Be sure to staple your graph to the handout!

CHEM 151 EMISSION AND ABSORPTION SPECTROSCOPY Winter 2009 Fill-in Prelab attached (p 13) Stamp Here Name Partner Lecturer Be sure to staple your graph to the handout! Introduction Date You have now observed that when certain materials are heated they give off light. A burner on an electric stove is an everyday example. Gases, as well as solids, will emit light when heated. For example, a little table salt (sodium chloride) or baking soda (sodium bicarbonate) sprinkled in a flame will excite the electrons and will impart a yellow color to the flame; a copper wire will impart a green color to the flame where it is hot enough to excite the electrons in the copper. The electrons jump from lower levels to higher levels as a result of the amount of heat added. When the electrons fall back down to the lower levels, some form of light is given off. Electrical discharges will bring about the same effect in gases as the flame does with solids. You have observed this in neon advertising signs where an electrical discharge through a tube containing neon, argon or helium has caused the tube to give brightly colored light. The color you observe is a mixture of all of the colors emitted. Many solutions are colored also. Copper (II) solutions and nickel (II) solutions have characteristic blue and green colors respectively. The color of a solution observed is simply a mixture of the colors of the different wavelengths of light that are not absorbed. PART A Emission Spectroscopy This emission of light is consistent with the theory you have learned in lecture about quantized energy levels. The electron is excited then falls back to a lower level, giving off a photon of light with an energy equal to the difference in the energy of the electron before and after its fall. (See Figure 1) E=hν or E=hc/λ Figure 1 Bohr model of the atom, with principal quantum levels n = 1 to n = 4 shown http://cse.ssl.berkeley.edu/img/em_bohr.gif #8 Spectroscopy Rev W09AEM Winter 2009 Page 1 of 13

Oscillating electric charges, electromagnetic disturbances and nuclear reactions can produce the electromagnetic radiation shown in Figure 2. Electron transitions in excited atoms can produce such electromagnetic radiation as X-rays, ultraviolet rays, visible rays and infrared waves. This experiment will only involve studies in the visible range of the spectrum. x-rays high energy high frequency short wavelength x-rays Wavelength (λ) in meters 10-11 10-9 -7 10 10-5 -3-1 3 10 10 10 10 ultraviolet visible light infrared rays (heat) violet blue green yellow orange red Hertzion waves microwaves low energy low frequency long wavelength 400 500 600 700 nanometer Figure 2 The Electromagnetic Spectrum Before we discuss the energy emission patterns of elements further, we need to first discuss the properties of visible light. Ordinary visible light (called white light) contains all the wavelengths of light in the 380 nm to 770 nm range. This white light can be broken up into its component colors using a spectroscope containing a prism or diffraction grating (See Figure 2). Since all wavelengths of light are present, the light is split up into a continuous spectrum. Each wavelength of light present in the spectra contains all visible wavelengths, thus it is called a continuous spectrum. And each color of light has a specific energy associated with it. The relationship between energy and wavelength can be determined as follows: E = hυ c = λυ Therefore υ = c/λ And E = hc/λ Let us also examine what the variables in the equations stand for Where E = energy of the photon h = Planck s constant: 6.626 x 10-34 J sec (Pronounced Nu) υ = frequency of light for the photon, in sec -1 where c = speed of light: 2.998 x 10 8 m/sec (pronounced lambda) λ = wavelength, in m #8 Spectroscopy Rev W09AEM Winter 2009 Page 2 of 13

