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Energy & the Electron
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In the study of the atom, there are some
important concepts that can help you
develop your understanding of matter.
Ideas that we will investigate include:
 1. The Wave & Particle Nature of Light (energy)
 2. The Nature of Electrons in Atoms
 3. Bonding
 4. Trends in the Elements on the Periodic Table
Fundamentals, Properties and Relationships in Light
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Light is a form of energy and the light we see
is part of the Electromagnetic (EM) Spectrum
Light has wave properties, including:
 Wavelength (λ) – distance between consecutive
crests or troughs in waves (measured in meters)
 Amplitude – height of the wave
 Frequency (f) – number of crests or troughs
passing each second (measured in s-1 or Hertz)
 Speed (c) – for light being 3.0 x 108 m/s
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The physical features of light are shown
below:
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There is a relationship between these
properties of a wave. This is shown by the
formula:
c = fλ
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We can use this relationship to solve for
either one of the wave characteristics.

What is the wavelength of light with a
frequency of 3.44 x 109 Hz?
c = fλ
Rearrange to get: λ =
𝑐
ν
3.00 ×108 𝑚/𝑠
λ=
3.44 × 109 𝑠 −1
= 8.72 x 10-2 m

What is the frequency of light with a
wavelength of 4.53 x 10-6 m?
c = fλ
Rearrange to get: f =
𝑐
λ
3.00 × 108 𝑚/𝑠
f=
4.53 × 10−6 𝑚
= 6.62 x 1013 Hz
Light’s Alter Ego
Light has both wave and particle properties (a
dual nature)
 Why? Well….
The wave model does not explain the observations
of why heated objects will only emit certain
frequencies of light at a given temperature.
 Max Planck (1856-1947) proposed that there
needed to be a minimum amount of energy that
can be gained, or lost, by an atom (this energy is
called a quantum)
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Planck determined a relationship for energy
and the observations made:
Equantum = hf
Where h = 6.626 x 10-34 J·s
Theory states that matter can only absorb or
emit energy in whole number multiples of hf
(1 hf, 2 hf, 3 hf, ...) i.e. No partial multiples

We can use the formula, E = hν to solve for:
 1. Energy of a particle, or the
 2. Frequency of the particle

Remember that h is a constant!
Example:
 What is the amount of energy in a particle that
has a frequency of 7.76 x 1014 Hz?
E = hf
E = (6.626 x 10-34 J·s)(7.76 x 1014 Hz)
= 5.14 x 10-19 J
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Example:
 What is the energy of a particle that has a
wavelength of 566 nm?
This question is little bit of a different take on the
problem, since it has two issues to overcome:
1. It provides λ instead of f, and
2. It states a wavelength in nm (nanometers)
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A nano- anything means that the measure is
actually a very small number.
As a metric prefix, it stands for 10-9.
So, if we have 600 nm, it basically means:
 600 x 10-9 m, or: 6.00 x 10-7 m
 Note: you can enter the value as 600 x 10-9 on your
calculator and it will work, but you are expected
to show the correct Scientific Notation value
when writing it down.
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So now that we have a way to fix the number,
what do we do with it, since we actually need
frequency?
Any ideas?
 We can change wavelength to frequency using
c=fλ.
c = fλ
f=
𝑐
λ
=
3.00
3.00 ×
× 10
1088 𝑚/𝑠
𝑚/𝑠
−7 𝑚
5.66
566 ×
×10
10−9
𝑚
= 5.30 x 1014 Hz
Remember that 566 x 10-9 = 5.66 x 10-7

Now that we have the value for ν, we can
solve for E:
𝐸 = ℎ𝑓
= (6.626 × 10−34 𝐽 ∙ 𝑠)(5.30 × 1014 𝐻𝑧)
= 5.31 × 10−19 𝐽
 So for radiation with a wavelength of 566 nm, the
energy of a particle is 5.31 x 10-19 J.
Understanding the Atom and the Electron
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A neat property of the elements is that each
element has a unique, what is called,
“emission spectrum”.
In a nutshell, an emission spectrum is a
pattern of light radiation that is produced by
an element after it has received energy (for
example, being heated).
This is an “Absorption Spectrum”
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With the understanding that light has behaviour
of both a particle and a wave, we can start to
understand the emission spectra of atoms.
One in particular, hydrogen (shown below)
The theory of Planck and Einstein states that
there are only certain allowable energy levels or
states.
The lowest allowable state is called the ground
state.
Recall that it was stated
that light’s properties
could not be explained
entirely by the Wave
Model.
 What was the evidence?
 A phenomenon was
known at the time where
a certain frequency of
light shined on a metal
surface will cause the
“photoelectric effect”
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This effect results in the release of electrons
from the metal.
Due to this apparent ability of light to cause
the ejection/excitation of electrons from a
metal, Albert Einstein proposed that light
existed as bundles of energy called
“photons”
Elements have the ability to absorb certain
amounts of energy.
 When the atoms of an element absorb enough
energy, they become “excited”. In this state it is
actually the electrons that become excited.
 When these electrons release this energy to go
back down to “ground” state, they release it in
the form of radiation (light).
 Each element will display a particular emission
spectrum, so we can actually identify elements
by their emission spectrum.

