Ch. 13 Electrons in Atoms

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Ch. 13 Electrons in Atoms

• Ch. 13.1 Models of the Atom
• Ch. 13.2 Electron Arrangement in Atoms
• Ch. 13.3 Physics and the Quantum
• Mechanical Model

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Ch. 13.1 Models of the Atom

• Evolution of Atomic Models
• John Dalton atomic theory
• JJ Thomson plum pudding model
• Ernst Rutherford - nuclear atom
• Niels Bohr planetary model
• Fixed energy levels
• Quantums of energy
• Schrodinger quantum mechanical model

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Ch. 13.1 Models of the Atom

• The quantum mechanical model
• Primarily mathematical model
• Restricts energy of electrons
• Estimates probability if finding an electron in a
  certain position
• 90 probability

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Ch. 13.1 Models of the Atom

• Atomic orbitals
• Designates energy levels of electrons with
  principal quantum numbers (n)
• n1, n2, n3, n4 and so on
• The principal energy levels have a specific
  number of sublevels
• The sublevels are designated s, p, d, and f
• Sublevels have specific numbers of orbitals
• The orbitals have specific shapes that correspond
  to the path of the electrons
• Nodes are areas where the probability of finding
  electrons is low

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Ch. 13.2 Electron Arrangement in Atoms

• Electron Configurations
• The ways in which electrons are arranged around
  the nucleus of an atom
• Three rules explain how to find the electron
  configuration of atoms
• Aufbau principle e- enter the lowest energy
  level first
• Pauli exclusion principle an atomic orbital may
  hold at most 2 e- (with opposite spin)
• Hunds rule when e- occupy orbitals of equal
  energy, one e- enters each orbital until all the
  orbitals contain one e- with parallel spins
• Exceptional electron configurations Cr, Cu

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Ch. 13.3 Physics and the Quantum Model

• Light and atomic spectra
• Electromagnetic radiation
• Amplitude height of wave from origin
• Frequency measured in hertz (Hz)
• Wavelength distance between crests (nm)
• Speed speed of light (a constant, c)
• v c/l
• Frequency speed/wavelength
• Visible spectrum (ROY G BIV)
• Atomic emission spectra (fingerprint)

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Ch. 13.3 Physics and the Quantum Model

• The quantum concept and the photoelectric effect
• Max Planck
• Amount of energy absorbed or emitted is
  proportional to the frequency of the radiation
• E hv (energy Plancks constant x frequency)
• h 6.6262 x 10-34 Js
• Albert Einstein
• Light can be described as quanta of energy
  (photons)
• Dual nature of light (waves and particles)
• First to explain the photoelectric effect
• Energy must meet a threshold value to eject e-
  from the surface of metal

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Ch. 13.3 Physics and the Quantum Model

• An explanation of atomic spectra
• When energized, e- move from the ground state to
  an excited state (n 2,3,4, etc)
• The same amount of energy that was absorbed is
  then emitted as light and the e- falls back down
  to a lower energy level
• Results in Lyman, Balmer and Paschen series
• An upper limit for the frequency of light emitted
  exists because a very excited e- will escape the
  atom
• Bohrs model of the atom was not able to explain
  bonding and was eventually replaced

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Ch. 13.3 Physics and the Quantum Model

• Quantum mechanics
• De Broglies equation predicts that all matter
  exhibits wave-like motions
• True for all matter, but motion depends on the
  size of the object
• Wavelengths for objects visible to the naked eye
  are not measurable
• Wavelengths for extremely small objects are
  measurable
• Heisenbergs uncertainty principle
• It is impossible to know exactly both the
  position and velocity of a particle at the same
  time
• The more precise the measured velocity, the less
  precise the position, and vice versa