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2 Matter, Energy, and Light

All matter, life included, is made up of atoms. Here we discuss the inner workings of atoms.  This will be relevant as we work to understand the chemistry and biology of life. Life also needs a source of energy to thrive. This energy can come in different forms, from sunlight to energy released when chemicals react. Light is a type of pure energy. The light and other kinds of radiation that reach us from objects in the universe encode a wide range of information about what those objects are like and how they work. If we can decipher this code and read the messages it contains, we can learn an enormous amount about the cosmos without ever having to leave Earth or its immediate environment.

Learning Objectives

By the end of this chapter, you will be able to:

  • Describe atoms and how they relate to chemical elements
  • Understand the properties of the subatomic particles electrons, protons, and neutrons
  • Explain what ions and isotopes of an atom are
  • Describe the cosmic abundance of different elements
  • Understand that there are different types of energy and that energy is conserved
  • Explain the wavelength, frequency and energy of a wave or particle of light
  • Discuss the different parts of the electromagnetic spectrum
  • Explain how and why the light emitted by an object depends on its temperature
  • Explain the difference between the luminosity and brightness of an object
  • Explain how the total energy emitted differs for hot and cold objects
  • Describe how spectroscopy can provide information about an object, such as a star.
  • Discuss the difference between continuous, absorption and emission spectra.

Atomic Theory

Atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you can directly test yourself in a chemistry laboratory. Every piece of matter we interact with, whether living or non-living, is made out of atoms. Pure elements consist of only one type of atom, which has certain properties characteristic of that element.  As an example, Figure 1 shows a pure copper penny, which means the penny is entirely made up of copper atoms that are linked or bonded together. Because an element consists of only one kind of atom, an element cannot be broken down into a different substance.  All atoms of a given element have identical chemical properties.  Inversely, atoms of a given element differ in properties from atoms of all other elements.

The left image shows a photograph of a stack of pennies. The right image calls out an area of one of the pennies, which is made up of many sphere-shaped copper atoms. The atoms are densely organized.
Figure 1 – Copper atoms in a penny. A pre-1982 copper penny (left) contains approximately 3×1022 copper atoms (several dozen are represented as brown spheres at the right), each of which has the same chemical properties.

An atom is the smallest unit of an element that can participate in a chemical reaction. Atoms are neither created nor destroyed during a chemical reaction; they are instead rearranged to yield substances that are different from the substances present before the reaction (Figure 2). This property is captured in the law of conversation of matter.  Because the number and nature of atoms remains constant before and after a chemical reaction (it is only the configuration of atoms that changes), the total mass of matter before and after a chemical reaction also remains constant.

The left stoppered bottle contains copper and oxygen. There is a callout which shows that copper is made up of many sphere-shaped atoms. The atoms are densely organized. The open space of the bottle contains oxygen gas, which is made up of bonded pairs of oxygen atoms that are evenly spaced. The right stoppered bottle shows the compound copper two oxide, which is a black, powdery substance. A callout from the powder shows a molecule of copper two oxide, which contains copper atoms that are clustered together with an equal number of oxygen atoms.
Figure 2 – Atoms are rearranged in a chemical reaction. In this example, we start on the left with copper atoms (shown here as brown spheres) and oxygen (shown here as red spheres). These atoms react and rearrange (right) to form a compound containing copper and oxygen (a powdery, black solid).

Atomic Structure

We now understand that each atom is composed of a very small nucleus, made up of positively charged protons and uncharged neutrons, and that this nucleus is surrounded by a much larger volume of space in which negatively charged electrons exist.

The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10−10 m, whereas the diameter of the nucleus is roughly 10−15 m—about 100,000 times smaller. To contextualize this, if an atom’s nucleus was the size of a blueberry, the diameter of the atom would be roughly the size of a football stadium (Figure 3).

The diagram on the left shows a picture of an atom that is 10 to the negative tenth power meters in diameter. The nucleus is labeled at the center of the atom and is 10 to the negative fifteenth power meters. The central figure shows a photograph of an American football stadium. The figure on the right shows a photograph of a person with a handful of blueberries.
Figure 3 – A scaling of the size of an atomic nucleus. If the size of an atom corresponded to the size of a football stadium, then the nucleus would be the size of a single blueberry.

Properties of Electrons, Protons, and Neutrons

Atoms—and the protons, neutrons, and electrons that compose them—are extremely small. For example, mass may typically be measured in terms of grams (g); a classic M&M is about 1 g.  The mass of a single carbon atom is less than 2 × 10−23 g.  An electron is over a thousand times less massive, with a mass on the order of 1 × 10−26 g.  Protons and electrons also carry electric charge in very small amounts that are unwieldy to express using more typical units.

When describing the properties of atoms and their subatomic particles, we therefore make use of appropriately small units of measure.  Rather than grams, we use the atomic mass unit (amu), which is approximately the mass of one proton.  For electric charge, we use the fundamental unit of charge (e), which is equivalent to the electric charge on a single proton.  In other words, a proton has a mass of ~1 amu and a charge of +1 e. Table 1 summarizes the properties of these subatomic particles.

Table 1 – Properties of Subatomic Particles
Name Location Charge (e) Mass (amu) Mass
(g)
electron outside nucleus -1 0.00055 0.00091 × 10−24
proton nucleus +1 1.00727 1.67262 × 10−24
neutron nucleus 0 1.00866 1.67493 × 10−24

Characterizing Atoms

Recall that all atoms of a given element have identical and unique chemical properties. The chemical properties of an atom are a function of the subatomic particles contained with that atom.

A defining trait of an element is its atomic number (Z), the amount of protons in the nucleus of an atom of that element.  The atomic number determines the identity of the atom.  For example, helium has an atomic number of 2.  All helium atoms therefore have two protons, and all atoms with two protons are helium atoms.  This is true regardless of how many neutrons and electrons an atom has.

For neutral atoms, atoms that carry no electric charge, the positive charge from protons must be balanced out by the negative charge of electrons.  For example, we know that a helium atom contains two protons.  A neutral helium atom must therefore also contain two electrons.  For neutral atoms, this means that the atomic number also corresponds to the number of electrons in the atom.

An atom’s mass number (A) gives the total number of protons and neutrons in the atom.  Recall that an element is defined by its atomic number (Z), the number of protons in its atoms neuclei, regardless of the number of neutrons.  We can use the mass number to calculate the number of neutrons in an atom as the difference between the mass number and the atomic number, i.e. A – Z = the number of neutrons.  For example, if a helium atom, which has two protons, has a mass number of 4, we can calculate that this helium atom has 2 neutrons (4-2).