As can be seen, a number of energy transitions may occur. The relationship between several of the energy transitions of the element hydrogen and its line spectra are shown in the chart on the wall in the lab. In this experiment you will learn how the known or experimentally determined spectral properties (their characteristic emission line spectra) of elements can be used to identify what elements are emitting the spectra. However, because these transitions vary in energy, only certain transitions fall within the visible spectrum region. Thus emission line spectra are termed discontinuous spectra. For hydrogen, only the Balmer series of transitions occur. The Balmer series is the name given for transitions that result in visible light being emitted by the ion or atom (e.g. ROYGBIV). The emission spectrum of hydrogen contains four visible lines. This series of transitions is unique to the element hydrogen. Transitions in the visible range give rise to color in a compound. However, the color observed is not the same as the color that is absorbed. As energy hits a compound, the full spectrum of light comes into each molecule. Only certain photons/wavelengths of light are absorbed, the rest are transmitted and are seen by the eye. The color that is seen in the complement of the color absorbed. If the colors of the rainbow are placed in a circle, complementary colors are opposite one another on this color wheel. 720 nm Examples: 400 nm V R O 650 nm If the compound is blue (B), then the color of the photon absorbed is orange (O). 470 nm B G Y 590 nm If the compound is red, the color of the photon absorbed is green. 520 nm PART B Absorption Spectrophotometry In the following discussion you will learn how the intensity of colors such as these can be used for quantitative analysis by measuring the absorbance of a specific wavelength of light by a solution. The basic idea is simple. The greater the concentration of a colored species the more intense the color and the greater the absorbance of the complimentary color, see your text in the Coordination Chemistry chapter for a color wheel. For example, the concentration of copper (II) which is blue in an unknown solution can be estimated by comparing its color intensity with a series of solutions having known concentrations. One can simply view tubes containing the same amounts of solution against a white background for a qualitative comparison. We use an instrument with a photocell for the analysis of colored solutions. Such an instrument is called a spectrophotometer. Figure 3 below is a block diagram showing the instrument in its simplest form. The technique is called spectrophotometry. As you probably expect, the use of a photocell is much more accurate than the human eye for measurements of light intensity. Another important consequence of photoelectric measurement is that the technique can be extended into the ultraviolet and infrared regions where the human eye does not work. Thus, thousands of substances that appear colorless to the eye can, in fact, be analyzed with a spectrophotometer. #8 Spectroscopy Rev W09AEM Winter 2009 Page 3 of 13

The current produced by a photocell is called the photocurrent the greater the light intensity the greater the current. The light source has multiple wavelengths and the monochrometer filters out all but one. For our equipment, the colorimeter allows you to select certain specific wavelengths. Light Source Monochrometer Photocell Wavelength Analyte solution in absorption cell Figure 3 Block Diagram of a Spectrophotometer M illiammeter Detector Of particular importance is the fact that light absorption is a function of the wavelength of light. A solution can be nearly transparent at one wavelength and completely opaque at a shorter or longer wavelength. Consequently, it is essential to use a light beam with a known wavelength. Light having a single wavelength or narrow range of wavelengths is said to be monochromatic. An optical device that accepts a continuous range of wavelengths from the light source and allows only one wavelength to pass is called a monochrometer. Spectrophotometers have a built in monochrometer, usually based on a prism, diffraction grating or interference filter that allows the operator to select any desired wavelength within the range of the instrument. In the following discussion it is always assumed that measurements are made using monochromatic light. In practice it is more convenient to use two matched cells. One, called the reference cell, is filled with solvent and used to measure the initial intensity, P o. The other, called the sample cell, holds the solution to be analyzed. Because the characteristics of the two cells might not be exactly the same, a cell correction is often applied. Transmittance, T, is how much of the light goes through the solution and is a fraction that varies from zero to one. It is measured by comparing the solution to be measured (analyte) to the reference (the solvent in this case). Often it is expressed as a percentage and given the symbol %T. Absorbance, A, is how much of the light is absorbed by the sample compared to the reference. The amount of light absorbed by a solution will depend on the length of the light path through the absorption cell, b. The longer the cell path the more light is absorbed. The amount absorbed also depends on the concentration of colored species, c. The relationship between the various factors is T = 10 -αbc where A = αbc and therefore T = 10 -A or % T/100= 10 -A. ε is a constant at a given wavelength of light. Thus, if the transmittance is known the absorbance can be calculated at the same wavelength. And, of course, if the absorbance is known the transmittance can be calculated. Also, if you take a close #8 Spectroscopy Rev W09AEM Winter 2009 Page 4 of 13

