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Electronic Structure
Wave Nature of Light
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Electromagnetic Radiation
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Gamma Rays, Visible Light
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Moves Through Vacuum at 3.00x108m/s (c)
Wavelength (λ, m) = distance between
successive peaks or troughs
Frequency (f, s-1) = how often a wave passes
through a particular point
c = f · λ, or v · λ
Wavelength Practice
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The brilliant red colors seen in fireworks are due
to the emission of light with wavelengths around
650nm when strontium salts such as Sr(NO3)2
and SrCO3 are heated. Calculate the frequency
of red light of wavelength 6.50x102nm.
A FM radio station broadcasts electromagnetic
radiation at a frequency of 103.4MHz. Calculate
the wavelength of this radiation. (1MHz=106s-1)
Quantize Energy and Photons
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Wave model explains much of the behavior of
light but not all:
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Black body radiation – Emission of light from hot
objects
Photoelectric Effect – Emission of electrons from
metal surfaces
Emission Spectra – Emission of light from excited
atoms
Black Body Radiation
Hot Objects and Quantization of
Energy
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When objects are heated they emit light
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Red hot (cooler) → white hot (hotter)
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Light only emitted at certain wavelengths
Max Planck declared that energy can only be
emitted or absorbed in packets (quanta, photon)
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E= h · v
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h = Planck's Constant 6.63x10-34J·s
Energy emitted at whole number multiples of hv
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Think walking up and down stairs
Why don't we notice this? Why does energy
Photoelectric Effect and Photons
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When light shines on an object, electrons are
emitted
Light has to have a specific energy, frequency
and wavelength in order for e- to be emitted
Photons are absorbed
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Too little energy – nothing happens
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Just right amount – electrons are emitted
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A little too much – electrons are emitted and excess
used as kinetic energy
Calculation Practice
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The blue color of fireworks is often achieved by
heating copper (I) chloride to about 1200°C.
Then the compound emits blue light having a
wavelength of 450nm. Calculate the frequency
and quantum of energy that is emitted at
4.50x102nm by CuCl.
Spectra
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Radiation emitted from a source contains
various λ's
When separated into its different λ's a spectrum
is formed
Two types of spectrum
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Continuous
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Line
Continuous Spectra
Line Spectra
Bohr Model
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Assumed electrons orbit nucleus in circular
patterns
To move between levels energy is absorbed or
emitted
Ground State – electron at lowest energy
Excited State – electron is at higher energy
state
Bohr model only accurately explains hydrogen
Wave Behavior of Matter
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Lights is both a wave and a particle
Louis de Broglie believed matter could have
wave properties
Applied idea to electrons
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λ=h/m·v
Works for all matter so why don't we observed
this in our everyday lives?
Uncertainty Principle
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If matter can act as a wave we should be able
to calculate position and velocity
Heisenberg determined that we cannot know
both position and velocity of subatomic matter.
Solving Schrödinger's equation gives use
probabilities of location.
The solutions correspond to the orbitals
Quantum Numbers
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n – principle quantum number
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Whole numbers – 1,2,3,....
l – azimuthal quantum number
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From 0 to n-1
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Determines shape of orbital
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0 → s = shape
1 → p = principle
2 → d = diffuse
3 → f = fundamental
Quantum Numbers Cont.
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ml – magnetic quantum number
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Goes from -l to l including 0
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Determines orientation of orbital
ms – spin quantum number
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Two values +1/2 and -1/2
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Determines spin of electron
No two electrons can have the same four
quantum numbers – Pauli Exclusion Principle
Possible Quantum Numbers
Quantum Numbers Example
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Which of the following sets of quantum numbers
are not allowed? For each incorrect set, state
why it is incorrect.
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n = 3, l = 3, ml = 0, ms = -1/2
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n = 4, l = 3, ml = 2, ms = -1/2
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n = 4, l = 1, ml = 1, ms = +1/2
Quantum Numbers Practice
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Which of the following sets of quantum numbers
are not allowed? For each incorrect set, state
why it is incorrect.
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n = 2, l = 1, ml = -1, ms = -1
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n = 5, l = -4, ml = 2, ms = +1/2
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n = 3, l =1, ml = 2, ms = -1/2
Quantum Number Practice
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What is the designation for the subshell with n =
5 and l = 1? How many orbitals are in this
subshell? Indicate the values of ml for each of
these orbitals.
Representation of Orbitals
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S – orbital is a sphere
Representation of Orbitals
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P Orbital
Representation of Orbitals
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D Orbital
Atoms with Multiple Electrons
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Shapes of orbitals
remain the same
Energy of orbitals
varies
For a given n, energy
increases as l
increases
Electron Configuration
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Governed by three rules
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Pauli Exclusion
Hund's Rule – for orbitals with same energy, lowest
energy is attained when the number of electrons
with the same spin is maximized
AUFBAU – Energy shells are filled from lowest
energy to highest energy
Orbital Diagram
Where Orbitals are Filled
Filling Order
Electron Configuration Example
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Write the electron configuration for the following
elements:
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Silicon
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Chromium
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Iodine
Electron Configuration Practice
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Write the electron configuration for the following
elements:
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Copper
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Sulfur
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Tin
Exceptions to Filling Rules
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Chromium
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Copper
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Silver
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Molybdenum
Homework
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2, 6, 10, 14, 20, 22, 26, 46, 52, 54, 60, 68
Electron Configurations