Transcript chpt7QuantTheory

Chapter 7
Quantum Theory of the Atom
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Electromagnetic Radiation
Electromagnetic wave
 A wave of energy having a frequency within
the electromagnetic spectrum and
propagated as a periodic disturbance of the
electromagnetic field when an electric
charge oscillates or accelerates.
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Electromagnetic Radiation
Electromagnetic wave
 wavelength
 frequency
 amplitude
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Electromagnetic Radiation
nl= c
where
n => frequency
l => wavelength
c => speed of light
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Wave Nature of the Electron
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Electromagnetic Spectrum
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Line Emission Spectrum
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Line Spectrum
A spectrum produced by a luminous gas or
vapor and appearing as distinct lines
characteristic of the various elements
constituting the gas.
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Emission Spectrum
The spectrum of bright lines, bands, or
continuous radiation characteristic of and
determined by a specific emitting substance
subjected to a specific kind of excitation.
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Ground State
The state of least possible energy in a physical
system, as of elementary particles. Also
called ground level.
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Excited State
Being at an energy level higher than the
ground state.
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Photoelectric Effect

the emission of electrons by substances,
especially metals, when light falls on their
surfaces.
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Photoelectric Effect
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Quantum Mechanics
Quantum theory
 the theory of the structure and behavior of
atoms and molecules.
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Black Body Radiation
http://www.cbu.edu/~mcondren/C11599/BBvis.mov
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Electromagnetic Radiation
Ehi - Elo = hc/l
where E => energy
h => Planck's constant
c => speed of light
l => wavelength
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Photons
The quantum of electromagnetic energy,
generally regarded as a discrete particle
having zero mass, no electric charge, and an
indefinitely long lifetime.
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Dispersion of White Light
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The Atomic Spectrum of
Hydrogen
and the Bohr Model
Bohr Model for the Hydrogen Atom
mnr = nh/2p
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Bohr Atom
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Bohr Model
E = -B/n2
where n => quantum number
1, 2, 3, 4, 5, 6, 7, etc
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Bohr Model
for hydrogen
ground state: n = 1
excited state: n > 1
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Bohr Model
E = (-2.179 X 10-18 J/part.)
(6.022 X 1023 part./mole)
(1 kJ/103 J)/n2
= (-1312 kJ/mol)(1/n2)
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Electron Transition in a Hydrogen
Atom
Lyman series => ultraviolet
n > 1 ==> n = 1
Balmer series => visible light
n > 2 ==> n = 2
Paschen series => infrared
n > 3 ==> n = 3
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Line Spectra
See CHEMWORKS software
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Absorption Spectrum

Light shinning on a
sample causes
electrons to be excited
from the ground state
to an excited state
 wavelengths of that
energy are removed
from transmitted
spectra
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Knowing diamond is transparent, which
curve best represents the absorption
spectrum of diamond (see below)?
A, B, C
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According to the energy diagram below for
the Bohr model of the hydrogen atom, if an
electron jumps from E1 to E2, energy is
absorbed
emitted
not involved
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Heisenberg, Werner
1901–76, German physicist
1932 Nobel Prize in physics
A founder of QUANTUM MECHANICS, he
is famous for his uncertainty principle,
which states that it is impossible to
determine both the position and momentum
of a subatomic particle (such as the
electron) with arbitrarily high accuracy.
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Heissenberg Uncertainty
Principle
“it is impossible to determine both the position
and momentum of a subatomic particle
(such as the electron) with arbitrarily high
accuracy”
The effect of this principle is to convert the
laws of physics into statements about
relative, instead of absolute, certainties.
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Orbitals

region of probability of finding an electron
around the nucleus
 4 types => s p d f
 maximum of 2 electrons per orbital
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Pure Atomic Orbitals
s
p
d
f
shape
# of orbitals / energy level
spherical
1
dumbbell
3
complex
5
very complex
7
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Atomic Orbitals, s-type
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Atomic Orbitals, ptype
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Atomic Orbitals, d-type
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Atomic Orbitals, f-type
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Shapes of Orbitals
http://www.colby.edu/chemistry/OChem/DEMOS/Orbitals.html
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