This diagram shows the symbol for helium, “H e.” The number to the upper left of the symbol is the mass number, which is 4. The number to the upper right of the symbol is the charge which is positive 2. The number to the lower left of the symbol is the atomic number, which is 2. This number is often omitted. Also shown is “M g” which stands for magnesium It has a mass number of 24, a charge of positive 2, and an atomic number of 12.
Figure 4 — Examples of a helium (He) and magnesium (Mg) atom characterized by its element’s two-letter symbol, the mass number as a left superscript, the atomic number as a left subscript (sometimes omitted), and the charge as a right superscript.

Ions

Recall that neutral atoms contain the same number of positively charged protons and negatively charged electrons; the charges from the protons and electrons cancel each other out so that the atom has a net zero electric charge.  An ion is an atom that does carry an electric charge because there are either more electrons or less electrons than there are protons.  The electric charge of an atom is define as follows:

Atomic charge = number of protons − number of electrons

Atoms typically acquire charge by gaining or losing electrons. An atom that gains one or more electrons electrons will carry a negative charge.  These negatively charged ions are called anions.  Positively charged atoms, called cations, have lost one or more electrons.

As an example, consider the helium (He) atom characterized in Figure 4.  Recall that helium has an atomic number of 2, which indicates that all helium atoms have two protons.  A neutral helium atom therefore has two electrons.  If a helium atom gains one electron, it will have three electrons total (2+1) and become an anion with a charge of -1 e.  If a helium atom loses one electron, it will have one electron left (2-1) and become a cation with a charge of +1 e.  The helium atom in Figure 4 has a charge of +2 e, meaning it has lost both of its electrons.

 

Isotopes

The number of protons in an atom defines the type of element it is; if the number of protons in an atom changes, the type of atom it is also changes.  The number of electrons can change and defines the charge of the atom.  What about the number of neutrons?

During the early 1900s, it was discovered that an element could have atoms with different masses.  Despite the different masses of these atoms, they were otherwise indistinguishable from one another.  The difference in masses corresponded more closely to the mass of neutrons and protons, which have similar masses, than the mass of electrons, which are less massive by about a factor of 2,000 (see Table 1).  It was therefore deduced that these atoms have different masses because they possess different numbers of neutrons.

Such atoms are called isotopes—atoms of the same element that have differing amounts of neutrons.

An isotope of an element is specified by the mass number (A) written as a superscript to the left of the element symbol (see Figure 4).  For example, magnesium naturally occurs in the form of three different isotopes that have mass numbers of 24, 25, and 26.  These isotopes would be written as 24Mg, 25Mg, and 26Mg respectively.  These isotope symbols are read “element, mass number.”  For instance, 24Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.”

 

Unstable Isotopes

Not all isotopes of an element are stable.  Unstable elements undergo a process known as radioactive decay, in which the subatomic particles in a an atom changes.  These unstable elements are described as radioactive.  As unstable nuclei decay, they often change from one isotope into another and sometimes into different elements depending on whether the number of neutrons or protons changes in the nucleus.  The rate of radioactive decay is commonly characterized by half-life, the amount of time in which approximately half the number of atoms will have decayed.

Let’s consider as an example uranium-238, which can also be written [latex]_{92}^{238}\text{U}[/latex], meaning it has 92 protons and and 238-92=146 neutrons.  Uranium-238 has a half-life of 4.5 billion years.  If we had a sample of 12 uranium-238 atoms, then after 4.5 billion years we expect that 6 of those atoms will have undergone radioactive decay.   In other words, over 4.5 billion years, each uranium-238 atom has a half/half or 50% chance of decaying.

Concept Check

We encountered alpha particles earlier in this chapter. How did Ernest Rutherford, Hans Geiger, and Ernest Marsden use alpha particles to understand the components of the atom?

When uranium-238 decays, it emits an alpha particle; an alpha particle consists of two protons and two neutrons.  Because it gives rise to an alpha particle, this process is known as alpha decay and can produce alpha rays.

Recall that uranium-238 has 92 protons and and 238-92=146 neutrons.  With the emission of an alpha particle, a uranium-238 atom loses two protons and two neutrons.  This decay results in an atom with 90 protons and 144 neutrons.  Because the number of protons has changed, the nature of the atom has also changed.  An atom with 90 protons, and therefore an atomic number of 90, is a thorium (Th) atom.  The mass number of this atom is the combined number of protons and neutrons it has, 90+144=234.  Therefore, we see that when undergoing radioactive decay, a uranium-238 atom has decayed into a thorium-234 atom.

 

The process of radioactive decay is happening around us all the time.  If a radioactive nucleus decays into another radioactive nucleus, then another decay can happen.  This chain of reactions is called a “decay chain”. Eventually, the resulting nucleus may be stable and no longer decay. Most of the nuclei in the natural world appear to be stable, but there is a possibility that all nuclei will eventually decay if given enough time. We do not yet know if  “stable” nuclei are just extremely long lived. The process of radioactive decay is key to measuring the ages of rocks and fossils, as we will see later.

Cosmic Abundance of Elements

After the Big Bang, there were only a few elements in existence: hydrogen, helium, and a sprinkling of lithium. As described in the section “Assembling the Periodic Table,” we know that the majority of the remaining elements, including most of those which make up the Earth, were formed at the cores of stars.  These elements were then distributed throughout the cosmos through their violent and explosive deaths.

We have used spectroscopy to identify the chemical composition of other stars, giant molecular clouds, and even the gas around and between galaxies.  These measurements allow us to calculate the abundance of various elements and isotopes.  We see that the universal abundance of elements match what is expected from the processes that create them.

These observations of abundances throughout the universe suggest that the laws governing chemistry and physics that we observe on Earth also operate the same way on any other planet at any other location in our galaxy and beyond. Therefore, if we can understand the chemical origins of life on Earth and in our solar system, we would gain insights into when, where, and how life might arise on other worlds.

Figure 5 shows the relative abundances of different elements in the Solar System. What are the most common elements?

Figure 5 – Relative abundance of elements in the Solar System. Note that this is a log scale so an element that is one major tick mark higher than another is ten times more abundant. Thus, hydrogen and helium make up more than 99% of all atoms in the cosmos as they are more than two major tickmarks above the next most abundant elements: oxygen and carbon.
Concept Check

All life on Earth is based on four elements: hydrogen (H), carbon (C), nitrogen (N) and oxygen (O). How common are each of these elements in the solar system, according to Figure 5?

Why might life not use the element helium (He) for life? Explain your answer!

Energy

Different types of energy are all around us — the electrical energy that powers our technology, the chemical energy that is stored in the food we eat, and the heat energy we feel when when walking in sunlight. These different types of energy can be changed, or converted, from one type to another. For example, the chemical energy in food is converted to energy that allows us to think and move around. We will see other examples of energy transformations in astrobiology, such as the gravitational energy that is converted into heat when a giant cloud of gas and dust collapses and eventually forms stars.

