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Title: Energy and Electrons

1
EnergyandElectrons
Part One
Pisgah High School M. Jones
Rev 11/05/08
2
Goals for this unit Know about
1. the lines in the hydrogen spectrum, and Bohrs
atomic theory,
2. the arrangement of electrons in atoms, and the
shape of the periodic table,
3. energy diagrams and electron configurations,
4. valence electrons and dot diagrams.

3
Light and Energy
4
Some background terms and concepts
5
Frequency
• Symbol f or f (or n Greek nu)
• Units Hertz (Hz) or 1/sec
• Pronounced reciprocal seconds
• Historically, frequency was in units of cycles
per second, but this made too much sense.

6
Wavelength
• Symbol l (Greek lambda)
• Units meters (m)
• l of radio waves in meters to micrometers. l of
light in nanometers (10-9 m).
• Wavelengths of light can be measured with a
spectroscope.

7
The sine curve is often used to represent wave
motion.
• Look at the following graphs.
• What is the relationship between frequency and
wavelength?

8
O n e w a v e l e n g t h
9
One wavelength
10
One wavelength
11
One wave-length
12
What is the relationship between frequency and
wavelength?
• As the frequency increases, the wavelength
decreases.

13
The relationship between frequency and wavelength
can be represented by
f frequency l (lambda) wavelength
14
The relationship between frequency and wavelength
can be represented by
Frequency is inversely proportional to wavelength
15
The relationship between frequency and wavelength
can be represented by
k f l
Frequency times wavelength equals a constant
16
For electromagnetic energy, the equation is
c f l
c is the speed of light
c 3.00 x 108 m/sec
17
The equation can also be written as
c ln
c speed of light l wavelength n
frequency
18
Light is part of the electromagnetic spectrum
19
Electromagnetic Spectrum
Longer wavelength Lower frequency Lower energy
Shorter wavelength Higher frequency Higher energy
20
Electromagnetic Spectrum
400 nm
700 nm
Longer wavelength Lower frequency Lower energy
Shorter wavelength Higher frequency Higher energy
21
The visible spectrum was discovered by
• Dr. Roy G. Biv

Red Orange Yellow Green Blue Indigo Violet
700 nm Lower energy
400 nm Higher energy
22
Electromagnetic waves carry energy.
23
The energy in an electromagnetic wave is
directly proportional to the frequency,
inversely proportional to the wavelength.
24
Energy in an EM wave
Small wavelength large energy
Large wavelength small energy
25
Next, a demonstration
• Look at sunlight and fluorescent light through
the spectroscope.
• Use the spectroscope to look at the light coming
from various gas discharge tubes.
• Record the colors and their order in the spectrum
of hydrogen.

26
Stop Complete the observation of atomic spectra,
then continue.
27
What did you see?
Hydrogen Spectrum
28
And now something completely different, some
history.
29
J. J. Thomson discovered the electron in 1897.
Thomson suggested the plum pudding model with
many, many electrons throughout the atom.
30
J. J. Thomson discovered the electron in 1897.
Thomson suggested the plum pudding model with
many, many electrons throughout the atom.
31
Ernest Rutherford explained that atoms had a
small, dense, positively nucleus.
Rutherford suggested in 1911 that electrons might
exist outside the nucleus in a planetary
arrangement.
Hantaro Nagaoka Japan, 1904
32
Enter Niels Bohr
Niels Bohr, with a brand new PhD in physics from
U. Copenhagen, went first to Thompson, then to
Rutherford to study in 1912.
Bohr was there right after the gold foil
explanation was published.
Bohr also knew about Rutherfords planetary
model of the electrons.
33
Niels Bohr
While at Manchester, Niels Bohr studied the
spectra of elements, and how these might relate
to the internal structure of atoms.
In 1913 he proposed a structure of the atom with
Rutherfords nucleus and electrons in discrete
energy levels.
34
Niels Bohr
Niels Bohr knew of the work done by Max Planck
and incorporated it into his atomic theory.
He used this new quantum theory to help explain
how energy could be absorbed and emitted.
35
Max Planck
• In 1900 Max Planck introduced an unusual idea.
• Energy exists in packets, or quanta.
• The beginnings of Quantum Theory.
• A quantum of energy is the smallest amount of
energy possible.
• Energy exists only in multiples of these quanta.

