Unlike most other scientists, astronomers cannot bring their test material into the laboratory for controlled experiments. To learn about our universe, we rely upon our knowledge of the nature of light: how it is produced, how it travels through a vacuum and through media, how it behaves when we "capture" it with a telescope. This lesson gives a survey of what we know about radiation, spectra, and telescopes.
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After completing this lesson, you should be able to
All of the information we get from the Universe, with the exception of a few meteorites and the rocks brought back from the Moon, comes from light. We must understand light, electromagnetic radiation—oscillating electric and magnetic fields—to understand the information it brings with it. We must understand how atoms work, how the electrons absorb and emit radiation. We must have confidence that we understand the periodic chart of elements. If we see characteristics of elements billions of light years from us that are identical to the characteristics of elements here on Earth, then those elements must be be present billions of light years away. The chemistry and physics that operates here on Earth operates the same everywhere.
Boynton's Program 8 contains a compact description of what we need to know about light. These notes will supplement that information and give you a better idea of what concepts of light are important to us. The following 2 links give a history about our knowledge of light, from the notes of Plato and Aristotle to Newton. Certainly the multitude of colors we detect and the interest in how the eye was structured and how it worked was of great interest for centuries. Quantitatively, Isaac Newton and James Maxwell gave us the fundamentals upon which we currently base our knowledge.
http://www.colorsystem.com/projekte/engl/01pyte.htm
http://www-scf.usc.edu/~kallos/light1.htm
To fully understand where light comes from, how it is produced, we must start with a brief summary of what, exactly, atoms are:
This view of the atom is quite different from that given in usual astronomy texts. The classical view of the atom, where electrons
are viewed as "mini-planets" orbiting a nucleus "sun" is now over 70 years
out of date. Even though we need to stick to a superficial view of the quantum
atom, it is not that much harder to understand. Plus, quantum physics is extremely
fascinating and tests all that we know about our world.
When describing the orbitals of an electron in an atom, we can state only the probability of where we might "find" it in the atom; we will pick the hydrogen atom in picturing what we mean. The figure depicts the hydrogen atom, nucleus (single proton) at the center, and the electron cloud surrounding the nucleus. As the electron gains energy, either through a collision with another electron or by absorbing a photon, the charged cloud occupies a larger volume. The higher density shows where there is a greater probability of finding the electron. Note that in the more excited state, there is still some probability that the electron will be found closer to the nucleus. (There is even a state where there is a small probability that the electron will be found in the nucleus.) If the electron gains enough energy, it will leave the proton and we say that the atom is ionized. (Atoms with multiple electrons can have multiple ionized states.) Note that the figure is not to scale: the proton would be much too small to see.
We can describe the properties of light:
Wavelength, frequency, and the speed of light are related by:
c = wavelength * frequency
The whole electromagnetic spectrum consists of (from highest energy, highest frequency, shortest wavelength . . . to . . . lowest energy, lowest frequency, longest wavelength): gamma rays, x-rays, ultraviolet, visible, infrared, microwave, radio.
The wavelengths range from a picometer (0.000000000001 m) for gamma rays, to 10,000 m for radio. (Note, however, that both ends of the scale are theoretically unbounded.)
| EM Radiation | Frequency (Hz) |
Wavelength (meters) |
Typical Astronomical Sources |
|---|---|---|---|
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Lowest frequency Lowest energy |
Longest Wavelengths Greater than 1 cm (0.01 meter) |
Active galaxies Quasars Alien civilizations |
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1 cm-1 mm (approximately) | Cosmic Microwave Background Interstellar molecules | |
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1 mm-700 nm | Interstellar dust | |
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700 nm-400 nm | Stars Interstellar gas | |
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400 nm-10 nm | Hot stars | |
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10 nm-0.1 nm | Black hole accretion disks Pulsars | |
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Highest frequency Highest energy |
Shorter than 0.1 nm (0.0000000001 m) Shortest wavelength |
Supernova explosion Mergers of neutron stars or black holes? |
For a wonderful and informative presentation of light, check out Imagine the Universe: Electromagnetic Radiation. The images in the above table were obtained from that site.
It's important to remember the kinds of electromagnetic radiation and whether they represent high energy (high frequency, short wavelength) or low energy (low frequency, long wavelength), or somewhere in between. Let's take a look at another representation of the full spectrum:
Almost everything we know about our cosmos comes from light: either given off by the object or reflected by the object. We talk about light coming to us from across the Universe. Light is the propagation of electric and magnetic fields. In a vacuum, it travels at an astounding 300,000 kilometers per second (or, more precisely: 299,792.458 km/sec). The changing electric field generates a magnetic field, and the changing magnetic field generates an electrical field—nature's most perfect dynamo-electromagnet.
