Transcript Bohr Model of the Atom

Electrons in Atoms
Chapter 5
What were the early steps in the
development of atomic theory?
• John Dalton – Billiard Ball Theory.
Atom was indivisible.
• J.J. Thomson – Plum Pudding Model.
Atom was composed of smaller
particles, including negative electrons.
• E. Rutherford – Nuclear Model.
Rutherford Model
• Nucleus contains all positive charge &
most of the mass.
• Nucleus is very small in volume - only
1/10,000th of atomic diameter.
• Electrons occupy most of the volume of
the atom.
Later Models
• N. Bohr – Planetary Model
• Schrodinger – Wave Mechanics
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Cathode Ray Tubes & Discovery of Electrons
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Movie of cathode ray tube
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Rutherford’s Experiment - 1911
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Discharge Tubes
(Neon Lights)
Problems with the Rutherford
Model
• Why doesn’t the electron spiral into the
nucleus?
• How are the electrons arranged?
• Why do different elements exhibit different
chemical behavior?
• How are the atomic emission spectra
produced?
Atomic Emission Spectra
• Put some gas in a glass tube and apply a
voltage across the ends. Produce light.
• The color of the light depends on the
gas in the tube.
• Every element produces its own unique
color.
The emission
spectrum of an
element is the
set of
frequencies (or
wavelengths)
emitted.
Atomic Emission Spectra: Line spectra
Why is the emission spectra
useful?
• We can use it to determine if a given
element is present in a sample.
(Identification!)
• We can use it to learn fundamental
information about atomic structure.
• Neon lights
Quick light review
• Energy of light = E = h
• In the visible spectrum,
– Blue light: shorter wavelength, higher
frequency, more energy
– Red light: longer wavelength, lower
frequency, less energy
Interlude: Electromagnetic
Radiation
All electromagnetic
radiation has a velocity
of 3.00 X 108 m/sec
c = 
Emission & Absorption Spectra of
Elements
Each line in the bright line spectrum
represents 1 electron jump.
The blue lines have shorter wavelength and
higher frequency, so they are light from
“bigger” jumps.
The red lines have longer wavelength and
lower frequency, so they are light from
“smaller” jumps.
Bohr Diagram
• Shows all the electrons in orbits or
shells about the nucleus.
n=3
n=2
n=1
E3
E2
E1
Bohr’s Model
•Electrons travel only in
specific orbits.
•Each orbit has a definite
energy.
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• The energy of the
electron must match the
energy of the orbit. So
the electron is only
allowed to have some
energies, not any energy.
NEW!
Bohr assigned a quantum number, n, to each orbit.
Bohr’s Model
•The orbit closest to
nucleus is the smallest
& has the lowest
energy. It has n = 1.
•Outer orbits hold
more electrons than
inner orbits.
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•The larger the orbit,
the more energy
associated with it.
Bohr’s Model
•Atoms emit radiation
when an electron jumps
from an outer orbit to
an inner orbit.
•Atoms absorb energy
when an electron jumps
from an inner orbit to
an outer orbit.
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Bohr Applet
•Outer orbits
determine atom’s
chemical properties.
Bohr Diagram
n=3
n=2
n=1
* Made-up Numbers!!!
E3 = 23*
E2 = 15*
E1 = 5*
What is the energy
change of the electron if
it moves from E1 to E3?
It must absorb 18 units of
energy.
From E2 to E1?
It must release 10 units of
energy.
Bohr Model
• Energy is absorbed when the electron
moves to a higher orbit, farther from
nucleus. Endothermic process.
• Energy is released when the electron
drops to a lower orbit, closer to nucleus.
Exothermic process.
Hydrogen Atom
Flame Tests
• Volume 2, CCA
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Na
Sr
Cu
Animation
The energy levels
get closer together
away from the
nucleus.
Larger orbits can hold
more electrons.
Emitted Light
• The energy of the emitted light, E = h,
matches the difference in energy
between 2 levels.
• Again, we don’t know the absolute
energy of the energy levels, but we can
observe how far apart they are from
each other.
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Tiger Graphic – potential energy
Tiger Graphic – electron orbits
Potential Energy
A ladder is often used
as an analogy for the
energy levels of an
atom.
But it’s a little bit
different – How?