look at the scale on a spectrophotometer and note that one can read either transmittance or absorbance, whichever is desired. Beer s Law A = εb c Step 1 or A = α b c Intensity = P= i P i Intensity = = Po o Monochromatic atic Pure Pure solvent solvent Step 2 Step Photocell Photocell Photocurrent = = i i o Intensity = P i i Intensity = P Solution Photocurrent = = i i Figure 4 Beer s Law is the starting point for nearly all spectrophotometric analyses. Thus, a graph of absorbance versus concentration is a straight line that, in an ideal situation, passes through the origin. This is a much easier relationship to utilize than a graph of transmittance against concentration, which is a curved line. Once this graph has been drawn one can immediately find an unknown concentration of the analyte (sample) simply by measuring its absorbance and reading the concentration off the calibration curve. The linear graph also helps to identify incorrect data that would not be so easily detected in a curved graph. The proportionality constant in Beer s Law ε, (epsilon), is called the molar absorptivity. The molar absorptivity is used when the concentration of the light absorbing species is given in moles per Liter. This constant is specific to each different compound, and gives a measure of how much compound is needed to create a brightly colored solution. Compounds with large absorptivities (>10 3 ) require small amounts of compound, relative to a molecule with a low absorptivity. Often when dealing with unknown compounds or polymers the molecular weight is unknown so it is impossible to calculate a molar concentration. In such cases the concentration is expressed in grams per liter and the proportionality constant is called the gram absorptivity, α (alpha), in this case Beer s Law is written A = αbc, where c is the concentration in grams per Liter and b is the path length of the cell (often 1 cm). The most common method of quantitative analysis is to prepare a series of solutions having known concentrations, measure their absorption, and construct a calibration curve such as the figure below. An unknown concentration can then be found simply by measuring the absorbance and reading the concentration from the calibration curve. If the absorbance is too high to measure accurately, a quantitative dilution must be made. Student data is represented on the next page. #8 Spectroscopy Rev W09AEM Winter 2009 Page 5 of 13

Figure 5 A Beer s Law Plot, or Calibration Curve. Absorbance vs Concentration 1 0.8 0.6 0.4 0.2 0 0.0000 0.0500 0.1000 0.1500 0.2000 0.2500 Concentration (g/l) EXPERIMENTAL PART A Procedure for Emission Spectroscopy-Flame Tests 1. Obtain a Bunsen burner. 2. Obtain a tray containing bottled salt solutions of Na +1 (as NaCl and NaHCO 3 ), Li + (LiCl), K + (KCl), Ca 2+ (CaCl 2 ), Ba 2+ (Ba(NO 3 ) 2 ) and Cu 2+ (CuCl 2 ), an unknown, and 4 cotton swabs. You do not need to transfer the solutions use them in the bottles provided. Each cotton swab can be used exactly twice once on each end! 3. Obtain 20 30 ml of water in a small beaker. This will be used to cool the cotton swabs after each test. 4. Dip one cotton swab into each salt solution for at least 5 seconds to saturate the swab. 5. Light the Bunsen burner and adjust the burner for a small flame with the gas valve at the bottom of the burner. 6. Place the tip of a saturated swab in the flame just long enough to observe and note the color of the salt but do not catch the cotton swab on fire. Place the used swab into the beaker of water to cool or extinguish any flames. 7. Perform this flame test for all of the salts, and the unknown. 8. Record the flame color for each of the salts, the unknown letter and the flame color for the unknown in the table. For each salt you should record the color as well as observations. Note things such as the size of the flame, if it smells, if it sparks, etc... 9. Return the tray of salt solutions to the dispensing area, dispose of the cooled swabs in the trash, and rinse the water down the drain. #8 Spectroscopy Rev W09AEM Winter 2009 Page 6 of 13

