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Chapter 5
“Electrons in Atoms”
Section 5.1 –
Models of the Atom
• OBJECTIVES:
• Identify the inadequacies in the Rutherford
atomic model.
• Identify the new proposal in the Bohr model of
the atom.
• Describe the energies and positions of electrons
according to the quantum mechanical model.
• Describe how the shapes of orbitals related to
different sublevels differ.
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Ernest Rutherford’s Model
• Discovered dense positive
piece at the center of the
atom- “nucleus”
• Electrons would surround
and move around it, like
planets around the sun
• Atom is mostly empty
space
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Ernest Rutherford’s Model
• These were all significant
discoveries, but a few
questions still went
unanswered…
• Why is there so much
empty space? Why don’t
the electrons just fall into
the nucleus, after all,
Electrons are negative &
Protons are positive?
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Niels Bohr’s Model
• Neils Bohr helped answer this question:
• He said electrons around the nucleus:
– Move like planets around the sun.
– Move in specific circular paths, or orbits,
at different levels.
– An amount of fixed energy separates one
level from another.
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Niels Bohr’s model
• We call these levels where electrons reside
Energy levels of an atom
• They work something like a ladder:
Who has more gravitational potential energy?
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The Ladder Analogy
the similarities…
A ladder:
• You can’t stand in
between the rungs,
only on them.
• When you step up to
the next rung on the
ladder you gain
potential energy.
• When you step down
you loose potential
energy.
An atom:
• Electrons can’t exist
between energy
levels.
• When electrons
absorb energy they
can move to higher
energy levels.
• When an electron
moves down to a
lower energy level it
emits energy.
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The Ladder Analogy
the similarities…
A ladder:
• When I move between
rungs I experience a
change in energy that is
a specific and definite
amount.
• That amount is
determined mainly by
my height from the
ground.
An atom:
• An electron requires a
specific and definite
amount of energy in
order to move to a
higher energy level.
• That amount is
determined by the
energy level it is moving
from and moving to.
• When an electron drops
to a lower energy level
it emits a specific and
definite amount of
energy
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The Ladder Analogy
the differences…
A ladder:
• The energy I posses
on the ladder is
gravitational
potential energy.
• I acquire greater
potential energy
after moving up the
ladder.
An atom:
• The energy an
electron has is
electromagnetic
energy.
• An electron must
have the energy
prior to moving up to
higher energy level,
it is this additional
energy that allows it
to move.
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The Ladder Analogy
the differences…
An atom:
A ladder:
• If an electron were
• If I were to jump
down to a lower rung,
to jump down to a
my energy would
lower energy level,
decrease because
its energy would
some of my energy
decrease because
would be converted
some of its energy
to kinetic energy
would be released as
(energy associate
light.
with motion).
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The Ladder Analogy
the differences…
A ladder:
• Rungs of a
ladder are
usually evenly
spaced.
An atom:
• Energy levels in
an electron are
NOT separated
by equal
amounts of
energy
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A Diagram of Energy Levels
High
energy
Low
energy
• Each level has a
certain amount of
energy associated
with it
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Let’s see how this works:
http://www.visionlearning.com/library/
module_viewer.php?mid=51
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Niels Bohr’s model
• A Quantum of energy is the amount of
energy (a packet or chunck) required to
move an electron from one energy level
to another
• Since the energy of an atom is NEVER
“in between” there must be a quantum
leap in energy.
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And then came…
The Quantum Mechanical Model
Good Grief!
Another Model
of the atom
!*#@?
• We know now that Bohr’s model, although
very close, is not completely accurate.
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The Quantum Mechanical Model
• The quantum mechanical model is the
currently accepted model of the atom.
• In the next section of this chapter when
we look at how electrons for specific
elements are arranged in the atom, we will
use the Bohr model. It is accurate and
sufficient for the work we’ll be doing.
• However, you are required to know some
characteristics of the currently accepted
model and so we will review them now.
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The Quantum Mechanical Model
• Energy levels are NOT circular paths.
• Energy levels are areas where there is a
high probability of finding an electron
that we say is “in” that energy level.
• We draw:
In reality:
●
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http://universe-review.ca/I15-53-quantum.jpg
The Quantum Mechanical Model
• The nucleus is found inside a blurry
“Electron Cloud”
• Think of a fan blades spinning
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The Quantum Mechanical Model
• Each energy level is made up of orbitals.
(sub-levels)
• If an energy level has 8 electrons in it, all 8
electrons don’t have the same probability of
being found in the same place.
• 2 of those electrons may spend most of their
time…
• Another 2 here…
Another 2 here…
Another 2 here…
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The Quantum Mechanical Model
• There are more orbitals with different & more
complicated shapes, but if we tried to illustrate
what an atom with many electrons in many
orbitals would look like…
• This would be a
simplified version21
Before we go into any more detail, lets do
a quick review of the historical
development of the atom
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Section 5.2 –
Electron Arrangement in Atoms
• OBJECTIVES:
• Describe how to write the electron
configuration for an atom.
• Explain why the actual electron
configurations for some elements
seem ‘out of order’
• Describe what type of electron
configuration makes an atom stable.
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Principal Quantum Number
• Generally symbolized by “n”, it denotes
the energy level in which the electron is
located.
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• Maximum number of electrons
that can fit in an energy level:
2n2
n=1:
2(12) = 2 … energy level 1 holds 2 en=2:
2(22) = 2(4) = 8 … energy level 2 holds 8 e-
n=3:
2(32) = 2(9) = 18 … energy level 3 holds 18 en=4:
2(42) = 2(16) = 32 … energy level 4 holds 32 en=5:
2(52) = 2(25) = 50 … energy level 5 holds 50 e26
Electron Configuration
• Energy levels with the lowest energy are
filled first.
