LAB_Spectroscopy(1)

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Feb 20, 2024

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Spectroscopy 1 Atomic Emission v051623_WST Objective: The student will observe the atomic emission spectra of several light sources and apply quantum mechanical concepts to them. An unknown element will be identified by comparing its emission spectrum to known spectra. Light, Matter, and Quantum Levels In the eighteenth century, a Scottish physicist Thomas Melvill reported that when he placed different substances in a flame, they emitted light of different colors. For example, he noted that when table salt was passed through a flame, the resulting emitted light was bright yellow which was a result of sodium (Figure 1). This concept was something that had been known for several centuries before Melvill, but what was important about his discovery is that he found that when the emitted light was passed through a prism and a small slit, there was a band of colors, or a spectrum, that was created that was unique to the specific element in the flame. Figure 2 shows the atomic emission spectrum of sodium. Throughout the nineteenth century, this idea of different elements producing a unique spectrum was utilized as a way of identifying what elements were in substances and was even used to discover new elements. J.J. Balmer, a teacher from Switzerland, made an important discovery in 1885. Balmer was working with hydrogen and its atomic spectrum and he knew that hydrogen emitted visible light at 4 different wavelengths. Through trial-and-error, he discovered that there was a pattern that the emissions followed that could be represented by the Balmer-Rydberg equation, ଵ ஛ = R ൤ ଵ ௡ ೑ మ − ଵ ௡ ೔ మ ൨ [1] where n final = 2 for lines appearing in the visible light region, n initial is a whole number greater than 2, R is the Rydberg constant (1.097 x 10 -2 nm -1 ), and  is the wavelength of the emission line in nanometers. See that the units of R reflect the reciprocal wavelength (1/  ) in Equation 1. Over the next thirty years, the picture of the atom would begin to take shape. Work by J.J. Thompson, E. Rutherford and other scientists showed that the atom was made of positively charged protons and neutral neutrons found within the nucleus, and that the nucleus was surrounded by a negatively charged electron cloud. There were several models that tried to describe atomic structure, but one of the most important models was proposed by Niels Bohr in 1913. Observing the atomic spectrum for hydrogen and the work done by Rydberg and Balmer, Bohr suggested that the hydrogen atom Figure 1 Figure 3 Figure 2

Spectroscopy 2 was made up of a nucleus with an electron “orbiting“ that nucleus. He also indicated that the electron could only occupy specific energy levels and that it would not be found between energy levels. The Bohr model of the atom can be seen in Figure 3. Bohr proposed that the electron in the hydrogen atom would exist in a stable, lowest energy or “ground state” orbital shown on Figure 4 as n = 1. However, when the atom is irradiated by light, the electron absorbs that energy (where the energy of the incoming light matches the difference in the energy level between the orbitals) which results in the electron being excited to a higher energy orbital. The electron would then relax to a lower, more stable energy level. For a more in-depth discussion of atomic spectroscopy and the Bohr model, see Sec. 2.3 and 2.5 of Tro 2 nd , Edition. The energy difference between the energy levels can be determined using the equation:  E = h  = E photon emitted or absorbed [2] Rearranging the speed of light equation (c =  ) to solve for frequency, the speed of light divided by the wavelength can be substituted in equation 2 for the frequency to give the equation: ୦ୡ ஛ [3] In equations 2 and 3, h is Planck’s constant (6.626 x 10 -34 J . s),  stands for the frequency (1/seconds), c is the speed of light (2.998 x 10 8 m/s), and  is the wavelength of the light in meters. Bohr found that using his model he could accurately predict the wavelengths of the lines on the hydrogen emission spectrum that Balmer had observed and described using the Rydberg-Balmer equation. He also determined that if an electron fell from an orbital higher to the level n = 2, the energy emission would be observed in the visible region of the spectrum. The problem with Bohr’s model of the atom and the Rydberg-Balmer equation was that they could not accurately predict the wavelengths of spectral lines for atoms with more than one electron due to the fact that this model and equation do not consider electron-electron repulsions that exist in elements with more than one electron. Though the Bohr model has its fallacies, the idea of different quantized energy states is applicable to all atoms and molecules. Diode Lasers Laser pointers are an example of a device called a semiconductor diode laser. Diodes are solid-state devices in which a junction is constructed between an electron-rich material (N-type), and an electron- deficient material (P-type). This is illustrated in Figure 5. The N-type material can be thought of as a source of electrons, while the P-type material can be viewed as having lots of vacancies or “holes” waiting to be filled by electrons. As in the hydrogen atom, the energy levels of the electrons and holes are quantized. In a diode laser, the energy level of the electrons in the N-type material is higher than the energy level of the holes in the P-type material. The difference in these energies is called the “bandgap”. Dependent on the materials chosen for the particular diode, the bandgap is adjusted by the manufacturer to produce the desired color for the laser. When a voltage is placed across the junction, electrons move from the N-type side of the junction to the P-type side. When the electron “falls” across the bandgap in Figure 4

