Lecture Outlines
Chapter 4
Astronomy Today,
6th edition
Chaisson
McMillan
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Chapter 4
Spectroscopy
Units of Chapter 4
4.1 Spectral Lines
4.2 Atoms and Radiation
The Hydrogen Atom
4.3 The Formation of Spectral Lines
The Photoelectric Effect
4.4 Molecules
4.5 Spectral-Line Analysis
Information from Spectral Lines
Spectrascope
Radiation can be analyzed with an instrument
known as a spectroscope.
Research instruments called spectrographs,
or spectrometers, used by professional
astronomers are similar but more complex
than a spectroscope.
4.1 Spectral Lines
Spectroscope: Splits light into component
colors
prism is used with a lens
Edwinn
Hubble
SPECTROSCOPY
•  = the study of the properties of light
that depend on Wavelength.
•  3 Types of Spectra:
– Continuous
– Emission
– Dark-Line (absorption)
Continuous Spectrum
= produced by a incandescent solid, liquid,
or gas under high pressure.
-Uninterrupted band of color
-Light from a “common lightbulb”
-Incandescent = to emit light when hot
If you put a certain gas in a glass jar and pass
an electrical discharge through it, the gas will
begin to heat up and glow – that is it will emit
radiation.
Instead of a continuous spectrum (all
wavelengths), the radiation will be only certain
wavelengths out of the spectrum (called
emission lines when viewed through a
spectroscope), which are unique to each
element.
Emission Spectrum
“Bright Line”
= produced by a hot (incandescent) gas under
low pressure.
4.1 Spectral Lines
Emission lines:
Single frequencies
emitted by
particular atoms
Scientists have accumulated extensive
catalogs of the specific wavelengths at which
many different hot gases emit radiation.
The particular pattern of light emitted by a gas
of a given chemical composition is known as
the emission spectrum of the gas.
4.1 Spectral Lines
Emission spectrum can be used to
identify elements
4.1 Spectral Lines
Absorption spectrum: If a continuous spectrum
passes through a cool gas, atoms of the gas will
absorb the same frequencies they emit
When sunlight is split by a prism, at first glance it
appears to produce a continuous spectrum.
However, closer scrutiny with a spectroscope shows
that the solar spectrum is interrupted vertically by a
large number of narrow dark lines.
We now know that many of these lines represent
wavelengths of light that have been removed
(absorbed) by gases present either in the outer
layers of the Sun or in Earth’s atmosphere.
These gaps in the spectrum are called absorption
lines.
Absorption Spectrum
“Dark Line”
= produced when white light is passed through
a cool gas under low pressure.
- A continuous spectrum with dark lines running
through it.
Solar absorption lines (from our Sun) were first
observed in 1802, then observed in greater detail
10 years later.
There are now over 600 absorption lines cataloged,
now referred to as Fraunhofer lines.
Similar lines are known to exist in the spectra of all
stars.
60+ elements identified
Scientists in the lab are able to produce these
same absorption lines by passing a beam of a
continuous spectrum of light through a cool gas.
They have discovered a connection between
emission and absorption lines:
the absorption lines associated with a given gas
occur at precisely the same wavelengths as the
emission lines produced when the gas is heated.
4.1 Spectral Lines
An absorption spectrum can also be used to
identify elements. These are the emission and
absorption spectra of sodium:
Sodium
Oxygen
Lines depend on the gas that produces them.
Oxygen
Absorption
Emission
Kirchoff’s Laws
The analysis of the ways in which in matter
emits and absorbs radiation is called
spectroscopy.
In 1859, Gustav Kirchoff formulated three
spectroscopic rules, now known as Kirchoff’s
laws:
Kirchoff’s Laws
1. A luminous solid or liquid, or a sufficiently
dense gas, emits light of all wavelengths and
so produces a continuous spectrum of
radiation.
Kirchoff’s Laws
2. A low-density, hot gas emits light whose
spectrum consists of series of bright emission
lines that are characteristic of the chemical
composition of the gas.
Kirchoff’s Laws
3. A cool, thin gas absorbs certain wavelengths
from a continuous spectrum, leaving dark
absorption lines in their place, superimposed on
the continuous spectrum.
Once again, these lines are characteristic of the
composition of the intervening gas – they occur at
precisely the same wavelengths as the emission
lines produced by that gas at higher temperatures.