The amount of the energy in the universe is always the same and it is continuously being changed from one form into another. This is the essence of the law of conservation of energy. Other quantities in nature, such as mass and momentum, are also conserved.

Though energy cannot be created or destroyed, it can be transformed into different forms. The two main forms of energy are kinetic energy and potential energy. Kinetic energy is the energy of motion. When an object is moving more quickly, it has more energy. A fast-moving car has more energy than a slow moving car. Thermal energy is actually a form of kinetic energy since higher temperature is really a measure of the average speed of atoms and molecules — the higher the temperature, the greater the kinetic energy of the atoms and molecules.

Potential energy is the energy that is associated with different positions in space. When you throw a ball high into the air, it has more gravitational potential energy than when it is held in your hand. When an object moves from an area of high potential energy to low potential energy, that energy will be transformed into another form. For example, when an apple falls to the Earth, it will speed up gaining kinetic energy. But potential energy does not need to be related just to gravity. For example, the energy stored in chemicals is in the form of chemical potential energy which can be determined based on the position of different atoms bonded to each other by electric forces of attraction.

A more recently discovered form of energy is the energy associated with mass itself. Einstein’s famous equation [latex]E=mc^2[/latex] tells us, that mass-energy equivalence is a fundamental feature of the universe. Any object that has mass therefore also contains energy — this type of energy is called rest-mass energy. Inside the core of stars, nuclear fusion converts some of the rest-mass of two protons into energy. This process occurs an astounding 1038 times every second and is the source of the energy from the Sun that we receive on Earth.

Light

A great deal of the energy that life on Earth uses comes in one way or another from the Sun’s light, although geothermal energy from inside the Earth can also be harnessed. Light from the Sun, or any star, is an example of electromagnetic radiation. Here, the words light and radiation are synonymous, and both are a form of energy. Radiation can sometimes be thought of as dangerous and some types are (for example, gamma rays and X rays), but radiation is a neutral term and is synonymous with energy.

One of the more bizarre aspects of light is the way that it acts, specifically the fact that it can behave like a wave and a particle. In the 17th century, the nature of light was debated and both Isaac Newton and Christiaan Huygens tried to explain it. Newton experimented with lenses and prisms and believed that light was a stream of particles, which he called corpuscles. Huygens, on the other hand, believed that light was a wave that traveled outward in all directions from a source, much like the ripples on a pond that expand outward through the water after a stone is dropped into it. Building upon pioneering work by James Clerk Maxwell and Max Planck, Albert Einstein proposed in 1905 that light can indeed behave both ways, thus solidifying the concept of wave-particle duality, one of the tenets of quantum mechanics. The wave and particle nature of light has now been experimentally verified countless times.

Light as a Wave

Light carries energy and information from one place to another, and this energy is carried in the form of electromagnetic (EM) waves. Other familiar types of waves in nature include sound waves and water waves. All three of these types of waves have some differences and some similarities.

Photo of a frog sitting in a shallow pool of water. A concentric series of waves spreading out from the center where the frog’s song has impacted the water.
Figure 6 – Making Waves. Water waves created when a frog jumps into a pond. The waves move outward away from the frog, and the distance between each wave peak is called the wavelength. Credit: OpenStax Astronomy 2e, CC BY

Water and sound waves are mechanical waves and require a medium to travel in. For sound waves, the medium is air and the medium is water for water waves. If you screamed on the surface of the Moon, which has no atmosphere, nobody would hear you as the sound waves have no medium to travel through. (The director Stanley Kubrick got it right in the film 2001: A Space Odyssey.) EM waves do not require water or air or any medium to travel though – they can travel through empty space. This was such a disturbing idea to nineteenth-century scientists that they actually made up a substance to fill all of space—one for which there was no evidence—just so light waves could have something to travel through: they called it the aether. Today, we know that there is no aether and that EM waves have no trouble at all moving through empty space (as all the starlight visible on a clear night must surely be doing).

The speed of a sound wave depends on the medium through which it is traveling. A sound wave moves faster through water than it does through air, for example. However, all electromagnetic waves move at the same speed in empty space (the speed of light—approximately 300,000 kilometers per second, or 300,000,000 meters per second, which can also be written as 3×108 m/s), which turns out to be the fastest possible speed in the universe.

Now for some similarities. All waves are a kind of repeating phenomenon. Whether it is the up-and-down motion of a water wave or the changing electric and magnetic fields in a wave of light, the pattern of disturbance repeats in a cyclical way. Thus, any wave motion can be characterized by a series of crests and troughs. Moving from one crest through a trough to the next crest completes one cycle. The horizontal length covered by one cycle is called the wavelength (λ). Ocean waves provide an analogy: the wavelength is the distance that separates successive wave crests. An example of water waves created when a frog leaps into a pond is shown in Figure 6.

Figure 7 – One cycle of a wave. Credit: OpenStax Astronomy 2e, CC BY

We can also characterize different waves by their frequency, which is the number of wave cycles that pass by per second. If you count 10 crests moving by each second, for example, then the frequency is 10 cycles per second (cps). In honor of Heinrich Hertz, the late nineteenth physicist who, inspired by physicist James Clerk Maxwell’s work, discovered radio waves, a cps is also called a hertz (Hz). Take a look at your radio, for example, and you will see the channel assigned to each radio station is characterized by its frequency, usually in units of kHz (kilohertz, or thousands of hertz) or MHz (megahertz, or millions of hertz).

Figure 8 – Wavelength and frequency for different colors of light. Credit, NASA (modified), public domain
Concept Check: Wavelength and Frequency

Figure 8 is a schematic showing the wavelengths and frequencies for different colors of light. Red light is at the top and its wavelength (λ) is shown as the distance between two peaks, where the peaks are marked by the vertical black lines. Blue light is on the bottom; notice that it has a shorter wavelength than red light (the length between the two peaks is noticeably shorter).

The frequency of a wave is a measure of how many waves pass by in one second. Let’s imagine that the amount of time that elapsed for each of the waves in Figure 8 is 1 second. Looking at red light, two full waves can pass by in this 1 second period, so we say the frequency is 2 cycles per second, or 2 Hz. (Actually, a bit more than two red light wave cycles can pass by, about 2.5, but we will say 2 full waves to keep this example simple.)

  1. Which color of light has the highest frequency? How about the lowest frequency?
  2. Which color of light has the shortest wavelength?
Show Answer
  1. Looking at Figure 8, the most crests pass by in one second for blue light (6 Hz) while the fewest pass by for red light. Thus, blue light has the highest frequency and red light has the lowest.
  2. Since blue light has the highest frequency, it also has the shortest wavelength. It is easy to see in Figure 8 that λ is shortest for blue light.