Results for Flame Tests: be descriptive on color observations. Be cautious about using the same color to describe different compounds Salt observed (Li + ) /LiCl Color of Flame and Observations (K + ) /KCl (Ca +2 ) /CaCl 2 (Ba +2 ) /Ba(NO 3 ) 2 (Cu +2 ) /CuCl 2 (Na +1 ) /NaCl (Na +1 ) /NaHCO 3 Unknown 1. Record the unknown letter on the sample bottle 2. What is the chemical identity of your unknown sample? 3. How does the flame test for NaCl compare with that of NaHCO 3? Explain your answer. (e.g. are the flames the same color, are they as intense as each other, are they the same size, do either of them spark?) 4. Did you notice any differences in the flame test for the different salts? yes there is a difference no, there is no difference 5. Do you think the anion (-) or the cation (+) accounts for the differences? (HINT: examine your results for all the species that contain the Cl -1 ion - examine the species that only contain the Na +1 ion which have the same color flame?) the cation determines the color of the flame the anion determines the color of the flame 6. Utilizing the emission chart for various elements posted on the wall, determine the identity of the following unknowns given the wavelengths of light emitted. a) 410.1 nm, 434.1 nm, 486.1 nm, 656.3 nm b) 405.0 nm,435.5 nm, 546.0nm, 577-580nm, 615.1nm, 690.8nm, 734.5nm. #8 Spectroscopy Rev W09AEM Winter 2009 Page 7 of 13

PART B: Procedure for the Absorption Spectrophotometry The solutions are all made from a food dye and may be disposed of down the drain. 1. Select one of the prepared dye stock (known) solutions; red, blue or yellow. Pour approximately 40 ml of the stock dye solution using a small beaker. On the data sheet, record the concentration of your stock dye solution, in g/l (you can find this written on the bottle). Also note the appropriate wavelength for your dye (you will set your colorimeter to this wavelength). Pour approximately 40 ml of distilled water into another small beaker. 3. Now you will make four more solutions, by dilution, so that you will have four solutions of the same color with different absorbencies spread in the range from 0.1 to 0.9 A.U. Refer to your handbook for directions for using a pipet; you may want to practice pipetting water first. If you have never pipetted before, or are unsure how, ask your lab instructor to demonstrate this technique to you. 4. Label five clean, dry test tubes 1-5. Your fifth sample will be your original solution. Pipet (use a dry GLASS graduated pipet - NOT plastic) 2.00, 4.00, 6.00, and 8.00 ml of the dye solution into the test tubes 1-4 respectively. With a second glass graduated pipet, deliver 8.00, 6.00, 4.00, and 2.00 ml of distilled water into test tubes 1-4 respectively. Note that the total volume in each is 10.00 ml. Thoroughly mix each solution with a glass stirring rod, taking care to rinse and dry the stirring rod between solutions. In the 5 th test tube, transfer 10.0 ml of the stock dye solution (no water added!) 5. Calculate the concentration in each solution after the dilution using C 1 V 1 = C 2 V 2. In this equation, C 1 represents the solution that you started with. V 1 represents how much of that solution that you used, C 2 is your final (diluted) concentration, and V 2 is your total volume (dye + water). Record your concentrations in the data table on page 10 of the lab handout. Check your calculated concentration values with the lab instructor before analyzing the solutions on the computer! 6. You are now ready to create your calibration curve. Take your five solutions, along with a wash bottle and a waste beaker to a computer with a colorimeter. There should be a plastic cuvette with your colorimeter by your computer. Make sure that your cuvette is clean and free of scratches. 7. Log into the computer using your student log-in (you can obtain one from the library). 8. Open Logger Pro on the computer. 9. Click on File (on the Menu bar), then Open. Click on the Chemistry with Computers folder. There you will find the LoggerPro files. Open Exp 11 Beer s Law in the Chemistry with Computers folder. 10. Prepare for data collection. You will need to make sure the units on the x-axis are in g/l. If they are not, change them to the appropriate unit: this is done by double clicking on the word concentration in the data table on the left side of the screen. In the same dialogue box, click the options tab to change the number of digits displayed in the Table window for concentration. Change the number of sig figs based on your calculated concentrations of your dyes Click done. 11. You are now ready to calibrate the Colorimeter. Prepare a blank by filling your cuvette 3/4 full with distilled water. To correctly use a Colorimeter cuvette, remember: All cuvettes should be wiped clean and dry on the outside with a kimwipe. Handle cuvettes only by the top edge of the ribbed sides. #8 Spectroscopy Rev W09AEM Winter 2009 Page 8 of 13