• An electron entering an atom will move into
the energy level with the lowest energy (as
long as it is not already filled.)
• What energy level gets filled first?
– Energy level 1, also written n=1
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Electron Configuration
– The closer an energy level is to its maximum
number of electrons, the more stable the atom is.
– Example: Energy level 1 can hold 2 electrons.
Which is more stable? Explain why – use space in notes.
Helium Atom
Hydrogen Atom
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Electron Configuration
• Hydrogen: a violently
• Helium is a very
explosive gas
stable gas – not
reactive under
normal circumstances
The Hindenberg &
Hydrogen?
Hydrogen Balloon
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An Example of Electron Configuration
- Phosphorous
• Phosphorous (P) has 15 electrons
• Draw a nucleus
• Fill up level 1
– 2 electrons
• Fill up level 2
– 8 electrons
nucleus
• Fill up level 3
– 5 electrons
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An Example of Electron Configuration
• Cl has 17 electrons
- Chlorine
• Draw a nucleus
7
• Fill up level 1
8
– 2 electrons
2
• Fill up level 2
– 8 electrons
nucleus
• Fill up level 3
– 7 electrons
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An Example of Electron Configuration • Cu has 29 electrons
• Draw a nucleus
• Fill up level 1
Copper
1
– 2 electrons
19?
18
8
– 8 electrons
2
• Fill up level 2
• Fill up level 3
– 19 electrons?
– No…Why?
– Only 18 e- fit
nucleus
• Fill up level 4
– 1 electrons
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An Example of Electron Configuration • Ca has 20 electrons
• Draw a nucleus
• Fill up level 1
– 2 electrons
• Fill up level 2
– 8 electrons
• Fill up level 3
– 10 electrons?
– No…Why?
– Up to 18 e- can fit…
Calcium
10?
8
2
nucleus
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Overlapping Amounts of Energy
High
energy
Energy level: 1 2 3 4 5 6 7
Low
energy
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35
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Energy levels that are farther from the
nucleus are closer together
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Back to Calcium
• Cu has 20 electrons
• Draw a nucleus
• Fill up level 1
– 2 electrons
• Fill up level 2
– 8 electrons
• Fill up level 3
2
8
8
2
– 8 electrons
• Fill up level 4
– 2 electrons
nucleus
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Energy Level Diagrams
Draw energy level diagrams for:
• Li
Energy levels:
Nucleus:
• Na
•
#e
•
1cm
radius
• K
• 1 = 2cm radius
• colored red
• Rb
• 2 = 4cm radius
• #p+
• Be
• 3 = 5cm radius
• #no
• B
• 4 = 6cm radius
• C
• 5 = 7cm radius
• N
• O
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Questions
• What do all the elements in the 1st
column of the periodic table have in
common?
• What do all the elements in the 2nd row
of the periodic table have in common?
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Section 5.3:Physics and the
Quantum Mechanical Model
OBJECTIVES:
• Identify the source of atomic emission
spectra.
• Describe the relationship between the
wavelength and frequency of light.
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Light
• The light that we see with our eyes is just a small
part of a large spectrum of electromagnetic
radiation
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“R O Y
Low Frequency
G
B I
V”
High Frequency
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Wavelength Longer
Wavelength, Frequency and Energy
• Frequency (ƒ ) and Wavelength (λ)
have an inverse relationship
– As one goes up the other goes down.
• Frequency (ƒ ) and Energy have a
direct relationship
– As one goes up, so does the other.
ƒ _____ ; λ _____; E ______
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Wavelength, Frequency and
Energy
• Different frequencies of light are
different colors, the whole range is called
a spectrum.
• The color of the light indicates its
wavelength, frequency and Energy
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Atomic Spectra
• White light is made up of all the colors
of the visible spectrum.
• Passing light through a
prism bends it. The λ
determines how much it is bent.
Each color, with its
own wavelength bends
at its own angle.
• Result: the light separates
into its component colors
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If the light is not white
• Recall from earlier in chapter:
• when electrons drop to a
lower energy they release
energy as light
• Each element gives off its own
characteristic colors
• When we pass the light given
off by an atom through a
prism, we see only the
component colors present
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Atomic Emission Spectrum
• The characteristic color sequence given off
by an element is called its Atomic Emission
Spectrum
• Can be used to identify the atom.
• This is how we know what stars are made
of.
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Why does each element have
its own spectrum?
How does this Happen?
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Ground State
• When we write electron configurations, as
in the previous section, we are writing the
configuration for the lowest energy.
• In this configuration, the atom is said to
be in its ground state
• The ground state is stable
• Atoms tend toward ground state
configurations.
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Ground State
• Let’s look at a hydrogen atom, with only one
electron, and in the first energy level
• This is the electron configuration with the
lowest energy for hydrogen, we call it
_____________________
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Changing the energy
• Heat, electricity, or light can move the
electron up to different energy levels.
The electron is now said to be “excited”
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Changing the energy
• The excited state is higher energy, less stable.
• The atom will tend toward its ground state.
• As the electron falls back to the ground state,
it gives the energy back as light
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Changing the energy
• They may fall down in specific steps
• Each step has a different energy
• The further they fall, more energy is released
and the higher the frequency
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Changing the Energy
• Remember, every amount of energy
corresponds to a certain frequency,
wavelength and therefore color
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• These are called
the atomic
emission
spectrum
• Unique to each
element, like
fingerprints!
• Very useful for
identifying
elements
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