Flame Emission Spectroscopy

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Flame Emission Spectroscopy

The light emitted from glowing gases is characteristic of the elements in that gas, and the brightness of each band of light is directly proportional to the quantity of that element in the glowing gas. While you are probably familiar with sodium vapor lamps and with ordinary fluorescent tubes for lighting. While the light from overhead fluorescent tubes seems white to our eyes, that from the sodium vapor in many parking lots and along highways is distinctly yellow-orange and objects within range have a rather eerie cast about them. This lifeless color from sodium lamps is due to the fact that they are essentially monochromatic - only one wavelength of light is given off (actually two very closely associate wavelengths; see below).

Just as sodium emits these two very close wavelengths of light, all other elements can emit one to many distinctive wavelengths, many of which are not visible to the human eye as they are in the infra-red or ultraviolet. In the studies on this page, we will only consider a few elements that give emissions within the visible portion of the photoelectric spectrum.

How these various spectral bands form is a matter of extreme importance because they were among the first experimental proofs of quantum mechanics done by Erwin Schrödinger, who explained their highly discrete, narrowness somewhat as follows. First, because a light wave's frequency is proportional to its energy, he saw the different colored bands as representing discrete energy levels with nothing between them - and that 'nothing-between-them' is the basis of the notion of quantum mechanics - sort of like integers have no fractions between them.

But how does all this come about? When an element is vaporized and energized (heated), one or two of its least tightly bound electrons are evaporated off "into outer space" as far as the small nucleus was concerned. However, whenever another passing free electron came close enough, it would fall back towards that nucleus. But it would fall back like something falling down stairs as momentary stops were made from step to step - with no halfway points. Thus whenever that electron fell from - say - step 5 to step 4, it would release the characteristic kinetic energy of that particular drop. If it then fell from step 4 to step 3, it would emit another discrete amount of light energy. The interesting thing is that the energy levels between upper steps is higher than those between lower steps. Schrödinger elegantly demonstrated this by merely having his onlookers observe the flames of three salts: LiCl, NaCl and KCl. (Chlorine emits outside of the visible range, so only the spectra of the metals are seen.) Lithium is the smallest atom and emits bright red light; sodium emits yellow light; and potassium emits lavender light. (It is useful to think of the evaporated electrons as being those participating in valence.)

We spoke of Schrödinger's demonstration as being 'elegant.' Elegance in science is doing something very simply with obvious results that pertain to some really great principle. Certainly, electron orbital theory and quantum mechanics qualify for greatness. As for simplicity, the methods needed are very simple indeed. Actually one might conceive of a really simple example - place a glass rod in a faint blue flame and see the yellow fire come from the rod. That is the sodium flame. Look at it with a spectroscope, and you will see that the yellow light consists of two very bright and narrow bands. But to add other elements, we need to use other set-ups. Two of them are given here: "Lump on a Burner Method," next, and "Loop in a Fire Method".

"Lump on a Burner Method"

1. How to make a lump of a salt: you make a pile of salt and add a drop of water as shown to the right.
2. Finally you have some lumps made of different salts:
3. You will need a "Meeker" Burner. Notice the grid a the top opening of the burner, and you MUST have a gas that gives an almost invisible flame as shown here.

Here is a spectrum emitted from a glowing wire such as in a light BULB - incandescent emission.

If you put your lumps of salts on different burners, you will see variously colored flames - emission flames.

You are now going to use your spectroscope to see really what colors are being emitted and mixed to form the color your eyes see. But first you must focus your spectroscope. Do so by looking at any FLUORESCENT light. Usually these tube-shaped lamps have mercury vapor inside. As the mercury gas is energized, it emits characteristic wavelengths of light which then hits the white powder on the inside of the glass tubes. This causes the powder to fluoresce yielding an even wider number of colors, the mix of which appears to our eyes as white. You can see all this happening with your spectroscope. You see the bright mercury emission lines in the red, green and blue. The dimmer broad yellows are fluorescence. Try to make your emission lines as narrow as possible. Once you have focused your spectroscope, continue to the next paragraph.

Were you to use a spectroscope to look at the emission spectrum of sodium, you would see the bright yellow emission lines. (If you have a high quality spectroscope that is properly focused, you might be able to see the TWO closely spaced emission lines. These are due to the slightly different energies of the paired electrons in the "evaporated" orbital. They are commonly called the sodium "D-lines."


If you take your focused spectroscope outside and look at the sky, you will see the solar emission spectrum. Notice that superimposed on the incandescent spectrum are many narrow bands. These are absorption bands. Can you find the ones for sodium? In fact every known element has its "negative" absorption lines on the sun. Long ago, several solar bands could not be paired up with any known element on earth, so they figured they represented a new element, and so named it after the sun - Helium after Helios (sun). Later small amounts of He were found in natural gas - especially those coming from a few wells in Siberia and near Helium Drive in Amarillo, TX.

"Metal Loop in the Fire Method"

If you do not possess "Meeker" type burners, you may proceed with ordinary Bunsen burners or propane plumbers torches. However, you will have to find some way to support your "elements" in the flame. This poses the problem that your support will also glow and emit light. You should try to find a support that does not produce light visible to the eye, and has other properties amenable to doing the job. It is suggested that the students initially try various supports in the shapes of metal loops (), so as to acquaint the students with a little of the creativity involved in the invention of tools to do a job. Easily obtainable wires from hardware stores are of steel, copper and aluminum. (You might also want to try tungsten, which can be most easily obtained inside of large wattage incandescent light bulbs.) Use 16 gauge or narrower wire to prevent so much heat travelling up the wires as to burn the fingers holding them at the other end six inches or more away. The small loop is most easily made with the needle-nose electrician's pliers in your toolbox.

The first action the students should take is to burn off any coatings that may be on the wires. As many flames are hot enough to melt some of the types of wire (hint: an undesirable trait), one should not roast the wire in the flame indefinitely. For an example observations of one type of wire: A copper loop is placed in the flame. Initially an orange flame is made, but this dissipates into near invisibility. Suddenly greenish sparks fly and shortly thereafter the wire wilts and melts. It is the portion of time of 'near invisibility' that is desired for future tests. After the students try out wires of different metals, they will find that undesirable traits are: (1) the wire itself makes too much emission flame of its own, and (2) the heat melts the wire to easily. Thus the students select the best choice: a wire that has little emission of its own, and holds up fairly well in the flame.

Stick small "flags" of double-overed masking tape onto the handles of each loop. Upon these flags write the names of the principle elements being tested (Na, K, etc.) The reason for doing this is because once a loop has been used - say - for sodium, it will always glow the sodium color intensely. Thus a loop cannot be used for several different salts.

Dip the labelled loop into its appropriate salt solution. Hold the loop high over the flame to allow the water to evaporate rapidly without boiling and "spattering." Once dry, move the loop into the flame. Almost immediately, the salt's emission color should be seen. At this point, look at the colored flame through the spectroscope. (It helps to have the room darkened.) Shown is an example of a copper loop coated with a small amount of NaCl along with a control or unused loop. (Yes, these can be used with Meeker burners also! They just don't burn nearly as brightly as the lumps mentioned in the top portion of this page. Besides it takes two people to manipulate the loops and spectroscope simultaneously.)

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