Transcript Quantum Numbers - MullisChemistry

Atomic Structure
Summary of Contributions
• Max Planck
– A hot object emits energy in small, specific amounts
called quanta.
• Albert Einstein
– E = mc2
– Light behaves as both a wave and a particle.
– Each particle of light carries a photon ( a quantum of
energy).
• Niels Bohr
– Created a model of the atom that showed a single
electron of hydrogen orbits the nucleus only in
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allowed orbits with a fixed energy.
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Atomic Structure
Summary of Contributions,
continued
• Werner Heisenberg
– Uncertainty principle: It is impossible to
determine simultaneously both the position and
velocity of an electron.
• Erwin Schrödinger
– Helped lay the foundation for quantum theory
with an equation that treats electrons like
waves.
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Electromagnetic Spectrum
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Solar Radiation: Infrared
• 45% of solar radiation is in the infrared (IR) region of the
electromagnetic (EM) spectrum.
• IR causes molecules to vibrate or rotate faster. Faster
movement of molecules is perceived by us as higher
temperature in any material that received the energy.
• CO2 and H2O vapor absorb most of the sun’s IR radiation.
• Some of the higher-energy solar radiation will be absorbed and
then reflected back—it is “re-radiated” away from the surface
of the Earth and is captured in the Earth’s atmosphere.
Solar Radiation: Ultraviolet (UV)
• 9% of solar radiation is in the UV part of the EM
spectrum. Based on the actual wavelength, UV is
divided into 3 types.
• UV-A (>320 nm): Lowest energy and is most
likely to be absorbed by a layer of ozone high in
the atmosphere.
• UV-B (280-320 nm): Much of this is also
absorbed by the ozone layer. Causes sunburn.
• UV-C (<280 nm): Highest energy and is absorbed
in the atmosphere. Can be used to kill viruses and
bacteria.
Wave Description of Light
• Electromagnetic radiation is a form of energy that
exhibits wavelike behavior as it travels through space
• Wavelength (λ)
– Distance between corresponding points on adjacent waves.
– Unit: nm,cm,m
• Frequency (ν)
– Number of waves that pass a specific point in a given time
– Unit: Hz or waves/sec
Recall that Speed = Distance/time (m/sec)
• Speed of light (c)
C=λν
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Behavior of Light
• Photoelectric effect
– The emission of electrons when light shines on
the metal
– Scientists found that below a certain frequency,
no electrons were emitted.
– Light also behaves as a particle: Since hot
objects do not emit em energy continuously,
they must emit energy in small chunks called
quanta.
• Quantum
– Minimum quantity of energy that can be gained
or lost by an atom
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Light as a particle and a wave
Planck and Einstein
• Max Planck: Relationship between quantum of
energy and wave frequency
• Planck’s constant h = 6.626 x 10-34 J-s
E = hν
E is energy, ν is frequency
• Albert Einstein: Established dual wave-particle
nature of light 1st
– Einstein explained PE effect by proposing that EM
radiation is absorbed by matter only in whole numbers
of photons.
– Electron is knocked off metal surface only if struck by
one photon with certain minimum energy.
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Quantum Theory
• Ground state: An atom’s lowest energy state
• Excited state: Higher potential energy than ground
state.
• Photon: A particle of electromagnetic radiation
having zero mass and carrying a quantum of energy
(i.e., packet of light)
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Niels Bohr links hydrogen’s
electron with photon emission
• Bohr proposed that an electron circles the
nucleus in allowed orbits at specific energy
levels.
– Lowest energy is close to nucleus
• Bohr’s theory explained the spectral lines
seen in hydrogen’s line emission spectrum,
but it did not hold true for other elements.
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Electron Configuration: The Rules
• Aufbau principle
– An electron occupies the lowest energy orbital that will
receive it.
• Pauli exclusion principle
– No two electrons in the same atom can have the same set
of 4 quantum numbers.
– Therefore, electrons can pair in an orbital as long as their
spins are opposite.
• Hund’s rule
– Each of the orbitals at a particular level have one electron
before any of them can have two electrons.
– All single electrons in the orbitals at a particular level
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have the same spin.Mullis Chemistry Holt Ch.4
Label and color the electromagnetic spectrum (See p.92).
Position the energy and wavelength as shown.
700 nm
400 nm
Increasing Energy and Frequency(ν )
104m
10 cm
103 nm
400-700nm
10nm
1 nm
10-2 nm
Wavelength, (λ)
(Larger wavelengths to left)
Include the following:
Visible spectrum (ROYGBIV)
Gamma Rays
Radio Waves
X- rays
Microwaves
Infrared (IR)
Ultraviolet (UV)
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Vocabulary
Photon_______________________________
__________________________________
Quanta_______________________________
__________________________________
Excited state of an electron
_____________________________________
Ground state of an electron
_____________________________________
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Indicate meaning (and value if a constant) by
each symbol.
λ____________________________________
ν____________________________________
c___________________________________
(This is a constant)
h____________________________________
(This is a constant)
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Write the main contribution of
each scientist for this topic.
Bohr
Aufbau
Pauli
Hund
Einstein
Heisenberg
Planck
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