All solutions should be free of bubbles. The cuvette cap should be securely snapped on to the cuvette. Always position the cuvette with its clear sides facing the white mark inside the sample holder. If you place it in the colorimeter such that the ribbed sides face the beam of light it will not read your samples accurately. Holding the cuvette by the upper edges, place it in the cuvette slot of the Colorimeter and close the lid. 12. Set the wavelength on the Colorimeter to the appropriate wavelength for your dye, by using the arrow keys on the colorimeter to toggle through the available wavelengths. Press the CAL button, and wait for the light to stop flashing. The colorimeter is now calibrated. 13. You are now ready to collect absorbance data for the five standard solutions. You must use only one cuvette! Click Collect. Empty the water from the cuvette. Using the most dilute solution first, rinse the cuvette twice with ~1-mL amounts of the solution and then fill it 3/4 full. Dislodge any bubbles in the solution and cap the cuvette. Wipe the outside with a kimwipe and place it in the Colorimeter. After closing the lid, wait for the absorbance value displayed on the monitor to stabilize. Then click Keep, type in the concentration in the edit box, and press the ENTER key. The data pair you just collected should now be plotted on the graph. 14. Discard the cuvette contents into your waste beaker. Repeat the process of rinsing and filling the cuvette in order of increasing concentration of your dye solutions. Measure the absorbance of each solution in a similar fashion, entering the appropriate concentration with each kept data point. 15. When you have finished collecting data for all of your known solutions, click Stop. 16. Record the absorbance and concentration data pairs that are displayed in the Table window in your lab handout on page 10) 17. Rescale the x-axis of your graph by clicking the Autoscale button. 18. Examine the graph of absorbance vs. concentration. To see if the curve represents a direct relationship between these two variables, click the Linear Regression button,. A best-fit linear regression line will be shown for your five data points. This line should pass near or through the data points and the origin of the graph. You must choose y= mx + b NOT why = Ax 19. Title and rescale the graph appropriately (Absorbance vs. Concentration is not an appropriate title). Include the color of the dye used in the title. Click on File, Print Graph (this prints the graph only), with a regression line and the equation of the regression line displayed. Part C: Determining the concentration of an unknown dye solution. 1. Obtain a 10 ml sample of the appropriate single-color unknown. 2. Add the unknown to the cuvette and allow the colorimeter to read its absorbance. You can do this with your calibration plot still showing on the screen. Do NOT hit collect, you simply need to record this absorbance value in the data table in your lab handout. 3. Dispose of all dye solutions down the drain (it is just food coloring ) #8 Spectroscopy Rev W09AEM Winter 2009 Page 9 of 13

Original concentration Color of dye solution Wavelength Show a sample calculation for dilutions (C 1 V 1 = C 2 V 2 ): Pay attention to sig figs!! Show a sample calculation for % T of your unknown dye %T A = - log 100 Pay attention to sig figs. Using the acquired absorbance on the computer, calculate the % T for your unknown dye Instructor initials for Concentration: Data Table for Spectrophotometry Test Tube Dye Volume (ml) Water Volume (ml) Total Volume (ml) Concentration (g/l) % Trans (T) Absorbance (A) 1 2.00 8.00 2 4.00 6.00 3 6.00 4.00 4 8.00 2.00 5 Original solution 10.00 0.00 Unknown -------- -------- ------- See your slope and calculations. #8 Spectroscopy Rev W09AEM Winter 2009 Page 10 of 13

