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Chapter 4: Glow in the Dark
Introductory Activity
List as many things as you can think of
that “glow”
What do you have to do to make these
“glowing” things “glow”?
Glow in the dark
This chapter will introduce the chemistry
needed to understand how glowing things
work
Section 4.1: History of Atomic Theory
Section 4.2: Atomic Structure
Section 4.3: Electron Structure
Section 4.4: Periodic table
Section 4.5: Periodicity
Section 4.6: Light
Section 4.7: Light and matter
Glow in the dark
Is based on
Electron
structure
Gives off
Changes in produce
Is a part of
Light
Which can
be determined using
Periodic
Table
Atomic
structure
Arranged to show
Is based on
Atomic
theory
Periodicity
Section 4.1—Development of
Atomic Theory
Dalton’s Atomic Theory
John Dalton’s theory based on experiments in early 1800’s
 All matter is made of tiny particles “atoms”
 Atoms cannot be created, divided, destroyed or
changed into other types of atoms
 Atoms of the same element have identical
properties
 Atoms of different elements have different
properties
 Atoms of different elements combine in wholenumber ratios to form compounds
 Chemical changes join, separate or rearrange
atoms in compounds
Cathode Ray Tubes
A cathode ray is a ray of light traveling in a vacuum (no other particles inside)
The ray travels from one metal plate to another as the plates are connected to
electricity
Cathode ray
Metal plate (cathode)
releases stream
Metal plate (anode) to
which stream travels
Cathode Ray Tubes & Charge
In the late 1800’s, JJ Thomson put charged plates outside the tube
Negatively charged plate
-
+
Positively charged plate
Ray is deflected
away from negative
plate and towards
positive plate
It made no difference what type of metal he used in the tube—all material produced
this stream that curved towards the positive charge
Thomson’s conclusions
 The evidence from Thomson’s work showed that
there was something negatively charged in
atoms
 Since all types of metal produced the same
result, the negative charge is in all types of
atoms
 Since atoms were overall neutral, if there was a
negative charge there had to also be a positive
charge
 In 1897, Thomson announced that the rays were
electrons and they had a negative charge
Theories change
 Thomson’s evidence showed Dalton’s idea of
solid, uniform atoms was incorrect.
 Eugene Goldstein conducted experiments to
label the positive part “protons” and determined
it has the same charge as the electron (with
opposite sign) but is 1837 times heavier!
 Thomson developed the “plum pudding” model.
Since most of us aren’t familiar with plum pudding, you
can think of it as a chocolate cookie dough theory
Thomson’s Theory
The “chips” are the negative
electrons.
The “dough” is the positive
portion
The “chips” are stationary and
don’t move within the “dough”
Remember, officially this theory
is called “plum pudding” but it’s
the same idea!
Gold Foil Experiment
Hans Geiger performed experiments in the
early 1900’s where he bombarded very
thin gold foil with radioactive particles
(alpha particles “”)
They expected these relatively heavy
particles to go through the atoms with a
small deflection
What happened in the experiment?

Gold foil
What did he see?
Most of the alpha particles passed straight
through with no deflection
These particles did not run into anything
Some did deflect slightly
These particles ran into something much
smaller than themselves
A few were reflected back the direction
they came from
These particles ran into something very dense
What did that mean?
Atoms are mostly empty space
Electrons (the smaller particles) were the
cause of the small deflections
There must be a small area of the atom
with most of its mass (the protons) that
caused the reflections.
He called this small, dense area the nucleus
A third particle
The protons and electrons could explain
the charges of the various parts of the
atom
They could not explain the total mass of
the atoms
Neutrons were proposed in 1920’s but not
confirmed until 1932 by James Chadwick
Neutrons had mass similar to protons and no
charge. They were located in the nucleus
More changes to the theory
Niels Bohr performed experiments with
hydrogen atoms & light
He determined that electrons are in levels
according to how much energy they have
and that only certain energy amounts were
allowed.
