition

• The wave function of the electron must have a
  wavelength that allows a whole number of wave
  cycles to fit around the orbit.

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Results

• Possible orbits have radii that are whole number
  multiples of the smallest possible radius.

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Energy Levels

• In a hydrogen atom, the energy of an orbiting
  electron is determined by its orbit number
  (principal quantum number).

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Emission Spectra

• When an electron moves from one orbit to another
  of lower energy, it emits a photon whose energy
  is the difference between those levels.

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Atomic Spectra

• Every element emits a unique set of EM
  frequencies.
• These frequencies are determined by the energy
  changes as electrons move from one allowed orbit
  to another.

Hydrogen Spectrum
Hydrogen has four visible lines.
Lithium Spectrum
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Absorption Spectra

• To move into an orbit with a higher energy, the
  electron must absorb a photon whose energy is
  equal to the difference in the levels.

Absorption Spectrum of Hydrogen
Absorption Spectrum of Mercury