Sample Emission Experiment: Line Spectra of Elements

This sample emission experiment is courtesy of MeasureNet, an Ocean Optics OEM partner that produces data collection and analysis networks for high-quality electronic measurements in chemistry laboratories. This experiment will become part of a lab manual being prepared for the University of Cincinnati. Though the experiments described here mention MeasureNet products, the same setups can be accomplished with Ocean Optics spectrometers and accessories.

The bright glow of neon in advertising signs is a striking example of a common phenomenon: many elements emit colored light when their atoms are excited (provided with excess energy) by an external source, such as an electric current or heat. When we look at a neon sign we perceive just one color, bright red-orange. However, the light emitted by excited neon atoms is actually polychromatic; that is, it consists of many colors.

When polychromatic light reflects from or passes through a diffraction grating it is dispersed into many rays, each of a single wavelength or color. The array of colors is called a spectrum (plural, spectra). When light emitted by excited atoms passes into a spectrometer through a narrow slit, a series of discrete lines (lines, because that is the shape of the slit) of color is produced. Hence, the spectrum is called a line spectrum; it is also called an atomic emission spectrum, since it is produced by emission of light from excited atoms. Each element has a unique atomic emission spectrum with different numbers and colors of lines. For example, neon's atomic emission spectrum consists of dozens of closely spaced lines, many of which are red or yellow. In contrast only a few widely spaced lines are seen in hydrogen's atomic emission spectrum.

The colors seen in fireworks are also due to atomic emission spectra of various elements. For example, a brilliant yellow light is produced by Na and red by Sr. Many elements are added as their salts, such as SrCl₂. The Sr²⁺ cations (other cations in other salts) are reduced to neutral atoms through chemical reactions in the flame. The atoms are excited and emit their characteristic colors. This behavior can be put to practical use: the light emitted can be used to signal the presence of a particular element. Given the elements that may be present in an unknown mixture, you could observe the light emission from each element separately and compare it with the light emission from the unknown mixture. For example, if Ne and Ar are both present in a sample, the emission spectrum will contain lines from both elements; if only Ne is present, then only lines from Ne will be seen, and so on.

Two ways of exciting atoms are employed in this experiment. One is to introduce a solution containing a salt of the metal ion (such as Na⁺) into a burner flame. The cations are reduced to neutral atoms (such as Na); the atoms are excited in the flame and emit their characteristic colors. Among the ions readily identified by the characteristic colors of the emission spectra of their atoms in flames are Li⁺, Na⁺, K⁺, Cs⁺, Ca²⁺, and Sr²⁺.

The second method requires a sample of an element (not a salt) sealed at low pressure in a glass tube equipped with electrodes. Excited atoms are produced when an electric discharge is passed through the tube. (That's how a neon sign works.) Only elements that are gases, such as Ne or Ar, or are fairly easily volatilized, such as Hg or Na, can be excited in this way.

One way to observe emission spectra is simply to use the naked eye. In addition to this, still employing our eyes for light detection, we will use a simple spectroscope in which emitted light is dispersed by a grating. Finally, the emitted light can be directed into a spectrometer. In this case a plot of signal intensity as a function of wavelength is obtained. Wavelengths for which high intensity is seen correspond to the wavelengths emitted by the sample.

In this experiment we will use atomic emission spectra in two ways. First, since the wavelengths of emitted light are related to the electronic arrangement inside the atom, we will make calculations that relate the spectrum of the H atom to its internal structure. Second, we will use emission spectra to identify which elements are present in unknown samples.

Relating Hydrogen's Spectrum and Electronic Structure

The information in this section is provided as background for your analysis of the emitted wavelengths in the H atom spectrum. A light wave is characterized by its wavelength, l , and frequency, n . Its velocity, c, is the product of these.

c = l n (8-1)

The energy, E, of a photon is directly proportional to its frequency

E = hn (8-2)

with the proportionality constant, h, Planck's constant, equal to 6.626 ´ 10^-34 J× s. The relationship between wavelength and energy is obtained by combining Equations 8-1 and 8-2:

(8-3)

An atomic emission spectrum contains light of only certain specific wavelengths. This means that the excited atoms only emit light of certain specific energies. In contrast, white light has all energies of visible light.

