Transcript Chapter 5 Sections 1,3

5.1 Revising the Atomic Model >
Chapter 5
Electrons In Atoms
5.1 & 5.3: Evidence for a New
Atomic Model
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5.1 Revising the Atomic Model > Development of Atomic
Models
Thomson Model
In the nineteenth century, Thomson described
the atom as a ball of positive charge containing
a number of electrons. Plum Pudding Model
Bohr Model
After Rutherford's discovery, Bohr proposed
that electrons travel in definite orbits
around the nucleus.
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Rutherford Model
In the early twentieth century, Rutherford
showed that most of an atom's mass is
concentrated in a small, positively charged
region called the nucleus. Nuclear Model
Quantum Mechanical Model
Modern atomic theory describes the
electronic structure of the atom as the
probability of finding electrons within
certain regions of space.
5.1 Revising the Atomic Model > Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
• It explained only a few simple properties of
atoms.
• It could not explain the chemical properties
of elements.
For example, Rutherford’s model
could not explain why an object
such as the iron scroll shown here
first glows dull red, then yellow,
and then white when heated to
higher and higher temperatures.
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5.1 Revising the Atomic Model > Energy Levels in Atoms
The Bohr Model
Niels Bohr (Danish, 1885-1962) proposed that
an electron is found only in specific circular
paths, or orbits, around the nucleus.
This eventually leads to
the idea of energy
levels in the electron
cloud.
Each possible electron
orbit in Bohr’s model has
a fixed, definite energy.
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5.1 Revising the Atomic Model > Energy Levels in Atoms
The Bohr Model
• Electrons are confined to specific energy levels,
but can change levels under certain conditions
• In order to move from one level to another, an
electron must absorb or release a certain
quantum or amount of energy.
• Electrons cannot exist between energy levels.
• The higher energy levels are closer together
within the electron cloud.
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5.1 Revising the Atomic Model > Light and Atomic
Emission Spectra
The Bohr Model & Atomic Emission Spectra
When atoms absorb energy, their electrons
move to higher energy levels. These electrons
lose energy by emitting light when they return
to lower energy levels.
Colors of light correspond to specific amounts
of energy.
Measuring the colors of the light allows the
amount of energy that was absorbed by the
electron to be calculated.
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5.1 Revising the Atomic Model >
The Bohr Model and Electron Transitions Illustrated
Electron absorbs a
quantum of energy
to shift levels...
Energy is released
as a photon with a
definitive color as
e’s return to
orginial level.
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5.1 Revising the Atomic Model > Light and Atomic
Emission Spectra
The Electromagnetic Radiation Spectrum
shows all forms of energy (radiation) over a broad
range of wavelengths.
Low energy
( = 700 nm)
Frequency  (s-1)
3 x 106
102
Wavelength  (m)
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High energy
( = 380 nm)
increase
3 x 1012
3 x 1022
10-8
10-14
decrease
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5.1 Revising the Atomic Model > Light and Atomic
Emission Spectra
The frequency () and wavelength () of
light are inversely proportional to each
other. As the wavelength increases, the
frequency decreases.
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5.1 Revising the Atomic Model >
 Frequency and Energy are directly
proportional to each other.
 In the visible spectrum, red light has
the longest wavelength and the lowest
frequency. It also has the least
amount of energy of all visible light.
 Purple light has the highest frequency
and therefore the highest amount of
energy.
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5.1 Revising the Atomic Model > Light and Atomic
Emission Spectra
Atomic Emission Spectra
A prism separates light into the colors it contains.
White light produces a rainbow of colors.
Less
Energy
More
Energy
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R ed
O range
Y ellow
G reen
B lue
I ndigo
V iolet
Screen
Light
bulb
Slit
Prism
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5.1 Revising the Atomic Model > Light and Atomic
Emission Spectra
Atomic Emission Spectra
The distinctive pattern of colored lines produced when
electricty moves through a gas form of an element is
the Atomic Emission Spectra. This is unique for
each element, and therefore is an intensive,
identifying property.
Helium lamp
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Slit
Prism
Screen
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5.1 Revising the Atomic Model >
Bohr’s Theory & Atomic Emission Spectra
1. An electron is normally in its low energy state or ground
state.
2. When the electron absorbs a certain amount of energy
or a “quantum”, it will “jump” to a higher level of energy
or its “excited” state.
3. This new state is unstable for the electron and so it
releases this excess energy as a photon of visible light
EMR, and returns to the ground state.
4. Therefore, Atomic Emission Spectral lines (colored lines)
are unique, identifying properties of atoms. These lines
show up as atoms release the quanta of energy (in the
form of photons of light) generated by electron
movement.
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5.1 Revising the Atomic Model > Light and Atomic
Emission Spectra
Atomic Emission Spectra
• The energy absorbed by an electron for it to
move from its current energy level to a higher
energy level is identical to the energy of the
light emitted by the electron as it drops back to
its original energy level.
• The various colors with their wavelengths added
together make up the atomic emission
spectrum of an element.
• No two elements have the same emission
spectrum!!!
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5.1 Revising the Atomic Model >
Problems with the Bohr Model
1. Could only explain the spectrum of Hydrogen
2. A charged electron revolving around a charged
nucleus SHOULD eventually lose energy and fall into
the nucleus (“the collapsing atom problem”), but it
doesn’t…
3. Violated the Heisenberg Uncertainty Principle, which
states that its impossible to know both the position
AND the speed of an electron at any given moment.
