Light, Spectra, and Matter
Why will we spend so much time discussing the
electromagnetic spectrum?
Not easy to visit astrophysical objects
(the Sun, planets, other stars) and make
direct in situ measurements
We rely on remote sensing of EM
radiation.
Tells us the temperature and
composition
This gives us important clues to
the origins of these objects.
Light properties
There are many properties of light that can be measured and
quantified. Each of these properties will have its own units.
•
•
•
•
•
•
Energy
Power
Intensity
Wavelength
Frequency
Speed
Which of the following is
NOT a unit of energy?
1.
2.
3.
4.
Joule
Kilowatt
Kilowatt-hour
Electron volt
Which of the following is
NOT a unit of energy?
1.
2.
3.
4.
Joule
Kilowatt
Kilowatt-hour
Electron volt (eV)
BTU = British thermal unit
1 eV = 1.6 × 10–19 J
Intensity
Intensity is a measure of
how much power from a
light source is distributed
over an area.
Its units are Watts per
square meter (W/m2),
which is also written W m-2
Power
Intensity =
Area
Solar Intensity
The Solar Energy Output is 4 x 1026 W.
How much of that hits us?
When the Sun is directly over head, it delivers the equivalent of
22 × 60 watt light bulbs over each square meter (m2) of
ground!!!
This amount, 1368 W m-2, is known as the solar constant
How is solar energy delivered from the Sun to
the Earth?
As light!!!!
Light travels as an electromagnetic wave
Electromagnetic Wave: propagating wave of electric and
magnetic fields that oscillate perpendicular to
each other and the direction of propagation
In a vacuum, wave propagates with speed = 3.00 x 108 m/s
(cosmic speed limit)
Electric
field
Magnetic field
Wave Properties
speed (v): how much distance the wave moves per unit
time for an EM wave the speed is always the speed of light.
In other words, v = c = 3 x 108 m/s.
frequency (f or ν (nu)): number of peaks that pass a
location in a given time (units: Hertz (Hz) = 1/s = s-1)
Wave Properties
wavelength (): distance between two consecutive peaks
(units: km, m, cm, mm, m, nm…)
amplitude: height of the wave (or depth of the trough);
related to intensity but we won’t use it
Wave Properties
speed (v): how much distance the wave moves per unit
time
(for an EM wave v = c = 3 x 108 m/s)
frequency (f): number of peaks that pass a location
is a given time (units: Hertz (Hz) = 1/s = s-1)
wavelength (): distance between two consecutive peaks
(units: km, m, cm, mm, m, nm…)
These three properties are related:
f 
c

If wavelength is 10 m and frequency is 100 Hz (oscillations /
seconds), what would be the speed of the wave?
1.
2.
3.
4.
10 m/s
1 m/s
1000 m/s
100 m/s
If wavelength is 10 m and frequency is 100 Hz (oscillations /
seconds), what would be the speed of the wave?
1.
2.
3.
4.
10 m/s
1 m/s
1000 m/s
100 m/s
If wavelength is 10 m and frequency is 100 Hz
(oscillations / second), what would be the speed of
the wave?
f =
v
l
v= fl
= (100 s-1 ) ´ (10 m )
= 1000 m/s
The Photon
Light behaves like both a particle and a wave!
Photon: smallest bundle of light energy
(i.e., a particle of light)
Photons carry light energy:
1. A photon’s energy is proportional to frequency
(Eph  f).
2. A photon’s energy is inversely proportional to
wavelength (Eph  -1).
E ph = hf =
hc
l
Plank’s constant (h) = 6.602 x 10-34 Js
Matter actually a wave too!
• All matter exhibits particle and
wave properties (DeBroglie, 1921)
• For ordinary objects, the wave
nature of matter is much too small
to measure
– The wavelength of a baseball
moving at 80 mph would be
about 10-34 meters
• But for small particles, this is wave
nature of matter is measurable
– The wavelength of an electron is Electron diffraction pattern
showing its wave nature
about 10-10 meters
The Visible Spectrum:
How is a difference in the frequency or wavelength of
light observed?
For visible wavelengths  COLOR
How does a prism work?
• Dispersion: Speed of light in
the prism (glass or plastic)
depends on the frequency
(color)
• Refraction: Change in speed
of light causes a change in its
direction
• Result: Blue changes
direction most since its
speed is the lowest inside the
prism. And red changes
direction least since its speed
is highest inside the prism.
Red
Orange
Yellow
Green
Blue
Indigo
Violet
R
O
Y
G
B
I
V
Herschel Thinks Outside the Box:
In 1800 William Herschel made a discovery when he tried to
determine the temperature of light.
• He noticed that a thermometer
recorded energy from the Sun`s
spectrum even when placed beyond
the red end of the visible rainbow.
•He called this emission Calorific
Rays and it was the first discovery
that light had colors invisible to the
human eye.