The energy of the electron in the hydrogen atom is quantized; only certain values of the energy are possible. These energy values are given by

(8-4)

R_H is the Rydberg constant (2.18 ´ 10^-18 J), and n is called the principal quantum number.

The energy emitted by an excited atom is equal to the difference between two of these possible levels of energy:

(8-5)

where E_II > E_I.

For a hydrogen atom, the wavelength of light emitted by an excited atom is related to the principal quantum numbers of the two energy levels between which a transition occurs by

(8-6)

where n_II > n_I.

As an example of the use of Equation 8-6, consider an electron in hydrogen's fifth energy level (n_II = 5) losing energy in a transition to hydrogen's first energy level (n_I = 1):

(8-7)

The calculated value of l can be compared to the experimental value to test whether n_I and n_II have been correctly chosen. The energy difference can then be determined by using the value of l in equation 8-5.

(8-8)

The energy difference between hydrogen's 5^th and 1^st energy levels is 2.091 ´ 10^-18 J; when this electronic transition occurs, the emitted photon has wavelength = 9.50 ´ 10^-8 m, or 95.0 nm.

Procedure

A. Observations Using Spectroscope and Spectrometer

• Gas discharge tubes (H, He, Ne)
• White light source
• Power source
• Hand-held spectroscope
• MeasureNet station
• Network spectrometer

Qualitative Observations of Spectra

1. Before proceeding to the observation of line spectra from excited atoms, it is instructive to examine the white light emitted by an incandescent filament. Obtain a hand-held spectroscope, direct it at the small white light bulb provided, and observe the spectrum of white light. Record your observations. What color is on the right? What color is on the left? What colors, in what order, lie in between? Do some colors appear to be more intense than others? Knowing that red light has a wavelength of about 700 nm and violet of about 400 nm, does the wavelength in the spectroscope increase from right to left or from left to right?

2. A board containing gas discharge tubes will be set up on the lab bench. Turn on the power to this board. Use the hand-held spectroscope to view the spectra of the elements in the first three gas discharge tubes: H, He, and Ne. Record your observations about each spectrum. How many lines of what colors are present?

Quantitative Measurement of the Hydrogen Spectrum Using the Network Spectrometer

3. Rather than having each student record the H atom spectrum individually, your teaching assistant will choose someone to help measure the spectrum and print out copies for everyone on your network. Two spectra will be recorded. The first will be taken with the most intense line in the spectrum still on scale. The second will be taken with the light intensity adjusted to be as high as possible. Some features will be off scale in the second spectrum, but it will be possible to identify additional, weaker H atom lines. Your TA will also help you to determine the positions of the lines seen in the spectrum by examining the raw data at the PC of your network. Carefully label the lines in your spectra with the wavelengths determined this way. It may be possible to observe minor lines in the second spectrum which are due to traces of elements other than hydrogen in the tube. You will identify and ignore any such lines in your analysis (see below).

B. Flame Tests

• Meker burner
• Wire coil
• Watch glass
• 50-mL beaker
• MeasureNet Station equipped with spectrometer

• 0.5 M Ca(NO₃)₂(aq); calcium nitrate
• 0.1 M KNO₃(aq); potassium nitrate
• 0.1 M NaNO₃(aq); sodium nitrate
• Conc. HCl solution
• 0.5 M CsCl(aq); cesium chloride
• 0.5 M LiNO₃(aq); lithium nitrate
• 0.5 M Sr(NO₃)₂(aq); strontium nitrate
• Unknown solutions (1, 2, or 3 cations)

You will work with all the students on your network to obtain emission spectra for all of the atoms to be studied in this experiment. Each pair will be assigned one known solution. The spectra obtained will serve as a set of standards for you to use as you identify the constituents present in unknown solutions you will be assigned. (If two student pairs record the spectrum of the same standard solution, the better result ("cleaner" spectrum) will be used as the standard.) All of the students on the network should gather in the vicinity of the spectrometer and record their observations during the collection of the standard spectra.

