1
A Brief History
2
Lucretius
On the nature of the Universe (55 BC)
• The light and heat of the sun; these are
composed of minute atoms which, when they
are shoved off, lose no time in shooting
right across the interspace of air in the
direction imparted by the shove.
http://www.gap-system.org/~history/HistTopics/Light_1.html
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Electromagnetic Radiation
4
5
Properties of Light
• Travels in straight lines. Casts shadows
• Carries energy. Heats objects it strikes.
• Travels fast. No delay when turning on
a light switch.
• Comes in different colors.
6
Christiaan Huygens
1678
7
http://www.launc.tased.edu.au/online/sciences/physics/diffrac.html
Huygens’ Principle
• Throw a rock in a quiet pool, and
waves appear along the surface
of the water.
• Huygens proposed that the
wavefronts of light waves
spreading out from a point source
can be regarded as the
overlapped crests of tiny
secondary waves.
• Wavefronts are made up of tinier
wavefronts—this idea is called
Huygens’ principle.
8
Huygens’ Principle
Every point of a wavefront may be considered the
source of secondary wavelets that spread out in all
directions with a speed equal to the speed of
propagation of the waves.
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Huygens’ Principle
• Plane waves can be generated in
water by successively dipping a
horizontally held straightedge into
the surface
• As the width of the opening is narrowed, less of the incident
wave is transmitted.
• The spreading of waves into the shadow region becomes
more pronounced.
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Characteristics of a Wave
11
Wavelength (λ)
12
Light has the properties of a wave.
wavelength
wavelength
(measured
from
(measured
from
peak totrough
peak) to trough)
13
Frequency (ν)
14
Frequency is the number of wavelengths
that pass a particular point per second.
15
Diffraction
16
Diffraction
17
Thomas Young
18
Constructive Interference
19
Destructive Interference
20
The interference of water is a common sight. In some places, crests
overlap crests; in other places, crests overlap troughs of other waves.
Interference patterns
of overlapping
waves from two
vibrating sources.
Thomas Young’s original drawing of a two-source interference pattern.
The dark circles represent wave crests; the white spaces between crests
represent troughs. Constructive interference occurs where crests overlap
crests or troughs overlap troughs. Letters C, D, E, and F mark regions of
destructive interference.
Interference Patterns
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Young’s Experiment 1802
http://www.acoustics.salford.ac.uk/feschools/waves/diffract3.htm
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Young’s Experiment 1802
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28
Max Planck
29
• In 1901, Planck developed the quantum
theory of light.
• Light is not continuous.
• Light is a stream of distinct units of energy
called quanta.
• The amount of energy on one of these units
depends on the frequency.
• Light is absorbed or emitted in multiples of
whole quanta.
• Quanta is like an atom of light called a
photon
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ROYGBIV
31
“I think it is safe to say that
no one understands quantum mechanics. Do not keep saying to
yourself, if you can possibly avoid it, ‘But how can it be like that?’
because you will go ‘down the drain’ into a blind alley from which
nobody has yet escaped. Nobody knows how it can be like that.”
Richard Feynman
1918-1988
The Manhattan Project
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The Electromagnetic Spectrum
33
visible light X-rays
Infrared light
410-7
102
Radio waves
1
10-2
10-4
Microwave Infrared
710-7
visible
104
10-8
10-10
Ultraviolet X-rays
10-12
Gamma
rays
Wavelength meters
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35
The Bohr Atom
36
Niels Bohr, a Danish physicist,
in 1912-1913 carried out research
on the hydrogen atom
37
A young man working with Rutherford in 1911 was Niels Bohr. They
created a halfway successful model of the atom called the RutherfordBohr model. They imagined a nucleus at the center with electrons
orbiting around it. Twenty-five years later, Bohr described his
collaboration with Rutherford...
If, twenty-five years ago, I had the good fortune to give a modest
contribution to this development, it was, above all, thanks to the
hospitality I then, as a young man, enjoyed in the famous laboratories
of England. In particular, I think with grateful emotion of the unique
friendliness and straightforwardness with which Rutherford, in the
midst of his unceasing creative activity, was always prepared to listen
to any student behind whose youthful inexperience he perceived a
serious interest.
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• At high temperatures or voltages,
elements in the gaseous state emit light
of different colors.
• When the light is passed through a
prism or diffraction grating, a line
spectrum results.
