What is Light?
electromagnetic radiation
• a form of energy
• travels in waves
• exists in increments called photons (a
packet of light)
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Electromagnetic radiation ranges from very low energy waves (TV and radio waves) to very
high energy waves (X-rays and gamma rays).
This broad continuum of electromagnetic energy is referred to as the electromagnetic
spectrum.
In the middle of this spectrum is a small range
of radiation that our eyes can detect, and that
we perceive as visible light. This narrow
range of radiation is the visible spectrum.
If we look more closely at the visible
spectrum, we can see that it is made up of
the colors of the rainbow: red, orange,
yellow, green, blue, and violet.
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Types of Electromagnetic Radiation
• Classified by the Wavelength
– Radiowaves = l > 0.01 m.
• Low frequency and energy.
– Microwaves = 10-4m < l < 10-2 m.
– Infrared (IR) = 8 x 10-7 < l < 10-5 m.
– Visible = 4 x 10-7 < l < 8 x 10-7 m.
• ROYGBIV.
– Ultraviolet (UV) = 10-8 < l < 4 x 10-7 m.
– X-rays = 10-10 < l < 10-8 m.
– Gamma rays = l < 10-10.
• High frequency and energy.
4
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Describing Electromagnetic Waves
amplitude – The height of one wave
wavelength (λ) – The length of one wave (nm length)
frequency (ν) – The number of waves per second
1 wave/second = 1 hertz (Hz)
10,000 Hz
IR
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10,000/s
10,000 s-1
UV
Describing Electromagnetic Waves
Wavelength and the frequency are inversely related to each other. This means that as the
wavelength decreases, the frequency increases, and vice versa.
wavelength
frequency
inversely related
c = ln
speed of light = wavelength x frequency
m
s
=
m x
1
s
c = speed of light = 3.00 x 108 m/s
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A beam of green light has a wavelength of 500 nm.
What is the frequency of this light?
c = ln
c = 3.00 x 108 m/s
l = 500 nm = 500 x 10-9 m
1 nm = 10-9 m
n = ?
c
= n
l
3.00 x 108 m/s
500 x 10-9 m
= n
units: 1/s = Hz
6 x 1014 Hz =
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n
Tro's "Introductory Chemistry", Chapter 9
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8
The energy of light depends on its frequency and wavelength.
longer wavelength
lower frequency
lower energy
shorter wavelength
higher frequency
higher energy
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Energy of a photon:
akg-images / bilwissediti
We can relate the energy of a single photon to
the frequency using the relationship E = h nu,
where E is the energy (measured in Joules), nu
is the frequency, and h is Planck’s constant,
which has a value of 6.63 x 10-34 Jˑs.
E = hν
energy
frequency
Planck’s constant
= 6.63 x 10-34 J.s
ν = c/ l
This constant is named after Max
Planck, a German physicist who
was instrumental in developing the
theories that relate light, energy,
and electron structure.
E = hc/l
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A photon has a frequency of 7.50 x 1014 Hz. What is the wavelength of
this light? What color is this light? What is the energy of the photon?
c = ln
c
n = l
3.00 x 108 m/s
=
14
7.50 x 10 /s
4.00 x 10-7 m =
= 400 nm
E = hn
l
E = (6.63 x 10-34 J.s)(7.50 x 1014/s)
l
violet
E = 4.97 x 10-19 J
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Summary
• Light is a form of electromagnetic radiation
• We describe light by its
• frequency (n)
• wavelength (l)
• energy (E)
• c = ln
• E = hn = hc/l
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Color, Line Spectra, and the Bohr Model
Lets explore the
relationship between light
and electronic structure.
Fireworks: The explosive
powder in fireworks often
contains metals that
produce distinctive colors
as the mixture burns.
(c) anterovium / depositphotos.com
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The Electromagnetic Spectrum
• Light passed
through a prism is
separated into all
its colors. This is
called a
continuous
spectrum.
• The color of the
light is determined
by its wavelength.
14
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Color
• The color of light is determined by its wavelength.
– Or frequency.
• White light is a mixture of all the colors of visible light.
– A spectrum ROYGBIV
– Red Orange Yellow Green Blue Indigo Violet.
