 # 10/2/2013 The Development of Atomic Models

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5.1
The timeline shoes the development of atomic
models from 1803 to 1911.
5.1
5.1
The timeline shows the development of atomic
models from 1913 to 1932.
The Development of Atomic Models
What was inadequate about Rutherford’s atomic
model?
1,
5.1
Rutherford’s atomic model could not explain
the chemical properties of elements.
Rutherford’s atomic model could not explain
why objects change color when heated.
The Bohr Model
What was the new
proposal in the Bohr
model of the atom?
Bohr proposed that an
electron is found only
in specific circular
paths, or orbits, around
the nucleus.
Niels Bohr
7 October 1885 – 18 November 1962
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5.1
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.
5.1
Like 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.
Lewis- The total number of electrons
represented in a Lewis structure is
equal to the sum of the numbers of
valence electrons on each individual
atom. Non-valence electrons are not
represented in Lewis structures.
Bohr The Bohr Model has an
atom consisting of a small,
positively-charged nucleus
orbited by negatively-charged
electrons.
The Quantum Mechanical Model
What does the quantum mechanical model
determine about the electrons in an atom?
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.
5.1
. Austrian physicist Erwin Schrödinger
(1887–1961) used new theoretical
calculations and results to devise and solve
a mathematical equation describing the
behavior of the electron in a hydrogen
atom.
The modern description of the electrons
in atoms, the quantum mechanical model,
comes from the mathematical solutions to
the Schrödinger equation
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5.1
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.
5.1
In the quantum mechanical model, the
probability of finding an electron within
a certain volume of space surrounding
the nucleus can be represented as a
fuzzy cloud. The cloud is more dense
where the probability of finding the
electron is high.
Different atomic orbitals are denoted by
Atomic Orbitals
How do sublevels of principal energy levels
differ?
An atomic orbital is often thought of as a
region of space in which there is a high
probability of finding an electron.
Each energy sublevel corresponds to an
orbital of a different shape, which
describes where the electron is likely to
be found.
letters. The s orbitals are spherical, and
p orbitals are dumbbell-shaped.
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Aufbau Principle
Representation of s, p, d atomic orbitals
Four of the five d orbitals have the same
shape but different orientations in
space.
Hydrogen Energy Levels
The s and p types of sublevel
Hydrogen Orbitals
Why does an H atom have so many orbitals and
only 1 electron?
An orbital is a potential space for an electron.
Atoms can have many potential orbitals.
s, p, d, f orbitals named for sharp, principal,
diffuse and fundamental lines on spectra.
Further orbitals designated alphabetically
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The number of electrons allowed in each of
the first four energy levels are shown
here.
s
p
f
f
d
f
g
g
d
g
g
Electron Spin
Close examination of spectra revealed
doublets
Need one more property to determine how
electrons are arranged
Spin – electron modeled as
a spinning like a top
Spin is the basis of magnetism
Hydrogen Energy Levels
The s and p types of sublevel
The numbers and kinds of atomic orbitals
depend on the energy sublevel.
1.
Rutherford's planetary model of
the atom could not explain
any properties of elements.
the chemical properties of
elements.
the distribution of mass in an atom.
the distribution of positive and
negative charges in an atom.
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Bohr's model of the atom proposed
that electrons are found
embedded in a sphere of positive charge.
in fixed positions surrounding the
nucleus.
in circular orbits at fixed distances from
the nucleus.
orbiting the nucleus in a single fixed
circular path.
2.
Electron Configurations
What are the three rules for writing the
electron configurations of elements?
The ways in which electrons are arranged
in various orbitals around the nuclei of
atoms are called electron configurations.
Three rules—the aufbau principle, the
Pauli exclusion principle, and Hund’s
rule—tell you how to find the electron
configurations of atoms.
3.
What is the lowest-numbered principal
energy level in which p orbitals are found?
1
2
3
4
Aufbau Principle
According to the aufbau principle,
electrons occupy the orbitals of lowest
energy first. In the aufbau diagram
below, each box represents an atomic
orbital.
Aufbau Principle
If this rock were to
tumble over, it would end
up at a lower height. It
would have less energy
than before, but its
position would be more
stable. You will learn that
energy and stability play
an important role in
determining how electrons
are configured in an atom.
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Different atomic orbitals are denoted by
letters. The s orbitals are spherical, and
p orbitals are dumbbell-shaped.
First, determine how many electrons are in the
atom. Iron has 26 electrons.
Arrange the energy sublevels according to
increasing energy:
1s 2s 2p 3s 3p 4s 3d …
Fill each sublevel with electrons until you have
used all the electrons in the atom:
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
The sum of the superscripts equals the atomic
number of iron (26)
Electron Configurations
and the Periodic Table
• The periodic table can be used as a guide for
electron configurations.
• The period number is the value of n.
• Groups 1A and 2A have the s-orbital filled.
• Groups 3A - 8A have the p-orbital filled.
• Groups 3B - 2B have the d-orbital filled.
• The lanthanides and actinides have the f-orbital
filled.