36
Albert Einstein
• In 1905 Einstein used Max Plancks idea of quanta
to explain the photoelectric effect.

Gave credibility to quantum theory.
37
Albert Einstein
Einstein explained with the Quantum Theory what
could not be explained with classical physics.
A photon is a packet of light, or quantum of
energy.
A photon of just the right energy can knock an
electron out of an atom.
38
Niels Bohr
Bohr said that electrons could exist only in
certain discrete energy levels, and
that electrons can only change energy levels
when they absorb or give off a certain amount of
energy. (1913)
39
Hydrogen atom
40
Hydrogen atom
Electrons can exist at this level,
41
Hydrogen atom
or in this level,
42
Hydrogen atom
or in this level,
43
Hydrogen atom
but not in between the levels.
44
Hydrogen atom
Unless the electron is absorbing energy, or
Giving off energy
45
Niels Bohr
When high voltage is connected to the hydrogen
discharge tube, a bluish light is given off.
When observed through a diffraction grating,
specific lines of color are observed.
Why ???
46
Niels Bohr
The electron in a hydrogen atom gains energy from
the electricity passing thru the tube and
the electron moves up to a higher energy level.
47
Niels Bohr
The electron in the excited state is unstable.
The electron drops to a lower energy level, and
gives off light of a certain energy and
wavelength.
48
Dont forget
and
The wavelength is inversely proportional to the
energy
49
Hydrogen atom
50
Hydrogen atom
The electron absorbs energy and
energy
51
Hydrogen atom
the electron is elevated to the next energy
level
52
Hydrogen atom
a lower energy level
53
Hydrogen atom
Light of a particular wavelength is given off
54
A line in the hydrogen spectrum is produced when
an electron moves from higher energy level to a
lower one.
Each line has a wavelength and color that
corresponds to the difference in energy between
the two levels.
55
Dont forget
and
The wavelength is inversely proportional to the
energy
56
Energy and l
DE is large
Shorter l
DE is small
Longer l
57
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
58
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
Suppose an electron is in level 5 and drops to
level 2.
59
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
Then purple light with a wavelength of 434 nm
will be emitted.
60
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
Suppose an electron is in level 4 and drops to
level 2.
61
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
Then you get blue-green light with a wavelength
of 486 nm
62
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
Suppose an electron is in level 3 and drops to
level 2.
63
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
Then red light with a wavelength of 656 nm will
be emitted.
64
5 4 3 2 1
Hydrogen atom
The energy levels are numbered
The colors of the visible lines come from the
energy given off by electrons moving from higher
energy levels down to level 2.
65
Dont forget
and
The wavelength is inversely proportional to the
energy
66
The Hydrogen Spectrum
Each color represents the transitions of
gazillions of electrons in gazillions of H atoms
going from higher energy levels to the second
energy level.
67
There must be a lot more lines in the spectrum of
hydrogen.
There can be transitions among seven energy
levels.
Why cant we see them?
68
Transitions in the H-spectrum
7 6 5 4 3 2 1
Transitions to the first energy level produce
ultraviolet lines.
69
Transitions in the H-spectrum
7 6 5 4 3 2 1
Transitions to the second energy level produce
visible lines.
70
Transitions in the H-spectrum
7 6 5 4 3 2 1
Transitions to the third and higher energy levels
produce infrared lines.
71
Bohr successfully calculated the wavelengths of
all the transitions in the hydrogen spectrum.
72
But only the hydrogen spectrum.The spectra of
elements with more than one electron could not be
accurately predicted.
73
Regardless of its shortcomings and the
modifications that were later applied, Bohrs
model of the atom was the first successful
attempt to make the internal structure of the
atom agree with spectroscopic data.
Asimov, 1964
74
The Arrangement of Electrons in the Quantum
Mechanical Model of the Atom
75
The Modern View of the Atom
1. A small, dense positively charged nucleus which
contains protons and neutrons.

76
The Modern View of the Atom
1. Electrons which exist outside of the nucleus at
1. various distances from the nucleus, and at
2. various energy levels.

77
The Electrons
1. The electrons can have both a mass, as does
matter, and a wavelength, as does light energy.

78
The Electrons
1. The electrons themselves are not little solid
spheres in orbit around the nucleus, but exist as
a fog of half-energy, half-matter. The
electrons can behave as either matter or energy,
depending on the experiment.

79
Energy Levels
1. Based on the ideas of Bohr, the electrons are
located
• in major energy levels,
• in energy sublevels within major energy levels,
• in orbitals within each sublevel.

80
The energy levels are like an organizational
81
Quantum Numbers
1. Each electron in an atom has a set of four (4)
quantum numbers.
2. The quantum numbers are like an address name,
street, city, state.
3. Pauli exclusion principle no two electrons in
the same atom can have the same set of quantum
numbers.

82
Quantum Numbers
Address Quntm. Sym. What it tells
State Principal n Major energy level
City Azmuthal L Sublevel
Street Magnetic ML Orbital
Name Spin MS Which e- in orbital
Luckily, we will only deal with the Principal
Quantum Number
83
Whats coming next?
1. 2n2, and the shape of the periodic table
2. Energy levels and sublevels
3. s, p, d, f and the periodic table
4. Orbitals, spin energy diagrams
5. e- config., valence e-, dot diagrams