EXPERIMENT 6 INTRODUCTION TO SPECTROSCOPY

EXPERIMENT 6 INTRODUCTION TO SPECTROSCOPY INTRODUCTION Much of what we know about the structures of atoms and molecules has been learned through experiments in which photons (electromagnetic radiation visible light, microwaves, ultraviolet or infrared radiation, radio waves, etc.) are emitted or absorbed by the atoms or molecules. The energy of a photon is related to its frequency,, and wavelength,, according to E photon c hv h (6-1) where h is Planck s constant, and c is the speed of light. The energy of an emitted or absorbed photon corresponds to the change in energy the atom or molecule experiences. E photon E E E (6-2) Whether photons are absorbed or emitted is correlated with the type of energy change the atom or molecule is undergoing. Thus, for example, a molecule can be raised to an excited electronic state by absorbing a visible or ultraviolet photon. A molecule already in an excited electronic state can return to the unexcited, or ground state by emitting a visible or ultraviolet photon. The energies of photons in this portion of the electromagnetic spectrum correspond to the differences between the ground and excited electronic states. Careful analysis of the details of the radiation absorbed or emitted as a function of wavelength (the absorption or emission spectrum), coupled with the formulation of physical models to interpret and explain them, has provided a wealth of detailed information about atoms and molecules. In addition to the structural information that can be gained, studies involving the absorption and emission of electromagnetic radiation have proven to be extremely useful in other practical ways. For example, the specific wavelengths of visible and ultraviolet radiation emitted by atoms and ions in a flame or in an electrical discharge can provide an unambiguous means of identification. In fact, a number of elements were first discovered in this way, when previously unknown emissions were observed. Spectroscopic measurements are now routinely employed in the analysis of chemical samples. While measurement of the wavelengths emitted or absorbed can provide a convenient means for qualitative analysis of samples (i.e., what is present), measurement of how much light is emitted or absorbed can be used for quantitative analysis (i.e., how much of a substance is present). Carrying out such measurements is referred to as spectroscopy. Although quantitative experiments can be performed using various regions of the electromagnetic spectrum, one of the most useful is the visible portion, sometimes in combination with the ultraviolet region. In this experiment we will be working with visible electromagnetic radiation, ordinary light. final initial

In order to study the emission and absorption of visible light, we will make use of an instrument known as a spectrometer. A simplified drawing of the main components of our spectrometer is shown in FIGURE 6-1. A spectrometer is an instrument that accomplishes two main tasks. First, it disperses, or spreads out, the light entering it into all of the wavelengths or colors present. This can be done with either a prism or a diffraction grating. Our spectrometer uses a grating. Second, it provides a signal proportional to the intensity of the light of each wavelength. It does this by directing the dispersed light onto a detector, which provides the electrical signal. In our spectrometer the detector consists of an array of 2048 tiny diodes arranged in a straight line and positioned so that the dispersed light is spread from one end of the array to the other. Therefore we actually have 2048 tiny, individual detectors, and each one has light of a slightly different wavelength, or color, falling on it. mirrors Whatever type of experiment we are carrying out, the signal from the spectrometer is always just a set of values, one from each of the tiny diodes, indicating the intensity of the light reaching them. (In actual practice we reduce the amount of data to be handled by averaging the signals from the diodes in adjacent pairs, thereby obtaining 1024 values from the original 2048.) The diodes are more sensitive to red light than blue light, so signals in the blue end of the spectrum will be somewhat reduced compared with the red end. However, the signal for any wavelength is proportional to the intensity of the light of that wavelength. Thus, for example, no matter what the color, the signal will double if the intensity of the light of that color is doubled. light in When we use the spectrometer to measure an emission spectrum (a plot of light intensity vs. wavelength), we simply direct the light emitted by the sample (gas in a discharge tube, flame, etc.) into the spectrometer. The set of values we get from the spectrometer can then be examined to see what wavelengths of light were emitted. In this experiment, we will examine an emission spectrum only from a qualitative emission standpoint, where only the wavelengths, and not the intensities, of the emitted light are important, so the variation of detector sensitivity with wavelength will not affect its usefulness. When we use the spectrometer to measure an absorption spectrum (a plot of absorbance vs. wavelength), the situation is quite different. We use a lamp to supply light of all wavelengths throughout the visible region and use the spectrometer to determine the extent to which light of each wavelength is absorbed. The physical arrangement is shown in FIGURE 6-2. The sample is contained in a cuvet, a small container with clear windows. I 0, is the intensity of the light incident on the sample as a function of wavelength, ; in other words, it is the light intensity coming from the light source at a particular wavelength. I is the intensity of the light remaining after passing through the sample. grating blue red diode array detector Figure 6-1. Simplified Spectrometer Diagram