1. The instructor will set up the spectrometer to measure the light emitted by samples introduced into a flame. Unlike the procedure for recording the emission spectrum of fluorescent lights, the emission of light from the flame is not steady. It will persist only for a fairly brief interval and will vary in intensity. You will therefore need to work together with your partner to coordinate the introduction of the sample into the flame with the collection of the spectral data. After a little practice it should become fairly easy to have the spectrometer ready (zeroed) and press the SAMPLE button at the appropriate moment. (Your instructor will demonstrate the proper technique.)

2. Set up your workstation for normal emission spectroscopy.

3. Clean the metal coil by dipping it in concentrated HCl and heating it in the burner flame. (If you are following someone else on the spectrometer, it may also be helpful to first carry out steps 4-6 below with distilled water as the "sample" in order to reduce any contamination remaining from the preceding sample.)

4. Place a few drops of your assigned standard solution on a clean watch glass, and position the watch glass with the liquid as close as possible to the air inlets at the base of the burner.

5. Heat the metal coil in the burner flame, then plunge it into the liquid on the watch glass. Some liquid will vaporize and some of the vapor will be carried into the flame through the air inlets. Observe the characteristic flame color of your sample and also any pertinent features of the color. (Does the color appear immediately when the sample is introduced into the flame? Does it last for a long time or short time? Is it a light or dark shade?) Record these observations in your notebook. You should observe the flame more than once. The yellow color produced by Na emission will appear to some extent in all flames. This is because Na⁺ impurities are introduced from glass storage bottles and are virtually impossible to avoid.

6. Record the emission spectrum. (Be sure you have already zeroed the spectrometer.) This requires two pairs of hands: one student should introduce the sample into the burner with the heated coil, and the other should press the SAMPLE button at the appropriate moment.

7. Return to your workstation to download and plot the spectrum. If the printout of the spectrum is acceptable, your TA will send you back to your workstation to print out enough copies for the entire group. Each time you receive a printed spectrum from anyone on your network, immediately label it with the identity of the solution that produced it.

8. Whether you are personally collecting the spectrum or not, you should remain in the vicinity of the spectrometer to record your visual observations of the flame test for each known solution.

9. Obtain your unknown samples from your TA. Immediately record the identification numbers of the unknowns. Record the emission spectrum for each unknown.

Results

A. Observations Using a Spectroscope and Spectrometer

1. Qualitatively describe each spectrum you observed in this section, addressing the points suggested in the Procedure section.

2. Present the two computer plots of the H atom emission, making sure each has an appropriate title. You should have labeled each line in the spectrum with the wavelength determined with the aid of the TA. For any lines in the visible region of the spectrum, identify each by its color.

3. Use Equation 8-6 to identify the energy levels, n_I and n_II, between which the transition occurs for the two strongest emission lines. This equation cannot be solved in a straightforward algebraic fashion; instead you must try possible solutions. For each measured value of l , pick values for n_I and n_II and evaluate the right-hand-side of Equation 8-6. If your value for this term and your value for l agree (within a small experimental error), these values of n correspond to the energy levels between which a transition occurred. If not, try another set of n values.

In deciding what n values to try, keep in mind the following:

(a) All the n values will be fairly small. As n increases, energy levels get closer together; therefore, large consecutive values of n give small values for D E = E_II - E_I and large values for l . The lines produced by these transitions are of lesser energy (longer wavelength) than visible light. If n_I and n_II differ widely, D E is large and l small. The lines for these transitions are of greater energy (shorter wavelength) than visible light.

(b) The difference between n_I and n_II will be less for lines toward the red (low energy) end of the spectrum and greater for lines toward the violet (high energy) end.

(c) Begin with the simplest combinations: n_I = 1 and n_II = 2; n_I = 1 and n_II = 3; etc. If none of these values matches a value of l from the experimental spectrum, try values of n_I = 2 and n_II = 3; n_I = 2 and n_II = 4; etc. Show an example of your determination of n_I and n_II for one of the lines.