Max Capacity of Bohr Orbits
Orbit
Max # of Electrons
1
2
2
8
3
18
4
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n
2n2
Ground State
• Each electron is in the lowest energy
orbit available. Lowest energy state of
an atom.
• The Bohr configurations in the
reference table are ground state
configurations.
Excited State
• Many possible excited states for each atom.
One or more electrons excited to a higher
energy level.
• We can give electron configurations for the
excited states as well as the ground state.
• You need to recognize excited state
configurations.
Excited State Configurations
• For the smaller elements, it’s easy.
– First level holds 2.
– Second level holds 8.
– If upper levels fill before the first or
second is full, it’s an excited state
configuration.
Some Excited States for Li
• Ground State of Li is 2-1.
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2-0-1
2-0-0-1
1-2
1-1-1
1-1-0-1
Each possible configuration
still has 3 electrons, but
now one electron has been
bumped up to a higher level.
Many other possibilities!
Excited State Configurations
•
Which is an excited state configuration
of Fluorine, F?
•
Fluorine is element 9.
a) 2-7 Matches g.s.
b) 2-8 10
c) 2-6
d) 1-8
8
Excited State Configurations
•
Which is an excited state
configuration of S?
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S is element 16.
a) 2-7-6 15
b) 2-8-6 G.S.
c) 2-7-8 17
d) 2-7-7
Excited State Configurations
Determine which of the following is an excited
state configuration of Manganese, Mn.
Manganese is element 25.
a) 2-8-13-1
24
c) 2-7-14-2
b) 2-8-13-3
26
d) 2-8-13-2 Matches g.s.
Excited State Configurations
• For the larger elements, the best thing
to do is compare the given configuration
with the ground state configuration in
the reference tables.
– Note: The configuration must have the
correct number of electrons for that
element.
– If the configuration matches the one in the
reference table, it’s ground state.
– If it does NOT match, it’s excited state.
Electron Transitions
• If an electron gains or absorbs a
specific amount of energy, it can be
excited to a higher energy level.
• If an electron loses or emits a specific
amount of energy, it moves down to a
lower energy level.
Emitted Light
• The energy of the emitted light, E = h,
matches the difference in energy
between 2 levels.
• Turns out, we don’t know the absolute
energy of the energy levels, but we can
observe how far apart they are from
each other.
Hydrogen has 1 electron, but it can make
many possible transitions.
Potential Energy
A ladder is often used
as an analogy for the
energy levels of an
atom.
But it’s a little bit
different – How?
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Tiger Graphic – potential energy
Tiger Graphic – electron orbits
Absorption & Emission
• We cannot easily detect the absorption
of energy by the electron.
• We can easily detect the emission of
energy by the electron. We can see the
photon that is kicked out.
Success of Bohr’s Model
• Bohr’s model could predict the
frequencies in the hydrogen emission
spectrum.
• Predicted correct size of H atom.
• Unfortunately, it didn’t work for
anything with more than 1 electron.
Which principal energy level of an
atom contains an electron with
the lowest energy?
a)
b)
c)
d)
n=1
n=2
n=3
n=4
What is the total number of
occupied principal energy levels in
an atom of neon in the ground
state?
a)
b)
c)
d)
1
2
3
4
What is the total number of fully
occupied principal energy levels in
an atom of nitrogen in the ground
state?
a)
b)
c)
d)
1
2
3
4
What is the total number of
electrons in a completely filled
fourth principal energy level?
a)
b)
c)
d)
8
10
18
32
Which atom in the ground state
has five electrons in its outer
level and 10 electrons in its
kernel?
a)
b)
c)
d)
C
Cl
Si
P
Kernel?
• Everything EXCEPT the valence
electrons.
• = the nucleus + ALL the INNER SHELL
electrons.
Which electron configuration
represents an atom in an excited
state?
a)
b)
c)
d)
2-8-2
2-8-1
2-8
2-7-1
Which electron configuration
represents an atom of Li in an
excited state?
a)
b)
c)
d)
1-1
1-2
2-1
2-2
The characteristic bright-line
spectrum of an atom is produced
by its
a)
b)
c)
d)
Electrons absorbing energy
Electrons emitting energy
Protons absorbing energy
Protons emitting energy