lamp I 0, s a m p l e I spectrometer Figure 6-2. Measurement of light absorption as a function of wavelength. There are two steps in an absorption experiment: first, light intensity is measured after the light passes through the reference cuvet, which is identical to the sample in every way except that it does not contain the absorbing molecules. If the absorbing species is present as the solute in an aqueous solution, as is the case in this experiment, the reference cuvet is simply filled with water, which does not absorb light in the visible region. At wavelengths where the sample does not absorb, I is equal to I 0,. Next, the reference cuvet is replaced by the sample cuvet, which contains a solution of the absorbing species, and the light intensity is measured once again after it passes through the sample. At the wavelengths where the sample absorbs, I is less than I 0,. At this point, we could, if we wished, make a plot of intensity vs. wavelength for the light after it passes through the reference (I 0, ) and after it passes through the sample (I ). A plot showing both of these spectra on the same set of axes is shown in FIGURE 6-3. However, we are usually more interested in how much light is absorbed by the absorbing species rather than how much light passes through the sample. The absorbance at any particular wavelength, A is related to I 0, and I by A I 0, log (6-3) I The spectrometer calculates A at each wavelength where the intensity measurements were made, and generates an absorption spectrum. An example is given in FIGURE 6-4. (The absorbance values were obtained from the intensity measurements in FIGURE 6-3.) The value of A has been found to depend on three things. First, the nature of the absorbing molecules determines what particular wavelengths of light will be absorbed, and to what extent. Second, the longer the path the light travels through the sample, the greater the fraction of light absorbed. Third, the greater the concentration of absorbing molecules, the greater the fraction of light absorbed. These three factors combine quantitatively in the following manner (Beer s Law): A b C (6-4) where is the molar absorptivity, characteristic of the absorbing molecules, b is the path length of the light through the sample (in other words, the width of the cuvet), and C is the molar concentration of the absorbing species. The most important routine use of absorption measurements is for the determination of the concentration of absorbing molecules. Since, as EQUATION 6-4 shows, the absorbance is directly proportional to the concentration of absorbing molecules, it is the quantity we use to determine the concentration. A series of

solutions of known, varying concentration are prepared, and the absorption spectrum of each solution is obtained. For each spectrum, the absorbance value is determined at a particular wavelength. Generally, it is best to select a wavelength where the absorbance is at a maximum, since this will minimize the relative error of the measurement. A plot of absorbance vs. concentration (a Beer's Law plot) can then be prepared. Because the The concentration of an unknown solution of the same absorbing species can be determined by obtaining the absorption spectrum of the unknown solution, determining the absorbance at the same wavelength used to make the Beer's Law plot, and then using the slope of the Beer's Law plot to calculate the concentration OBJECTIVES to measure the emission spectrum of the fluorescent light fixtures in the laboratory ceiling and correlate the spectrum with the colors observed visually to measure the absorption of visible light by a solution of an absorbing substance (a food dye) and examine how the intensity plots result in the calculated absorption spectrum to prepare a plot of absorbance (at the wavelength of its maximum absorption) vs. concentration for a series of solutions (Beer s Law plot), and determine the molar absorptivity,, of the absorbing species at that wavelength to measure the absorption spectrum for a sports drink and use the Beer s Law plot to determine the concentration of food dye in the sports drink.

Figure 6-3. Absorbance Experiment Raw Data: Reference and Sample Signals Figure 6-4. Absorbance calculated from data in Figure 6-3.

