Electronic Structure of Atoms

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• Electronic Structure of Atoms
• Chemistry involves the study of the interactions
  of atoms and molecules
• These interactions are at the outer regions of
  atoms and molecules
• The outer part of atoms and molecules contain the
  electrons
• Thus, a study of the details of how the electrons
  are organized in atoms is very important to
  understanding chemistry
• Many of the phenomena weve discussed are
  associated with how the electrons in atoms are
  arranged
• The arrangement of the elements in the periodic
  table
• The stoichiometry of ionic and molecular
  compounds
• The geometric arrangement of atoms in molecules
• Many physical properties of substances
• The chemical properties of substances
• Quantum Theory gives the current picture of how
  electrons are arranged in atoms and molecules
• Experimental basis of quantum theory comes from
  how light interacts with matter and how
  subatomic particles behave when they move at high
  speeds.

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• Electronic Structure of Atoms
• Wave Nature of Light
• One viewpoint about light is that it has a wave
  nature
• There is an oscillating electric field
  perpendicular to an oscillating magnetic field
• Both fields oscillate perpendicular to the
  direction of travel
• The wave travels through a vacuum at velocity c
  3.00 x 108 m/s
• The wavelength, l,is the distance between two
  successive points on the wave of maximum
  amplitude
• The frequency, n, of light is the number of
  oscillations the wave makes in 1 s.
• nl c

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• Electronic Structure of Atoms
• Wave Nature of Light
• Various parts of the Electromagnetic spectrum
  have different wavelengths
• Radio frequencies l0.1 m
• Microwaves 0.001 m l 0.1 m
• Infrared frequencies 10-6 m l 10-3 m
• Visible light 400 x 10-9 m l 750 x 10-9 m
• Ultraviolet frequencies 10-8 m l 350 x 10-9 m
• X-rays 10-11 m l 10-8 m
• g-rays l
• Scientists often use units other than meters for
  wavelength
• 1 Angstom (Å) 10-10 m for X-rays
• 1 nm 10-9 m 0.1 Å for UV and visible
• 1 ?m 10-6 m for IR
• 1 cm 10-2 m for microwaves
• m for radiowaves

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• Electronic Structure of Atoms
• Frequency-wavelength conversions
• Example What is the frequency of yellow visible
  light having wavelength 598 nm?

• Standing waves, as opposed to travelling waves,
  have endpoints that are fixed in space
• Standing waves have 2 or more nodes - points
  along the wave with zero amplitude
• If a is the distance between the end nodes, the
  only possible wavelengths occur for

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a
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• Electronic Structure of Atoms
• The Quantum Nature of Energy
• The explanation of amplitude distribution of the
  frequencies of light emitted by black body
  radiators and the temperature effect on this
  distribution was explained by Max Planck in
  1900.
• Planck postulated that energy exists in fixed
  quantities or quanta.
• The energy of a quantum is
• E hn h6.63 x 10-34 J s Plancks Constant
• energy is produced or consumed only in
  integer multiples of hn
• Example What is the energy of 1 quantum of UV
  light having l 200 nm and how much energy is
  in one mole of these quanta?

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• Electronic Structure Atoms
• The Photoelectric Effect
• When light is incident on some substances,
  including some metals, electrons can be emitted
  from the surface
• The surface is negatively charged and there is
  nearby a positive electrode to capture the
  emitted electrons.
• For any such photoemissive material, light below
  a particular frequency will not stimulate the
  emission of electrons no matter how bright the
  light.
• Albert Einstein explained this phenomenon
• He assumed light consisted of energy packets
  called a photon.
• The energy of the photons is E hn.
• In order to expel the electron, a certain amount
  of the photons energy is used to overcome the
  attractive forces - the work function, w, - of
  the electron for its solid matrix.
• If the photon does not have enough energy - n is
  too small - the electron will remain in the
  metal.
• Photo electrons will have kinetic energy Eke hn
  - w
• This creates an apparent paradox - it seems
  impossible to reconcile the wave
  point of view and the particle theory as
  explanations for the nature of light

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• Electronic Structure of Atoms
• Line Spectra light emitted by energetic atoms in
  the gas phase
• The light is produce only at discrete wavelengths
  unlike light from a continuous source that
  produces a continuous rainbow of colors
• H atom spectrum

l, nm
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Electronic Structure of Atoms Balmer showed a
mathematical relationship between n and n for the
visible lines Additional line spectra are known
for other parts of the electromagnetic spectrum
which have a similar mathematical
description Rydberg Equation
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• Electronic Structure of Atoms
• Bohrs Model of the H atom
• Postulated that the electron moved in a circular
  orbit about the proton
• Postulated that the electrostatic force of
  attraction was counterbalanced by the
  centrifugal force associated with the electron
  moving in the circular orbit.
• Postulated that the electron could have only
  certain energies in the atom the energy of an
  electron in an orbit is quantized
• This turns out to be the same as fixing the radii
  of the orbits
• The electron can only have an allowed energy
  state
• Bohr showed the energies of these orbits are
• where n is called a quantum number representing a
  different electron orbit
• Rhc 2.18 x 10-18 J

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• Electron Structure of Atoms
• Bohrs Model of the H Atom
• The radius of each orbit increases with n2
• for n1, we have the smallest Bohr Orbit
• for n2, we have the next orbit which is 4x as
  big as the n1 orbit
• Notice the negative energies calculated by the
  Bohr equation
• for n1, E-Rhc
• for n2, E-3Rhc
• for n, E0
• Bohr postulated that light is emitted when an
  electron goes from an orbit with a high n value
  to a lower n value
• where i refers to the initial state and f to the
  final state

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Electronic Structure of Atoms Energy Level
Diagrams
O
n
-9Rhc
n4
-8Rhc
n3
IR
-3Rhc
n2
vis
Energy
n1
-Rhc
UV
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n