The Bohr Model
It consists of the nucleus with protons &
neutrons and electrons in concentric orbits
(circles) outside the nucleus
The circle closest to the nucleus contains
the lowest energy electrons
The first level can hold 2 electron, then the
next two levels can each hold 8 and then
levels farther out can hold 18.
Pictures of the Bohr Models
Electron
Proton
Neutron
Hydrogen-1
Helium-4
Lithium-6
Use of the Bohr Model now
We no longer believe electrons are in
concentric circles, but this is still a
convenient way to show energy levels on
2-dimensional paper
Modern Atomic Theory
In the 1920’s, Bohr’s research lead the
way for the study of quantum mechanics
(the study of tiny particles)
Modern atomic theory uses calculus
equations to show how the subatomic
particles act as both particles and waves
These equations show the most probable
location of electrons in the atom (known as
atomic orbitals)
Section 4.2—Atomic Structure
What are atoms?
Atom - smallest piece of matter that
has the chemical properties of the
element.
What’s in an atom?
An atom is made of three sub-atomic particles
Particle
Location
Mass
Charge
Proton
Nucleus
1 amu =
1.6710-27 kg
+1
Neutron
Nucleus
1 amu =
1.6710-27 kg
0
Electron
Outside the
nucleus
0.00055 amu
9.1010-31
kg
-1
1 amu (“atomic mass unit”) = 1.66  10-27 kg
What gives an atom its identity?
What makes an atom “carbon” as opposed
to “oxygen”?
Every atom has a different number of
protons.
The number of protons determines the
identity of the atom
The atomic number shows the number of
protons.
Atomic number = protons
The Nucleus & Mass
Since the nucleus has protons & neutrons,
and the mass of each one is 1 amu…
The mass of the nucleus (in amu’s) is the
number of protons + neutrons
Since electrons have relatively no mass
(0.054% of one proton or neutron), we
don’t need to worry about them when
determining mass of an atom
Mass # = protons + neutrons
Charges
Protons have a positive charge
Electrons have a negative charge
Neutrons have no charge
Overall charge = protons + (-1)×electrons
Charge = protons - electrons
How do we show information
about an element?
Element symbols
Element Symbol
1 or 2 letters, found on
the periodic table
Mass number
# protons + # neutrons
Charge
A
C
X
Z #
# protons - # electrons
(assumed to be “0” if
blank)
Atomic number
# of protons
Number
How many atoms do you
have?
Example: Element symbols
Element Symbol
O = Oxygen
Charge
-2
Mass number
16
16
-2
O
8
Atomic number
8
Number
Assumed to be “1” if blank
Let’s Practice
Example:
Fill in the
missing
values
Symbol
Name
Atomic
#
Magnesium-25
Mass
#
Charge Proton Neutron Electron
+2
82
126
82
Isotopes
What are isotopes?
Isotopes - n. Atoms of the same
element with a different number of
neutrons
Some isotopes are radioactive—but not
all…many are quite stable!
Isotopes Example
Mass # = 2 amu
Mass # = 1 amu
Hydrogen-1
Hydrogen-2
 If they have different number of neutrons, and
neutrons have a mass of 1 amu…
 Then isotopes of the same element will have
different masses!
 But because their protons are the same, they
are the same element!
Identifying Isotopes
Isotopes can be differentiated by their different
mass numbers in the element symbol
12
C
Carbon-12
13
C
Carbon-13
Or by the mass number following their name.
Mass Number versus Atomic Mass
Mass Number
Average Atomic Mass
# of protons + # of neutrons
Average of actual masses
Always a whole number
Not a whole number
For one specific isotope
only
Weighted average of all
isotopes
Is not found on the periodic
table
Is found on the periodic table
Calculating Average Atomic Mass
Average atomic mass is a weighted average (it
takes into account how often each isotope occurs).