Spectroscopy 3 this way, a photon is emitted in a manner analogous to that of the electron relaxing between quantum states in the hydrogen atom. Like the emissions from the hydrogen atom, this photon has a fixed energy defined by the bandgap, and therefore a particular wavelength. This happens many times, and the resulting photons are focused into a coherent, monochromatic beam that occurs at a frequency associated with the bandgap of the laser. (There’s actually a bit more that goes into getting all these photons into a coherent beam, but those design considerations aren’t necessary for our objectives here.) Absorbance and the Perception of Color We’ve seen that atoms are capable of absorbing and emitting light at specific wavelengths, but molecules are also capable of absorbing and emitting light. It’s this light absorption by molecules that allows us to sense many of the colors we see on an everyday basis. For example, the reason we perceive grass as being green is due to the molecules of chlorophyll within the grass. As seen in Figure 6, chlorophyll absorbs light in the red and violet regions of the visible electromagnetic spectrum. The primary wavelength of light that is not absorbed is around 500 nm which corresponds with the green region of the spectrum. This particular wavelength of light is reflected back to the eye and we see that the grass is green. The same idea can be applied to transparent colored material such as stained glass. The color that we see arises from the colors that are transmitted through the material and not absorbed. In both cases the absence of the absorbed color in the light that reaches our eye is the reason for the color we perceive. An effective qualitative method of determining the color of a substance that will be observed considers an artist’s color wheel (Figure 7). The color absorbed and the color observed tend to be complementary. For chlorophyll the color absorbed is red and the color observed is green. As you can see from the color wheel these two colors are complementary, across from each other on this simplified diagram. Figure 7. Color wheel relating color with wavelength of light. Figure 6. Absorption spectrum of chlorophyll Figure 5

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Spectroscopy 4 Procedure: All measurements should be recorded using the correct number of significant figures appropriate to the equipment being used. All calculated values should be recorded to the correct number of significant digits. Carefully read Sec. E.3-4 of Tro 2 nd Edition to re-familiarize yourself with significant digits and being consistent with them in mathematical operations. IF THE SPECTROVIS PLUS IS NOT ALREADY ON, TURN IT ON NOW TO ALLOW THE INSTRUMENT TO WARM UP BEFORE USE. If you have not already done so, download the “Spectral” app to your phone or tablet and then pair your phone or tablet to your SpectroVis Plus. Safety:  Wear safety goggles. Wash your hands thoroughly before leaving the lab.  Unknowns in the lab are known to be harmful if swallowed. I. Blackbody Radiators, the Continuum Spectrum, and Absorption Characteristics of Colored Films 1. Turn on the incandescent lamp and examine the emission spectrum using the spectroscope. The filament in the lamp is a blackbody radiator (an object emitting light as a result of its temperature) and will emit what is known as a continuum spectrum. Record your observations. 2. Collect the emission spectrum of the incandescent lamp using the optical probe and the SpectroVis Plus. Save this spectrum as a .csv file. 3. While observing the emission spectrum through the spectroscope, record your observations on what happens when the red, blue, and green filters are placed between the lamp and the entrance slit of the scope. How does each of the colored films alter the appearance of the continuum spectrum emitted by the lamp? 4. Repeat this process only collect the spectra using the SpectroVis Plus. These spectra represent the residual light reaching the instrument after having passed through the colored filter. Wavelengths absorbed by the filter do not reach the instrument. Save these spectra as .csv files. II. Emission Spectra of Diode Lasers NOTE: Do not look directly into the laser pointer!!! 1. Shine the laser pointer on a white piece of paper. 2. Point the optical probe of the SpectroVis Plus spectrometer at the spot on the paper and collect the emission spectrum for the laser. 3. Identify the wavelength of the light emitted by the laser. 4. Repeat this process for the other laser pointer.