Part D. Analyzing the data: (review your pre-lab questions before answering Q2-5!) 1. Record the value for the concentration of your unknown, as read off of your calibration curve. Please print out the graph, and with a pen, mark the absorbance of your unknown on the best fit line with an X. Then draw a vertical line (using a ruler!) down to the x-axis where you will guesstimate your concentration by reading your graph. This number is the concentration that you have read off your graph. Pay attention to sig figs as the concentrations on your x-axis given by the computer may not have the correct sig figs!! Don t forget your units c read = 2. Using your calibration curve and the equation of best fit line, identify the slope for your data. Careful with sig figs!! Don t forget your units here! slope = 3. From this value, determine the value for α for your dye. (Hint: what is the size of the cuvette? See prelab #2). Don t forget your units!! 4. Write out a generic equation for Beer s Law, substituting the numerical value of αb that you have obtained from the slope of your line. Don t forget to include your units!! 5. Calculate the concentration of the unknown in g/l, using the equation above. Show your calculation. Don t forget your units!! Pay attention to sig figs!! α = A = c calculated = 6. Do the values you obtained in question 1 and 5 differ? By how much? (calculate a % difference!) Can you list a reason why they might be different? (hint: What is the y-intercept and what should it be based on the Beer s Law?) #8 Spectroscopy Rev W09AEM Winter 2009 Page 11 of 13

Homework Questions: can be done at home, need not be finished to get final stamp 1.) Look up the energy-wavelength relationship, write the equation. As the wavelength increases, does the energy of transition increase or decrease (circle the appropriate answer)? equation if λ increases, E increases decreases Look up the energy-frequency relationship, write the equation. What happens to the frequency as the energy increases (circle the appropriate answer)? equation if υ increases, E increases decreases 2.) Using the data gathered from your absorbance vs. concentration data and the information given in the lab handout, answer the following questions a. What was the color of your dye? b. Therefore, what color of light was absorbed? c. What is the energy associated with this absorbed light? SHOW YOUR WORK!! Pay attention to units and sig figs. E = 3.) Given the following wavelengths, circle which has the largest energy of transition? a. 450 nm b. 686 nm c. 545 nm 4.) Given the following frequencies, circle which has the largest energy of transition? a. 2.45 x 10 9 s -1 b. 8.68 x 10-9 s -1 c. 6.43x10 13 s -1 Don't forget to attach your graph with the equation of the best-fine line shown on the graph itself. #8 Spectroscopy Rev W09AEM Winter 2009 Page 12 of 13

Stamp: Prelab Questions Read the handout before attempting to answer the questions below. 1. Beer s Law states that A = αbc or A = εbc. Both ε and α have to do with absorptivity. If you were performing a spectrophotometry experiment (like you are today!), under what specific conditions would you want to use ε? Under what specific conditions would you want to use α? 2. Beer s Law is A= αbc, where A is the absorbance of the sample, α is the gram absorptivity, b is the path length of the cell in cm, and c is the concentration in g/l. A plot of A vs. c gives a straight line where the y-intercept ideally should pass through the origin. Therefore, A = (αb)c + 0 should be the equation of the line. a.) The slope of the line is given by which variable(s): a) α b) αb c) A d) c b.) The slope of a Beer s Law graph is 0.0098 L/g. Using your answer for question 2a, solve for α if the path length (b) of the cell is 1.0 cm. Be sure to give the units. α = c.) If a sample has an absorbance of 0.82, what is the concentration (you will also need the information from 2b to solve this problem)? (note: absorbance is unitless) c = 3. Given the color wheel in this handout and the fact that a solution absorbs wavelengths opposite the observed color, what color is absorbed for the following solutions. What wavelength does the color have? a) A Ni +2 solution appears green color absorbed = λ absorbed = nm b) A Cu +2 solution appears blue color absorbed = λ absorbed = nm 4. You must perform a dilution of a purple dye. You take 3.0 ml of a 0.0876 g/l stock solution and dilute the dye with 17.0 ml of water. What is the concentration, in g/l, of this new solution? SHOW ALL WORK! (see page 8 in the handout for help) #8 Spectroscopy Rev W09AEM Winter 2009 Page 13 of 13