Actual mass
(not mass
number)
“Sum of”
Average
=
atomic
mass
(
Abundance of
isotope
 Mass of )
isotope
What fraction of the time is that
isotope present?
Example of Finding Avg Atomic Mass
Example:
Find the atomic mass of
chlorine if
Chlorine-35 has a mass of
34.969 amu
and Chlorine-37 has a
mass of 36.966 amu and is
present 24.22% of the time.
This chart summarizes
the information in the
problem:
Remember that percents add up to
100.
So they said the second isotope is
present 24.22% of the time.
This means that the first isotope is
present 100-24.22 = 75.78% of the
time
Isotope
Mass
Percent
Decimal
1
34.969 amu 75.78
0.7578
2
36.966 amu 24.22
0.2422
Avg Mass  0.7578  34.969amu  0.2422  36.966
= 35.45 amu
(this is what’s on the periodic table for Cl!)
Section 4.3—Electron
Structure
The Electron Hotel
The story of the Electron Hotel
Shopping Center
Parking Garage
Restaurant
A man built an hotel for electrons with a restaurant next door.
But he was making so much money that he decided to add on with some more
rooms and a parking garage.
He still had high demand and decided to add on some more rooms and a
shopping center.
He used the last space he could to put some rooms above the shopping center.
How the Electron Hotel Fills
Shopping Center
Parking Garage
Restaurant
This man had some very strange ideas about how to run his hotel. He insisted
four things:
• The lowest possible must be used first (actually it was the fire inspector
that insisted on this one)
• There can only be one person in a room until all rooms at that level
have someone
• No more than 2 people to a room
• When two people are in a room, they must be of opposite sex
If 8 people come to the hotel, where would he put them?
Another Example
Shopping Center
Parking Garage
Restaurant
This man had some very strange ideas about how to run his hotel. He insisted
four things:
• The lowest possible must be used first (actually it was the fire inspector
that insisted on this one)
• There can only be one person in a room until all rooms at that level
have someone
• No more than 2 people to a room
• When two people are in a room, they must be of opposite sex
If 21 people come to the hotel, where would he put them?
You Try
Shopping Center
Parking Garage
Restaurant
This man had some very strange ideas about how to run his hotel. He insisted
four things:
• The lowest possible must be used first (actually it was the fire inspector
that insisted on this one)
• There can only be one person in a room until all rooms at that level
have someone
• No more than 2 people to a room
• When two people are in a room, they must be of opposite sex
If 42 people come to the hotel, where would he put them?
Where do electrons really live?
Electron Clouds
They don’t live in a hotel…They are in
the area outside of the nucleus where
the electrons reside.
Electron Clouds
Electron Hotel
Electron
cloud
Which section
of the hotel
Principal
energy levels
The electron cloud is
made of energy levels
Which floor
Subshells
Energy levels are
composed of subshells
Which room
Orbitals
Subshells have orbitals.
Subshell versus Orbital
Subshell – A set of orbitals with equal
energy
Orbital – Area of high probability of the
electron being located.
Each orbital can hold 2 electrons
Energy increases
Types of Subshells
Subshell
Begins in
energy level
Number of
equal energy
orbitals
Total number
of electrons
possible
s
1
1
2
p
2
3
6
d
3
5
10
f
4
7
14
Electron Configuration
What are electron configurations?
They show the grouping and position of
electrons in an atom.
The number and configuration of electrons
determines how something glows…so it’s
important to know “where the electrons live” for
an atom!
Electron configurations use boxes for orbitals and
arrows for electrons.
Energy and Subshells
6p
6s
5p
5d
4f
4d
5s
4p
3d
4s
3p
3s
2p
Energy
2s
Subshells are filled from
the lowest energy level to
increasing energy levels.
Does this look familiar? Electron Hotel!
1s
Aufbau Principle
The first of 3 rules that govern electron configurations
1
Aufbau Principle: Electrons fill subshells (and
orbitals) so that the total energy of atom is the
minimum
What does this mean?