Spectroscopy 5 III. Atomic Spectrum of Hydrogen 1. With the power supply unplugged and the power switch in the OFF position, insert the hydrogen discharge tube into the power supply. 2. Plug in and turn on the power supply. 3. Set up the spectroscope by moving the slit toward the center of the discharge tube while being careful not to touch the tube. 4. Record the colors of the emission lines that are visible through the spectroscope. 5. Collect the emission spectrum of the hydrogen lamp using the optical probe and the SpectroVis Plus. 6. Record the precise wavelengths of each of the emission lines that are detected. (You will need these wavelengths later for your calculations.) Were any additional lines visible using the SpectroVis Plus? Hint: There should be two emission lines in the violet part of the spectrum. 7. Turn off the power supply and allow the discharge tube to cool. IV. Identification of Unknown Gas 1. Select one of the discharge tubes for your unknown and place it into the power supply as you did with hydrogen. 2. Turn on the power supply and observe the line spectrum through the spectroscope. Record the number of visible lines, and their colors. 3. Collect the emission spectrum of your unknown gas using the SpectroVis Plus. Record the precise wavelengths of each emission line observed.

Spectroscopy 6 Pre-Laboratory Assignment 1. An electron in a lithium atom moves from the 2p orbital to the 2s orbital with a ΔE of 2.96 x 10 -19 J. When the transition occurs, energy equal to ΔE is released in the form of a photon. What is the wavelength of the light that is emitted? 2. Does the emission line determined question1 fall in the visible region of the electromagnetic spectrum? If so, what color is the light that is emitted? 3. Beta-carotene is an example of a molecule that absorbs strongly in the visible region of the spectrum. Given the absorption spectrum of beta-carotene shown below, what are the significant wavelengths that are transmitted? What color would a solution containing this molecule appear to be?

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Spectroscopy 7 Data Sheet I. Blackbody Radiators, the Continuum Spectrum, and Absorption Characteristics of Colored Films 1. Observations of the emission spectrum of an incandescent bulb: 2. What does this spectrum resemble? Hint: The sun is also a blackbody radiator. 3. What wavelengths of light are present in this continuum spectrum? Which are missing? 4. Observations of the continuum spectrum when red, green, and blue films are placed between the spectroscope and the light source. 5. Import the .csv files for the unfiltered continuum spectrum and the three spectra obtained using the colored films into an Excel spreadsheet (a total of four files) and make line plots with intensity as the y-variable and wavelength in nanometers as the x-variable. Attach these spectra to your data sheets to turn in. A good way to make these plots efficiently and highlight their

Spectroscopy 8 differences is to overlay them on the same set of axes. A correct plot will have both axes labelled, including specifications of the units. Give your spectrum a descriptive title. 6. What wavelengths were absorbed by each of the colored filters? How is the perceived color of the filter related to the colors (wavelengths) that are absorbed? II. Emission Spectra of Diode Lasers 1. How do the diode laser spectra differ from the continuum spectrum emitted by the incandescent lamp? 2. Wavelength emitted by the red laser in nm: _______________________ Wavelength emitted by the green laser in nm: _______________________

Spectroscopy 9 III. Atomic Spectrum of Hydrogen Note: some people’s eyes are not sensitive to the photons of the highest energy purple line, so some may only see 3 visible lines instead of all 4. Be sure a note the units of the wavelength drawn from spectroscope. Wavelength ( ) Line color         IV. Identification of Unknown Gas Unknown ID: _____________ Provide a general sketch of the emission spectrum (below) observed for your unknown; approximate number of lines and colors; relationship of line colors to the color perceived by the eye. Especially note particularly intense lines, as these few can greatly help you fingerprint your unknown.