FLOW CHART FOR LAB ACTIVITY Emission Spectrum of Fluorescent Light Instructor will obtain spectrum and print copies for class Observe fluorescent light through handheld spectroscope and correlate colored lines to emission spectrum peaks. Label peaks with corresponding colors and submit to TA. Absorption Spectrum of Food Dye Stock Solution Obtain spectrum Save spectrum (using file code 001) Preparation of Dilution Sample Prepare 10.00 ml of diluted sample Repeat for 2 nd 3 rd, 4 th dilution samples Absorption Spectrum of Dilute Solution Obtain spectrum Save spectrum Absorption Spectrum of Sports Drink Obtain spectrum Save spectrum Print Spectra Open Absorbance Series program at computer Enter station number and select spectrum files Enter wavelength and print plot of absorption spectra

EQUIPMENT NEEDED MeasureNet spectrometer on network volumetric flask, 10 ml, with cap Mohr (graduated) pipet, 5 ml cuvets with caps (2) volumetric pipet, 3mL hand-held spectroscope disposable plastic pipet pipet pump fluorescent light source CHEMICALS NEEDED aqueous stock solutions of food dyes FD&C blue #1 and FD&C red #40 (concentrations on containers) Sports Drinks PROCEDURE A. Emission Spectrum of Laboratory Fluorescent Lights Note: To save time, your instructor may decide to carry out steps A1 through A6 for you and print out the emission spectrum for everyone on the network. If that is the case, you should begin with step A7 after you receive your copy of the spectrum. Setting up the Workstation A1. Press the MAIN MENU button on the workstation. (Make a note of your station number you will need it at the spectrometer.) Select SPECTROSCOPY, then EMISSION(1), and then press DISPLAY. Zeroing the Spectrometer Note: Pay attention to the spectrometer display. The first line shows the selected station and the type of spectrum to be measured. It also shows which operations remain to be performed. For example, if the letters Z and S are displayed, the ZERO scan and the SAMPLE scan are to be made, in that order. If Z, R, and S are displayed, the ZERO, REFERENCE, and SAMPLE scans are to be made, also in that order. After each scan is made, its letter will disappear. If you accidentally try to do something out of order, an error message will appear, and you will have to begin again. The display will also indicate when a scan is being made, when data are being transmitted to the controller, etc. Always wait until it has completed whatever it is doing before pressing a button. A2. At the spectrometer, press the button labeled STATION NUMBER, then your station number, then ENTER. The spectrometer display should indicate your station number, and it should show that your station is set up for emission by displaying STA# EMI-VIS Z S on the first line. A3. Cover the end of the fiber optic cable from the spectrometer with a folded tissue or paper towel to prevent any light from entering, and press the button labeled ZERO on the spectrometer. This will allow the spectrometer to compensate for any small, nonzero signals from any of the elements of the detector in the absence of light. (The Z on the display will disappear.) Collecting Data A4. The fiber optic cable will be mounted in a clamp on a ring stand, pointing toward the ceiling lights. You need to adjust the intensity of the light reaching the spectrometer so that the signal is maximized without saturating the detector, that is, without exceeding the light intensity to which it is capable of responding.

To do this, press the spectrometer button marked INTENSITY CHECK. While observing the display on the spectrometer, make slight adjustments in the direction in which the fiber optic cable is pointing. The spectrometer is scanning repeatedly and displaying the maximum intensity value found in each scan. Try to get this to be reasonably close to the maximum value of 4095, say 3500 or greater. If the message SATURATED appears, the intensity is too great. Once you have a satisfactory intensity, be careful not to disturb the position of the fiber optic cable. A5. To record the emission spectrum, press INTENSITY CHECK again to leave the intensity check mode, then press the spectrometer button labeled SAMPLE. The S on the display will disappear. Return to your station, where, in a few seconds, the spectrum will appear on your station display. A6. Press the FILE OPTIONS button on the station, select PRINT standard, enter the number of copies to be printed (probably 2, since you are working in pairs), and press ENTER. When your data have been sent to the printer, the message SELECT A FUNCTION will appear. The copies of your spectrum will be printed on the printer attached to the computer on your network and will be labeled at the top with your station number. Visual Comparison A7. A principal goal of this experiment is for you to correlate the emission spectrum recorded by the spectrometer with the colors of the emitted light. Take your printed emission spectrum with you to the location on the bench of the simple plastic spectroscope. View the ceiling lights through this spectroscope and compare the features in the emission spectrum with your visual observations. (Consult the instructor if you need help in using the spectroscope.) Label the most intense features on your emission spectrum with the colored lines you see (you should see at least 5 lines). The crude wavelength scale in the spectroscope itself, along with the following very approximate scale, may be useful in making the assignments. Turn the labeled plot in to your TA before you leave lab for the day. ultraviolet violet blue green yellow orange red infrared 400 nm 500 nm 600 nm 700 nm B. Absorption Measurement Stock Solution COMMON MEASUREMENT MISTAKES TO AVOID In order to obtain high-quality, reliable absorption spectra, you must be careful to avoid a number of common mistakes. The most typical mistake, which will affect your results significantly, is to position the cuvet incorrectly in the cell holder. The cuvets have two clear sides and two frosted sides; if the cell is positioned so that the light passes through the frosted sides, the absorption will be too high or too low, depending on whether the reference cell or sample cell was the one positioned incorrectly. If any of your spectra are above or below zero at the end, you should redo the spectrum, making sure to align the cuvets properly. Still other mistakes to avoid include having a smudge or fingerprint in the light path on one of the cuvets. Another possibility is the presence of an air bubble in the light path, or one of the cells may not be positioned all the way in the bottom of the cell holder or may have been tilted from the vertical. In short, to achieve good results, you must be careful in each measurement to have clean cuvets, properly positioned and oriented, with no air bubbles to interfere with the light path.