4.1 Spectral Lines
Kirchhoff’s Laws:
•  Luminous solid, liquid, or dense gas
produces continuous spectrum
•  Low-density hot gas produces emission
spectrum
•  Continuous spectrum incident on cool, thin
gas produces absorption spectrum
4.1 Spectral Lines
Kirchhoff’s laws illustrated:
4.2 Atoms and Radiation
Existence of spectral lines required new
model of atom, so that only certain amounts
of energy could be emitted or absorbed
Bohr model had certain allowed orbits for
electron
1
2
3
This description of the atom contrasts sharply
with the predictions of Newtonian mechanics,
which would permit orbits with any energy, not
just at certain specific values.
In the atomic realm, such discontinuous
behavior is the norm.
In the jargon of the field, the orbital
energies are said to be quantized.
The rules of quantum mechanics,
the branch of physics governing the
behavior of atoms and subatomic
particles, are far removed from
everyday experience.
4.2 Atoms and Radiation
Existence of spectral lines required new
model of atom, so that only certain amounts
of energy could be emitted or absorbed
Bohr model had certain allowed orbits for
electron
In Bohr’s original model, each
electron orbital was pictured as
having a specific radius, much like a
planetary orbit in the solar system.
However, the modern view is not so
simple. Although each orbital does
have a precise energy, the orbits are
not sharply defined, as indicated in
the figure. Rather, the electron is
now envisioned as being smeared
out in an “electron cloud”
surrounding the nucleus.
4.2 Atoms and Radiation
Emission energies correspond to energy
differences between allowed levels
Modern model has electron “cloud” rather than
orbit
We do not currently have a physical
theory that can predict the actual
location of the electron. Instead, we
can only speak of the probability of
finding it in a certain location within
the cloud. The average distance
from the electron cloud to the
nucleus is commonly referred to as
the “radius” of the electron’s orbit.
When a hydrogen atom is in its
ground state, the radius of the orbit
is about .05 nm.
4.2 Atoms and Radiation
Emission energies correspond to energy
differences between allowed levels
Modern model has electron “cloud” rather than
orbit
Compare diagrams: Classical atom
vs. Modern Atom
Atoms do not always remain in their
ground state. An atom is said to be
in an excited state when an electron
occupies an orbital at a greater-thannormal distance from its parent
nucleus. An atom in such an excited
state has a greater-than-normal
amount of energy.
4.2 Atoms and Radiation
Energy levels of the hydrogen atom, showing
two series of emission lines:
The energies of
the electrons in
each orbit are
given by:
The emission
lines correspond
to the energy
differences
The excited state with the lowest
energy (closest to ground state) is
called the first excited state, that
with the second-lowest energy is the
second excited state, and so on.
4.2 Atoms and Radiation
Energy levels of the hydrogen atom, showing
two series of emission lines:
The energies of
the electrons in
each orbit are
given by:
The emission
lines correspond
to the energy
differences
An atom can become excited in one
of two ways: by absorbing some
energy from a source of
electromagnetic radiation or by
colliding with some other particle –
another atom, for example. However,
the electron cannot stay in a higher
orbital forever (about 10
nanoseconds only); the ground state
is the only level where it can remain
indefinitely.
Radiation as Particles
Because electrons can exist only in
orbitals having specific energies,
atoms can absorb only specific
amounts of energy as their electrons
are boosted into excited states.
Likewise, atoms can emit only
specific amounts of energy as their
electrons fall back to lower energy
states.
Thus, the amount of light energy
absorbed or emitted in these
processes must correspond
precisely to the energy difference
between two orbitals.
4.2 Atoms and Radiation
Energy levels of the hydrogen atom, showing
two series of emission lines:
The energies of
the electrons in
each orbit are
given by:
The emission
lines correspond
to the energy
differences
The atom’s quantized energy levels
require that light be absorbed and
emitted in the form of distinct
“packets” of electromagnetic
radiation, each carrying a specific
amount of energy. We call these
packets photons. A photon is, in
effect, a “particle” of
electromagnetic radiation.
4.2 Atoms and Radiation
Energy levels of the hydrogen atom, showing
two series of emission lines:
The energies of
the electrons in
each orbit are
given by:
The emission
lines correspond
to the energy
differences
The idea that light sometimes
behaves not as continuous wave,
but as a stream of particles, was
proposed by Albert Einstein in 1905
to explain a number of experimental
results (especially the photoelectric
effect) then puzzling physicists.