• When an object absorbs some of the wavelengths of white light
while reflecting others, it appears colored.
– The observed color is predominantly the colors reflected.
Tro's "Introductory Chemistry", Chapter 9
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15
Flame tests
observe colors emitted by different metal ions
A wire is dipped in a solution containing metal ions.
The ions give off a characteristic color when heated.
For example, when a wire is dipped in a solution
containing calcium, the flame burns bright orange.
When dipped in a solution of copper, the flame burns
bright green.
Photo credits: GIPhotoStock/Science Source
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Gas lamps also produce unique colors:
We see a similar effect in gas lamps,
like those used in neon signs.
These lamps produce light by
passing an electric current through a
tube filled with a gas such as
helium, neon, argon, krypton, or
xenon.
Richard Megna / Fundamental Photographs
Like the metals in a flame test, each gas in a lamp produces a
characteristic color.
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Emission Spectrum
Unique to each
element; only
certain
wavelengths of
light are given off.
18
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Each element produces a unique line spectrum.
He
Li
Kr
Scientists often use line spectra as
“fingerprints” to identify elements.
For example, analyzing light from the
sun shows fine lines that correspond
to the spectral lines of hydrogen and
helium.
From this, they know that the sun and
other stars are composed largely of
these two elements,
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Revell Introductory Chemistry, © 2018 Macmillan Learning
Early 20th Century:
• Dense nucleus surrounded by electrons
• Photoelectric effect: light causes atoms to eject electrons
This effect showed that
light energy was
somehow connected to
electron structure.
But how did this relate to
the line spectra?
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The Bohr Model (1913)
•
•
•
•
Electrons orbit the nucleus.
Only certain orbit energies are “allowed”.
Electrons can jump between levels.
Light is absorbed or released when
electrons jump.
• Ground state: all electrons in lowest
possible levels.
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• To help think about the Bohr model, imagine
standing on a staircase.
• You can jump from one step to another, but
you can’t hover between steps – you will
always land on one of the steps.
• Similarly, electrons only occupy specified
energy “steps”.
• Electrons can absorb energy to
move to a higher level (step),
or release energy to move to a
lower level.
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releases energy
(blue light)
• When the atom gains energy, the
electron leaps from a lower energy orbit
to one that is further from the nucleus.
– However, during that “quantum leap” it
doesn’t travel through the space between
the orbits, it just disappears from the lower
orbit and appears in the higher orbit.
releases energy
(UV light)
absorbs energy
-
• When the electron leaps from a higher
energy orbit to one that is closer to the
nucleus, energy is emitted from the atom
as a photon of light—a quantum of
energy.
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In a H atom, four transitions produce visible light, resulting in
four spectral lines.
Notice that the smallest transition produces the lowest energy
light (red), while the biggest transition produces the highest
energy light (purple).
Other transitions also take place, but they produce radiation
that is either too high in energy (ultraviolet) or too low in
energy (infrared) for our eyes to detect.
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Bohr Model
Explained
• The hydrogen line spectrum (Framework of
electron structure)
• Some properties of main group elements
Did not explain
• More complex line spectra (Elements larger
than hydrogen)
• Properties of the transition elements
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The Quantum Model and Electron Orbitals
Bohr Model:
1913
• Erwin Schrödinger applied the Quantum Model:
mathematics of probability and
1920s-30s
the ideas of quantizing energy
to the physics equations that
describe waves, resulting in an
equation that predicts the
probability of finding an
electron with a particular
amount of energy at a particular
location in the atom.
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Heisenberg’s Uncertainty Principle
It is impossible to precisely
know the exact velocity and
location of a particle.
• The result is a map of regions in
the atom that have a particular
probability for finding the electron.
• An orbital is a region where we
have a very high probability of
finding the electron when it has a
particular amount of energy.
– Generally set at 90 or 95%.
Quantum mechanics: describes electrons
most probable locations
energies
Photo credits: Ted Kinsman/Science Source
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The wave nature of electrons
Tiny, fast-moving particles
also behave as waves.
Ted Kinsman/Science Source
This explains electron energy levels.
?