The electron configuration of an atom is a
shorthand method of writing the location of
electrons by sublevel.
We can use the periodic table to predict which
sublevel is being filled by a particular element.
The sublevel is written followed by a
superscript with the number of electrons in the
sublevel.
If the 2p sublevel contains 2 electrons, it is written
2p2
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Electron Configurations
and the Periodic Table
• The periodic table can be used as a guide for
electron configurations.
• The period number is the value of n.
• Groups 1A and 2A have the s-orbital filled.
• Groups 3A - 8A have the p-orbital filled.
• Groups 3B - 2B have the d-orbital filled.
• The lanthanides and actinides have the f-orbital
filled.
5th 2nd 6th 1st
Different atomic orbitals are denoted by
letters. The s orbitals are spherical, and
p orbitals are dumbbell-shaped.
The electron configuration of an atom is a
shorthand method of writing the location of
electrons by sublevel.
The sublevel is written followed by a
superscript with the number of electrons in the
sublevel.
If the 2p sublevel contains 2 electrons, it is written
2p2
The electron configuration of an atom is a
shorthand method of writing the location of
electrons by sublevel.
The sublevel is written followed by a
superscript with the number of electrons in the
sublevel.
If the 2p sublevel contains 2 electrons, it is written
2p2
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When we write the electron configuration of a
positive ion, we remove one electron for each
positive charge:
Na → Na+
1s2 2s2 2p6 3s1 → 1s2 2s2 2p6
When we write the electron configuration of a
negative ion, we add one electron for each
negative charge:
O
1s2 2s2 2p4
→
→
O21s2 2s2 2p6
Electron Configurations
and the Periodic Table
Pauli Exclusion Principle
According to the Pauli exclusion
principle, an atomic orbital may describe
at most two electrons. To occupy the
same orbital, two electrons must have
opposite spins; that is, the electron spins
must be paired.
Hund’s Rule
Hund’s rule states that electrons
occupy orbitals of the same energy
in a way that makes the number of
electrons with the same spin
direction as large as possible.
• The periodic table can be used as a guide for
electron configurations.
• The period number is the value of n.
• Groups 1A and 2A have the s-orbital filled.
• Groups 3A - 8A have the p-orbital filled.
• Groups 3B - 2B have the d-orbital filled.
• The lanthanides and actinides have the f-orbital
filled.
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http://wine1.sb.fsu.edu/chm1045/notes/Struct/EConfig/Struct08.htm
The two electrons in Helium represent the complete
filling of the first electronic shell. Thus, the
electrons in He are in a very stable configuration
For Boron (5 electrons) the 5th electron must be
placed in a 2p orbital because the 2s orbital is filled.
Because the 2p orbitals are equal energy, it doesn't
matter which 2p orbital is filled.
Orbital Filling Diagram
Ne 1s2 2s2 2p6
F 1s2 2s2 2p5
F- 1s2 2s2 2p6
Mg 1s2 2s2 2p6 3s2
Mg2+ 1s2 2s2 2p6
(10 electrons)
(9 electrons)
(10 electrons)
(12 electrons)
(10 electrons)
Notice – different elements can have the
same number of electrons
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Some actual electron configurations
differ from those assigned using the
aufbau principle because half-filled
sublevels are not as stable as filled
sublevels, but they are more stable
than other configurations.
Exceptions to the
aufbau principle are
due to subtle electronelectron interactions in
orbitals with very
similar energies.
Copper has an electron
configuration that is an
exception to the
aufbau principle.
Exceptional Electron Configurations
Why do actual electron
configurations for some elements
differ from those assigned using the
aufbau principle?
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1.
Identify the element that corresponds to
the following electron configuration: 1s22s22p5.
3.
Spectral lines in a series become closer
together as n increases because the
F
energy levels have similar values.
Cl
Ne
energy levels become farther apart.
atom is approaching ground state.
O
electrons are being emitted at a slower rate.
2.
N.
Write the electron configuration for the atom
1s22s22p5
1s22s22p3
1s22s1p2
1s22s22p1
3.
The electron configurations for some
elements differ from those predicted by the
aufbau principle because the
the lowest energy level is completely filled.
none of the energy levels are completely filled.
half-filled sublevels are less stable than filled
energy levels.
half-filled sublevels are more stable than some
other arrangements.
`
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Ground state – atoms with their electrons in
the lowest energy level possible; this lowest
energy state is the most stable.
Excited state – a higher energy state;
electrons jump to higher energy levels by
absorbing energy
Atom is less stable in an excited state; it will
release the extra energy to return to the
ground state
103
.
The configuration that corresponds to the
lowest electronic energy is called the ground
state.
Any other configuration is an excited state.
As an example, the ground state
configuration of the sodium atom is
1s22s22p63s1 , as deduced from the
Aufbau principle .
The first excited state is obtained by
promoting a 3s electron to the 3p
orbital, to obtain the 1s22s22p63p1
configuration, abbreviated as the 3p
level.