Electrons must fill the lowest available subshells and orbitals before
moving on to the next higher energy subshell/orbital.
Where did we see this “rule” in the Electron Hotel?
Hund’s Rule
2
Hund’s Rule: Place electrons in unoccupied
orbitals of the same energy level before doubling
up.
How does this work?
If you need to add 3 electrons to a p subshell, add 1 to each before
beginning to double up.
Where did we see this “rule” in the Electron Hotel?
Pauli Exclusion Principle
3
Pauli Exclusion Principle: Two electrons that
occupy the same orbital must have different spins.
“Spin” describes the angular
momentum of the electron
How does this work?
“Spin” is designated with an up
or down arrow.
If you need to add 4 electrons to a p subshell, you’ll need to double
up. When you double up, make them opposite spins.
Where did we see this “rule” in the Electron Hotel?
Determining the Number of Electrons
Charge = # of protons – # of electrons
Atomic number = # of protons
Example:
How many
electrons
does Br-1
have?
Writing Electron Configurations
1
Aufbau Principle: Electrons fill subshells (and orbitals) so that
the total energy of atom is the minimum
2
Hund’s Rule: Place electrons in unoccupied orbitals of the
same energy level before doubling up.
3
Pauli Exclusion Principle: Two electrons that occupy the same
orbital must have different spins.
Example:
Write the
boxes &
arrows
configuration
for Cl
Spectroscopic Notation
Spectroscopic Notation
Shorthand way of showing electron
configurations
The number of electrons in a subshell are
shown as a superscript after the subshell
designation
1s
2s
2p
3s
1s2 2s2 2p6 3s2 3p5
3p
Writing Spectroscopic Notation
1
Determine the number of electrons to place
2
Follow Aufbau Principle for filling order
3
Fill in subshells until they reach their max (s = 2, p = 6, d = 10,
f = 14).
4
The total of all the superscripts is equal to the number of
electrons.
Example:
Write
spectroscopi
c notation for
S
No charge written  Charge is 0
Atomic number for S = 16 = # of protons
0 = 16 - electrons
1s 2 2s 2 2p 6 3s 2 3p 4
Electrons = 16
2 + 2 + 6 + 2 + 4 = 16
Noble Gas Configuration
Noble Gases & Noble Gas Notation
Noble Gas – Group 8 of the Periodic
Table. They contain full valence shells.
Noble Gas Notation – Noble gas is
used to represent the core (inner)
electrons and only the valence shell is
shown.
Br
Spectroscopic
Noble gas
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5
[Ar] 4s 2 3d 10 4p 5
The “[Ar]” represents the core electrons and only the valence electrons are shown
Which Noble Gas Do You Choose?
How do you know which noble gas to use to symbolize
the core electrons?
Think: Price is Right.
How do you win on the Price is Right?
By getting as close as possible without going over.
Choose the noble gas that’s closest without going over!
Noble Gas
# of electrons
He
2
Ne
10
Ar
18
Kr
36
Xe
54
Noble Gas Notation Example
1
Determine the number of electrons to place
2
Determine which noble gas to use
3
Start where the noble gas left off and write spectroscopic
notation for the valence electrons
Example:
Write noble
gas notation
for As
Section 4.4—The Periodic
Table
History of the Periodic Table
Different scientists organized the elements
differently—this lead to confusion
In 1869, Dimitri Mendeleev designed a
periodic table based on atomic mass.