(“Ceiling Fans”)
SO…research continued…
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5.1 Revising the Atomic Model >
See Data from Unit 3 Lab 1—Flame
Tests!
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5.1 Revising the Atomic Model >
The Quantum Mechanical Model
• Austrian physicist Erwin Schrödinger (1887-1961)
developed a mathematical equation describing
the behavior of the electrons.
• Also believed that electrons travel in energy
levels.
• The modern description of the electrons in atoms,
the quantum mechanical model, came from the
mathematical solutions to the Schrödinger
equation.
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5.1 Revising the Atomic Model >
The Quantum Mechanical Model:
• Like the Bohr model, the quantum mechanical
model of the atom notes that electrons have
specifc amounts of energy.
• Unlike the Bohr model, however, the quantum
mechanical model does not specify an exact
path the electron takes around the nucleus
because the electron is too small and is
traveling at almost the speed of light.
• Only the probability of finding the electron in
an energy level can be determined.
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5.1 Revising the Atomic Model > The Quantum
Mechanical Model
• In the quantum
mechanical model, the
probability of finding an
electron within a certain
volume of space is called
the orbital.
• Combining all the
orbitals in all the energy
levels makes up the
electron cloud.
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5.1 Revising the Atomic Model > The Quantum
Mechanical Model
The Quantum Numbers:
What do the Quantum Numbers
determine about the electrons in an
atom?
• Describes the amount of energy and probable location
of an electron, much like an address
• Every electron in an atom is unique; each electron has
a different energy and therefore will have a different set
of quantum numbers.
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5.1 Revising the Atomic Model >
Example: Quantum Numbers of Helium
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Electron 1
Electron 2
•
•
•
•
•
•
•
•
n=1
l= 0
ml = 0
ms = + 1/2
n=1
l= 0
ml = 0
ms = - 1/2
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5.1 Revising the Atomic Model >
Schroedinger’s Quantum Numbers
Type of
Quantum
Number
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Symbol for
Quantum
Number
Gives the…
Possible
Values…
Principal
Quantum
Number
n
Energy level
1 7
Orbital
Quantum
Number
l
Sublevel & its
shape
s, p, d, f
Magnetic
Quantum
Number
ml
Direction of the
sublevel in the
orbital
s=1
p=3
Spin Quantum
Number
ms
Spin of the
electrons in the
orbital
+ 1/2
d=5
f=7
and -1/2
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5.1 Revising the Atomic Model > Atomic Orbitals
Four Types of Quantum Numbers
• The energy levels of electrons in the quantum
mechanical model are labeled by Principle
Quantum Number (n).
• These numbers are assigned the values n = 1  7.
• For energy levels greater than 1, there are several
orbitals with different shapes within the different
energy levels.
• These orbital shapes are known as sublevels,
and are designated by the Orbital Quantum
Number (l).
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5.1 Revising the Atomic Model >
Sublevel Shapes & Directions in Orbitals:
s-1 direction:
p-3 directions:
d-5 directions:
f-7 directions:
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5.1 Revising the Atomic Model >
In energy level 1, the s orbital is present.
In energy level 2, the s & p orbitals are present.
In energy level 3, the s, p, & d orbitals are present.
In energy levels 4-7, the s, p, d, & f orbitals are
present.
s orbitals use the least amount of energy; f orbitals
need the most energy.
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5.1 Revising the Atomic Model > Atomic Orbitals
The different directions are given by the ml Quantum
Number.
Having different orientations allows multiple electrons
to be in the same general area but to also stay away
from each other!
The Spin Quantum Number (ms ) of +1/2 and -1/2
shows that the 2 electrons within the same orbital
direction will spin in opposite directions so as to stay
away from each other.
These are shown using arrows:
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5.1 Revising the Atomic Model > Atomic Orbitals
Summarizing Quantum Numbers:
• The principal quantum number, n, always equals
the number of sublevels within that principal
energy level.
• The total number of orbital shapes in a principal
energy level is equal to n 2.
• A maximum of two electrons can occupy an
orbital.
• Therefore, the maximum number of electrons that
can occupy a principal energy level is given by the
formula 2n 2.
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5.1 Revising the Atomic Model >
Energy
Level
1
2
3
4
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Possible
Shapes
(Equal to n)
Number of
Orbital
Directions
in Sublevel
(n2)
Number of
Electrons in
Sublevel
Total per
Level
(2n2)
s
s
p
s
P
d
s
P
d
f
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5.1 Revising the Atomic Model >
Calculate the maximum number of
electrons in the 5th principal energy level
(n = 5).
The maximum number of electrons that can
occupy a principal energy level is given by the
formula 2n 2.
If n = 5, 2n 2 = 50.
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5.1 Revising the Atomic Model >
Quantum Number Review…
Quantum Orbital Shapes!!
The Quantum Number Rag!!
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5.1 Revising the Atomic Model >
Key Concepts
Bohr proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus.
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 of an atom.
Each energy sublevel corresponds to one or
more orbitals of different shapes, which
describe where the electron is likely to be
found.
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5.1 Revising the Atomic Model >
END OF 5.1
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