•These rays are known today as
Infrared light.
Herschel’s work  color is associated with a temperature
Visible light is just a small part of the electromagnetic (EM)
spectrum
Fraunhofer’s Surprise
In 1813, Joseph von Fraunhofer, the owner of a glass manufacturing
firm in Munich, made an even more interesting discovery.
Using a precision dispersing prism, he discovered that the
`solar blackbody` was cut by thousands of dark bands.
Fraunhofer’s Surprise
Fraunhofer tried to test whether this effect was real.
1) He tested with different optics.
2) He tested by looking at different objects (moon and planets).
Bunsen and Kirchhoff`s solution:
Robert Bunsen (Univ. of Heidelberg)
turned pyromania into one of the great
discoveries of modern physics. Bunsen
set fire to things in order to figure out
their elemental composition
A colleague there, Gustav
Kirchhoff, suggested using a prism to
break the light apart. They quickly
discovered (1860) that burning
substances produced light in narrow
bands with unique patterns.
Iron
Blueprint to Composition:
Bunsen and Kirchhoff`s trick was the key to finding out the
composition of anything from the light it produced.
Many of the lines they found had the same wavelength as those
of Fraunhofer`s dark bands. They were seeing the composition of
the Sun!
Kinds of Spectra:
Bunsen found that he could identify the signature of different
elements in the Fraunhofer spectrum of the Sun.
Why were Bunsen`s heated gas spectra composed of bright lines
while Fraunhofer`s exhibited a continuous spectrum with dark
bands?
Bunsen`s fires were stimulating light emissions in the hot gas.
So what are Fraunhofer`s bands?
Absorption by (and re-emission from) a cooler gas!
Types of Spectra
Continuous: black body radiation
continuous
Absorption: requires a cool object
in front of a hot background
(ex: Fraunhofer) discrete
Emission: requires a hot object
with a cool background
(ex: Bunsen) discrete
Spectroscopy is the use of light’s interaction with matter to
identify or characterize properties of matter.
Basic definitions:
• Element: a substance that cannot be broken
down by chemical means (defined by number
of protons)
• Atom: the smallest piece of matter that is still
an element
• Molecule: two or more atoms that are bound
together by chemical bonds
• Nucleus: the protons and neutrons bound
together at the center of an atom
Atoms and light
Why do elements have the `discrete` interactions that Bunsen
saw?
Why do different elements (and molecules) have different
interactions?
This has to do with the nature
of atoms and how they are put
together.
A brief history of the atom
First discussion of the nature of matter (~400 BCE):
• Leucippus, a Greek philosopher, all matter consisted “tiny and
indivisible bodies called atoms”.
• The word atom comes from the Greek word `atomos` (not
divisible).
• Democritus, another Greek philosopher, these atoms were not
all alike, but had different shapes and sizes to make different
matter.
• Opposed (Aristotle): The prevailing view that everything was
made up of four basic elements: earth, fire, air, and water, not
atoms.
• Views such as Aristotle`s dominated science for many
centuries, until the Renaissance.
Dividing the ‘indivisible’:
The plum pudding model
• J. J. Thomson discovers cathode rays
are made of electrons (he called them
‘corpuscles’ – 1897). Electrons are
shown to have a negative charge
• Thomson proposes model of the atom
(1904):
– Atom has smaller components
– Negatively charged corpuscles/electrons
(plums)
– Positive ‘soup’ to balance negative
charge (pudding)
Discovering “nothing”
Meanwhile, Ernest Rutherford (Cambridge) discovers two new
types of radiation emitted by uranium (1899):
1. Alpha particles (): later found to be the helium nucleus
2. Beta particles (): later called the “electron” by Thomson
In 1909, Rutherford fires alpha particles at gold foil.
Expected only small angle scattering
due to gold atoms’ “plum pudding”.
Positive Nucleus
Saw mostly no scattering with
occasional back scattering
Matter is mostly empty space!!!!
Negative
electron
Rutherford’s atom
1. Mass is highly concentrated in the positivelycharged nucleus at the center of the atom.
2. Electrons (negatively charged) “orbit” the
nucleus.
3. Lots of empty space in-between
Positive
4. Similar to today’s atom
–
–
Number of protons determine
the element identity
Number of electrons determine
the chemical properties of the atom
Nucleus
Negative
electron
The nucleus is not uniform
Rutherford (1918): Discovers the proton.
The proton is about 2000 times as massive
as the electron and has a positive charge,
exactly the same magnitude as the
electronic charge.
James Chadwick (1932): Discovers the
neutron. Neither positive nor negative, it
has about the same mass as a proton.
Nuclei are made up of protons and neutrons.
