Absorbance Spectrophotometry: Analysis of FD&C Red Food Dye #40 Note: there is a second document document that goes with this one! “2046 “2046 - Absorbance Spectrophotometry - Calibration Curve Procedure.” The second document has instructions for determining determining the slope and intercept of a linear data set. You will need to do this for the calculations in the second week of this experiment.
Introduction The purpose of this experiment is to determine d etermine the concentration of red dye in an unknown sample, provided by your instructor. Students will familiarize familiarize themselves with use of volumetric pipettes, volumetric flasks and the absorption spectrophotometer. A review of molarity calculations will be helpful in performing this lab! lab!
Discussion Absorban ce Spectr Spectr ophotometr y is a commonly used laboratory technique for is
determining the concentration of substances substances in solutions. In this experiment we will prepare several samples of FD&C Red #40, a commercial food dye, and use them to calibrate the absorbance absorbance spectrophotometer. Using the calibrated calibrated spectrophotometer, we will then analyze a sample of Kool-Aid brand drink mix to determine the quantity of Red #40 in the sample. Many compounds are colored due to their absorption of visible light. light. Our eyes are naturally capable of detecting many man y different wavelengths or energies of light, and we perceive these different wavelengths as different colors, as summarized summarized in the table below: Color violet blue green yellow orange red
Wavelength 380 – 450 450 nm 450 – 495 495 nm 495 – 570 570 nm 570 – 590 590 nm 590 – 620 620 nm 620 – 750 750 nm
So, for instance, light with wavelengths between 620 and 750 nm are perceived by the human eye as “red” light. “White” light, such as that produced by the sun or many artificial light sources, contains all of the visible wavelengths of light combined together; our eye perceives this as “white” but it is really a blend of all of the colors co lors of the spectrum (red, orange, yellow, green, blue, indigo and violet). When white light passes through a colored liquid, some of the wavelengths are absorbed abso rbed and some are transmitted. transmitted. The result is that that the liquid will appear colored to our eye. If a liquid absorbs all of the light in the violet, blue, green and yellow wavelength ranges, but allows orange and red light to pass through, then the liquid will appear orange-red 1
because its only transmitting orange and red light. All of the other colors are filtered out by the molecules in the liquid. On the other hand, a substance that absorbs in the wavelength range 500 -750 nm will appear blue or purple to our eyes, because all of the red, orange, yellow and green light have been absorbed. If a sample absorbs no light within the visible spectral range, then it will appear as a clear, colorless liquid. Water, for instance, absorbs such a small amount of visible light that our eyes generally cannot perceive the absorbance; unless you have a very large volume of water, it appears completely clear. (Please note that water is NOT truly clear it DOES absorb visible light to a small degr ee. That’s why it gets dark er as you move to the bottom of a pool of water, even if the water is clean). If a sample absorbs ALL of the visible wavelengths then it is opaque ; no visible light will pass through it at all. Many aqueous solutions of chemicals are transparent (meaning that they transmit light) but not colorless; they will absorb certain wavelengths and transmit others, resulting in a colored solution as described above. Aqueous solutions of copper ion, for instance, appear blue or green because they absorb light in the range from about 450 - 750 nm. Aqueous solutions of certain iron ions are orange because they absorb in the range from about 380 - 600 nm. To illustrate these ideas, consider the figure below. Light is shown on the left as incident on a sample and then emerging from the sample at the right. In general, Po, the radiant power of the incident light, will be larger than P, the power of the emerging light, because the sample will absorb some of the light. This is referred to as the absorbance of the solution. Not only is Po > P, but the colors (wavelengths) of the emerging light may be different than the incident light. sample Po
P
b
A simple and logical relationship exists between the absorbance of the solution and the concentration of the solution; namely, that the more sample molecules we place in the path of the incident light, the greater will be the absorbance of the light. Low concentrations will transmit more light (and the color will appear lighter or less intense), and high concentrations will transmit less light (and the color will appear darker or more intense). This relationship is the basis of Beer’s Law, usually expressed as follows: A = bc
(equation 1)