4. You should have found in the previous step that the two strongest lines in the H atom spectrum result from transitions having the same value of n_I, the quantum number for the lower energy level, but different values of n_II, the quantum number for the higher energy level. Use equation 8-5 to predict the wavelengths expected for other transitions also sharing this lower energy level, but starting from still higher levels (higher values of n_II). Examine your list of wavelengths to see whether any of these lines were observed in the experimental spectra.

5. There may be emission lines in your spectra that do not come from H atoms and therefore cannot be associated with quantum numbers as in step 3 above. A likely contaminant is oxygen, since the glass tube consists mainly of SiO₂. Examine this spectrum carefully to see whether any of the "extraneous" lines in the experimental H atom spectrum may be assigned to traces of oxygen present and therefore excluded from the analysis. (Don't forget that this spectrum also could have "extraneous" lines, that is, not due to oxygen. Just use it to try to identify the extraneous lines in the H atom spectrum.)

6. Use your experimentally determined wavelengths to calculate the energy emitted in each transition, Equation 8-3. Show a sample calculation. Present your results for all of the lines you observed in a table, using the following headings: l (nm); Source; Color; n_I; n_II; E (J). Source refers to the atom emitting the light, such as H or O. Clearly no n_I or n_II values can be included for any emission not arising from H atoms.

7. Construct an energy level diagram for H. Use graph paper (at least 10 squares/inch) or a computer plot so that you will have a reasonably accurate scale for the energy levels. Calculate the energy associated with each energy level using Equation 8-4. Along the energy scale axis (vertical), draw a horizontal line for each energy level involved in any of the transitions observed in your experimental hydrogen spectrum. (Note: since the energy values are all negative, the energy scale in your diagram should have a value of zero at the top.) Label each level with its quantum number. Draw arrows between energy levels indicating the transitions that produce the lines you observed in the spectrum. Label each arrow with the wavelength of the light that is produced by the transition.

1. On the plots of the standard spectra you and your classmates obtained using known samples, write the color (based on wavelength) of each line next to the corresponding signal. Discuss the observations you made concerning the flames from the standard samples.

2. Identify the cations present in each of your unknown samples by comparing its spectrum with the standard spectra collected by the entire group. As you do so, you will have to think just as an analytical chemist using a sophisticated, expensive, atomic emission apparatus would, keeping a number of potential problems in mind:

a. You must always consider the possibility of contamination (in the standards as well as in the unknowns). For example, there will usually be some signal from Na, so it is only considered to be genuinely present if its signal is quite intense. Furthermore, it is likely that the substances present in one sample could appear as contaminants in the next measurement. It may be necessary to repeat a measurement more than once, carefully cleaning the watch glass and the wire coil, to reduce contamination from this source. It may also be helpful to use distilled water as a "sample" between real samples in order to remove some of the possible contamination between samples.

b. Be sure to note that the emission lines in the spectra of some metals are relatively broad, while other metals have very narrow lines. The total amount of light emitted in a broad line is spread over a greater wavelength range, so the signal in the spectrum may not be very large at the maximum in the line. When a similar amount of light is confined to a narrow line, the signal will be much larger at the maximum in the line. Your eye responds to the total amount of light emitted, so both cases can produce bright flames. If an unknown sample contains a mixture of metals giving both types of emission, the lines may differ greatly in intensity. Don’t overlook the broader, apparently less intense emission in comparison with the narrow lines. Carefully compare your unknown plot with the plots for known solutions.

c. The spectrometer can "see" colors your eyes can’t, so even flames which show very little color may have intense emission outside the visible region (about 400 - 700 nm).

3. Report your results in summary form, including your reasoning in arriving at your conclusions as to which cations are present in each unknown. For each unknown, give the unknown number, list the cations present, and give your reasoning for your conclusions.

As a portion of your discussion, include answers to the following questions:

1. For the gas discharge tubes, what simple qualitative (not quantitative) correlation do you observe between the number of electrons in an atom and the number of lines in its spectrum?

2. Use Equation 8-6 to show that a transition between two energy levels with large n values results in the emission of light of high wavelength. (Hint: Choose similar values of 20 or higher for n_I and n_II to calculate the wavelength of emitted light.)

3. Use Equation 8-6 to show that a transition between energy levels with small n_I and n_II values results in the emission of light of low wavelength.

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