Atoms, elements, and isotopes
Atoms (below - periodic table (Mendeleev, Meyer, 1867))
– Nucleus
• Protons – number determines the element (atomic #)
• Neutrons – number determines the isotope (mass #)
– Electrons – number determines the chemical properties
Atoms, elements, and isotopes
Isotopes:
Atoms with the same number of protons but different numbers
of neutrons are called isotopes.
Isotopes have the same chemical properties, but different
masses, different emission spectra, and participate in different
nuclear reactions.
hydrogen (1H) deuterium (2H) tritium (3H)
p+
p+
eA stable isotope
of hydrogen – 99.98%
natural abundance
n p+
p+ n
p+
n
n
e-
A stable isotope
of hydrogen – 0.02%
natural abundance
helium (4He)
e-
n
e-
Radioactive isotope
of hydrogen
eNew element; not an
isotope of H
Bohr atomic model
Energy States:
Niels Bohr (Copenhagen Univ.),
based on Rutherford’s work,
suggested a quantized
structure of electronic orbits
in an atom (1913)
Bohr and Werner Heisenberg
later (1926) modify structure
to account for the wave
properties of electrons.
p+
E1
E2
E3
e-
Electron distances and
energies are discrete
values
Atoms and light
Electrons exist in `orbits`
(much like planets in the
solar system) that are stable
at specific separations from
the nucleus.
Energy States:
The distance from the
nucleus determines the
energy of the electron
(lower E is closer).
p+
E1
E2
E3
e-
The spacing of these
energy levels is not even.
E1E2 > E2E3 > E3E4
etc…
Atoms and light
So what does all of this have to do with
Bunsen and Franhofer lines?
Atoms and light
Energy States:
If you heat the atom up to
high enough temperatures,
the electron will jump to
higher orbits (higher
energy state).
p+
E1
E2
E3
e-
How does `heating` do
this? Collisions
Atoms and light
Energy States:
If you heat the atom up to
high enough temperatures,
the electron will jump to
higher orbits (higher
energy state).
p+
How does `heating` do
this? Collisions
E1
eE2
E3
Atoms and light
After a time, the electron falls
back to the lowest energy state.
Energy States:
A photon is given off.
p+
E1
E2
E3
e-
The energy of the
photon is exactly equal
to the energy difference
between the two energy
states.
Atoms and light: absorption
Process of emission is fully reversible.
Energy States:
Electron can absorb a
photon and jump to a
higher energy level.
p+
E1
E2
E3
e-
The energy of the
photon must be exactly
equal to the energy
difference between the
two energy states.
Atoms and light: absorption
Process of emission is fully reversible.
Energy States:
Electron can absorb a
photon and jump to a
higher energy level.
p+
E1
eE2
E3
The energy of the
photon must be exactly
equal to the energy
difference between the
two energy states.
Conservation of energy
The energy difference between electron orbital states is
exactly equal to the energy of the photon emitted or
absorbed.
E2 – E1 = h f
Where E1 and E2 are the energies associated with the
electronic orbital states, f is the frequency of light, and h
is Planck’s constant = 6.62 × 10–34 J•s
hydrogen energy level diagram
Quantized energy
• Different frequencies are perceived as different
colors
• Atoms of different elements have different
allowable energy level transitions and thus emit
and absorb different discrete colors.
• Example: Each line in the spectrum of iron is
different energy level transition
Iron
Types of spectra
Continuous: black body radiation
continuous
Absorption: requires a cool object
in front of a hot background
(ex: Fraunhofer) discrete
Emission: requires a hot object
with a cool background
(ex: Bunsen) discrete
Ionization:
e-
Energy States:
What happens if a very energetic
photon interacts with an atom?
p+
e-
Such a photon can give enough
energy to the electron that it can
escape the atom.
The amount of energy necessary to
do this is called the binding energy
of the atom.
Ionization:
e-
Energy States:
p+
e-
e-
When an atom absorbs light (or
thermal) energy greater than the
binding energy, the electron
escapes.
The atom is left with a
positive charge and is called
an ion.
Together, ions and free electrons
are called plasma.
Plasma is found in stars, space, and
parts of our atmosphere.
What about molecules?
Electronic
Energy States:
e-
Molecules are atoms that are
connected by bonds (electrons). At
a basic level a molecule will
behave similarly to an atom.
ep+
p+
p+
p+
Molecules also have
discrete electron energy
levels.
-
Like atoms, electrons in
molecules can absorb a
photon and move to a
higher energy level
What about molecules?
e-
Electronic
Energy States:
A photon with enough energy
can free an electron by
overcoming the binding energy.
p+
p+
p+
p+
This produces a
molecular ion.
ePlasmas can also contain
molecular ions.
Photo-dissociation
Electronic
Energy States:
Or overcome the molecular binding energy and break the
molecule up (photo-dissociation).
ep+
e-
p+
p+
e-
e-
p+
Spectrum of HCl, a diatomic molecule
Frequency increases right along the x-axis; intensity is the y-axis
We can use light spectra to identify molecules!