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where A is the absorbance, is a mathematical proportionality constant called the molar absorptivity , b is the pathlength of the light through the sample (as shown in the figure above), and c is the concentration of the absorbing species. Beer’s Law can be used to determine concentration of a compound in solution if we measure A and know both and b, or their product b. The law’s most common application in chemistry is to liquid-phase solutions; the concentration, c, is usually expressed in molarity (M) and the pathlength, b, is usually expressed in centimeters. The -1 -1 constant is specific to each wavelength and units of “M cm .” If you look at the units of absorbance in Beer’s Law, you can see that the value of A is unitless (or dimensionless). A spectrophotometer is a simple device that measures the amount of light which is absorbed at a particular wavelength; in terms of th e figure above, it compares P (the light which passes through the sample) to Po (the light which is incident on the sample). In theory, absorbances can range from zero (a sample which is totally transparent at the wavelength tested) to infinity (a sample which is completely opaque at the wavelength tested). In practice, however, absorbances above about 2 are not measurable; above an absorbance of 2, the sample transmits so little light that it is difficult to measure. (You will notice as you perform the experiment that if the reading rises above 1.999, the display will flash, indicating an error – readings above this value are not reliable). In very simple terms, the spectrophotometer measures how “dark” the color of a solution is – and Beer’s Law allows us to mathematically relate the “darkness” of t he color to the concentration of the dissolved solute. Look again at Beer’s law, above – you will notice that it has the form of a linear equation. The form of a linear equation is shown below, compared to Beer’s law: y = mx + intercept A = ( b)c + 0 You will notice that a plot of absorbance on the y-axis vs. concentration on the x-axis is predicted to be a linear plot, with a slope equal to b, and an intercept of zero. In absorbance experiments, this is referred to as the cali bration curve . The first step in performing an absorbance experiment is to determine the calibration curve, by preparing samples with known concentration (x-values) and measuring their absorbances (y-values). The calibration curve is a description or representation of the mathematical relationship between A and c. Thus, once the calibration curve is complete, you will be able to interconvert between A and c.
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For example, the graph below is a calibration curve for iodide ion, I . Several solutions of iodide ion were prepared, and their absorbances measured on the spectrophotometer:
Notice that the intercept of the y-intercept of the line is close to zero – this makes sense, because if the concentration the light-absorbing compound iodide is zero, then the absorbance should also be zero. In practice, the intercept will be a very small number, close to zero, but may not be exactly equal to zero. {Don’t be surprised when you do your calculations if your intercept is not zero!}. Also, the slope of the line has been calculated and is shown in the bottom left corner of the figure, along with the calculated intercept. Now that the calibration curve is known, the concentration of any iodide solution is easily determined by measuring its absorbance. For instance, if an unknown solution containing iodide has a measured absorbance of 0.400, then you can see from the graph that the concentration of iodide must be around 0.005 M; that’s the x-value that corresponds to a y-value of 0.40. Rather than approximating the concentration, however, we can precisely calculate it. Knowing that on this linear plot the slope is 79.12874, and the intercept is 0.0153, the relationship is: y = mx + b 1ABS = 79.12874 [I ] + 0.0153 If the absorbance is 0.400, the concentration of iodide is calculated to be 0.00486 M. Note that calculating the molarity gives a much more precise value than simply estimating it from the graph. This is why the slope and intercept of the line must be determined. The calculation process outlined in this example is very similar to the method you will use to determine the concentration of FD&C Red Dye #40 in an unknown sample.
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The compound which we will analyze in this experiment is known as FD&C Red Dye #40, commonly referred to as “Red 40” or “Allura Red.” It is found is hundreds (maybe thousands?) of food and drink products, including soft drinks, candies, processed fruit products, yogurt, ketchup and barbecue sauces, and much more. (Do an internet search for “Red Dye 40” or “Allura Red” if you would like to know more). The compound has a formula of C18H14 N2S2O8. You will receive a sample of Red 40 with a known concentration during the first week of this experiment, and use it to calibrate the instrument. During the second week, you will receive a sample of Kool-Aid powder, which contains Red 40. You will prepare a solution from the powder and determine the concentration of Red 40 in the solution. The overall goal of the experiment is to calculate % (by mass) of FD&C Red Dye #40 in the Kool-Aid powdered mix. NOTE: Your lab manual also has a discussion of absorbance spectrophotometry, in Appendix IV, under the heading “Light Absorption”; you can use that as a supplement to this document if you wish.
