Flame Tests of Selected Ions
When the electrons in the atom or ion are excited, for example with electricity or by being heated in a flame, the additional energy pushes the electrons to higher energy orbitals. These “excited” state electrons, however, are not stable, and they will eventually fall back in toward the nucleus and settle in the ground state (most stable configuration). When the electrons fall back down to a lower energy level and leave the excited state, energy is re-emitted in the form of a little packets of light energy called photons. We see this as colored light and we call it the element’s emission spectrum (literally, “emitted image”). Both the emission and absorption of energy are quantized – only certain energy levels are allowed.
The particular color or colors that we see depend on the wavelength of light emitted and the wavelength depends on the difference in energy between the excited and the ground states for the electron. In fact, the wavelength of the photon emitted matches exactly the energy difference between the excited energy level and the energy level to which the electron returns. The greater the energy difference between the excited and ground states, the shorter the wavelength of light emitted. The fact that only certain colors appear in an element's atomic emission spectrum means that only certain wavelengths of light are emitted. Each of these wavelengths is related to energy by the formula:
Ephoton = hc
λ
where E is the energy of the photon
h is Planck's constant = 6.63 x 10–34 J s
c is the speed of light = 3.0 x 108 m s–1
λ (“lambda”) is the wavelength of light in meters
Photons with wavelengths between 390 and 750 nm are visible to us. In this lab, we qualitatively study the colors produced by selected Group 1 and Group 2 metals, along with some transition metals species.
Materials:
CaCl2 , NaCl, KCl, LiCl, Sr(NO3)2, BaCl2, CuCl, CuSO4, wooden splints, distilled water
Safety:
Place wooden splints in a beaker of tap water before throwing them in the trash.
Wear goggles, aprons, and tie hair back.
Wash hands thoroughly with soap and water before leaving the lab.
Procedure:
1. Prepare two beakers of water, one with distilled water and one with tap water. Place eight wooden splints in the beaker with the distilled water.
2. Obtain the vials with salt samples.
3. Light your Bunsen burner.
4. Dip the soaked end of the wooden splint into one of the samples – you only need a tiny amount on the splint. Place it in the flame and observe the color.
5. Allow the splint to burn until the color disappears. Try not to get the salts on the Bunsen burner. Immediately place the burned end of the splint into the beaker with tap water to put it out.
6. Record the color of the flame in the data table below.Repeat steps 4 - 6 for the remaining salts.
7. Clean up: place the used wooden splints in the trash. Tightly screw the caps on the salt samples and put them back in the designated bin. Rinse and dry the beakers you used. Wash your hands.
Data Table
Metal Ion / Symbol w/ Charge / Color of Flame / Wavelength (nm) / Wavelength (m)Lithium
Sodium
Potassium
Calcium
Strontium
Barium
Copper (I)
Copper (II)
Color of Light / Wavelength of Light (nm)
Violet / 410.
Blue / 470.
Blue-green / 490.
Green / 520.
Yellow-green / 565.
Yellow / 580.
Orange / 600.
Red / 650.
Analysis Questions (answer on a separate sheet of paper and attach - Show all work – don’t forget units!)
1. Explain the difference between the excited state and the ground state electron configurations of an atom.
2. Which color of light in the visible spectrum is associated with the highest energy? The lowest energy? How can you determine this knowing only the wavelength?
3. What is the relationship between frequency and wavelength?
4. For Li+, Na+, K+, and Ba2+:
a. calculate the frequency of emitted light in Hertz
b. calculate the energy emitted with one photon in Joules
c. calculate the energy released when one mole of electrons returns to the ground state in
kJ mol–1.
5. Rank the Group 1A elements tested in order of increasing energy emission when 1 mole of electrons returns to the ground state. Explain this pattern in terms of changes in energy levels.
6. An excited electron falls from n = 4 to n = 2.
a. Calculate the energy change in Joules associated with this transition.
b. What is the wavelength in nm of the emitted electromagnetic radiation? What color is it?
c. Calculate the energy change in kJ mol–1.
7. An excited electron falls from n = 3 to n = 1.
a. Calculate the change in energy in Joules associated with this transition.
b. What is the wavelength in nm of the emitted light? How is this different from the electromagnetic radiation emitted in Question 6b? Could you detect this light with your eyes?
c. Calculate the energy change in kJ when one mole of electrons moves from the third principal energy level to the first principal energy level. Compare this to your answer to Question 6c.
d. Using Einstein’s famous equation, E = mc2, where m = mass, determine the mass of a photon with the wavelength calculated in part 7b.
8. A certain line in the spectrum of Hydrogen is associated with the electronic transition from the fifth energy (n = 5) level to the third energy level (n = 3).
a. In the electronic transition described above, does the atom emit or absorb energy? Justify your answer.
b. Calculate the wavelength of the photon associated with this spectral line. Account for the observation that the energy associated with the same electronic transition in the He+ ion is greater than that associated with the corresponding transition in the H atom.
9. When a chloride solution of an element is vaporized in a flame, the color of the flame is purple. What element could be in the solution?
10. Why is it important to use distilled water to pre-soak the wooden splints? Why can’t we just use tap water?