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The Quantum Model
Main Ideas:
• uncertainty principle
• wave nature of electrons
QM describes electrons by
• energy
• probable locations
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Energy Levels and Sublevels
1. Electrons occupy different energy levels.
• Level is identified by its principal quantum number, n (1, 2, 3…)
• Higher energy levels can hold more electrons
Level
1
2
3
4
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Electron
Capacity
2
8
18
32
Energy Levels and Sublevels
2. Each energy level contains one or more sublevels.
Sublevel
s
p
d
f
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Energy Levels and Sublevels
3. Each sublevel contains one or more orbitals.
Orbital: a region where electrons are most likely to
be found.
Sublevel
s
p
d
f
Number of
Orbitals
1
3
5
7
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Orbits: Bohr model
describes the pathways of
the particles orbiting the
nucleus.
Orbitals: describes the
region around the atom
where the electron is most
likely to be.
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Revell Introductory Chemistry, © 2018 Macmillan Learning
Energy Levels and Sublevels
4. Each orbital holds up to two electrons.
– Electrons have a magnetic field, called spin.
– Electrons with opposite spins pair together.
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Energy Levels and Sublevels
1.
2.
3.
4.
Electrons occupy different energy levels.
Each level contains sublevels.
Each sublevel contains orbitals.
Each orbital holds up to two electrons.
Sublevel
Number of
Orbitals
Electron
Capacity
s
p
d
f
1
3
5
7
2
6
10
14
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Level 1: s only
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Level 2: s + p
As you go to higher
levels, the orbitals
INCREASE in size.
A 1s orbital is
smaller than a 2s
orbital is
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Level 2: s + p
x
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Level 2: s + p
Sublevel
s
p
Number of
Orbitals
1
3
Total:
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Electron
Capacity
2
6
8
Level 3: s + p + d
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Level 3: s + p + d
Sublevel
s
p
d
Number of
Orbitals
1
3
5
Electron
Capacity
2
6
10
Total:
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18
Level 4: s + p + d + f
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Level 4: s + p + d + f
Sublevel
s
p
d
f
Number of
Orbitals
1
3
5
7
Total:
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Electron
Capacity
2
6
10
14
32
Summary
Energy sublevels follow a predictable pattern: Each new level adds a new sublevel, and each new sublevel has two more
orbitals (four more electrons) than the one before.
Above energy level four, the trend continues: Level 5 has five sublevels, level 6 has six sublevels, etc. However, even the
largest elements on the periodic table fit all of their electrons within the s, p, d, and f sublevels, so we don’t need to
worry about any sublevels beyond these four.
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Aufbau Diagram:
This diagram shows the relative energy differences
between the different energy levels and sublevels.
The s orbital is the lowest in energy, followed by the p,
d, and f.
As the energy levels get higher, they become more
closely grouped together.
As the energy levels branch out into more sublevels,
they actually overlap each other.
Ex: energy level 4s is actually lower than energy level
3d.
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Describing Electron Configuration
Quantum Model:
Energy levels – 1, 2, 3…
Energy sublevels – s, p, d, f
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Filling an Orbital with Electrons
• Each orbital may have a maximum of 2
electrons.
– Pauli Exclusion principle.
• Electrons spin on an axis.
– Generating their own magnetic field.
• When two electrons are in the same
orbital, they must have opposite spins.
– So their magnetic fields will cancel.
Tro's "Introductory Chemistry", Chapter 9
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50
Tro's "Introductory Chemistry", Chapter 9
51
Orbital Diagrams
• We often represent an orbital as a square and
the electrons in that orbital as arrows.
– The direction of the arrow represents the spin of
the electron.
Unoccupied
orbital
Orbital with
1 electron
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Orbital with
2 electrons
As a general rule, electrons
occupy the lowest available
energy levels- this is called
the ground state.
Aufbau Principle
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Hydrogen:
electron configuration: 1s1
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Helium:
1s2
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Lithium:
1s22s1
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Beryllium:
1s22s2
e-
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Boron:
1s22s22p1
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Hund’s Rule:
If empty orbitals of the same
energy are available, electrons
singly occupy orbitals rather
than pairing together.
Carbon:
1s22s22p2
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Li:
1s22s1
B:
1s22s22p1
Be:
1s22s2
C:
1s22s22p2
N:
1s22s22p3
O:
1s22s22p4
F:
1s22s22p5
Ne:
1s22s22p6
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What is the electron configuration of silicon?