Procedure Week 1, Part A: Solution Preparation Proper use of volumetric flasks : You will be preparing your dilutions in volumetric flasks, a type of glassware that is probably new to you. A volumetric flask has an advantage over a graduated cylinder because it is more precise; a graduated cylinder is precise to within one tenth of a milliliter (one place past the decimal), but a volumetric flask is precise to within one one-hundreth of a milliliter (two places past the decimal). Because we are measuring small quantities in this experiment, it’s best to use very precise equipment.
Volumetric flasks are very simple to use, but remember a few precautions: Be sur e to fi ll the fl ask exactly to the fil l l in e, and not above! This is the most common mistake people make when using volumetric flasks. There is a horizontal marking on the flask, which is the mark for 250.00 mL. If you overfill the flask, there is no way to correct that error… you have to discard the solution and start over! Be cer tain to use the ri ght stopper or l id f or your f lask! Some of the flasks have screw-top lids, and some will require the use of a rubber stopper. If you use a stopper that is too small, it will get stuck in the neck of the flask! M ake sur e you have the r igh t size ! It should fit snugly, but not go down into the neck of the flask. Af ter f il li ng the fl ask, mix the soluti on thoroughly . Simply invert the stoppered flask a few times to make sure the solution is thoroughly blended. In this portion of the experiment, you will p repare four solutions, and collect the data required to complete a cali bration curve , as described in the discussion above. A “stock” solution of FD&C Red Dye #40 will be prepared for you. The concentration of this solution is expressed as grams of dye/Liter of solution. Record this concentration, 5
which you can convert to a molarity (remember that Red 40 has a formula of C18H14 N2S2O8). Obtain about 125 mL of the stock Red 40 solution in a clean, dry beaker. Please do not take more than you need ! Any excess will be wasted. Using 5 mL, 10 mL and 25 mL volumetric pipettes, and volumetric flasks, prepare the following dilutions of the stock solution:
Standard #1: 10.00 mL of stock, diluted to 250.00 mL. Standard #2: 25.00 mL of stock, diluted to 250.00 mL. Standard #3: 30.00 mL of stock, diluted to 250.00 mL. (use 25 mL and 5 mL pipettes). Standard #4: 50.00 mL of stock, diluted to 250.00 mL. (use 25 mL pipette, twice).
Calculate the molarity of Red 40 in each of the above solutions. Recall that M1V1 = M2V2. These four solutions are referred to as “standards”. In chemistry lab, a “standard” solution is one for which you kno w the concentration. Week 1, Part B: Measuring Absorbance of the Standards Proper use of cuvettes : Test tubes look perfectly clear to our e ye, but the absorbance spectrophotometer is much more sensitive at detecting light absorbance than our eyes; glass test tubes do in fact absorb a small but measureable amount of light. Also, test tubes can vary in thickness (thicker glass will absorb more light, thinner glass will absorb less light), and sometimes there are small bubbles or impurities in the glass from its manufacture, which affect the transmission of light and would affect our readings.
To avoid these problems, absorbance measurements must be made in special containers called “cuvettes”. They look like test tubes, but are made of an optically pure substance (quartz) that absorbs very little light. They are also manufactured to higher specifications than test tubes, so that all cuvettes will absorb a minimum amount of light. Cuvettes are expensive and must be handled properly to avoid damaging them, and to ensure that your readings are accurate. Please remember a few important things about using cuvettes:
Please do not get the cuvettes mi xed in with the test tu bes in your drawer . When
you finish using a cuvette, rinse it with distilled water and place it on the drying rack provided next to the spectrophotometer. Cuvettes are very easil y scratched ! NEVER use paper towels or scrub brushes on a cuvette! Wipe them with soft tissues only (Kim-Wipes are available for this purpose). In the spectrophotometer, the light will pass through the bottom half of the cuvette. When handling the cuvettes, you should always pick them up by the top half; oils from your skin will leave residue on the surface – while this residue is not visible to the naked eye, it will affect the absorbance of the sample! For the same reason you
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should wipe the outside of the cuvette with a Kim-Wipe before placing it in the spectrophotometer. You will notice that the cuvette has a vertical white line at the top; and if you look at the sample container on the instrument, you will n otice that there is a black line at the front of the opening. When placing the cuvette in the instrument, the white line on the cuvette should align with the black line on the instrument. This ensures that you always put the cuvette in with the same orientation, so that your absorbance results are reproducible. When you take your measurements, the cuvette should be at least half-full.