Wavelength (λ) and frequency (f) are related because all electromagnetic waves travel at the same speed. The formula for this relationship can be expressed as follows: for any wave motion, the speed at which a wave moves equals the frequency times the wavelength. Waves with longer wavelengths have lower frequencies. Mathematically, we can express this as

[latex]c = \lambda f[/latex]

where the Greek letter for “l”—lambda, λ—is used to denote wavelength and c is the scientific symbol for the speed of light.

The electromagnetic wave model of light (as formulated by James Clerk Maxwell) was one of the great triumphs of nineteenth-century science. In 1887, when Heinrich Hertz actually made invisible electromagnetic waves (what today are called radio waves) on one side of a room and detected them on the other side, it ushered in a new era that led to the modern age of telecommunications. His experiment ultimately led to the technologies of television, cell phones, and today’s wireless networks around the globe.

Light as a Photon

By the beginning of the twentieth century, physicists had to reluctantly accept that sometimes light behaves more like a “particle”—or at least a self-contained packet of energy—than a wave. We call such a packet of electromagnetic energy a photon.

A photon carries a specific amount of energy. How much energy a photon has depends on its frequency. We can use the idea of energy to connect the photon and wave models. How much energy a photon has depends on its frequency. A low-energy radio wave has a low frequency, while a high-energy X-ray at your dentist’s office is a high-frequency wave. Among the colors of visible light, violet-light photons have the highest energy and red-light photons have the lowest.

Test whether the connection between photons and waves is clear to you. In the above example, which photon would have the longer wavelength as a wave: the radio wave or the X-ray? If you answered the radio wave, you are correct. Radio waves have a lower frequency, so the wave cycles are longer

The Electromagnetic Spectrum

Objects in the universe send out an enormous range of radiation, or light. Scientists call this range the electromagnetic (EM) spectrum, which are divided into a number of regions. The spectrum is shown in Figure 9, with some information about the waves in each part or band.

 

Figure 9 – The electromagnetic spectrum. This diagram of the EM spectrum shows the wavelength ranges for different types of radiation. Credit: https://jila.colorado.edu/~ajsh/courses/astr2030_19/electromagneticspectrum.html

Looking at the EM spectrum in Figure 9, we see that:

  1. Gamma rays have the highest energy and radio waves have the lowest energy
  2. Radio waves have the longest wavelength and gamma rays have the shortest wavelength
  3. Radio waves have the highest frequency and gamma rays have the lowest frequency

These properties of light are summarized by the beautifully simple equation that relates the energy of a photon to its frequency (or wavelength), where h is Planck’s constant:

[latex]E=hf[/latex]

This relationship shows the wave-particle duality of light, as the energy of photon (a particle of light) is directly related to its frequency (a wave property). Since h has a constant value, you can immediately get the energy of a particular color of light simply by knowing its frequency (or wavelength, which can be expressed as [latex]\lambda = c / f[/latex]).

Another property of light that has already been mentioned but is worth re-emphasizing is that all types of light travel at the speed of light. Since the speed of light is constant, all light travels at the same speed.

Worked Example: Frequency and Wavelength

Red light has a wavelength of about 650 nm and blue light has a wavelength of about 450 nm. Recall that nm is the abbreviation for nanometers.

What is the frequency of red light?

The frequency of any type of light is related to its wavelength as [latex]f = c / \lambda[/latex], where c is the speed of light.

To use this equation, we must be sure that the units are all consistent. If we use 3×108 m/s for the speed of light, then the wavelength must be in units of meters and the frequency will be in units of Hz (where 1 Hz = 1 cycle per second). So, we must first convert the wavelength from nanometers into meters. One nanometer is one-billionth of a meter: 1 nm = 10-9 m so it follows that 650 nm = 650×10-9 m. Expressed in scientific notation, the wavelength of red light is 6.50×10-7 m.

[latex]f = c / \lambda[/latex] = (3×108 m/s) ⁄ (6.50×10-7 m) = 461,538,461,538,461 Hz

That’s a huge number of cycles in one second! Expressing this in scientific notation, f = 4.62×1014 Hz.

The wavelength of visible light (like red and blue) is often reported in THz, where T is the prefix one trillion, or 1012.

Extra 1: Show that the frequency can also be expressed as 462 THz.

Show Answer

The frequency 4.62×1014 Hz can also be written as 4.62×102×1012 Hz

4.62×102 = 462 and 1012 Hz = 1 THZ, so f = 462 THz

Extra 2: Show that the frequency of blue light (with a wavelength of 450 nm) is 666 THz.

Show Answer

[latex]f = c / \lambda[/latex] = (3×108 m/s) ⁄ (4.50×10-7 m) = 6.66×1014 Hz = 666 THz

Types of Electromagnetic Radiation

Throughout our study of astrobiology, we will encounter all types of light. When we collect light using a camera or sensor on the Earth’s surface, that light must first travel through the Earth’s atmosphere before it reaches the ground. Depending on the type of light, some of it may be partially or completely absorbed by molecules in different parts of the Earth’s atmosphere. This means that some types of telescopes need to be above the Earth’s atmosphere in order to detect any light. Also, we will see that this absorbed light will show up as a kind of set of Earth’s chemical fingerprints in any observations that contain the Earth’s atmosphere; this is the basic idea behind biosignatures for life, and we will probe it more deeply in the exoplanet chapters.

The way that light is absorbed in Earth’s atmosphere for each part of the EM spectrum is shown in Figure 10.

Figure 10 – Absorption of light by Earth’s atmosphere. The Earth’s atmosphere stops most types of electromagnetic radiation from space from reaching Earth’s surface. This illustration shows how far into the atmosphere different parts of the EM spectrum can go before being absorbed. Only portions of radio and visible light reach the surface. Credit: OpenStax Astronomy 2e, CC BY

We highlight a few properties of each type of light that are of relevance to astrobiology, as well as what types of telescopes are used to study the universe in each type of light.

Radio Waves

Radio waves have the longest wavelengths of any type of radiation — they can range from a few millimeters all they way up to hundreds of kilometers (for reference, Switzerland is about 300 km across). Radio waves have very low energies and are all around us in our everyday lives. On Earth, cell phones, radio and TV transmissions, satellites and radar all emit radio waves and these pass right through our bodies. Radio waves are also produced in nature by a number of astrophysical objects, including stars and pulsars. The most abundant element in the cosmos — hydrogen — also naturally emits radio waves at a very specific wavelength of 21 cm (which corresponds to a frequency of 1420 MHz). Astronomers search for signals from advanced extraterrestrial civilizations that are close to the 21 cm radiation that hydrogen produces. Most radio waves reach the ground, as seen by the “radio window” in Figure 10. The lowest frequency radio waves do not make it to the Earth’s surface and this region of the EM spectrum remains unexplored.