Sample Preparation B1. Obtain about 25 ml (no more!) of the food dye stock solution assigned to you in a 50 ml beaker. Be Logging In sure to record the concentration listed on the bottle. Fill your two cuvets about three-fourths full, one of them with distilled water (reference) and the other with the stock solution (sample). B2. Press the MAIN MENU button on the workstation. Select OTHER, then LOG IN. Follow the instructions on the workstation screen to log in, using the 4-digit ID obtained from the Chem21 website. Check with your TA to make sure you are logged in properly. Setting Up the Station B3. Press the MAIN MENU button on the workstation. Select SPECTROSCOPY, then ABSORPTION(1), and, then press DISPLAY. Setting Up the Light Source and Sample Holder B4. A fiber optic cable is used to connect the tungsten/halogen light source and sample holder to the spectrometer. If this has not yet been done, ask your instructor for assistance. Make sure the light source is turned on. Zeroing the Spectrometer B5. At the spectrometer, press STATION NUMBER, your station number, then ENTER. The display will show STA# ABS-VIS Z R S on the first line, where # is your station number. B6. Insert the black cuvet to block the light from the light source, then press ZERO on the spectrometer. The Z will disappear. The display on the spectrometer will indicate when you may go on with the next step. It should not take longer than a few seconds. Collecting Data Reference B7. Remove the black cuvet and place your cuvet with distilled water (reference) in the cell holder (fully inserted), making sure that the arrow on the top of one side of the cuvet is facing the light source. If necessary, carefully wipe the sides of the cuvet with a Kimwipe to remove any dust or fingerprints, etc., before placing it in the cell holder. When you are ready to record the reference spectrum, press the spectrometer button labeled REFERENCE. The R will disappear from the display, which will indicate when you may go on to the next step. Collecting Data Sample B8. Place your sample cuvet in the cell holder, clean, properly oriented, and fully inserted. When you are ready to record the sample spectrum, press the spectrometer button labeled SAMPLE. The S will disappear, and the display will indicate when you may go on. Take your cuvets and return to your station. In a few seconds the absorbance curve calculated from the reference and sample spectra will appear on the display. B9. Evaluate the displayed spectrum on the workstation screen the maximum absorbance (not including the 350-450 nm region) should be ~1.0. If it is not, you may have inserted either the reference cuvet or your sample cuvet into the cell holder incorrectly. Redo steps B3-B8, making sure that the arrow on the cuvet is always facing the light source. If the spectrum appears to be satisfactory, press the FILE OPTIONS button and select SAVE to save your absorption spectrum on the PC attached to your