This way showed patterns in properties that
repeated across rows and similarities down
columns
He couldn’t find elements to fit all the
property trends, so he left holes
History of the Periodic Table
The holes he left were later filled in as
more elements were discovered
The modern periodic table is arranged by
atomic number rather than atomic mass
This caused a few “switches” in placement, but
overall is very similar to Mendeleev’s
Organization of the Periodic
Table
Groups and Periods
Periods
Rows are called
“periods”
Groups
Columns are called
“groups” or “families”
Information for Each Element
Most periodic tables give the following information,
but it can be in a different location
Atomic Number
Element Symbol
If there’s a second
letter, it’s lower-case
Atomic Mass
Number with decimals
Gives the mass for 1 mole
of atoms, in grams
6
C
Carbon
12.01
Whole number—
elements are ordered
by this on the periodic
table.
Element Name
Parts of the Periodic Table
The rows at the bottom
Most periodic tables are written with 2 rows at the bottom.
This is done to allow the font to be bigger on a piece of paper.
The rows at the bottom
Most periodic tables are written with 2 rows at the bottom.
This is done to allow the font to be bigger on a piece of paper.
But they really belong here!
Follow the atomic numbers on your periodic table to see it!
Electron Configurations and
the Periodic Table.
Configurations Within a Group
Look at the electron configurations for the Halogens
F
1s2 2s2 2p5
Cl
1s2 2s2 2p6 3s2 3p5
Br
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
I
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
All of the elements in Group 7 end with 5 electrons in a p subshell.
In fact, every Group ends with the same number of electrons in
the highest energy subshell
Configurations and the Periodic Table
s-block
p-block
d-block
s1 s2
p1 p2 p3 p4 p5 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f-block
f1
f2
f3
f4
f5
f6
f7
f8
f9 f10 f11 f12 f13 f14
How to remember the filling order?
1s
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6p
6s
4f
5d
7s
5f
6d
s subshells begin in level 1, so begin the s-block with “1s”
p subshells begin in level 2, so begin the p-block with “2p”
d subshells begin in level 3, so begin the d-block with “3d”
f subshells begin in level 4, so begin the f-block with “4f”
How to remember the filling order?
1s
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6p
6s
4f
5d
7s
5f
6d
To see the filling order of subshells, read from left to right,
top to bottom!
Note that this tool shows that the 3d energy level is filled
after the 4s energy level!
Where Does the Noble Gas Leave Off?
How do you know where to start off after using a
noble gas?
Use the periodic table!
1s
He
2s
2p
Ne
3s
3p
Ar
4s
3d
4p
Kr
5s
4d
5p
Xe
6p
Rn
6s
4f
5d
7s
5f
6d
The noble gas fills the subshell that it’s at the end of.
Begin filling with the “s” subshell in the next row to show
valence electrons.
Section 4.5—Periodicity
What is periodicity on the periodic table?
The predictable pattern by which
properties of elements change across
or down the periodic table.
There are always exceptions to these periodicity trends…each of
the trends is a “general” trend as you move across a period or
down a group.
Trend 1: Atomic Radii
What is atomic radius?
Half of the distance between the nuclei
of two bonded atoms.
H
H
Distance between nuclei
Atomic radius of hydrogen atom
Atomic Radii Trends
Decreases
Increases
Atomic Radii Trends Reasoning, Part 1
Why do atomic radii decrease across a period?
Moving left to right, the number of protons, neutrons and
electrons all increase.
e
e
e
n
p
n p
n p
e
Move across the
periodic table
Radius decreases
Lithium atom
e
e
pn
p p
p
n nn
Beryllium atom
As the # of protons electrons increase, the attraction between
the positive nucleus and negative electron cloud increases.
This attraction “pulls” in on the electrons.
e
Atomic Radii Trends Reasoning, Part 2
Why do atomic radii increase down a group?
Protons, neutrons and electrons are also added as you move
down a group.
e
e
e
+
Move down the
periodic table
e
e
Radius increases
e
e
e
e
e
e
e
Lithium atom
e
+
e
Sodium atom
However, the electrons are added in new energy levels.
The inner electrons “shield” the new outer electrons from the pull of
the nucleus, therefore it doesn’t pull in like the last slide.
Trend 2: Ionization Energy
What is Ionization Energy?