Microwaves

Microwaves are also low energy radiation and have sizes that range from about 1 mm up to 300 mm. The most familiar source of microwaves on Earth is probably microwave ovens. Microwaves are used extensively by satellites to monitor weather on the Earth (Doppler radar maps use microwaves) and for remote sensing of the Earth’s surface. The ancient radiation left over from the Big Bang also happens to fall into the microwave part of the EM spectrum. As seen in Figure 10, some microwave radiation makes it to the ground, but other parts are also absorbed by the Earth’s atmosphere.

Infrared

Infrared (IR) light is given off by anything with a temperature. Our eyes are only sensitive to visible light (some animals, such as vampire bats and goldfish, can see into the infrared), although we can still “see” this type of light by using an IR camera or night vision goggles. These types of sensors pick up differences in infrared intensity. Since most IR light gets absorbed by the Earth’s atmosphere, many IR telescopes are in space.

Images of the Carina Nebula taken in visible light (left) and infrared light (left) by the Hubble Space Telescope.
Figure 11 – Visible and infrared light comparison. Images of the Carina Nebula taken in visible light (left) and infrared light (left) by the Hubble Space Telescope. Credit: NASA, ESA and the Hubble SM4 ERO Team

IR light is of great importance in astronomy because it can penetrate, or travel through, dust that is opaque to visible light. This idea is best shown with a comparison of the same picture taken in visible light and infrared light. Figure 11 shows pictures of the Carina Nebula taken by the Hubble Space Telescope (HST) in visible light (left) and infrared light (right). You can see an enormous amount of stars and other objects in the infrared that are not seen in the visible, as infrared light travels right through the dust while visible light is absorbed by the dust.

The James Webb Space Telescope (JWST) detects IR light. IR light covers the wavelength range of about 1-1000 micrometers, where micrometers are abbreviated as μm and are often referred to as microns; for reference, E. Coli bacteria are 2 microns and a human hair is 100 microns. JWST can sense wavelengths in the range of 0.6-28 microns. One of the main goals of JWST is to study planetary systems and the origin of life. One way JWST is doing this is by studying the atmospheres of exoplanets to see if any molecules needed for life, as we know it, are present. This will be discussed in more detail in the chapters on exoplanets.

Visible Light

Human eyes evolved to detect wavelengths spanning the visible part of the EM spectrum, which covers 400 nm to 750 nm. Stars emit visible light and life on Earth would not exist without the light from the Sun. Microscopes and optical telescopes both take visible light and focus it into an image — the study of visible light is essential to astrobiology. For example, detailed pictures of the surface of Mars that show evidence for past liquid water are taken with visible cameras like Mastcam-Z on the Perseverance rover.

Ultraviolet

Ultraviolet (UV) light has a shorter wavelength than visible light and ranges from about 10-400 nm. Stars hotter than the Sun emit primarily UV light but the Sun emits some UV light, too. This is probably most apparent when you are outside on a sunny day and forgot to bring some sunblock — your skin will absorb some of the UV (specifically, UVA) light. The fact that stars emit UV radiation is important for many aspects of life, in addition to protecting your skin. UV light can actually break the bonds between the hydrogen and oxygen atoms in a water molecule in a process called photolysis and this can create ozone in the atmosphere, which in turn protects life from dangerous radiation that can damage cells.

UV light is emitted from a range of astrophysical sources, including sites of star formation and planetary aurorae. As seen in Figure 10, nearly all UV radiation is absorbed by the Earth’s atmosphere so UV telescopes are almost always in space, although some telescopes on high mountain tops, like the Keck telescopes atop Mauna Kea in Hawaii, have UV instruments. NASA’s SWIFT satellite is actively studying the universe in UV light, and the Hubble Space Telescope has been probing a wide range of UV sources since 1990.

X Rays

X rays are a high energy type of radiation, with short wavelengths of just 0.01-10 nm, and they can be absorbed by tissue and bones in our bodies. They can destroy DNA so great care is taken when X ray machines are used to take medical images. Fortunately for life, X rays are absorbed by the Earth’s atmosphere so they are not a concern on Earth’s surface (but are a great concern for astronauts traveling above the Earth’s atmosphere).

X rays are generated by high energy astrophysical phenomena such as supernova as well as the Sun’s corona. During a total solar eclipse, the very hot gas in the Sun’s corona can be seen. The composition of soil on Mars is studied using X ray instruments on Martian rovers, such as Spirit and Opportunity and the active Perseverance mission.

Gamma Rays

Gamma rays are the most energetic type of radiation in the cosmos and have wavelengths on the order of the size of the nucleus inside an atom. On Earth, gamma rays are produced naturally by lightning and radioactivity. Gamma rays are also created when two neutron stars merge, in addition to the gravitational waves that these events create. Gamma rays produced in neutron star mergers produce many of the heavy elements on the periodic table, such as gold and platinum.

Gamma rays do not make it to the Earth’s surface so gamma ray observatories need to be in space. The Fermi Space Telescope has been studying gamma rays on the sky since 2008 and has created a complete map of the sky in gamma rays. The surface composition of a planet can be studied using gamma ray instruments — the MESSENGER mission to Mercury carried a gamma ray spectrometer.

Multiple Wavebands

To fully understand an object, we ideally want to look at in every type of light that it emits. For many object, this can include all types of radiation, from radio waves to gamma ray. An example of this is shown in Figure 12 for the Crab Nebula:

Figure 12 – The Crab Nebula in different parts of the EM spectrum. Telescopes used to create the images: Very Large Array (radio), Spitzer (IR), Hubble Space Telescope (visible), Swift (UV), Chandra (X), Fermi (gamma). Credit: wikimedia, CC BY-SA

Understanding Radiation

Some astronomical objects emit mostly infrared radiation, others mostly visible light, and still others mostly ultraviolet radiation. What determines the type of electromagnetic radiation emitted by the Sun, stars, and other dense astronomical objects? The answer often turns out to be their temperature.

At the microscopic level, everything in nature is in motion. A solid is composed of molecules and atoms in continuous vibration: they move back and forth in place, but their motion is much too small for our eyes to make out. A gas consists of atoms or molecules that are flying about freely at high speed, continually bumping into one another and bombarding the surrounding matter. The hotter the solid or gas, the faster the motion of its atoms or molecules. The temperature of something is thus a measure of the average motion energy of the particles that make it up.