Data Acquisition : It is good practice to rinse the cuvette with some of the solution that is to be tested. Take a clean cuvette and fill it with a few mL of the first standard solution. Discard the solution; Red 40 is non-toxic and can go down the sink. Before using the spectrophotometer, the 0% T an d 100% T baselines must be adjusted. Your instructor will do this for you, but ensure that this procedure has been completed before proceeding. After rinsing, fill the cuvette at least halfway with the solution to be tested; place the sample in the instrument and ensure that the white line on the cuvette matches up with the black line on the spectrophotometer. Close the lid and record your absorbance. Absorbance readings should be three places past the decimal ; for the analog spectrophotometer, you will have to estimate the last decimal (as you would with a graduated cylinder or burette). The settings on the spectrophotometer are adjusted either by black knobs on the front of the instrument, or by buttons on the top (depending on which model of spectrophotometer you are using). Please do NOT adjust any settings; if you accidentally move the knobs or change settings, please inform your instructor so the spectrophotometer can be re-adjusted for use! Any Red 40 waste or excess can be disposed of in the sink. Part C: Plotting the Calibration curve (between week 1 and week 2) . Note: you should plot your calibration curve and have it ready when you arrive for week 2 of the lab! For instruction on how to plot and use your calibration curve, see the other document for this lab, “2046 - Absorbance Spectrophotometry - Calibration Curve Procedure”. (This can be downloaded from Dr. Rubini’s website, or you can get a copy from your instructor). Be certain to include the origin (concentration = 0, ABS = 0) as one of your data points. Week 2, Part D: Unknown Determination You will use absorbance spectrophotometry to determine the quantity of Red 40 in a sample of Kool-Aid drink mix. Make sure that you work with the same spectrophotometer this week as you did last week ! Not all of the instruments will have the same calibration curve.
Your instructor will furnish you with a sample of Kool-Aid drink mix. Carefully measure, into 250 mL beaker, 10 - 12 grams of the drink mix. Add about 150 mL of water - be careful not to add too much! - and stir to dissolve the powder. This may take a 7
few minutes to completely dissolve. When the powder is compl etely dissolved , transfer the solution to a 250 mL volumetric flask. Rinse the beaker several times with small portions of water, to ensure that all of the solution has been transferred to the volumetric flask. Then carefully fill the flask to the fill mark with water. Cap the flask and invert it several times to make certain the solution is thoroughly mixed. The solution of Kool-Aid that you have just made will be referred to on the data page as Solution 1. Solution 1 contains Red 40, but in a concentration that is too high to be accurately measured by the spectrophotometer – its red color will be too dark! (Remember from the discussion that absorbance values over 2 are not reliable). So we will dilute the solution again to bring it down to a range that is measurable. Pipette 50 mL of the Kool-Aid solution into another 250 mL volumetric flask and dilute it with water to the mark on the flask. As before, invert the flask several times to ensure a thorough mixing of the dye. This new solution will be referred to on the data page as Solution 2. Locate the SAME spectrophotometer that you calibrated last week . (It may not be in the same spot in the laboratory!). Check with your instructor to assure that the the 0%T and 100%T points have been set. Rinse a cuvette with a small amount of your second Kool-Aid mixture, and discard the rinse. Fill the cuvette at least halfway with your mix, and determine the absorbance of the sample.
Using the calibration curve which you have prepared, determine the concentration of Red 40 in the Kool-Aid solution. Any waste or excess Kool-Aid can be disposed of in the sink.
