Spectra Lab
C.O. – Students will explore the spectrum given off by elements due to e- changing energy levels.
L.O. Students will work in groups identifying and exploring the spectra of different elements.
When atoms ABSORB energy, electrons move into higher energy levels. These electrons lose that energy by emitting
light when they drop back to their original energy level.
Key vocabulary (chapter 5.3)
Atomic emission spectra –
Photons –
Louis de Broglie –
Clinton Davisson –
Materials –
Plastic spectrometer
Wire loop/wood stirrer
gas tubes (Ne, He, H2O etc…)
Bunsen burner
ionic compounds
camera
protective eyewear (required)
Background –
Every element has its own “fingerprint”. Since elements have electrons that can absorb and emit energy in certain
frequencies (color) those colors can be recognized in the electromagnetic spectrum. Humans can only see the visible
light spectrum of wavelengths which includes red (long), orange, yellow, green, blue, indigo, and violet (short). Many
animals and insects can see beyond those wavelengths and if you have the right tool (like infrared binoculars) you also
can see below or above the visible light spectrum. Our sun gives off a “full spectrum” of light, which means it emits all
frequencies. This is how we see a rainbow on days that have moisture in the atmosphere and the sun shining at the
same time. It requires a prism to separate the wavelengths and the varying sizes of water droplets provide the prism for
us.
Each element gives off a very specific color and spectrum of light, because of this, we can even look into space and tell
when other stars are moving towards or away from us (the Doppler effect). When we burn ionic compounds, the energy
absorbed by the atom is given off as light and specific colors as the electrons return to their original energy level.
Rutherford’s model didn’t account for this ability of electrons to absorb and release energy so it was the next step as
Neils Bohr added this phenomenon to the atomic theory. This is what we will test today. First you will look at a few
examples with a plastic spectrometer (YOU MUST BE CAREFUL WITH THIS TOOL!), then we will burn some ionic
compounds in a flame and record the results.
*TECHNOLOGY* - have at least one person at your table try to take pictures THROUGH the plastic spectrometer while
looking at the flame so you can have reference and a copy of that element or gas’s specific spectra “fingerprint”. When
using the spectrometer it should be oriented so that the numbers (4-7) are RIGHT SIDE UP and in the bottom left hand
corner. This should be obvious.
Example 1:The sun (full spectrum) DO NOT LOOK AT THE SUN. Point your spectrometer out the window. Without looking
directly at the sun, you should notice a spectra show up. Take a picture for reference, then using the scale below, color
the corresponding colors above the scale exactly as you see them. Some colors may blend to the next, be as accurate as
possible.
Example 2: Fluorescent bulb. Mercury (Hg) vapor at low pressure is placed inside a glass housing with phosphorous (P),
the white powdery stuff, lining the inside. As the Mercury vapor has electricity passed through it, it emits ultraviolet
light. The UV light reacts with the phosphorus particles, causing them to glow, producing light. Take a picture for
reference, then copy exactly what you see on the scale below.
SUN
FLUORESCENT BULB
Example 3/4/5/6: Using one of the vapor bulbs, turn off any other light source that you can. Using your phone, try to
take a picture of the element and its spectra. Use this as a reference as you are coloring. Be as accurate as possible on
the graph below.
Ex 3: ____
Ex 4: ____
Ex 5: ____
Ex 6: ____
Separating Ionic Compounds
Ions are made of (+) metal ions (cations) bonded with (-) non-metal ions (anions). Elements from group 1 and 2 on the
periodic table readily LOSE electrons and become positive to get to its closest full energy level, while group 16 and 17
elements readily GAIN electrons and become negative to fill up its outer energy level. After becoming ions with a charge,
they (cations and anions) stick together in compounds like NaCl (table salt) and other compounds. Group 1 elements
readily lose 1 e-, while group 2 elements readily lose 2 e-. Group 17 elements readily gain 1 electron. Today you will be
burning these compounds in liquid form and viewing the spectra of the METAL ions. Ionic compounds form with an
EQUAL number of positive and negative charges to equal an overall charge of “0”.
Procedure: You will need the ionic compounds, and protective eyewear (REQUIRED AT ALL TIMES), working Bunsen
burner and a metal loop or wooden stir stick. Fill in the data table below. If using the metal loops dip and burn off the
excess liquid in solution A (HCl – Hydrochloric Acid) to clean it and prepare for the next ionic solution. DO NOT MIX
SOLUTIONS. Be as accurate as possible. You may need to burn some of the compounds multiple times.
Solution
Chemical Formula
B
KCl
C
CuCl2
D
Na2(CO3)
E
Li2(CO3)
F
SrCl2
Color
Spectra colors
Conclusion Questions
1. Which metal ion were you observing the spectra for in solution B?
_________________. What is its Atomic # ____, Atomic mass ____ ,# of p+
____, # of neutrons ____, # of e- _____. If it readily loses 1 electron to
become an ion, write its isotope notation INCLUDING its ionic charge
___________. What would you assume is the charge of “Cl” in KCl? _____,
why? _______________________________________________________
2. Which metal ion were you observing the spectra for in solution C?
_________________. What is its Atomic # ____, Atomic mass ____, # of p+
____, # of neutrons ____, # of e- _____. If it loses 2 electrons, write its
isotope notation INCLUDING its ionic charge:
3. Which metal ion were you observing the spectra for in solution D?
_________________. What is its Atomic # ____, Atomic mass ____ , # of p+
____ # of neutrons ____, # of e- _____. If it loses 1 electron , write its isotope
notation INCLUDING its ionic charge:
4. Given what you now know about Na (from question 10) what would you
guess the overall charge of (CO3) is? _______________
5. Which metal ion were you observing the spectra for in solution E?
__________________. What is its Atomic # ____, Atomic mass ____ , # of p+
____ # of neutrons ____, # of e- _____, If it loses 1 electron , write its isotope
notation INCLUDING its ionic charge:
6. Which metal ion were you observing the spectra for in solution F?
__________________. What is its Atomic # ____, Atomic mass ____ , # of p+
____ # of neutrons ____, # of e- _____. If Chlorine gains 1 electron to
become negative, what overall charge do you think Strontium has? ______
Write strontium’s isotope notation INCLUDING its overall charge:
7. Write Sodium’s e- configuration AFTER it has lost its electron.
8. Write Chloride’s e- configuration AFTER it has gained its electron.
9. Write strontium’s e- configuration AFTER it has lost its electrons.