The energy needed to remove the
outermost electron.
Ionization Energy Trends
Increases
Decreases
Ionization Energy Trends Reasoning, Part 1
Why does Ionization Energy increase across a period?
Moving left to right, the radius of the atom decreases as more
protons pull on more electrons.
Move across the
periodic table
e
e
n
p
n p
n p
e
e
Radius decreases
IE increases
Lithium atom
e
e
pn
p p
p
n nn
e
Beryllium atom
When an atom is smaller, the electrons are closer to the
nucleus, and therefore feel the pull more strongly.
It is harder to pull electrons away from these smaller atoms.
Ionization Energy Trends Reasoning, Part 2
Why does ionization energy decrease down a group?
As you move down a group, the radius increases as more
electrons shells are added.
e
e
e
+
Move down the
periodic table
Radius increases
IE decreases
e
e
e
e
e
e
e
+
e
e
e
e
Lithium atom
Sodium atom
As the outer electrons (those involved in bonding) are farther
from the nucleus, they will feel the “pull” of the nucleus less.
It is easier to remove an electron from a larger atom.
Trend #3: Electron Affinity
What is Electron Affinity?
energy released when an electron is
added to an atom
Electron Affinity Trends
Increases
Decreases
Electron Affinity Trends Reasoning, Part 1
Why does Electron Affinity increase across a period?
Moving left to right, the radius of the atom decreases as more
protons pull on more electrons.
Move across the
periodic table
e
e
n
p
n p
n p
e
e
Radius decreases
EA increases
Lithium atom
e
e
pn
p p
p
n nn
e
Beryllium atom
When an atom is smaller, the electrons are closer to the
nucleus, and therefore feel the pull more strongly.
A smaller atom can handle an extra electron more easily as it can
be more “controlled” by the closer nucleus
Electron Affinity Trends Reasoning, Part 2
Why does electron affinity decrease down a group?
As you move down a group, the radius increases as more
electrons shells are added.
e
e
e
+
Move down the
periodic table
Radius increases
e
e
e
e
e
e
EA decreases
e
e
e
Lithium atom
e
+
e
Sodium atom
As the outer electrons (those involved in bonding) are farther
from the nucleus, they will feel the “pull” of the nucleus less.
The larger atom is less able to “control” a new electron added.
Ionic Charge & Radii
Review Some Definitions
Ion – atom with a charge.
Cation – positively charged ion.
Results from loss of electrons.
Anion – negatively charged ion.
Results from gain of electrons.
Ionic Radii—Cations
How does the radius of a cation compare to the parent
Atoms lose electrons to create positive ions
atom?
e
e
e
+
Lithium atom
Creating a cation,
losing electrons
Radius decreases
e
e
+
Li+ ion
When electrons are lost, there are now more protons than
electrons
Therefore, the protons have a greater “pull” on each of the
electrons.
Ionic Radii—Anions
How does the radius of an anion compare to the parent
atom? Atoms gain electrons to create negative ions
e
e
e
e
Creating an anion,
gaining electrons
e
+
e
e
e
e
e
e
e
e
+
Radius increases
e
e
e
Oxygen atom
e
e
O2- ion
When electrons are gained, there are now more electrons
than protons
Therefore, the protons have a weaker “pull” on each of the
electrons.
Let’s Practice
Example:
List Li, Cs
and K in
order of
increasing
Atomic radii
Example:
List Li, N
and C in
order of
increasing
Atomic radii
Ionization Energy
Electron Affinity
Ionization Energy
Electron Affinity
Section 4.6—Light
Light is Electromagnetic Radiation
Electromagnetic energy is energy that has
electric and magnetic fields
There are many types of Electromagnetic
Radiation…visible is just one type!