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Spectroscopy 10 Results 1. Attach your plotted spectra for the incandescent lamp and the colored films. Acceptable plots should have the axes labeled clearly and with units. Each plot should have a title. If desired, your plots can be overlaid on the same set of axes. 2. Calculate the frequencies and energies associated with the light emitted by the red and green lasers. Show clear work at the bottom of this page for the red laser. Red Laser Green Laser ν = ________________ (Hz) ν = ________________ (Hz) E = ________________ (J) E = ________________ (J)

Spectroscopy 11 3. Using the Rydberg equation [1], calculate the energy level from which the electron fell to produce the emission you observed from the hydrogen lamp. Remember: when visible light is produced, the electron falls to the 2 nd energy level. Show at least one sample calculation. Purple line 1 Purple line 2 Blue line Red line

Spectroscopy 12 4. Indicate the electron transition with arrows and associated energy for each of the wavelengths of light observed for hydrogen on the diagram below. Label the color of light observed with each transition arrow.

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Spectroscopy 13 5. Line spectra are like fingerprints. No two elements have the same line spectrum. Which of the following best matches your unknown? Helium Mercury Neon Unknown? _______________________________

Spectroscopy 14 Post-Laboratory 1. The Balmer-Rydberg equation only works perfectly for hydrogen, allowing the precise determination of n f and n i values for any emission observed from this element. (a) Sketch a model of a hydrogen atom and of a helium atom in the boxes below (see Fig. 3). Represent the nucleus with a small circle with a charge (+1 or +2). Represent electron(s) orbits with circles. Add electrons (e). Label the He electrons e 1 and e 2 . hydrogen atom helium atom What force(s) of attraction or repulsion are present in He that are absent in H, the reason the Balmer-Rydberg equation only works perfectly for H. 2. Astronomers use light collected from distance stars and galaxies to discern their elemental compositions. Such observations lead them to conclude that H is the most abundant atom by number and by mass across the whole universe , a remarkable assertion given the limitations in space exploration even today. Think about your observations of line spectra for H and other elements. If a distant galaxy or star contained H, what would you expect to see in visible spectra observed from these objects?

Spectroscopy 15 3. Manufacturers have two chemical approaches to making colored products: 1) Use a mixture of dyes, so that the combination gives the desired color, or 2) use a single dye of exactly the color one wants. Grape Crush soda is one such brightly colored product. The ingredient label for Graph Crush and its observed light absorption spectrum are listed below. (a) According to the ingredient label, which of the two approaches appears to be the way manufacturers make the purple color in Grape Crush. (b) Over what visible wavelengths might Blue 1 dye be expected to show high absorbance? (c) Consider your answer to (b) and your red dye spectrum. Is the Grape Crush spectrum consistent with the dye(s) used to make its purple color? Defend your answer.

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Spectroscopy 16 4. Red laser pointers were the first to be developed and represented the simplest application of the direct diode process described above. The development of diode lasers in other colors presented a more challenging engineering problem, and many of the diode lasers in other colors that are currently available utilize a more complex mechanism to produce the desired color. If you wanted to make a blue diode laser with a wavelength of 473 nm that utilized a direct diode process, what adjustment to the bandgap of the diode would be required relative to that of a red laser? 5. Take a close look at the energy you calculated for the photons emitted by the green laser pointer. Is this a large or a small amount of energy? Even low power laser pointers operate at an output of 1 – 5 mW, and can cause permanent damage to the human retina if exposed for too long. If we assume the green laser you used in this lab has a power output of 5 mW, what would the total energy deposited onto your retina be if you shined it into your eye for 1 second? How many photons would your retina receive in this one second?

LAB_Spectroscopy(1)

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