Furthermore, Einstein was able to
quantify the relationship between
the two aspects of light’s double
nature. He found that the energy
carried by a photon had to be
proportional to the frequency of the
radiation. (or inversely proportional
to the wavelength of radiation).
The constant of proportionality in
the preceding relation is now known
as Planck’s constant, in honor of the
German physicist Max Planck, who
determined its numerical value. It is
always denoted by the symbol h,
and the equation relating the photon
energy E to the radiation frequency f
is usually written
E = hf
Which can also be written
E = hc/λ
(from the wavespeed equation: c =
Like the universal gravitational
constant G and the speed of light in
a vacuum, c, Planck’s constant is
one of the fundamental physical
constants of the universe.
In SI units, the value of Planck’s
constant is a very small number:
h = 6.63 x 10-34 Joule seconds
Consequently, the energy of a single
photon is tiny. Even a very highfrequency gamma ray (the most
energetic type of radiation) has an
energy of just 7 x 10-12 J – about the
energy carried by a flying gnat.
Nevertheless, this energy is more
than enough to damage a living cell.
The basic reason that gamma rays
are so much more dangerous to life
than visible light is that each
gamma-ray photon typically carries
millions, if not billions, of times
more energy than a photon of visible
radiation.
The Photoelectric Effect
The equivalence between the energy
and frequency (or inverse
wavelength) of a photon completes
the connection between atomic
structure and atomic spectra. Atoms
absorb and emit radiation at
characteristic wavelengths
determined by their own particular
internal structure.
Because this structure is unique to
each element, the colors of the
absorbed and emitted photons – that
is, the spectral lines we observe –
are characteristic of that element
and only that element. The spectrum
we see is thus a unique identifier of
the atom involved.
4.2 Atoms and Radiation
The photoelectric effect:
•  When light shines on metal, electrons can be
emitted
•  Frequency must be higher than minimum,
characteristic of material
•  Increased frequency—more energetic
electrons
•  Increased intensity—more electrons, same
energy
4.2 Atoms and Radiation
Photoelectric effect can only be understood if
light behaves like particles
4.2 Atoms and Radiation
Light particles each have energy E:
Here, h is Planck’s constant:
Many people find it confusing that
light can behave in two such
different ways. Modern physicists
don’t yet fully understand why
nature displays this wave-particle
duality. As a general rule of thumb,
in the macroscopic realm of
everyday experience, radiation is
more usefully described as a wave,
whereas in the microscopic domain
of atoms, it is best characterized as
a series of particles.
4.3 The Formation of Spectral Lines
Absorption can boost an electron to the
second (or higher) excited state
Two ways to decay:
1.  To ground state
2.  Cascade one orbital at a time
4.3 The Formation of Spectral Lines
(a) Direct decay
(b) Cascade
4.3 The Formation of Spectral Lines
Absorption spectrum: Created when atoms
absorb photons of right energy for excitation
Multielectron atoms: Much more complicated
spectra, many more possible states
Ionization changes energy levels
4.3 The Formation of Spectral Lines
Emission lines can be used to identify atoms
4.4 Molecules
Molecules can vibrate and
rotate, besides having energy
levels
•  Electron transitions produce
visible and ultraviolet lines
•  Vibrational transitions
produce infrared lines
•  Rotational transitions
produce radio-wave lines
4.4 Molecules
Molecular spectra are much more complex
than atomic spectra, even for hydrogen:
(a) Molecular hydrogen
(b) Atomic hydrogen
4.5 Spectral-Line Analysis
Information that can be gleaned from
spectral lines:
•  Chemical composition
•  Temperature
•  Radial velocity
4.5 Spectral-Line Analysis
Line broadening can
be due to a variety of
causes
4.5 Spectral-Line Analysis
4.5 Spectral-Line Analysis
The Doppler shift may
cause thermal
broadening of spectral
lines
4.5 Spectral-Line Analysis
Rotation will also cause broadening of
spectral lines through the Doppler effect
Summary of Chapter 4
•  Spectroscope splits light beam into
component frequencies
•  Continuous spectrum is emitted by solid,
liquid, and dense gas
•  Hot gas has characteristic emission spectrum
•  Continuous spectrum incident on cool, thin
gas gives characteristic absorption spectrum
Summary of Chapter 4 (cont.)
•  Spectra can be explained using atomic
models, with electrons occupying specific
orbitals
•  Emission and absorption lines result from
transitions between orbitals
•  Molecules can also emit and absorb
radiation when making transitions between
vibrational or rotational states