39
Sir Isaac Newton’s crucial experiment, 1672.
he refracts white light with a prism, resolving
it into its component colors: red, orange,
yellow, green, blue and violet.
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41
He
Ne
Ar
Kr
Xe
42
Flame test colours
Ba
Li
Sr
Na
Cu
K
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Diffraction grating
• Composed of a large number of close, equally
spaced slits for analyzing light source
• Produced by spectrometers that disperse white
light into colors
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The lenses of “rainbow” glasses
are diffraction gratings. Looking
through them causes light to
separate into its color
components. The spectral
patterns seen from light sources
such as fireworks and neon signs
emit a distinct number of
discontinuous colors and do not
show all the colors of the
rainbow. The spectral patterns
from these light sources are the
atomic spectra of elements
heated in the light sources.
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Lamps of a chandelier seen through diffraction-grating glasses.
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Each element has its own
unique set of spectral emission
lines that distinguish it from
other elements.
These colored lines
indicate that light is
being emitted only at
certain wavelengths.
Line spectrum of hydrogen. Each line corresponds
to the wavelength of the energy emitted when the
electron of a hydrogen atom, which has absorbed
energy, falls back to a lower principal energy level.
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Incandescent bulb has a continuous spectrum
Hydrogen
Sodium
Mercury
The Bohr Atom
50
Electrons
revolve
An
electron
has a
around the
nucleus
in it
discrete
energy
when
orbits thatan
areorbit.
located
occupies
at fixed distances from
the nucleus.
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The color
When
an electron
of the
fallslight
from a higher
emitted
corresponds
energy level
to
to a lower
one
of the
energy
lines level
of the
a
hydrogen
spectrum.
quantum of
energy in the
form of light is emitted by
the atom.
52
Different lines of the
hydrogen spectrum
correspond to different
electron energy level
shifts.
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Light is not emitted
continuously. It is
emitted in discrete
packets called quanta.
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E1
E2
E3
An electron can have
one of several possible
energies depending on
its orbit.
55
• Bohr’s calculations succeeded very well
in
correlating
the
experimentally
observed spectral lines with electron
energy levels for the hydrogen atom.
• Bohr’s methods did not succeed for
heavier atoms.
• More theoretical work on atomic structure
was needed.
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Baseball Waves
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• In 1924 Louis De Broglie suggested
that all objects have wave properties.
– De Broglie showed that
the wavelength of
ordinary-sized objects,
such as a baseball, are too
small to be observed.
– For objects the size of an
electron the wavelength
can be detected.
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In this view, the electron is thought of not as a particle
located at some point in the atom but as if its mass and
charge were spread out into a standing wave surrounding
the atomic nucleus with an integral number of
wavelengths fitting evenly into the circumferences of the
orbits.
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(a) An orbiting electron forms a standing wave only when the
circumference of its orbit is equal to a whole-number multiple of
the wavelength.
(b) When the wave does not close in on itself in phase, it undergoes
destructive interference. Hence, orbits exist only where waves
close in on themselves in phase.
The circumference of the innermost orbit,
according to this picture, is equal to one
wavelength.
• The second has a circumference of two
electron wavelengths.
• The third orbit has a circumference of three
electron wavelengths, and so forth.
• The electron orbits in an atom have discrete radii because the
circumferences of the orbits are whole-number multiples of the
electron wavelength.
• This results in a discrete energy state for each orbit.
• This model explains why electrons don’t spiral closer
and closer to the nucleus, causing atoms to shrink to the
size of the tiny nucleus.
• If each electron orbit is described by a standing wave,
the circumference of the smallest orbit can be no smaller
than one wavelength.
– No fraction of a wavelength is possible in a circular (or
elliptical) standing wave.
• As long as an electron carries the momentum necessary
for wave behavior, atoms don’t shrink in on themselves.
Wire loop at rest.
For a fixed
circumference, only
some wavelengths are
self-reinforcing.
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Tacoma Narrows Bridge Collapse
"Gallopin' Gertie" Nov. 7, 1940
• http://www.youtube.com/watch?v=j-zczJXSxnw
Davisson & Germer 1927
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The diffraction on the left was made by a beam of
x-rays passing through thin aluminum foil. The
diffraction on the right was made by a beam of
electrons passing through the same foil.