Neils Bohr developed a model for an atom. He also
developed a quantum model of hydrogen that helps
explain the visible spectrum of hydrogen.
 Although hydrogen has only one electron, it can have
many different excited states.
 There is a different energy level corresponding to
each possible orbit around the atom (the lowest
energy level for the orbit closest to the nucleus)
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Bohr defined each orbit around an atom as
having a Quantum state or number. The closest
one having a value of n=1.
Orbit
Quantum
Number
Orbit Radius Corresponding
(nm)
energy level
Relative
Energy
First
n=1
0.0529
1
E1
Second
n=2
0.212
2
E2 = 4E1
Third
n=3
0.476
3
E3 = 9E1
Fourth
n=4
0.846
4
E4 = 16E1
Fifth
n=5
1.32
5
E5 = 25E1
Sixth
n=6
1.90
6
E6 = 36E1
Seventh
n=7
2.59
7
E7 = 49E1
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The emission spectrum that we do see is
only a part of what is released by the atom
of hydrogen (the visible part is called the
Balmer Series).
There are 2 other Series corresponding to
the ultraviolet range (Lyman) and infrared
range (Paschen), which we cannot see.
we can see what levels of electrons move
from and to using the diagram on the next
slide:
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A scientist in the mid-1920’s by the name of
de Broglie proposed that since waves can
display particle-like behaviour, then particles
can show wave-like behaviour.
His idea developed into the following:
λ= h
mv
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Where: m = mass in kg and v = velocity in m/s
What is the wavelength for a car of mass 910
kg and a velocity of 25 m/s?
2.9 x 10-38 m
1.
2.
3.
What is the frequency of green light with a
wavelength of 540 nm?
5.6 x 1014 Hz
How much energy in joules is there in light
with a frequency of 4.67 x 1017 Hz?
3.09 x 10-16 J
What is the wavelength of a Panther with a
mass of 85 kg and a velocity of 12 m/s?
6.5 x 10-37 m
Study of the Configuration of the Atom’s Electrons
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Heisenberg’s Uncertainty Principle states:
 Electrons are constantly moving.
 We cannot know both the precise location of an
electron around an atom and its speed.
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Instead, we have a region called an “orbital”
which indicates its most probable location.
Each orbital can carry, at most, 2 electrons.
There are 4 different types of orbitals:
 s, p, d, and f
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We use a special notation to write the
electron configuration for an element that
utilizes the s, p, d, and f orbital names.
Each orbital type has a unique
shape to account for the
additional electrons.
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This is how it works:
 We start with s orbitals and go up to 2
 We can then add p orbitals and go up to 6
 Next are the d orbitals that go up to 10, then f orbitals to
14
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Observe these examples:
 Starting with hydrogen, we know that it has only the 1
electron. Its electron configuration is 1s1.
 For helium, which has 2 electrons, its configuration is
1s2.
 For lithium, which has 3 electrons, its configuration is
1s22s1 (why?)
 For boron, which has 5 electrons, its configuration is
1s22s22p1
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Write the full electron configurations for the
following elements:
 Carbon
 Fluorine
 Sodium
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Elements we know, have certain abilities to
form chemical bonds.
The number of bonds is based on the number
of “valence” electrons for that element (the
electrons in the outermost shell or orbital).
We can illustrate this for each element by
drawing what is called an “electron dot
diagram”.

Here is how Dot Diagrams are drawn:
 1. Hydrogen
 2. Helium
 3. Lithium
H•
• He
•
• Li
 5. Boron
• Be
•
•
•
•
•
 4. Beryllium
We place electrons around the
symbol (in no particular order).
•B •
•
•
•
We add new electrons to a spot
where there is no electrons until
the element is surrounded.
We now add electrons to form
pairs.
We add electrons until
all electrons are paired.
In order to determine the number of valence electrons
that an element has, we need to use the electron
configuration for that element.
 The number of valence electrons is equal to the
number of electrons found in its highest orbital
(“principal quantum number”).