This motion at the microscopic level is responsible for much of the EM radiation on Earth and in the universe. As atoms and molecules move about and collide, or vibrate in place, their electrons give off EM radiation. The characteristics of this radiation are determined by the temperature of those atoms and molecules. In a hot material, for example, the individual particles vibrate in place or move rapidly from collisions, so the emitted waves are, on average, more energetic. And recall that higher energy waves have a higher frequency. In cooler materials, the particles have lower energy atomic and molecular motions and thus generate lower energy waves.

Blackbody Radiation and Wien’s Law

To further understand the relationship between temperature and light (EM radiation), we consider a type of ideal object called a blackbody. A blackbody absorbs all light that hits it; none of the incoming light is reflected away or passes through the object. The object heats up and gains energy; the object then loses this energy by emitting light at all wavelengths (recall that light is a type of energy). In other words, the object cools off until it reaches the same temperature as its surroundings — this is called thermodynamic equilibrium — and it does this by converting the energy it absorbed into energy of motion that is radiated in a specific way (this radiated energy is called thermal or blackbody radiation). A glowing ember of coal cools by radiating energy until it reaches the same temperature as its environment.

Any dense, solid object that has a temperature emits blackbody radiation; this includes people, animals, stove tops, light bulbs, stars and planets. For example, the Earth absorbs EM radiation from the Sun, heats up, and then emits radiation as mostly infrared light. Note that any objects that reflect light, such as a book with a red cover, are not blackbodies. Similarly, if light passes through the object, as with a glass table, it is also not a blackbody. In that respect, the Earth is not a perfect blackbody, since clouds and other particulate matter do reflect some sunlight.

No object in nature is a perfect blackbody but many opaque objects behave like a blackbody so it is a very good approximation and we can use some fairly simple math to learn more about the object. As mentioned above, blackbody radiation is not emitted at just one wavelength or one part of the EM spectrum, but covers all wavelengths. At first this may sound strange — since humans act like blackbodies, does that mean that our bodies emit dangerous X rays and gamma rays? No, that is not the case (I don’t have X ray eyes) and any blackbody emits the most intense radiation at one particular wavelength. Humans have a typical temperature of 98.6°F (or 310 K) and emit the most intense radiation at 9.3 microns (1 micron = 10-6 m, so this an infrared wavelength, meaning we radiate heat!). The peak wavelength of the radiation that a blackbody gives off depends only on its temperature. This means two very different objects with the same temperature emit identically with the same wavelength of peak intensity. The Earth is just a little bit cooler than a human at 288 K, and its most intense radiation is at 10.1 microns. This relationship between the temperature of a blackbody and its peak wavelength is called Wien’s Law and is:

[latex]\lambda_{\rm max} = \frac{ 0.0029}{T} \rm{~m}[/latex]

where the wavelength λmax is in meters and the temperature is in K (the constant 0.0029 has units of m × K). You can easily verify λmax for a human at 310 K: (0.0029 m K)/(310 K) = 9.3×10-6 m = 9.3 microns.

The thermal radiation emitted by any blackbody shows a similar shape in the intensity of light at different wavelengths. These shapes are called blackbody curves or thermal energy curves. Figure 13 shows this general shape for blackbody curves for objects of different temperatures. The shape is similar for all temperatures but notice that the intensity for each blackbody curve peaks at a different maximum wavelength.

 

Figure 13 – Wien’s Law Illustrated. This graph shows how much light is given off at each wavelength for objects at four different temperatures. The wavelengths corresponding to visible light are shown by the colored bands, with UV and IR to the left and right. Note that at hotter temperatures, more energy (in the form of photons) is emitted at all wavelengths. The higher the temperature, the shorter the wavelength at which the peak amount of energy is radiated (this is known as Wien’s law). Credit: OpenStax Astronomy 2e, CC BY

Notice in Figure 13 that the curves show that, at each temperature, the blackbody objects emit radiation at all wavelengths (all colors) but that the most energy is emitted at a peak wavelength. Physically, this peak corresponds to the average speed of atoms or molecules inside the object, which is a manifestation of its temperature.

The hotter the object, the shorter the peak wavelength — the object at 5500 K has a shorter λmax (530 nm) than the object at 2500 K (which has λmax of 1160 nm and is in the infrared portion of the EM spectrum). It makes sense, then, that hot objects give off a larger fraction of their energy at shorter wavelengths (higher energies) than do cool objects. You may have observed examples of this rule in everyday life. When a burner on an electric stove is turned on low, it emits only heat, which is infrared radiation, but does not glow with visible light. If the burner is set to a higher temperature, it starts to glow a dull red. At a still-higher setting, it glows a brighter orange-red (shorter wavelength). At even higher temperatures, which cannot be reached with ordinary stoves, metal can appear brilliant yellow or even blue-white.

We can use these ideas to come up with a thermometer for measuring the temperatures of stars. Because many stars give off most of their energy in visible light, the color of light that dominates a star’s appearance is an indicator of its temperature. If one star looks red and another looks blue, which one has the higher temperature? Because blue is the shorter-wavelength color, it is the sign of a hotter star. (Note that the temperatures we associate with different colors in science are not the same as the ones artists use. In art, red is often called a “hot” color and blue a “cool” color. Likewise, we commonly see red on faucet or air conditioning controls to indicate hot temperatures and blue to indicate cold temperatures. Although these are common uses to us in daily life, in nature, it’s the other way around.)

What about the Sun? The Sun has a surface temperature of 5800 K so Wien’s law immediately tells us that λmax is 500 nm. Which corresponds to green light. But the Sun doesn’t look green, it looks yellow or white! What’s going on?! This makes sense when you consider that the Sun is also emitting red light and blue light; these all mix together as a nearly white color. Pure purple stars are also never seen for similar reasons — the blue and violet mix into a more deep blue color.

Luminosity and Brightness

Wien’s Law shows that blackbodies, like stars, emit the most intense radiation at a specific wavelength, λmax. What exactly is this intensity a measure of?

Luminosity is the total amount of energy that an object (like a star) puts out each second. It has dimensional units of energy per second. In the same way that a 100 W bulb will always put out 100 Watts whether we are standing close or farther away, the luminosity of a star does not depend on our distance from it.

However, astronomers do not measure luminosity directly with a telescope; they measure brightness: the luminosity that is intercepted by a detector such as a photographic plate or a digital camera. If you imagine a spherical surface – a bubble – around a star, then the luminosity is the integrated (total) light from the surface of that bubble. The brightness, which is the luminosity per unit area, decreases as the surface area of the spherical volume increases. This is the same phenomenon that happens with expanding balloons. The balloon has a certain amount of material, usually latex or rubber. Analogous to luminosity, that amount of material is constant, no matter how much air is in the balloon. However, as the balloon expands, that constant amount of material is stretched over a larger surface area. The walls of the balloon get thinner and the amount of material per unit area decreases. The luminosity of a star is constant. The brightness that we measure depends on whether we are “up close” or far away from the star.

