Bohr Model of the Atom

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Bohr Model of the Atom

• Why are the emission spectra of elements not a
  continuous spectrum?
• In 1913, a Danish physicist named Niels Bohr
  tried to discover the answer, using hydrogen.
• Bohr proposed that hydrogen had only certain
  allowable energy states.

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Hydrogens Energy States

• The lowest energy state of an atom ground state.
• When the atom gains energy excited state.
• Bohr said hydrogen, even with only 1 e-, is
  capable of many different excited states and
  these are related to the orbit of the e- in the H
  atom.

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Hydrogens Energy States

• Bohr said that the single e- of hydrogen orbited
  the nucleus in only certain allowed circular
  orbits.
• The smaller the orbit, the lower the energy state
  (energy level) the larger the orbit, the higher
  the energy state (energy level)
• Bohr assigned a quantum number, n to each orbit.
  The one closest to the nucleus he numbered 1,
  the next 2, etc.

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An explanation of Hydrogens energy states

• In the ground state, (n 1) the atom does NOT
  emit energy. Normal state.
• But when excited, the e- jumps up to a higher
  energy level (n 2, or 3)
• When this excited e- drops back to the ground
  state, it emits a photon equal to the difference
  in energy between the 2 energy levels.

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Explanation, continued

• Because only certain orbits are possible, only
  certain frequencies of EM radiation can be
  emitted.
• Think of the energy levels as rungs on a ladder
  you can only go up or down on the rungs, NOT
  in-between.

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ORIGIN OF THE LINES IN THE HYDROGEN EMISSION
SPECTRUM
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The problem with Bohr

• Bohr explained hydrogens spectral lines almost
  perfectly.
• However, it did not explain any other element!
• It also did not explain other chemical behavior.
• And though his work was groundbreaking, later
  experiments proved his model fundamentally wrong.
  Electrons dont move around the nucleus in
  circular orbits.
• Alas, poor Bohr!

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Enter Louis de Broglie

• In the 1920s, French graduate student Louis
  deBroglie asked, if waves (light) can behave
  like particles, can particles (electrons) behave
  like waves?

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The de Broglie Equation

• deBroglie eventually derived an equation to
  explain the wave properties of particles.
• ? _h_
• mv
• ? wavelength of particle
• h Plancks constant
• mmass of particle
• v velocity

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The de Broglie Equation

• This equation proves that all moving particles
  have a particular wave property
• Why cant these waves be seen, like light?
• They simply are too small. Even sensitive
  instruments cant detect these waves.
• When large objects move, when their large mass is
  divided into the small number of Plancks
  constant, the resulting wavelength is very, very
  tiny.
• Electrons, on the other hand, have almost NO
  mass, so their wavelength is easily measured.

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Heisenberg Uncertainty Principle

• The de Broglie equation ultimately was proven
  correct.
• The next scientist to contribute to our
  understanding of the atom was German physicist
  Werner Heisenberg.
• Heisenberg said that its impossible to make a
  measurement on an object without disturbing it
  at least a little bit.

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Heisenberg Uncertainty Principle

• The smaller the object, the more it is disturbed
  by measurement.
• For an electron, even shining light on it will
  disturb its frequency and position.
• Therefore, Heisenberg concluded, it is
  impossible to know precisely both the velocity
  and position of a particle at the same time.
• Scientists of the time found it hard to accept,
  but later tests proved it true.
• It shows that the uncertainty of any electrons
  position is VERY large.

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Schrodingers Wave Equation

• In 1926, Austrian scientist Erwin Schrodinger
  further refined de Broglies equation.
• He developed a very complex equation that treated
  the hydrogen electron as a wave. It worked
  perfectly for not only hydrogen, but ALL
  elements.
• It is the basis for the current model for
  electrons in atoms, the quantum mechanical model.

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Basics of Quantum Mechanics

• The Schrodinger equation defines the basis of
  energy levels and possible locations of electrons
  in atoms.
• An electrons location is simply a probability of
  where it can be located at a given time.
• To make probability predictions easier, quantum
  mechanics assigns numbers to energy levels and
  shapes to sublevels.

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Basics of Quantum Mechanics

• Principal Quantum number (n) relative size
  energy of the atomic orbitals (where an electron
  can orbit the nucleus)
• As n increases, the orbital gets larger, the e-
  spend more time away from the nucleus, the
  energy level increases.
• Therefore, the Principal Quantum number
  represents Principal Energy Levels.
• The lowest energy level is assigned the number 1.
  Up to 7 levels are possible for hydrogen.

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Energy Levels
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Sublevels

• Principal Energy levels contain sublevels (also
  called orbitals).
• Energy level 1 has 1 sublevel, level 2 has 2
  sublevels, 3 has 3 and so on.
• The sublevels are named for their shapes, as
  follows sspherical pdumbbell d varying
  shapes (mostly lobes) and f varying shapes
  (like flower petals, almost)
• Each orbital can hold up to TWO electrons,
  maxium.

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