 For example: Sn (tin) has 50 electrons
 its electron configuration is:
[Kr]5s24d105p2
So Sn has
4 valence
electrons
•
• Sn
•
•
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Recall that ground state is the lowest possible
energy state for an element (and its electrons)
The arrangement of an element’s electrons is
dictated by three rules or principles:
 Aufbau Principle – each electron occupies the lowest
energy orbital possible.
 Pauli Exclusion Principle – two electrons may occupy
the same orbital as long as they have opposite “spins”.
 Hund’s Rule – states that electrons must fill empty
orbitals before pairing electrons of opposite spin.
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In the Orbital Diagram, we draw each orbital
in an element using a box:
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A single box represents the s orbital
A triple box represents the p orbitals
A set of 5 boxes represent the d orbitals
A set of 7 boxes represent the f orbitals
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When we draw an orbital diagram, we fill in
the boxes using arrows to represent the
electrons like we see below:
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Try drawing the diagram for nitrogen (atomic
number 7):
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Covalent chemical bonding is based on the
number of valence electrons that are available to
form that bond for the element.
We are used to elements having the ability to
form bonds like with carbon, where it can form
up to 4 bonds (one for each valence electron).
The standard rule for bond formation is to
complete what is called an “octet” (8 valence
electrons).
Another requirement that has been made
necessary is the need for free electrons in order
to form a chemical bond.
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How do orbitals affect bonding?
Consider the following compound: BeCl2
If you draw the box diagram for Be, you
normally would get this:
So you actually don’t have any free electrons
to form bonds.
 So how is Be able to form bonds?
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How it works for many elements is what is
called “hybridization of orbitals”.
In this process, an element creates free
electrons by forming a hybrid orbital.
This occurs by combining orbitals of the
same quantum number.
For BeCl2, we see this:
We move one electron
from the pair to the
available space in
the next orbital type
In the previous example, Be actually changes its
bonding orbital type to the combination of the
orbitals combined: sp
 The naming is based on the type of orbitals
combined and how many “boxes” are used in the
formation of the hybrid.
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 Other possible hybrid types: sp2, sp3, sp3d, sp3d2
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What kind of orbitals would we need for AlCl3?
 Draw the Box Diagrams for both the non-hybrid Al
and a hybrid Al.
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Each chemical molecule will have a particular
shape associated with it.
Hybridization of orbitals will cause the
formation of a variety of molecular shapes
that are very interesting:
Open your text to page 260.
 You will find a table showing a variety of chemical
molecules and their known shape.
 The type of hybridization does influence the
generated shape for that molecule.
Studying the Properties of the Elements
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We know that the elements vary in their
properties, however, the elements will display
similarities as well.
We know that elements within a Group will
have the same chemical properties (react the
same).
Trends are also seen as we move through the
elements both down a Group (column) as well
as across the Periodic Table.

We will look at the following properties:
 1. Atomic Radius
▪ This is the measure of the atom of an element from the
center of its nucleus to the edge of its electron cloud.
 2. Ionic Radius
▪ This is the measure of the ion-form of an element from
its center to the edge of the its electron cloud.
 3. Ionization Energy
▪ This is the measure of the amount of energy it takes to
remove an electron from the atom of that element.
 4. Electronegativity
▪ This is basically a relative measure of the strength of
attraction for electrons in that element.

The Trends:
 Atomic Radius – we see is that the radius of an
atom will:
▪ Increase in size as we move down a Group (atoms get
bigger, more electron levels/shells)
▪ Decrease in size as we move across the Period (due to a
higher core charge, pulls the electrons closer)
 Ionic Radius – the ionic radius will:
▪ Increase as we move down a Group (the number of
orbits increase just like in atoms).
▪ Decrease as we move across (same idea as before, core
charge increases), until will change the type of ion
(positive to negative), then the pattern resets.
▪ The key ideas here are that:
▪ 1. Negative ions are larger than their atom
▪ 2. Positive ions are smaller than their respective atom
 Ionization Energy pattern (cont’d)
▪ The highest ionization energies belong to the Noble
Gases (since they will NOT want to lose them).
▪ The lowest belong to the elements that typically form
positive ions (easily lose electrons).
 Ionization Energy will:
▪ Decrease as we move down a Group (electrons get
further away from the nucleus).
▪ Generally increase as we move across a Period (towards
the Noble Gases).
 Electronegativity – the strength of an atoms’
attraction towards electrons will generally:
▪ Increase as we move from left to right (Noble Gases
have no affinity for electrons).
▪ Decrease as we move from top to bottom (along a
Group).

What is the role of electronegativity?
 These electronegativity values will dictate what
kind of bond will form between the atoms.

Remember that there are 2 main types of bonds:
 Covalent
 Ionic

The way we use electronegativity is the
difference between the values:
 If the difference is 1.0 or less, it will be a covalent
bond.
 If the difference is 2.0 or more, it will be an ionic bond.
 If the difference is between 1.0 and 2.0, it will be
covalent with some polar character

Atomic Structure
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1. The Electromagnetic (EM) Spectrum
2. Energy of a Particle (Planck)
3. de Broglie Equation
4. Ground State vs. Excited Stated
5. Spectra of the Elements
6. Quantum Number
7. Electron Configuration
8. Electron Dot Diagrams
 8. Electron Box Diagrams
 9. Hybridization of Orbitals
 10. Trends in the Periodic Table
▪ Atomic Radius
▪ Ionic Radius
▪ Ionization Energy
▪ Electronegativity