This is shown in Figure 14, where the light from a star with luminosity L spreads out into spheres of increasing surface area as it travels away from the star. At a distance d from the star, we can measure the brightness ($B$) by dividing the total power emitted by the star (L) by the surface area of the sphere that the light has now spread into:

[latex]B = \frac{L}{4 \pi d^2}[/latex]

This idea—that the apparent brightness of a source (how bright it looks to us) gets weaker with distance in the way we have described—is shown in Figure 14 below. At point 1, the light is concentrated into one box. By the time the light reaches point 2, which is twice as far as point 1, it is now spread out into four squares.

Figure illustrating the inverse square law for light. At left is a spherical light source such as a star. Four arrows move radially outward from the source toward the right, with each arrow representing one corner of a square. At one unit away from the source (labeled 1), the arrows are close together to form a square one unit high, and thus one unit square. At the next labeled unit (labeled 2), two steps away from the source, the square is now two units high, and thus four units square. At this distance from the source the energy is 4 times less than at step 1. The final labeled step (labeled 3) is 3 units away, thus the square bounded by the arrows is now 9 units in area, and the energy is 9 times less than at step 1.
Figure 14 – Inverse Square Law for Light. As light radiates away from its source, it spreads out in such a way that the energy per unit area (the amount of energy passing through one of the small squares) decreases as the square of the distance from its source. Credit: OpenStax Astronomy 2e, CC BY
Concept Check: Brightness and distances of stars

You observe two stars, named Sol-2 and Sol-3, that have the exact same luminosity as the Sun (this means they have the same temperature, radius, and luminosity as the Sun). You measure the brightness of both stars with the same instrument and find that the light from Sol-2 is twenty-five times brighter than the light from Sol-3. Which star is closer to the Earth and by how much? Explain your reasoning.

Show Answer

Ans: Sol-2 is five times closer than Sol-3. Brightness drops as the distance squared, so the ratio of the distances to Sol-2 and Sol-3 is given by [latex]\sqrt{B_{Sol2}/B_{Sol3}}[/latex] or [latex]\sqrt{25}[/latex] = 5.

Let’s consider the luminosity of the Sun and the amount of energy from the Sun that reaches the Earth. This will be important to know when we study exoplanets and want to know how much light from the host star reaches the exoplanet. The luminosity of the Sun is 3.9×1026 W — in other words, it emits the same amount of power as 3.9×1024 100 W light bulbs (or 39 trillion trillion 100 W bulbs!). As the light travels away from the Sun, it is spread out more and more, as all of that light needs to cover spheres (bubbles) that have greater surface area. The distance from the Earth to the Sun is 151 million km, or 1.51×1011 m. Putting this value into the brightness equation, we find that we receive 1360 W/m2 of sunlight on the Earth. That means every 1 m2 patch on the Earth, which is about the size of a standard chess board, receives 1360 Joules of energy every second. Venus is closer to the Sun than the Earth and receives more energy for every square meter, 2600 W/m2, while Mars is further and receives only 600 W/m2.

Worked Example: Calculating the energy received from the Sun

How much radiation from the Sun is received at Saturn’s moon Titan? You can assume that Titan and Saturn are at the same distance from Earth.

We can find the value by using the relationship for brightness, luminosity and distance:

[latex]B = \frac{L}{4 \pi d^2}[/latex]

Here, L is the luminosity of the Sun and d is the distance of Titan from the Sun. The distance from the Sun to Titan (Saturn) is 1.45 billion km (or 1.45×1012 m).

B = (3.9×1026 W)/(4 π (1.45×1012 m)2) = 14.8 W/m2

This makes sense — Saturn is almost 10 times further from the Sun than the Earth (Earth is 1 AU and Saturn is 9.57 AU), so the value should be significantly lower than Earth’s. [In fact, you could also find this value by taking the ratio (1/9.57)2 and multiplying by 1360 W/m2]

Show mathematically that the Earth receives 1360 W/m2 from the Sun.

Show Answer

The only difference here is the distance: now it is the distance from the Earth to the Sun, which is 151 million km.

B = (3.9×1026 W)/(4 π (1.51×1011 m)2) = 1361 W/m2

Spectroscopy

Light from objects is filled with information. By taking the light from an object and splitting it apart into a spectrum, we can decode this light. Spectroscopy is a fundamental tool for astrobiology and allows us to detect exoplanets and study the chemical makeup of their atmospheres.

The most familiar example of spectroscopy is the rainbow of colors that is produced when white light is sent through a prism, as seen in Figure 1. A water droplet can also split up (or disperse) light and create a rainbow when a light beam passes through it. Both prisms and water droplets can serve as spectrometers or spectrographs — tools used to disperse light into a spectrum.

White light being sent through a prism to create a rainbow
Figure 15 – Different wavelengths of light disperse through a glass prism. Credit: NASA, ESA, Leah Hustak (STScI), public domain

Stars emit all types of electromagnetic radiation but emit the most intense light at a specific wavelength (see Wien’s Law). Note that infrared and ultraviolet light are also emitted by the Sun and are split out as seen in Figure 15 — our eyes just can’t perceive these types of radiation. The distribution of colors, which is naturally organized by wavelength (or frequency), is called a spectrum. There are three main types of spectra: continuous, absorption, and emission as shown in Figure 16. If the shape of the spectrum on the bottom left of Figure 16 looks familiar, it is because you have seen it before when learning about blackbody (thermal) radiation curves.

Figure 16 – Formation of different types of spectra: continuous, emission, and absorption. Credit: NASA / STScI, public domain

A continuous spectrum looks like the rainbow shown in Figure 15. However, sunlight that reaches the Earth actually is an absorption spectrum, not a continuous spectrum. Why? The Sun does emit a continuous spectrum from its core but some of that light is absorbed by different atoms as it passes through the Sun’s atmosphere. This is shown in Figure 16. On the left, the Sun emits a continuous spectrum at its surface. But that light passes through the atmosphere (the cloud of gas in the middle) and some of that light is absorbed by atoms in the gas. The spectrum on the right is the type of spectrum seen for the Sun, where the dark lines correspond to the wavelengths associated with the atoms or molecules that absorbed the light. Figure 17 shows an absorption spectrum for the Sun, with some of the atoms and molecules present in the Sun’s atmosphere labeled. This spectrum was taken from the Earth’s surface, so the O2 (molecular oxygen) lines are from Earth’s atmosphere.