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Absorbance Spectrophotometry: Data Page Name _________________________________________________________________ Lab partner _____________________________________________________________ Spectrophotometer number ____________ Week 1: Absorbance of Standards
The stock solution, which was provided, was prepared by dissolving ____________ grams of Red 40 in water to make _________ L of solution. Standard 1
Concentration of Red 40 (M)
Absorbance
2 3 4 Questions – be sure to show ALL CALCULATIONS on your calculation page!
1. Calculate the molarity of the standard Red 40 solution that was provided. The molecular formula for Red 40 is C18H14 N2S2O8. _____________________ M 2. Calculate the concentration of Red 40 in the four standards you prepared. Enter them in the table above. 3. The pathlength of the cuvettes we used in this experiment is 1.00 cm. Use the formula for Beer’s Law (equation 1 in the discussion) to determine the molar absorptivity of Red 40. (For the purpose of this question, you may assume that the intercept of the calibration curve is zero). Report your answer for all four standards; they should be very similar, although they may not be exactly the same.
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Data Page
(continued)
Week 2: Kool-Aid analysis
Mass of powder mix used: ____________________________ grams Absorbance of Unknown _____________________________ Slope of calibration curve ____________________________ Intercept of calibration curve __________________________ Questions – be sure to show ALL CALCULATIONS on your calculation page!
4. Based on the absorbance of the Kool-Aid, determine the concentration of Red 40 in Solution 2. ______________________ M 5. How many moles of Red 40 were in your Kool-Aid Solution 2? _____________________ moles 6. How many grams of Red 40 were in your Kool-Aid Solution 2? ____________________ grams 7. How many grams of Red 40 were in Solution 1? (This is a very simple question! What is the relationship between the amount of solute in solution 1 and solution 2?) ____________________ grams 8. Calculate the % of Red 40 (by mass) in the Kool-Aid powder mix. ______________________ %
Remember to submit a copy of your calibration curve with this report. 10
Absorbance Spectrophotometry: Advance Study Assignment Name ______________________________________________________ 1. A solution is prepared by dissolving 0.1896 grams of FD&C Red #40 in water to produce 2.500 Liters of solution. Given that the molecular formula of Red 40 is C18H14 N2S2O8, determine the molarity of this stock solution.
_________ M 2. 25.00 mL of the dark red stock solution described question #1 is transferred to a volumetric flask and diluted with water to 100.00 mL. a. how many moles of Red 40 are in the 25.00 mL sample that was transferred?
_________ moles b. After the dilution, what is the molarity of the Red 40 in the 100.00 mL of solution?
_________ M note: the preceding problem (#2) is a dilution calculation. A convenient formula for calculating concentrations after dilution is M 1V 1 = M 2V 2 , where M 1 and V 1 are the molarity and volume of the concentrated solution, respectively, and M 2 and V 2 are the molarity and volume of the dilute solution. 3. a. The solution described in problem #2 is placed in the spectrophotometer. Based on the calibration curve shown on the next page, what would you predict the absorbance to be? Note : like the example given in the discussion, the absorbance should be calculated, not just estimated from the graph. Notice that the slope and intercept of the calibration curve are provided in the figure on the next page.
Absorbance = _________
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--------------------------------------------------------------------------------------------------------The following calibration curve was produced for Red 40. The slope was calculated to be 4 2.1033 x 10 . The intercept of the calibration curve is 0.0039.
--------------------------------------------------------------------------------------------------------3b. A sample of powered drink mix with a mass of 0.2548 grams is dissolved in water to make 250.00 mL of solution. A sample of this solution is placed in the spectrophotometer and its absorbance is measured as 0.3110. Based on the calibration curve above, calculate the molarity of Red 40 in the unknown.
_________ M c. How many grams of Red 40 are in the drink mix solution?
_________ g d. What is the % by mass of Red 40 in the drink mix powder?
_________ % 4. A sample of powdered drink mix with a mass of 11.22 g is dissolved in water to make 250.00 mL of solution. 50.00 mL of the resulting solution is transferred to a new container and diluted to a total volume of 250.00 mL. How many grams of the drink mix powder are in the second solution? (hint: there is a very simple relationship between the concentrations of these two solutions... this is a basic one -step calculation!)
__________ grams 12