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Baseball Waves
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• In 1926 Erwin Schrödinger created a
mathematical model that showed
electrons as waves.
72
– Schröedinger’s work led to a new branch
of physics called wave or quantum
mechanics.
– Using Schröedinger’s wave mechanics,
the probability of finding an electron in a
certain region around the atom can be
determined.
– The actual location of an electron within
an atom cannot be determined.
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• Based on wave mechanics it is clear that
electrons are not revolving around the
nucleus in orbits.
• Instead of being located in orbits the
electrons are located in orbitals.
• An orbital is a region around the nucleus
where there is a high probability of
finding an electron.
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Energy Levels
of Electrons
75
TheAccording
wave-mechanical
model
of
the
atom
to Bohr the energies of
also predicts discrete principal energy
electrons in an atom are quantized.
levels within the atom
76
As n increases, the
energy of the electron
increases.
The first four
principal energy
levels of the
hydrogen atom.
Each level is
assigned a principal
quantum number n.
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Each principal energy level
is subdivided into sublevels.
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Within sublevels the electrons are found in
orbitals.
An s orbital is spherical in
shape (s from sharp).
The spherical surface
encloses a space where
there is a 90% probability
that the electron may be
found.
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An atomic orbital can hold a maximum of two
electrons.
An electron can spin in one
of two possible directions
represented by ↑ or ↓.
The two electrons that
occupy an atomic orbital
must have opposite spins.
This is known as the Pauli
Exclusion Principal.
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A p sublevel is made up of three orbitals.
Each p orbital has two lobes.
Each p orbital can hold a maximum of two
electrons.
A p sublevel can hold a maximum of 6
electrons.
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pz
The three p orbitals share
a common center.
py
px
The three p orbitals point in
different directions.
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A d sublevel is made up of five orbitals.
The five d orbitals all point in different directions.
Each d orbital can hold a maximum of two
electrons.
A d sublevel can hold a maximum of 10 electrons.
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Number of Orbitals in a Sublevel
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f Orbitals
85
Web Sites
• http://web.mit.edu/3.091/www/orbs/
• http://winter.group.shef.ac.uk/orbitron/
AOs/1s/index.html
87
Distribution of Sublevels by Principal Energy Level
Main
energy
Level (n)
Maximum
Sublevels (n)
Maximum Maximum
Orbitals Electrons
(n2)
(2n2)
1
1: s
12 = 1
2
2
2: s and p
22 = 4
8
3
3: s, p, and d
32 = 9
18
4
4: s, p, d, and f 42 = 16
32
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Atomic Models
John Dalton’s atom 1803
Bohr’s orbit model 1913
Thomson’s pudding model 1897
Rutherford’s nuclear model 1911
Schrödinger electron cloud model 1926
Atomic Structure of the
First 18 Elements
90
To determine the electronic structures of
atoms the following guidelines are used.
91
1. No more than two
electrons can occupy
one orbital and then
only if they have
opposite spin. This
is the Pauli
Exclusion Principle.
92
Wolfgang Ernst Pauli
93
1 s orbital
2 s orbital
2. Electrons occupy the lowest energy orbitals
available. They enter a higher energy orbital only
after the lower orbitals are filled. This is called the
Aufbau Principle.
3. For the atoms beyond hydrogen, orbital energies
vary as s<p<d<f for a given value of n.
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4. Each orbital on a sublevel is occupied by a
single electron before a second electron
enters. This is Attila the Hun’s Rule
(Hund’s Rule). For example, all three p
orbitals must contain one electron before a
second electron enters a p orbital.
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Friedrich Hund
1896-1997
96
Attila the Hun
97
Nuclear makeup and electronic structure of
each principal energy level of an atom.
number of protons
and of electrons
number
neutrons in the nucleus
in each sublevel
98
Electron Configuration Using
Spectroscopic Notation
Number of
electrons in
sublevel orbitals
Arrangement of
electrons within their
respective sublevels.
6
2p
Principal
Type of orbital
energy level
99
Orbital Filling
100
• In the following diagrams, boxes
represent orbitals.
• Electrons are indicated by arrows: ↑ or
↓.
– Each arrow direction represents one of
the two possible electron spin states.
101
Filling the 1s Sublevel
106
H
↑
1s1
Hydrogen has 1 electron. It will occupy the orbital of lowest
energy which is the 1s.