Figure 17 – The absorption spectrum for the Sun. Each black line indicates the presence of a specific atom in the Sun’s atmosphere. For example, we can see two sodium (Na) lines at 590 nm, which tells us that there are sodium atoms in the Sun’s atmosphere. Credit: continuous spectrum

What does it mean to say “the wavelengths associated with the atoms”? Every element on the periodic table has a different number of protons and thus a different structure. Hydrogen is the simplest atom, with just one proton and one electron orbiting the proton in a neutral atom. To get the electron out of the hydrogen atom entirely requires a certain amount of energy. If there is no energy supplied to the atom, it is stable and the electron is in what we call the ground state of energy. However, the electron can move to different energy levels above the ground state — imagine rungs on a ladder, with the ground state at the bottom and the top rung being the amount of energy needed to kick the electron out of the atom. Hydrogen has a unique set of energy levels and each level has an exact amount of energy. And recall that this energy has an exact frequency or wavelength given by [latex]E=hf = h c/\lambda[/latex]. When light particles (photons) with exactly one of those wavelengths encounters the atom, it is absorbed. And every element has a different set of energy levels, meaning a distinct set of its own characteristic wavelengths. We thus have a very powerful way of learning which elements are inside a star’s atmosphere.

Using the unique pattern of an element to identify whether it is inside of an atmosphere is similar to DNA barcoding used in forensics to identify a match to evidence at a crime scene. Figure 18 shows an example of some of the colors (wavelengths) for a few different elements. No two are the same and we can use these patterns to identify an element in a star or planet’s atmosphere. This same technique is used extensively in biochemistry and many other scientific fields. We can think of the unique emission lines associated with an atom or molecule as its chemical fingerprint.

Figure 18 – The “chemical fingerprints” for a few different elements. Every element has its own unique set of spectral lines. Astronomers can use these fingerprints to determine what atoms or molecules are in a star or exoplanet’s atmosphere. The units for the wavelength axis are called angstroms (Å), and 1 Å is equal to 10-10 m. How many angstroms are in 1 nanometer (nm)? Credit: OpenStax Astronomy 2e, CC BY

The middle part of Figure 16 shows the physical idea of how an emission spectrum is created. Let’s say you have a gas made from atoms of just one type of element. If you heat up that gas, the electrons can move up to higher energy levels and then drop back down, emitting light in the process. This is almost the same situation as the absorption spectrum on the left except we are not looking at the continuous source and the gas but instead just at the gas…so there is no continuous spectrum in the background, just the bright lines that correspond to whatever element is in the gas. The atom is emitting energy at the unique set of wavelengths (colors) for that particular element. Of course, the gas can contain many different atoms and molecules at the same time.

You can take a look at the fingerprints for some other elements using the simulation below. Click on any element to see its unique set of spectral lines. Notice that you can choose to see both absorption (top) and emission (bottom) spectra.

 

Key Concepts and Summary

Atoms are fundamental units in chemistry that explain the Periodic Table as a sequence with increasing numbers of protons. The nucleus of an atom is comprised of positively charged protons and uncharged neutrons. It is the number of protons in the nucleus that defines the element – isotopes of elements have different numbers of neutrons. In an electrically neutral atom, the number of electrons equals the number of protons. After the Big Bang, the only elements that emerged were hydrogen, helium, and traces of lithium. All other elements were forged in the cores of stars or during highly energetic explosions. The elements in the universe today are generally decreasing in abundance with increasing atomic mass (i.e., highest atomic mass elements are rarest), but this trend is modulated by even-odd atomic numbers, a signature of the fusion of alpha particles. The energy of light is our messenger about the universe. Different colors of light are simply different wavelengths and different wavelengths of light contain different amounts of energy. The longest wavelengths in the electromagnetic spectrum are radio waves and these have the lowest energy. The shortest wavelengths are high energy gamma rays. By observing the universe with detectors that are sensitive to different energies of light, we can learn about the energy output and the peak temperatures of the objects we detect. Because the brightness of a star or other celestial object decreases with distance, a measurement of distance is needed to back out luminosity (energy per second). The wavelengths of light encode the energy and chemical composition of stars and other celestial bodies – astronomers use spectrographs to disperse light collected at a telescope into a spectrum. Objects emit a continuous (sometimes called “blackbody”) spectrum with an intensity distribution that depends on temperature. When atomic transitions take place in cooler outer layers of stars, the spectrum can also have absorption or emission lines. Absorption lines occur when atoms absorb photons of light pushing electrons to a higher energy state. Emission lines occur when atoms release photons and the electrons cascade to a lower energy state.

 

Review Questions

Summary Questions

  1. What is an atom?
  2. How are atoms relate to chemical elements?
  3. What is the difference between an atom and a molecule?
  4. What three “subatomic” particles are found inside an atom? Describe the similarities and differences between them.
  5. How are different isotopes of a particular element distinguished? What are some reasons that certain isotopes are more common than others?
  6. Which elements are the most common and which elements are the least common in our universe? How does astrophysics help explain this?
  7. What is energy? Describe two different kinds of energy and give examples of each.
  8. What does it mean to say the energy is “conserved”?
  9. How are the frequency and wavelength of light related to each other?
  10. What is a photon?
  11. Which property of light determine how much energy a photon has?
  12. What are the seven main parts of the electromagnetic spectrum? For each band (part), give an example of a source of that type of radiation.
  13. How does the type of light emitted by an object depend on its temperature?
  14. What is the difference between the brightness of an object (such as a star) and its luminosity?
  15. How does the total amount of energy emitted differ for hot and cold objects?
  16. How does spectroscopy give information about an object? What happens when starlight passes through a prism?
  17. What are the differences between continuous, absorption and emission spectra? What are the similarities between all three?
  18. Why are the spectral lines for an element sometimes called “chemical fingerprints”?

Activities

  1. Building atoms. Open the “Build an atom” simulator (https://phet.colorado.edu/en/simulations/build-an-atom) and select the Atom option. Build the following:
    1. hydrogen, deuterium, tritium
    2. lithium (is this atom stable?)
    3. carbon, carbon-12, carbon-13
  2. Understanding blackbody radiation. Open the simulation at https://phet.colorado.edu/sims/html/blackbody-spectrum/latest/blackbody-spectrum_en.html.
    • The default is set to the temperature of the Sun. Click on the “labels” box to show different parts of the EM spectrum (UV, visible, and infrared). What part of the EM spectrum has the most intense light from the Sun.
    • Now set the thermometer to the temperature of the star Sirius. The y-axis now needs to be rescaled to see the value for the peak intensity, Adjust this by using the zoom out button. Once you can see the peak, click on the “Graph Values” box. What is the wavelength of the peak intensity light for Sirius?
  3. Energy flux. Calculate the solar flux (in W/m2) at the planet Mercury.
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Cosmic Origins Copyright © by Debra Fischer; Allyson Sheffield; Joshua Tan; Lily Ling Zhao; and Dawn Erb is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.