Electromagnetic energy travels in waves
at the speed of light “c” (3.0 × 108 m/s)
Types of Electromagnetic Radiation
400 nm
700 nm
Visible light
Cosmic Gamma
X- Ultraviolet
Rays
Rays
Rays
Rays
10-12
10-10
10-8
Infrared Microwaves
Rays
10-6
10-4
10-2
Radio waves
1
102
Electric
Power
104
106
Wavelengths (cm)
Wave Properties—Frequency
Frequency () is the number of times a
wave completes a cycle in one second
(cycles per second is “Hertz” or “Hz”)
Lower frequency
Higher frequency
Wave Properties—Wavelength
Wavelength () is the distance from trough
to trough of a wave (measured in meters
“m”)
wavelength
Relationship between wave properties
The shorter the wavelength, the higher the
frequency
The higher the frequency, the higher the
energy
The speed of light is equal to wavelength
(in meters) times frequency (in sec-1 or s-1
or Hz)
c = 
Visible Range
Wavelength increases
Frequency decreases
Energy decreases
400 nm
700 nm
Visible light
White light is made of all the colors…a prism can separate white light into a rainbow!
Light is Quantized
Light is quantized, which means it comes
in packets—you can only have certain
amounts of it
The “packets” are called “photons”
The Energy of a photon (in Joules, “J”) is
equal to Planck’s constant “h” times
frequency)
E = h where h = 6.63 × 10-34 Js
Section 4.7—Light & Matter
Electrons Absorbing Energy
Photon coming into atom collides with
electron. Photons are energy.
+
Electrons Absorbing Energy
Photon coming into atom collides with
electron. Photons are energy.
+
The electron is “excited” to a higher energy level with is
newly increased energy from absorbing the photon.
Excitation
The process of an electron absorbing a
photon of light (energy) and being
promoted to a higher energy level from its
“ground state”
And later…
The electron cannot remain in that excited state indefinitely
+
And later…
The electron cannot remain in that excited state indefinitely
+
Energy is released during relaxation
Relaxation
The process of an electron releasing a
photon of light (energy) and falling back
down to a lower energy level.
Energy of photon and levels jumped
The higher the energy of the photon, the
greater the electron jump!
A photon of UV light has more energy than
a photon of Infrared light
The UV photon would cause a higher energy
jump (jump up more levels) than the IR photon.
Total energy in = Total energy out
However much energy was absorbed must
be released again, but it can be released
in smaller packets
A high energy photon might be absorbed,
but two lower energy photons might be
released as the electron falls in a “stepwise” manner.
Photons must match energy changes
The energy of the photon must exactly
match the energy change of the electron.
If the photon is not an exact match, the
photon will pass through unabsorbed.
+
Measuring light absorption
Not all
colors
come out
All
colors
of light
go in
Sample of hydrogen
Hydrogen Spectrum
The black bars are the colors that a hydrogen atom absorbs.
The other colors pass through the atom un-absorbed.
Absorption of Molecules
Because the structure of a molecule is
much more complicated than a single
atom, they absorb regions of light rather
than single wavelengths.
Absorption spectrum of water
Picture from http://www.lsbu.ac.uk/water/vibrat.html
Ways of producing light
Fluorescence: visible light is absorbed and
visible light is emitted at the same time—
the relaxation happens very quickly after
excitation
Phosphorescence: Visible light is
absorbed and then a while later is
emitted—relaxation occurs after a period
of time
Ways of producing light
 Incandescence: Energy is put in from heat and
given off as visible light
 Chemiluminescence: Energy released during a
chemical reaction is absorbed to cause
excitation. Relaxation produces visible light
 Biolouminescence: Chemiluminescence that
occurs in a biological organism.
 Triboluminescence: Physical pressure or torque
provides energy for excitation. Relaxation
produces visible light.
What did you learn about
glowing things?
Glow in the dark
Is based on
Electron
structure
Gives off
Changes in produce
Is a part of
Light
Which can
be determined using
Periodic
Table
Atomic
structure
Arranged to show
Is based on
Atomic
theory
Periodicity