Atoms can move from one
configuration to another by absorbing
or emitting energy
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He showed that the energy of a quantum is
related to the frequency of emitted radiation
by the equation
Equantum=hv
Where E is energy,
h is Plancks constant and
v is frequency
Plancks constant is 6.626 x 10-34 J*s
J is the symbol for the joule
Check how as energy increases, v will
increase—(why violet light has more energy
than red light)
The photoelectric effect is when electrons
are emitted from a metals surface when a
certain frequency of light shines on the
surface.
The mystery concerns the frequency, and
therefore color of the light.
Too low a frequency, no matter how long or
intense it is will not eject electrons.
As long as the frequency is greater than
1.14 x 10 15 Hz, electrons will be ejected.
In 1900 Max Planck searched for an
explanation of why light was emitted from
heated objects.
His study lead to a conclusion that matter
can gain or lose energy only in small specific
amounts called quanta—a quantum is the
minimum amount of energy that can be lost
or gained by an atom.
Einstein proposed in 1905 that
electromagnetic radiation had both
waveleline and particlelike natures and called
these bundles of energy photons.—particles
of electromagnetic radiation with no mass
that carries a quantum of energy.
Einstein calculated that a photons energy
depended on its frequency.
E photon = hv
There has to be a minimum value to cause the
ejection of a photoelectron
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Neon advertising signs are
formed from glass tubes
bent in various shapes. An
electric current passing
through the gas in each
glass tube makes the gas
glow with its own
characteristic color. You
will learn why each gas
glows with a specific color
of light.
Light
How are the wavelength and frequency of
light related?
The wavelength and frequency of
light are inversely proportional to
each other.
The product of the frequency and
wavelength always equals a constant (c),
the speed of light.
The amplitude of a wave is the wave’s
height from zero to the crest.
The wavelength, represented by λ (the
Greek letter lambda), is the distance
between the crests.
The frequency, represented by ν
(the Greek letter nu), is the
number of wave cycles to pass a
given point per unit of time.
The SI unit of cycles per second is
called a hertz (Hz).
According to the wave model, light
consists of electromagnetic waves.
Electromagnetic radiation includes
radio waves, microwaves, infrared
waves, visible light, ultraviolet waves,
X-rays, and gamma rays.
All electromagnetic waves travel in a
vacuum at a speed of 2.998 × 108 m/s.
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Sunlight consists of light with a
continuous range of wavelengths and
frequencies.
When sunlight passes through a
prism, the different frequencies
separate into a spectrum of colors.
In the visible spectrum, red light
has the longest wavelength and the
lowest frequency.
The electromagnetic spectrum consists
of radiation over a broad band of
wavelengths.
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A prism separates light into the colors it
contains. When white light passes through a
prism, it produces a rainbow of colors.
c/ λ =v
When light from a helium lamp passes
Practice problems 1-4 p121 H and
p140—14 & 15 regs
through a prism, discrete lines are
produced.
The frequencies of light emitted by an
Atomic Spectra
What causes atomic emission spectra?
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.
element separate into discrete lines to
give the atomic emission spectrum of
the element.
Mercury
Nitrogen
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5.3
The three groups of lines in the
hydrogen spectrum correspond to the
transition of electrons from higher
energy levels to lower energy levels.
An Explanation of Atomic Spectra
How are the frequencies of light an atom emits
related to changes of electron energies?
5.3
In the Bohr model, the lone electron in the
hydrogen atom can have only certain specific
energies.
When the electron has its lowest possible
energy, the atom is in its ground state.
Excitation of the electron by absorbing
energy raises the atom from the ground
state to an excited state.
A quantum of energy in the form of light is
emitted when the electron drops back to a
lower energy level.
Quantum Mechanics
How does quantum mechanics differ from
classical mechanics?
In 1905, Albert Einstein successfully
explained experimental data by proposing
that light could be described as quanta
of energy.
The quanta behave as if they were
particles.
Light quanta are called photons.
In 1924, De Broglie developed an
equation that predicts that all moving
objects have wavelike behavior.
5.3
Today, the wavelike properties of beams
The light emitted by an electron moving
from a higher to a lower energy level has a
frequency directly proportional to the
energy change of the electron.
of electrons are useful in magnifying
objects. The electrons in an electron
microscope have much smaller
wavelengths than visible light. This allows
a much clearer enlarged image of a very
small object, such as this mite.
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Classical mechanics adequately describes
the motions of bodies much larger than
atoms, while quantum mechanics describes
the motions of subatomic particles and
atoms as waves.
2.40 109 Hz
2.40 1024 Hz
2.40 106 Hz
2.40 102 Hz
The Heisenberg uncertainty principle
states that it is impossible to know
exactly both the velocity and the
position of a particle at the same time.
This limitation is critical in dealing
with small particles such as electrons.
This limitation does not matter for
ordinary-sized object such as cars or
airplanes.
1.
Calculate the frequency of a radar wave
with a wavelength of 125 mm.
2.
The lines in the emission spectrum for an
element are caused by
the movement of electrons from lower to higher
energy levels.
the movement of electrons from higher to lower
energy levels.
the electron configuration in the ground state.
the electron configuration of an atom.
The Heisenberg Uncertainty Principle
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