This document provides an overview of spectroscopy and how it can be used to determine the composition of astronomical objects. It discusses how light interacts with matter on an atomic level, causing absorption and emission spectra that act as "elemental barcodes." The spectra are caused by electrons transitioning between quantized energy levels in atoms and emitting or absorbing photons of specific wavelengths. Measuring the absorption lines in a star's spectrum allows astronomers to identify the elements present in a star's atmosphere and determine its chemical composition, such as the fact that hydrogen and helium make up over 97% of the Sun's mass.

1. Astonishing Astronomy 101
With Doctor Bones (Don R. Mueller, Ph.D.)
Educator
Entertainer
J
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Scientist
Science
Explorer

2. Chapter 4 - Spectroscopy

3. The Nature of Light
 The electromagnetic (EM) wave
 A small disturbance in an electric
field creates a small magnetic field,
which in turn creates a small electric
field and so on…
 EM waves transport energy.
Particles of light (photons)
travel through space.
These photons have very
specific energies, that is,
light is quantized.
• Light is radiant energy.
• Travels very fast –
300,000 km/sec!
• Light can be described
as either a wave or a
particle.

4. The Effect of Distance on Light
• Light from distant objects
seems very dim
• Why? Because the photons are
losing energy?
• No – because the light is simply
spreading out as it travels from
its source to its destination.
• The farther from the source
you are, the dimmer the light
seems.
• We say that the object’s
brightness or amount of light
received from a source is
decreasing.
2
d4
OutputLightTotal
Brightness


This is an inverse-square law –
the brightness decreases as the
square of the distance (d) from
the source increases.

5. The Nature of Matter
• The atom has a nucleus at its
center containing protons
and neutrons.
• Outside of the nucleus,
electrons whiz around in
clouds called orbitals
– Electrons can also be
described using wave or
particle models
– Electron orbitals are
quantized – that is, they
exist only at particular
energies.
– The lowest energy orbital
is called the ground state,
one electron wave long.
• To move an electron from one orbital to the
next higher, a specific amount of energy
must be added. Likewise, a specific
amount of energy must be released for an
electron to move to a lower orbital.
• These are called electronic transitions.
Please insert figure 21.3

6. The Chemical Elements
• The number of protons (atomic number) in the
nucleus determines which element a substance is.
• Each element has a number of electrons equal to the
number of protons.
• The electron orbitals are different for each element,
and the energy differences between the orbitals are
unique as well.
• This means that if we can detect the energy emitted
or absorbed by an atom during an electronic
transition, we can tell which element the atom
belongs to, even from millions of light years away!

7. Absorption
• If a photon of exactly
the right energy
(corresponding to the
energy difference
between orbitals) strikes
an electron, that
electron will absorb the
photon and move into
the next higher orbital
– The atom is now in an
excited state
• If the photon is of
higher or lower
energies, it will not be
absorbed – it will pass
through as if the atom
were not there.
• This process is called absorption
• If the electron gains enough energy
to leave the atom entirely, we say the
atom is now ionized. It is an ion.
Please insert figure 21.6

8. Emission
• If an atom drops
from one orbital
to the next lower
one, it must first
emit a photon
with the same
amount of energy
as the orbital
energy difference.
• This is called
emission.
Please insert figure 21.4

9. Seeing Spectra
http://www.youtube.com/watch?v=QI50GBUJ48s
• Seeing the Sun’s
spectrum requires a
few special tools, but
it is not difficult.
• A narrow slit only lets a
little light into the
experiment.
• Either a diffraction
grating or a prism splits
the light into its
component colors.
• If we look closely at the
spectrum, we can see
lines, corresponding to
wavelengths of light that
were absorbed.

10. The Bohr Model:
http://www.youtube.com/watch?v=45KGS1Ro-sc
• The Danish physicist Niels Bohr developed a model of the hydrogen
in the early 1900s that describes its possible energy levels.
• The Bohr model of hydrogen describes the electron as orbiting at a set
of discrete radii.
• If we number the energy levels starting from the ground state as n = 1,
2, 3 and so on, then the electron’s orbital radius is r = 0.053 × n2
nanometers (nm). (1 nm = 10−9 m).
• The energy En of level n can be described by the formula:
E1 (Hydrogen) = - 2.18 x 10-18 J/n2 = -13.6 eV/n2
n = 1 is the ground state quantum number and E1 is the ground
state energy.
1 eV = 1.60 x 10-19 J

11. Hydrogen Energy Levels
E1 = -13.6 eV/12 = -13.6 eV
E2 = -13.6 eV/22 = -3.4 eV
E3 = -13.6 eV/32 = -1.5 eV
E4 = …
.
.
E∞ = 0

12. Emission Spectra
http://www.youtube.com/watch?v=8TJ2GlWSPxI&feature=related
• Imagine that we heat up a
gas of atoms.
• Collisions among the atoms
cause electrons to jump to
higher orbitals, hence higher
energy levels.
• Collisions can also cause the
electrons to fall to lower levels,
emitting photons of energy:
E = hυ = hc/.
• If the electron falls from
orbitals 3 to 2 in hydrogen, the
emitted photon will have a
wavelength of  = 656 nm. Red
• If the electron falls from orbital
4 to orbital 2, the emitted
photon will have a wavelength
of  = 486 nm. Blue
• We can monitor the gas, counting the number
of photons at each wavelength. If we graph
this data, we’ll see an emission spectrum.

13. Emission spectrum of hydrogen
http://csep10.phys.utk.edu/astr162/lect/light/absorption.html
http://www.youtube.com/watch?v=OJzW2RoZq1Y&feature=related

14. Atomic Spectra: Elemental Fingerprint
• Every element has its own
spectrum. A fingerprint so to
speak. Note the differences
between hydrogen and helium
spectra below.

15. Absorption
Spectra
• What if, instead of hot
hydrogen gas, we had a
cloud of cool gas between
us and a star?
• Photons of an energy that
corresponds to the
electronics transitions in
hydrogen will be
absorbed by electrons in
the gas.
• The light from these
photons is effectively
removed from the
spectrum.
• The spectrum will have
dark lines where the
missing light would be:
An Absorption Spectrum:
The elemental barcode!

16. Types of Spectra
• Kirchoff’s Laws:
– If the source emits light that is
continuous and all colors are
present, we say that this is a
continuous spectrum.
– If the gas molecules are moving
rapidly (i.e., high temperature),
the atoms will emit characteristic
frequencies of light. This is an
emission-line spectrum.
– If the molecules of gas are cool,
they will absorb light of a
characteristic frequency as it
passes through. This is an
absorption line spectrum.

17. Spectra of Astronomical Objects
Please insert figure 24.7

18. Composition of the Sun, from Spectroscopy
Element Atomic Number Relative
Number of
Atoms
Percent by Mass Energy to Remove
Outermost Electron
Hydrogen 1 1012 70.6% 2.179 × 10−18 J
Helium 2 9.77×1010 27.4% 3.940 × 10−18 J
Carbon 6 3.63×108 0.31% 1.804 × 10−18 J
Nitrogen 7 1.12×108 0.11% 2.328 × 10−18 J
Oxygen 8 8.51×108 0.96% 2.182 × 10−18 J
Neon 10 1.23×108 0.18% 3.454 × 10−18 J
Silicon 14 3.55×107 0.07% 1.306 × 10−18 J
Iron 26 4.68×107 0.18% 1.266 × 10−18 J
Gold 79 10 0.000000014 1.478 × 10−18 J
Uranium 92 <0.3 <0.0000000057 0.992 × 10−18 J