Here's the strangest thing about light: it can behave either as a wave or as a particle. How it behaves depends on what it is traveling through, or how you are trying to measure it, among other things. For example, light enters the lens of an eye as a wave, and is refracted and focused on the retina. There, the light acts as particles, transmitting their energy to the rods and cones, stimulating them to send the nerve impulses to our brain. Light travels through a vacuum and through our atmosphere as a wave. When we hold our solar-powered calculator to the light, it then behaves as a particle, generating an electrical current when it strikes the solar energy cell. Light behaves as a wave when refracted through raindrops forming a rainbow.
As a wave, the characteristics of light are described by its wavelength, velocity (speed), and frequency. These three properties are related by:
(from Imagine
the Universe)
As a particle, light behaves as tiny packets of energy called "photons" (do not confuse with "protons"). Photons at radio frequencies have the lowest energy—that's why we can live with radio waves passing by us and through us continuously. Photons at gamma ray frequencies have the highest energy—that's why you do not want to meet up with an atomic bomb or a nuclear meltdown. Gamma ray photons destroy life. The energy of a photon is related to its frequency by:
E = hv
E is the energy of the photon; h is another fundamental constant of the Universe called Planck's constant (of order 7 × 10-34 in units of energy * time!); and v is the Greek letter nu representing frequency. One can see by this relationship that the higher the frequency of the light, the higher the energy contained in its photons. Light as a wave and light as a particle are intimately related, and one cannot think of light without considering both of these characteristics.
(Thought questions: if infrared, microwave, and radio waves are all forms of electromagnetic radiation, why can't we operate our calculators by holding them up to a stove or our cell phones or towards some nearby radio towers? In a pinch during a nuclear war, would gamma rays work to run our calculators?)
A continuous spectrum is what we normally think of as all of the colors or the rainbow. Any object having a temperature above absolute zero, emits thermal radiation. Humans, with a temperature of about 310 Kelvin, radiate at infrared wavelengths. If the thermal radiator is a perfect radiator, we call it a blackbody, and it will emit radiation over a range of wavelengths and form a continuous spectrum.
See http://www-astro.phast.umass.edu/courseware/vrml/bb/ for an interactive activity on blackbody radiation.
Light is emitted or absorbed when electrons of an atom change their energy states. We think of electrons as jumping from one energy state to another. A photon is emitted when an electron jumps from a high energy state to a lower energy state. The photon must somehow get rid of that energy and it does so by emitting energy in the form of electromagnetic radiation. (Think of gravitational potential energy: if you are 3 meters up on a ladder, you have more potential energy than when you are standing on the ground. If you were to jump down to the floor, you would lose that potential energy.) In order for an electron to jump to a higher energy state, it needs to somehow acquire enough energy to do so. It can get that energy at the expense of another electron by bumping into it (not on purpose, of course), or it can absorb the energy contained in a photon. (In our ladder analogy, you had to take in energy in order to be able to climb the ladder.) The more energy an electron absorbs, the higher the energy state.
A curious thing about atoms and their electrons: the electrons exist only in discrete energy states, and these energy states are different for the atoms of each element in the periodic chart. Each element of the periodic chart has a unique spectrum. Study and compare the emission lines of the various elements shown below: hydrogen, helium, neon, mercury. How do the patterns of the emission lines differ? How do you think astronomers use this knowledge of the patterns and colors to determine if certain elements are present in stars? The different spectra of stars are often compared to individuals having different fingerprints. Why is this a good analogy? Because every element's signature spectrum is unique. Not only that, but the spectra of that same element ionized one or more times are different. Isotopes of an element have different spectra. As part of the lab for this lesson, you will be closely examining the spectra of a few different elements.
| Element | Spectrum |
|---|---|
| Argon |  |
| Helium |  |
| Mercury |  |
| Neon |  |
| Sodium |  |
Let's take some of the concepts introduced above and work with them so we get a better understanding. We will just be dealing with the hydrogen atom. It's complicated enough, and the helium atom—with just two electrons—is so much more complicated that it would be impossible to describe here. We can use the Bohr model of the atom only for hydrogen. The above link will take you to a more detailed mathematical explanation. It is important that you try to follow the logic.