network and to the Chem21 website. When asked to enter a three-digit number, enter 001. Press ENTER to save the file. Check with your TA to make sure the spectrum looks OK and the file has been uploaded to Chem21 properly. C. Absorbance Measurements Beer s Law Plot In this part of the experiment you will record the absorption spectra of four solutions you prepare by diluting the stock solution from Part B of this experiment. Preparing the Diluted Solutions Before preparing the diluted solutions, practice pipetting with the Mohr pipet using distilled water until you can measure out volumes to the nearest 0.1 ml by controlling the flow with the pipet pump. Measure the position of the bottom of the meniscus with respect to the graduation marks on the pipet. A good way to measure out 2.0 ml, for example, is to fill the pipet to the 0.0 ml mark and drain it to the 2.0 ml mark. Before using the pipet for the stock solution, rinse it with a small portion of the stock solution, allowing it to drain into a waste beaker. Obtaining Absorbance Spectra of Diluted Solutions C1. Pipet 2.00 ml of stock solution into the 10-mL volumetric flask. Using a disposable pipet, add distilled water to the flask to bring the bottom of the meniscus to the mark on the neck of the flask. Prepare this solution as carefully as you can, since the quality of your Beer s Law plot will be determined to a large extent by how accurately the concentrations are known. Cap the flask and invert it several times to mix the diluted solution. C2. Use a small amount of the diluted solution to rinse out the sample cuvet, and then fill the cuvet about three-fourths full with the solution and cap it. (For each absorbance measurement you will use the same reference cuvet as in Part B, containing distilled water, and the same sample cuvet for each of the solutions.) C3. Obtain the absorbance spectrum for the solution as in B4-B9 above. Press the FILE OPTIONS button on the station, select SAVE, using any 3-digit number (other than 001) for creating a file name. Be sure to write down the 3-digit code you used to save the data along with the volume of stock solution used to make the solution. Do not discard any of your solutions until you are completely satisfied C4. Repeat step C1 this time using 4.00 ml of the stock solution. Obtain and save an absorption spectrum for this solution as in steps C2-3, using a different 3-digit code when saving the spectrum. C5. Repeat steps C1-3 for two more solutions, using 6.00 ml and 8.00 ml of the stock solution, respectively. Obtaining Absorbance Spectra of Sports Drink C6. Obtain ~5 ml of a sports drink (it must be the same color as the solutions you prepared). Rinse the 10 ml volumetric flask with distilled water, then use the 3 ml volumetric pipet to deliver 3 ml of the sports drink into the volumetric flask. Dilute to the mark with distilled water, and obtain the absorption spectrum as in steps B4-B9 above. Save the absorption spectrum (with a different 3-digit number). Determining Absorbance Values at Selected wavelength C7. You will now determine the absorbance values at a selected wavelength for each of the five known solutions and your sports drink. At the PC, select the program Absorbance Series. Enter your station

number and select the six filenames you used when saving the absorption spectra (the 3-digit numbers you entered when saving the data will be the last three numbers in the filename). Check to make sure the spectra all have approximately the same baseline (the absorbance value at high wavelengths); if some of the spectra have a baseline that is too high or too low, obtain a new spectrum. Select a wavelength that corresponds to a point on each curve near the absorbance maximum, and enter this value in the appropriate box. Print your results: a plot with all six absorption spectra in one graph and a list of the absorbance values at the chosen wavelength. Waste Disposal All solutions may be poured down the sink. RESULTS A. Emission Spectroscopy 1. As a part of your report, submit the emission spectrum of the fluorescent light, giving the plot an appropriate title. (What parameters does the plot include? What sample was being investigated?) Label each of the most intense features with its corresponding color. Turn this in to your instructor before you leave lab for the day. B. Absorption Spectra The spectrometer does not measure absorbance directly; rather it measure the light intensity after the light passes through the reference cell (I 0, ) and the light intensity after it passes through the sample containing the absorbing species (I ) and then calculates A at all wavelengths. In this part you will demonstrate your understanding of how the absorbance values are obtained. 1. Log on to the Chem21 website and access the exp. 6 lab report. 2. Using the absorbance and intensity plots linked to the report sheet corresponding to the data obtained from the absorbance spectrum of the stock solution, use the cursor to determine I 0,, I, and A at two chosen wavelengths. Select wavelengths at which the absorbance is >0.1. 3. Calculate A from the values of I 0, and I obtained above. (hint: look at eq. 6-3) Enter these values in the appropriate boxes on the REPORT SHEET. C. Absorbance Measurements Beer s Law Plot 1. Use the instructions at the end of the handout to create your Beer s Law plot using Excel. Do this before entering any of your data onto Chem21. You will need to calculate the concentration of each of your diluted solutions; M 1 V 1 = M 2 V 2 will be useful here. 2. From your data and EQUATION 6-4, determine the molar absorptivity,, for each solution at the wavelength used for this plot. Note: the path length b through the cell is 1.00 cm. Determine the average molar absorptivity for the substance you investigated. 3. A Beer s Law plot of absorbance vs. concentration gives a straight line passing through the point (0,0). This has the form y = mx, where m is the slope of the line. Thus, using EQUATION 6-4, A b C, a plot of absorbance (y-axis) vs. concentration (x-axis) will give a straight line where the slope = εb. Note this very important consequence when collecting experimental data: if the data fits a straight line, we can often obtain information regarding the phenomenon being observed from the slope of the line. Use the slope of linear regression line from your Beer s Law plot to determine the molar absorptivity for the substance you investigated. Note: when a printout lists a value in scientific notation, but is not capable of displaying

