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Bohr Model
 Nucleus: Neutrons and
Protons
 Orbits: Electrons
 We know both specific
energy and location of
each electron
Nucleus
Energy
Levels
 Electrons orbit the nucleus
in certain fixed energy
levels (or shells)
Bohr Model
 Bohr’s Atomic Model
of Hydrogen
 Bohr - electrons exist in
energy levels AND
defined orbits around
the nucleus.
 Each orbit corresponds
to a different energy
level.
 The further out the
orbit, the higher the
energy level
Bohr’s Model
 The Photoelectric Effect
 Light releases electrons
 Not all colors work
 Atomic Emission Spectra
 Hydrogen gas emitted
specific bands of light
 Bohr’s calculated
energies matched the IR,
visible, and UV lines for
the H atom
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Electromagnetic Radiation
 Photoelectric Effect – There is a
minimum frequency to eject the electron
Electromagnetic Radiation
 Photoelectric
Effect
 Only explained by “energy
packets” of light called a
quantum
 Quantum - minimum
amount of energy that can be
gained or lost by an atom
 Photons are massless
particles of light of a certain
quantum of energy
 Based on the frequency and
wavelength of the photon
Bohr’s Model
 Excited electrons
 Energy added to atom –
electrons “jump” up
energy levels
 When the atom relaxes electron “falls” to lower
energy levels and emits
photon
 Bohr Model of hydrogen
 Reference Sheets!!!!!
Electromagnetic Radiation
 Atomic Line
Spectra
 Electrons in an atom add
energy to go to an “excited
state”.
 When they relax back to
the ground state, they
emit energy in specific
energy quanta
Electromagnetic Radiation
 These observations suggested that electrons
must exist in defined energy levels
First, the electron absorbs energy
and jumps from the ground state
to an excited state
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Next, the excited electron
relaxes to a lower excited
state or ground state
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Electromagnetic Radiation
 Wave Nature of Light
 Property of Waves
 Frequency
No. of waves per
second
 Wavelength
▪ Distance between
corresponding points
in a wave
 Amplitude
▪ Size of the wave peak
▪
Electromagnetic
Radiation
 Mathematical Relations
C=λf
• C = speed of light = 3.0 x 108 m/s
• λ (lamda) = wavelength (m)
• f = frequency (Hz or s-1)
 This is how we know what color light is
emitted!
Frequency is inversely proportional
to wavelength
If λ increases, f decreases.
If f increases, λ decreases.
Speed of the wave is always constant at
3.0 x 108 m/s.
Electromagnetic Radiation
 Particle
Nature of
Light
 Wave nature could not
explain all observations
(Planck & Einstein)
Photoelectric Effect

E = hf
When light strikes a metal
electrons are ejected
Atomic Line Spectra
▪
When elements are
heated, they emit a unique
set of frequencies of visible
and non-visible light.
Other Scientists Contributions
 de Broglie
 Heisenberg
 Modeled electrons as
waves
 Heisenberg Uncertainty
Principle: states one
cannot know the position
and energy of an electron
 Electrons exist in
orbital’s of probability
 Orbital - the area in
space around the nucleus
where there is a 90%
probability of finding an
electron
Other Scientists Contributions
Schrödinger
 Schrödinger Wave
Equation - mathematical
solution of an electron’s
energy in an atom
 Quantum Mechanical
Model of the atom –
current model of the atom
treating electrons as waves.
Quantum Mechanical Model
 Nucleus: Neutrons and
protons
 Orbitals: region in space
surrounding the nucleus
where there is a 95%
probability of finding an
electron.
 We know either energy or
location of each electron.
Reflection
 Who contributed to the modern model of the
atom? How is it different from Bohr’s?
 Why do atoms give unique atomic line
spectra?
 What are ground and excited states?