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Ch. 5
Atomic Models
• Dalton –
• Thomson –
• Rutherford –
Previous models could not
explain the chemical properties
of elements.
• Why do different elements give off different
colors when heated?
• Why does the color change when more heat
is added to an element?
Bohr Model
• Bohr proposed that an electron is found
only in specific circular paths, or orbits
around the nucleus.
• Electrons are restricted to fixed
(quantized) orbits
• Electrons can jump orbits by absorbing or
emitting a photon that has a specific
wavelength of energy.
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The Bohr Model
Each possible electron orbit in Bohr’s
model has a fixed energy.
–The fixed energies an electron can
have are called energy levels.
–A quantum of energy is the
amount of energy required to
move an electron from one energy
level to another energy level.
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The
Bohr
Model
.
1Like the rungs of the
strange ladder, the
energy levels in an atom
are not equally spaced.
The higher the energy level
occupied by an electron,
the less energy it takes to
move from that energy
level to the next higher
energy level.
The Quantum Mechanical Model
–The quantum mechanical
model determines the allowed
energies an electron can have
and how likely it is to find the
electron in various locations
around the nucleus.
The Quantum Mechanical Model
• The propeller blade has the same probability
of being anywhere in the blurry region, but
you cannot tell its location at any instant. The
electron cloud of an atom can be compared to
a spinning airplane propeller.
Modern Model of the Atom:
Quantum Mechanical Model
Erwin Schrodinger developed an equation
to describe the energy of electrons
“I don't like it, and I'm sorry I
ever had anything to do with it.”
(Erwin Schrodinger talking about Quantum
Physics)
1887-1961
Atomic Orbital: a region of space in
which there is a high probability of
finding an electron. (S, P, D, F)
S sublevel: sphere shaped
Each S sublevel holds 1 orbitals
P sublevel: dumbbell shaped
Each P sublevel holds 3 orbitals
D sublevel: clover leaf shape
Each D sublevel holds 5 orbitals
F sublevel: complex shape
Each D sublevel holds 7 orbitals
(Don’t worry, I’ll explain this later)
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Atomic Orbitals
.
The numbers and kinds of atomic orbitals depend on
1 the energy sublevel.
5• The number of electrons allowed in each of
. the first four energy levels are shown here.
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Electron Arrangement in Atoms
Chapter 5 section 2
Electron Configuration
• The way electrons are arranged in various
orbitals around the nuclei of an atom
• Example: Arsenic
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3
Or (orbital diagram)
1s __ 2s __ 2p __ __ __3s___ 3p__ __ __
4s__ 3d __ __ __ __ __ 4p__ __ __
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Electron Configurations
• Orbital Filling Diagram
3 RULES FOR WRITING THE
ELECTRON CONFIGURATION
Aufbau Principle
Electrons occupy the lowest energy
level first.
• The first energy level is the lowest energy
level
Pauli Exclusion Principle
• An atomic orbital may describe and hold
at most two electrons.
• To occupy the same orbital electrons must
have opposite spins; clockwise and counter
clockwise
• Example 1s___
Hund’s Rule
• Electrons occupy orbitals of the same
energy level so that one electron enters each
orbital until all the orbitals contain one
electron with the same spin.
• Example 3p __ __ __
Exceptions to the Electron
Configurations
• Chromium
• 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1
• Copper
• 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1
Section 3
Physics and the Quantum
Mechanical Model
c=λf
• The wavelength and frequency of light are
inversely proportional to each other.
• c = the speed of light 3.00 x 108 m/s
• λ = wavelength in meters
• f = frequency in Hertz (s-1 or cycles/second)
All electromagnetic waves travel
in a vacuum at 3.00 x 108 m/s
What color
of light has
the highest
frequency?
• The amount of radiation emitted at each
wavelength depends on the temperature of the
object. Hot objects emit more of their light at
short wavelengths, and cold objects emit more
of their light at long wavelengths.
700 nm
400 nm
• The amount of light produced at each
wavelength depends on the temperature of the
object producing the light.. Solid objects
heated to 1,000 degrees C appear red but are
putting out far more (invisible) infrared light.
Max Planck analysis
• The color of light from a hot glowing solid
varies with temperature.
• There is a relationship of energy in a solid
and the wavelength of light emitted
• Energy varies by small whole numbers
• Any frequency of light should emit photons.
(not so)
Einstein
• Light travels in packets of energy (photons)
in waves
• Light energy can be both particles and
waves
• Quantum – minimum amount of energy that
can be lost or gained by an atom.
Planck’s Equation
(with the help of Einstein)
The energy of a photon is described by the
following equation.
E=fh
• E: joules
• f : frequency
• h : Planck’s constant
6.626x 10-34 j-s
Each frequency of
light has its own
specific energy
per photon and no
other.
Atomic spectra
• Neils Bohr (1913) explains…
• When atoms absorb energy, electrons move
into higher energy levels.
• These electrons then lose energy by
emitting light when they return to lower
energy levels.
Bright line spectrum
Consists of several distinct lines
of color, each with its own
frequency.
Every element has a different
bright line spectrum.
Modern Model of H Atom
Improving Bohr’s model
Louis deBroglie (1923)
Experimentally confirmed
Einstein and Planck’s theory
that particles (photons) could
Have properties of waves.
Won the Nobel Prize in 1929
The work done by de Broglie
and Schrodinger became the
branch of physics known as
Quantum
Mechanics
It is often stated that of all the theories proposed in this century, the silliest is
quantum theory. Some say the only thing that quantum theory has going for
it, in fact, is that it is unquestionably correct. - R. Feynman
Heisenberg Uncertainty Principle
Werner Heisenberger
It is impossible to know exactly
both the velocity and position of
a electron (particle) at the same
time.
Orbitals and Energy Levels
Every energy level has a specific
number of orbitals.
Energy level
# of orbitals
2
n
n
1
1
2
4
3...
9...
4 quantum numbers used to
describe an electron in an orbital
1. n = principle quantum number.
It describes the size of the
orbital.
2. l = designates the shape of the
orbital. s = spherical
p = double-lobed
4 quantum numbers used to
describe an electron in an orbital
3. m = describes the orbitals
orientation (around x-axis,
y-axis, or z-axis).
4. Designates the spin of the
electron.