superscripts, the symbol E is often used to represent a power of 10. For example, a slope value of 1.67 x 10 2 would be displayed as 1.67E+02. You must click on the Graph Data link on the report sheet before entering your slope value, or you will receive an error message. 4. Using the slope value of your Beer s Law plot, determine the concentration of food dye in your diluted sample. Since A b C and slope = εb (from above), then Aλ = (slope)c. Thus, the concentration of your diluted sample can be found by A Concentrat ion of food dyein diluted sportsdrink (6-6) slope provided the absorbance value from the spectrum of the unknown solution is obtained at the same wavelength as those absorbance values used to generate the slope of the line from your Beer s Law plot. 5. Remember that you diluted 3.00 ml of your sports drink to 10.00 ml before obtaining the absorption spectrum. Use this information to calculate the concentration of food dye in your original undiluted sports drink sample.

EXPERIMENT 6 REPORT SHEET Name: Date: Partner: PART B. ABSORPTION SPECTRA Wavelength, nm I 0, I A (measured) A (calculated) PART C. BEER S LAW PLOT Substance: Wavelength used: nm Sample number Volume of stock solution, ml Concentration, mol/l Absorbance Molar Absorptivity (1) stock solution xxxxxxxxxxxxxxx xxxxxxxxxxxxxxx (2) first diluted solution 2.0 (3) second diluted solution 4.0 (4) third diluted solution 6.0 (5) fourth diluted solution 8.0 Average Determination of food dye concentration in sport drink Sports drink flavor Absorbance of diluted sports drink Concentration of food dye in diluted sports drink Concentration of food dye in undiluted sports drink

Changes in Procedure: Laboratory Notes Introduction to Spectroscopy Part A: Your TA will record the emission spectrum and print out a copy for you. You need only to do Step A7 in this part of the experiment. Turn in the labeled spectrum to your TA before you leave the lab room today. Parts B and C Sample Assignments: even stations: FD&C Blue #1 odd stations: FD&C Red #40 Note that only printouts you will obtain are the emission spectrum of the fluorescent light and the printout of all the absorption spectra using the Absorbance Series program. All individual absorption spectra should be saved, noted printed. When preparing your Beer's Law plot, be sure to set the y-intercept of the trendline equal to 0. 5/12

Instructions for Preparing Beer s Law Plot for Exp. 6 Excel 2007 Instructions 1. Open Excel. In cell A1, type the column heading Concentration, M. In cell B1 type the column heading Absorbance. 2. Select columns A through B: click on the heading for Column A (the shaded cell with the letter A), and while holding the mouse button down, drag the mouse to the heading for column B. Click on the line separating the headings for columns A and B a line width window should pop up. Change the width of columns A and B to 15.00. 3. Enter the concentration and absorbance data for the stock solution and each of the four diluted solutions (not the unknown solution). 4. Select all of the data in columns A and B by clicking on cell A1, dragging across to B1, then dragging down to B6. 5. Select Insert from the menu bar. From the Graph section, select Scatter, then Scatter with only Markers. 6. Draw a best fit line through your data by selecting Trendline. Select More Trendline Options, make sure Linear is selected for the regression type, then select Display equation on chart and Set y-intercept = 0. Click on Close. 7. Enter the slope value onto the online report sheet.