14 e- total
1s22s22p63s23p2
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Examples:
What is the electron configuration for:
1. Phosphorus
2. Potassium
3. Iron
4. Selenium
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Revell Introductory Chemistry, © 2018 Macmillan Learning
Describing Electron Configuration
(part 2)
valence level: The highest
occupied electron energy level
• Up to 8 electrons in valence level
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Argon: (18 e-)
1s22s22p63s23p6
valence level
Potassium: (19 e-)
1s22s22p63s23p64s1
valence level
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Noble Gases have Filled Valences
1s2
1s22s22p6
Octet Rule:
An atom is stabilized by
having its highest-occupied
(valence) energy level filled.
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
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Electron Configurations for Larger Atoms
inner electrons
Sodium:
Phosphorous:
Chlorine:
Noble gas
notation
1s22s22p63s1
[Ne]3s1
1s22s22p63s23p3
[Ne]3s23p3
1s22s22p63s23p5
[Ne]3s23p5
1s22s22p6 = [Ne]
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Electron Configurations for Larger Atoms
not involved with bonding
inner
electrons
Titanium: [Ar]4s23d2
valence
+ d, f sublevels
essential to bonding
outer
electrons
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Write the electron configuration for selenium using the noble gas shorthand.
Identify the inner electrons, the outer electrons, and the valence electrons.
1s22s22p63s23p64s23d104p4
[Ar]
inner
outer
[Ar]4s23d104p4
valence
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Electron Configurations for Ions
What is the electron configuration of a sodium atom?
What is the electron configuration of a sodium ion with a +1 charge?
full
configuration
noble-gas
shorthand
species
Symbol
sodium atom
Na
1s22s22p63s1
[Ne]3s1
sodium ion (+1 charge)
Na+
1s22s22p6
[He]2s22p6 or [Ne]
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Electron Configurations for Ions
What is the electron configuration of an oxide ion,
which is an oxygen ion with a charge of -2?
species
symbol
full
configuration
oxygen atom
O
1s22s22p4
oxide ion (-2 charge)
O2-
1s22s22p6
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noble-gas
shorthand
[He]2s22p4
[He]2s22p6 or [Ne]
Many ions form noble gas configurations
O: 1s22s22p4
Na: 1s22s22p63s1
O2–: 1s22s22p6
Na+: 1s22s22p6
Ne: 1s22s22p6
These are isoelectronic
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Revell Introductory Chemistry, © 2018 Macmillan Learning
Electron Configuration
and the
Stock Montage/Getty Images
Periodic Table
Sovfoto/Getty Images
Revell Introductory Chemistry, © 2018 Macmillan Learning
Lithium
[He]2s1
(3 electrons):
Sodium
[Ne]3s1
(11 electrons):
Potassium
[Ar]4s1
(19 electrons):
Photo credits from top: SPL/Science Source; SPL/Science Source; Andrew Lambert Photography/Science Source
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Fluorine:
[He]2s22p5
Chlorine:
[Ne]3s23p5
Bromine:
[Ar]4s23d104p5
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The row indicates the
highest occupied
electron energy level.
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The column gives the
outermost electron
configuration.
s1
1
s2
p6
p1
p2
p3
p4
p5
f10 f11
f12
f13 f14
2
3
d1
d2 d3
d4
d5
d6
d7
d8
d9 d10
f1
f3
f4
f5
f6
f7
f8
4
5
6
7
f2
f9
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Revell Introductory Chemistry, © 2018 Macmillan Learning
Revell Introductory Chemistry, © 2018 Macmillan Learning
What is the outermost electron configuration for sulfur?
3p
4
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Write the configuration for the highest-energy occupied sublevel
for potassium, phosphorus, and iron.
K: 4s1
P: 3p3
Fe: 3d6
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Write the electron configuration for aluminum.
How many valence electrons does aluminum have?
[Ne] 3s2 3p1
3 valence electrons
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C
Si
Ge
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The column gives
the outermost
electron
configuration.
The row indicates
the highest
occupied electron
energy level.
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Revell Introductory Chemistry, © 2018 Macmillan Learning