Transcript Quantum Theory

Quantum Theory
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Schroedinger’s Cat
Place a cat in a box
Also place a radioactive
isotope and a vial of poison
The isotope decays once per
hour
If the particle triggers a
Geiger counter, the cat dies
If the Geiger counter is not
triggered, the cat lives
Seal the box and wait an hour
What happened to the cat?
Electromagnetic Spectrum
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The speed of light©
is 3.00 x 108 m/s
Photoelectric Effect
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The emission of electrons
from a metal when light
shines on the metal
Light had to be a certain
minimum frequency for
electrons to be emitted
Wave theory of light said that
any frequency of light should
have worked
This led to the concept of
light as a particle
Light as a particle
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Max Planck (1900)
Hot objects emit light and
other forms of
electromagnetic radiation,
but not continuously, as
expected
Instead, it is emitted in small,
specific amounts called
quanta.
E=hv
H=6.626 x 10 -34 J*s
This number is called
Planck’s constant
Dual Nature of Light
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Albert Einstein (1905)
Light behaves as a wave when it
travels through space
Light behaves as a particle when
interacting with matter
Even though it isn’t quite right to
do so, you can think of light as a
stream of particles that travels as a
wave
A photon is a massless bundle of
light
E photon=hv
Some metals hold electrons more
tightly than others and require
higher frequencies to move
electrons
Emission Spectra
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Ground State
Lowest energy state of an electron
Excited State
Higher energy state
Energy is absorbed to move from
ground to excited states and is
released as light(photons) when
returning to the ground state
Specific patterns of light are
emitted for any given element
Continuous spectrum-continuous
range of em light(rainbow)
(Bright) Line emission spectrumOnly certain wavelengths of light
are seen
The Bohr Model
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Niels Bohr (1913)
Allowed for electrons to have
orbits
Electrons have fixed energies
Lower energy-closer to nucleus
Higher energy-farther from nucleus
Electrons can gain energy to raise
to the next energy level and release
the same energy when falling to the
ground state
This model works well for
hydrogen but not for other
elements
Electrons as Waves
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Louis de Broglie (1924)
Suggested that electrons
could be considered as
waves confined to the
space around a nucleus
Diffraction- bending of
light around edges
Wave Interference- when
waves overlap
Quantum Theory
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Werner Heisenberg
(1927)
The Uncertainty
Principle
One cannot
simultaneously know
the position and
velocity of an electron
Quantum Theory
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Erwin Schroedinger (1926)
Wave Equation- The
quantization of an
electron’s energies is an
outcome
Quantum Theorymathematically describes
wave properties of
electrons and other small
particles
Atomic Orbitals and Quantum
Numbers
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Quantum Numbersproperties of atomic
orbitals and
properties of
electrons in orbitals
Like the address of an
electron or a seat in a
stadium or theater
Principal Quantum Number
(n) Main energy
level(shell)
1,2,3
K,L,M
This would be like the
section on a ticket for
a stadium seat
Angular Momentum Quantum
Number
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shape of an orbital
0,1,2,3
s,p,d,f
Like the row in a stadium
or theater
Magnetic Quantum Number
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m
Orientation of an orbital
around the nucleus
+1 to -1
Like finding your seat in
a stadium or theater
Spin Quantum Number
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Ms
Fundamental spin state of an
electron
+1/2 or -1/2
A single orbital can hold up to a
maximum of 2 electrons of
opposite spin
This would be which direction
you are facing in a theater or
stadium seat
Some Quantum Theory Rules
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Pauli Exclusion Principle- No two electrons
have the same set of four quantum numbers.
Or An orbital within a sublevel can contain up
to 2 electrons of opposite spin
Hund’s Rule- Each orbital within a sublevel
receives an electron of positive spin before any
can receive an electron of negative spin