Download Lab 5

Astronomy 115
Name(s):
Lab 5: Chemistry and spectroscopy
Purpose: To build physical and symbolic chemical representations of matter. To
demonstrate the use of spectroscopy to infer properties about matter.
Note: The first part of this lab references Lab 4 questions.
Chemistry
Chemistry is the study of matter and its changes, using the rules of physics. Since ours is
a universe of matter, you will need to know a little about these rules of how matter
changes. Atoms are the smallest unit of matter, but most matter is made of
molecules. Therefore, you will need a set of simplified rules to describe bonding
between atoms to form molecules. Obtain a “ball-and-stick” molecular model kit from
the cart.
The rules:
• White balls represent hydrogen atoms
• Red balls represent oxygen atoms
• Black balls represent carbon atoms
• Blue balls represent nitrogen atoms
• The plastic sticks represent chemical bonds (which are merely a
pair of shared electrons between atoms); length doesn't matter
• For a molecule to be "happy" (i.e., have all of its bonding
requirements satisfied), all holes must be filled with bonds
1. Using two white balls and one plastic stick as a bond, make a model of the simplest and
most common molecule in the universe — the hydrogen molecule. Write the formula for
the hydrogen molecule by finding the elemental symbol for hydrogen (hint: see the
periodic table of elements poster near the clock) and then adding a subscript number
that indicates how many atoms are bonded together.
2. So a hydrogen atom can only bond to one other atom, maximum. What is the maximum
number of atoms that carbon could bond to (hint: count holes)? That oxygen could bond
to? That nitrogen could bond to?
3. Why do you suppose there are no helium atoms in the kit? Hint: Helium is a “noble”
gas.
4. Create the following molecules: water (H2O), ammonia (NH3) and methane (CH4)
and carbon dioxide (CO2). Draw the molecules, as best as you can, below.
At high temperatures, like in a star or a nebula, chemical bonds (the sticks) all break,
leaving atoms free. Atoms contain three principal subatomic particles: the neutron, the
proton and the electron.
More rules:
• Atoms of an element all contain the same number of protons, given by the atomic number
on the periodic table; if the number of protons changes in an atom, it becomes a different
element.
• Neutral atoms contain the same number of protons and electrons.
• The total number of protons and neutrons in an atom is called the isotope number, and
distinguishes different isotopes of the same atom.
5. Quick review of atomic structure; fill in the missing information in the table below:
Atom
(isotope)
hydrogen-1
Symbol
hydrogen-2
(deuterium)
hydrogen-3
(tritium)
helium-3
2H
helium-4
4He
carbon-12
12C
1H
Number of
protons
1
Number of
neutrons
0
Number of
electrons
1
3H
3He
Moreover, electrons from within an atom can be ejected (ionized). The resulting atom is
called an ion, and will contain fewer electrons than protons, resulting in a net positive
charge; this type of ion is called a cation. Meanwhile, the electron that was ionize can
stabilize around a different atom, which will then have more electrons than protons,
resulting in a net negative charge – this is an anion.
Ions therefore have an overall non-zero electrical charge and are thus more reactive in
general than atoms. The overall charge is written as a superscript to the right of the element
symbol, so, for instance, a calcium ion that has two fewer electrons than protons would be
written: Ca2+.
Confusingly, an old astronomical notation for cations uses Roman numerals to indicate
various numbers of missing electrons, but this numbering system starts with (I) for neutral
atoms. For instance, calcium (I) is neutral calcium, but calcium (III) is the Ca2+ ion given
above.
6. Fill in the missing information in the table of ions:
Ion
H (neutral)
Astronomy
# of protons
notation of ion
H (I)
1
# of neutrons
# of electrons
0
1
H+
He+
He2+
Ca (II)
Chemistry and spectroscopy
7. Explain the observation of the UV graph, specifically the number of “peaks” you see.
Note: there are some stars that emit in these wavelengths, which require UV
spectroscopes to “see”.
8. Obtain and observe the helium gas lamp spectrum. Do hydrogen and helium have
distinct, different emission wavelengths (the hydrogen lamp will be set out as
well)? Let’s generalize; will emission wavelengths be useful for distinguishing different
elements?
9. Thus, the solar spectrum (the spectrum of the Sun) is a mystery indeed. How can the
Sun’s spectrum be so continuous (as opposed to the fluorescent bulb, which had a
discrete spectrum)? That is, what does the continuous nature of the solar spectrum tell
you about the Sun? Hint: I have said that the Sun is made of hydrogen and helium. Is it
made only of hydrogen and helium?
10. Examine the solar spectrum (on the “Spectra of Different Elements” chart).
However, it is not as continuous as I had implied. What is the origin of the “missing”
colors of the Sun’s continuous spectrum? Hint: what kind of spectroscopy (emission
or absorption) leads to this kind of spectrum?
11. What is the chemical formula for water? What are the chemical symbols for
hydrogen gas and oxygen gas? So what is the difference between water vapor, and a
mixture of hydrogen and oxygen gas?
12. Examine the spectrum of the water vapor lamp, and the oxygen lamp. Are there
any emission lines (colors) that are in the water vapor spectrum that are not already in
the hydrogen or oxygen spectra? If so, list the wavelengths in nanometers. If not, explain
what happened to the water vapor to yield such similar spectra as the elements.
13. Therefore, is visible light bright-line (emission) spectroscopy a good way to detect
elements in a glowing object? Is visible light bright-line spectroscopy a good way to
detect chemical compounds?
On the next page is a table of stellar spectra; each “row” represents the emissions from a
particular star. The star’s abbreviated name is given to the left of its spectrum; for
instance, the first star is “10 Lacerta”, which is the tenth brightest star in the
constellation Lacerta. The star’s spectral classification is given to the right.
Of interest to us are the apparently light vertical “lines” that appear at irregular
intervals throughout each spectrum. These spectra are actually negatives — that is, most
of the spectrum should be light-colored, and the lines should appear dark.
14. Do the lines therefore represent light emission or light absorption?
15. Based on your answer above, what is the cause of these “lines”? Note that your
answer will depend on whether you believe these are emission lines or absorption lines!
16. List an element that appears in both uncharged (neutral) and charged (ionic) form in
a spectrum. Do they have the same wavelength absorption?
17. Of course, even ions can be heated. Recall that the Sun was not only made of
hydrogen and helium. So, in addition to the source you cited to supply the “missing”
colors of the Sun’s spectrum in question 9, what other source of emission do stars have?
18. Annie Jump Cannon’s spectral classification system was based on these lines.
The table on the next page shows some unclassified stellar spectra. Based on the table of
classified stellar spectra on the previous page, please classify the unclassified stars. You
need not worry about the Arabic numeral; just put down a letter.
Star
Alpha
Stellar
class
Beta
Gamma Delta
Epsilon Eta
Theta
Iota
Omega
The tick marks at the top of the spectra give the identities of the lines – in other words,
what element/compound is responsible for the line. The letter is the chemical symbol of
the element or compound; a Roman numeral I or a Greek letter indicates the neutral
(not ionized) version of the element/compound. The Roman numeral II indicates a
positively charged ion (1+) of that element/compound. The Roman numeral IV
indicates a triply positively charged (3+) ion of that element/compound. Finally, the
four digit Arabic numeral represents the wavelength in angstroms of that particular
line.
Note: 1 angstrom = 0.1 nm