He
↑↓
1s2
Helium has two electrons. Both helium electrons occupy the
1s orbital with opposite spins.
107
Abbreviated Electron Configurations
Use the symbol of the nearest preceding noble
gas to represent the electron configuration of the
core electrons.
Phosphorus: 1s2 2s2 2p6 3s2 3p3
Core
Electrons
[Ne] 3s2 3p3
Valence
Electrons
108
Filling the 2s Sublevel
109
Li
↑↓
↑
1s
2s
1s22s1
[He]2s1
The 1s orbital is filled. Lithium’s third electron will enter the
2s orbital.
Be
↑↓
↑↓
1s
2s
1s22s2
[He]2s2
The 2s orbital fills upon the addition of beryllium’s third and
fourth electrons.
110
Filling the 2p Sublevel
111
B
↑↓
↑↓
1s
2s
↑
1s22s22p1
[He]2s22p1
2p
Boron has the first p electron. The three 2p orbitals have the same
energy (degenerate). It does not matter which orbital fills first.
C ↑↓
↑↓
1s
2s
↑
↑
1s22s22p2
[He]2s22p2
2p
The second p electron of carbon enters a different p orbital than the
first p electron so as to give carbon the lowest possible energy.
N ↑↓
↑↓
1s
2s
↑
↑
2p
↑
1s22s22p3
[He]2s22p3
The third p electron of nitrogen enters a different p orbital than its
112
first two p electrons to give nitrogen the lowest possible energy
.
O ↑↓
↑↓
1s
2s
↑↓ ↑
↑
2p
1s22s22p4
[He]2s22p4
There are four electrons in the 2p sublevel of oxygen. One of the
2p orbitals is now occupied by a second electron, which has a spin
opposite that of the first electron already in the orbital.
F
↑↓
↑↓
↑↓ ↑↓ ↑
1s
2s
2p
1s22s22p5
[He]2s22p5
There are five electrons in the 2p sublevel of fluorine. Two of the 2p
orbitals are now occupied by a second electron, which has a spin
opposite that of the first electron already in the orbital.
113
Ne ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
1s22s22p6
There are 6 electrons in the 2p sublevel of neon, which fills the
sublevel.
114
Filling the 3s Sublevel
115
Na ↑↓
↑↓
↑↓ ↑↓ ↑↓
↑
1s
2s
2p
3s
1s22s22p63s1
[Ne]3s1
The 2s and 2p sublevels are filled. The next electron enters the
3s sublevel of sodium.
Mg ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
↑↓ 1s22s22p63s2
2
[Ne]3s
3s
The 3s orbital fills upon the addition of magnesium’s twelfth
electron.
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117
118
Electron Structures and
the Periodic Table
119
In 1869 Dimitri Mendeleev of Russia and
Lothar Meyer of Germany independently
published periodic arrangements of the
elements based on increasing atomic
masses.
Mendeleev’s arrangement is the precursor
to the modern periodic table.
120
Dimitri Mendeleev
121
122
Glenn Seaborg
1912-1999
123
Period numbers correspond
Horizontal rows are
to the highest occupied
called periods.
energy level.
124
Elements
with
similar
Elements
in
the
B groups
groups
in
A
Groups are
numbered
properties
arenumerals.
organized
are
designated
transition
are
designated
with
Roman
in groups or families.
elements.
representative
elements.
125
Some A groups have
Alkaline earth metals family names.
Alkali metals (H not included)
Chalcogens
Halogens
Noble gases
126
127
128
TheForchemical
A family elements
behaviortheand
valence
properties
electron of
elements
configuration
in a family
is the same
are associated
in each column.
with the
electron configuration of its elements.
129
With the exception of helium which has a filled s
orbital, the nobles gases have filled p orbitals.
130
d orbital numbers are 1 less
than dthe
period
number
orbital
filling
Arrangement of electrons
according to sublevel being filled.
131
f orbital numbers are 2 less
than the
period
number
f orbital
filling
Arrangement of electrons
according to sublevel being filled.
132
Period number corresponds with the
highest energy level occupied by
electrons in that period.
133
The
Thegroup
elements
numbers
of a family
for thehave
representative
the same
outermost
elements electron
are equal
configuration
to the total number
except that
of
outermost
the electrons
electrons
are inindifferent
the atoms
energy
of the
levels.
group.
134
135