If you take a look at deep-sky images online, in particular the eta Carinae, horsehead, and Orion nebulae, you will see red, glowing gas. That gas is hydrogen, and it looks read because there are electrons jumping from energy level 3 to energy level 2 and emitting a photon having a wavelength of 656.3 nm. Now, the electrons are also jumping from higher energy levels to energy level 2, they all have multiple energy levels available to them, remember, but the red wavelengths dominate. Take a look at the image shown below of the Eagle Nebula. Hydrogen gas here is getting energy from the nearby, extremely hot stars that is precisely the right amount to excite electrons from energy level 2 to higher energies. Then, the electrons jump back to energy level 2, and the energy is emitted as photons. Where we are observing the thin cloud of hydrogen gas without a continuous source behind it (that is, a star) we will see an emission spectrum containing all of the Balmer lines. If there are places where we must look through the thin cloud of gas in order to observe a star, we will see an absorption spectrum since the cloud will be cooler than the blackbody radiator behind it.
A word of caution must be introduced here when it comes to interpreting the colors seen in the images in textbooks, online, and in these notes: colors can be enhanced and misrepresented for clarity (and confusion, unfortunately). For example, take a look at the Eagle Nebula and the region around the constellation of Orion in the following table. Since we have no "color" for infrared, scientists use the nearest thing: red (usually). The red in the Eagle Nebula represents hydrogen atoms at a temperature of approximately 10,000 K, while the red in the Orion image represents dust at temperatures only a fraction of that.
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M16, The Eagle Nebula, taken through filters isolating emission of http://physics.uwyo.edu/misc/GradsByGrads/wiro_pics.html [accessed 27 December 2007] |
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The constellation Orion in visible light (left) and infrared (right, as seen by IRAS). One team is studying IC 2118, or the "Witch Head Nebula," highlighted in the lower right. http://www.spitzer.caltech.edu/Media/happenings/20051208/ [accessed 27 December 2007] |
One more area of confusion is related to the energies of the electrons in atoms and the emission or absorption lines the energy jumps produce, and the amount of energy and maximum wavelengths we measure as a result of the blackbody temperature of the star. A physical process can produce lines in the high-energy part of the spectrum (ultraviolet, x-ray, and gamma ray). The corona of the Sun has a temperature of over 1 million degrees and so we will see emission lines of hydrogen that are from the Lyman series mentioned above because at that temperature there is a lot of high-energy photons that can make an electron jump up from energy level 1. But, the corona of the Sun is definitely not a blackbody. We would have to do spectroscopy with a telescope, though, that can detect ultraviolet radiation, and also one that needs to orbit above the Earth's atmosphere. On the other hand, a very hot star—one with a blackbody temperature of 30,000 Kelvin—radiates most of its energy at ultraviolet wavelengths, around 100 nm. These stars also need a telescope above the atmosphere to detect most of their light, but the detectors or the purpose of the research may be different.
See Boynton, Cosmic Perspective, Program 8.
Q: Where in an atom would you find electrons? Protons? Neutrons? Summarize the modern view of an atom, using hydrogen as an example.
We will be talking about different wavelengths and different regions of the electromagnetic spectrum throughout the remainder of this course. It is advantageous for you to be able to recognize what energies we are considering.
Q: Which of the following options orders light from the longest wavelengths to the shortest correctly?
Q: After you have checked the answer, have we ordered the light from the lowest energies to the highest energies, or from highest to lowest?
Q: What do we mean when we talk about light being a particle and light being a wave?
Q: Given these examples, is light going to behave as a particle or as a wave?
Q: The spectrum
for the hydrogen atom is shown here in emission and in absorption. There are
no differences in the wavelengths of the lines between the emission spectrum
and the absorption spectrum. Why not? Should there be?
Q: What is a blackbody? How does the energy emitted by a blackbody depend on its temperature? How does the "peak" wavelength—where the maximum of the blackbody curve occurs—depend on the temperature? (Look up Wien's law, expressed in nanometers.)
Q: Assume the Sun radiates as a blackbody with a temperature of 5800 degrees Kelvin. Using Wien's law: peak wavelength (nm) = 2.9 x 106 / temperature, calculate the wavelength at which it radiates most of its energy. What part of the spectrum does this wavelength lie? Does radiation at these wavelengths reach the surface of the Earth? How does this wavelength correlate to the range of wavelengths that our eyes are capable of seeing?
Q: Assume you are observing at microwave wavelengths, 1.9 mm to be exact. Convert this to nanometers and calculate the corresponding temperature.
