Introduction to Spectroscopy A Rainbow of Colors光谱颜色的彩虹的介绍

Student Workbook Pg. 1 of 52 STUDENT WORKBOOK TABLE OF CONTENTS: Introduction to Spectroscopy: A Rainbow of Colors Experiment……….………………………………………………..……..3-10 Pre-Lab Questions…………………………………..……….7 Data & Observations…………………………………………7-8 Results & Discussion………………………………..……….9 Advanced Questions………………………………..……….10Accuracy and Precision: Quantitative Analysis of Aspirin Experiment …………………………………………………..………..…11-19 Pre-Lab Questions…………………………………..……….16 Data & Observations…………………………………………16-17 Results & Discussion………………………………..……….18 Advanced Questions………………………………..……….19How an Acid-Base Indicator Works & Determination of its Equilibrium Constant Experiment ……………………………………………………….……...20-27 Pre-Lab Questions…………………………………..……….24 Data & Observations…………………………………………24-25 Results & Discussion………………………………..……….26 Advanced Questions………………………………..…….….27Kinetics of the Reaction of Methylene Blue with Ascorbic Acid Experiment ……………………………………………………….……...28-35 Pre-Lab Questions…………………………………..……….32 Data & Observations…………………………………………32-33 Results & Discussion………………………………..……….34 Advanced Questions………………………………..………..35LeChatlier’s Principle: Shifting Equilibrium of an Acid-Base Indicator Experiment ……………………………………………………………...36-42 Pre-Lab Questions…………………………………..……….39 Data & Observations…………………………………………39-40 Results & Discussion………………………………..……….41 Advanced Questions………………………………..….…….42Quantification of Tartrazine in Mt. Dew: A Beer’s Law Experiment Experiment …………………………………………………………….....43-51 Pre-Lab Questions…………………………………..……….48 Data & Observations…………………………………………48-49 Results & Discussion………………………………..……….50 Advanced Questions………………………………..………..51 Pg. 2 of 52 Introduction to Spectroscopy: A Rainbow of Colors Experiment for the i-LAB? Spectrophotometer Introduction Spectroscopy is the study of the interaction of light (or electromagnetic radiation) with matter, and it is one of the most important tools used by scientists to study atoms and molecules. This experiment uses a visible spectrometer to investigate food dyes and aids our understanding of how we perceive color. Light or electromagnetic radiation (EMR) is a form of energy which is made up of perpendicular electric and magnetic fields. When EMR interacts with matter, it can be absorbed, reflected, refracted (bent), or pass through it. How EMR interacts with the matter is dependent on the type of EMR and the chemical and physical properties of the material, so scientists can study this interaction to learn about the properties of the material. Scientists classify EMR based on its energy, wavelength, or frequency. The energy of one photon or particle of light can be related to its frequency and wavelength using Equations 1 and 2. E = energy (J)-23h = Planck’s constant (6.626*10 J/s)E = hν(Equation 1)-1ν = frequency (Hz or s)8c = speed of light (3.0*10 m/s) E = hc/λ (Equation 2)λ = wavelength (m) EMR consists of all radiation, spanning from the low energy radiowaves through the ionizing energy of gamma rays. The visible portion of the EMR spectrum, which is the focus of this lab, includes all EMR that can be detected by the human eye and includes wavelengths ranging from 380 to 750 nm. The entire EMR spectrum is shown in Figure 1. As illustrated in this figure, only a very small portion of this radiation can be perceived as light by the eye. Increasing Wavelength (λ) in nm 400500600700 Visible spectrum Increasing Frequency (ν) 24222018161412108642010101010101010101010101010ν (Hz) FMAMγ raysX raysUVIRMicrowaveLong wave radiosRadio waves λ (m)-16-14-12-10-8-6-4-20246810101010101010101010101010 Increasing Wavelength (λ) Figure 1. Entire spectrum of electromagnetic radiation (EMR) with inlay of visible region. Pg. 3 of 52 It is thought that color is perceived by three different types of cones in the human eye: one tuned to longer wavelengths (reds), one to medium wavelengths (greens), and one to shorter wavelengths (blues). Visible light excites one or more types of cones, and the brain combines the signals from all excited cones into a perceived color. An object appears blue when the cones of the eye sensitive to shorter wavelengths are excited by incident blue light. A blue object reflects mostly blue light and absorbs most of the other colors of light. White is detected when all the colors are reflected equally, and black is perceived when all the colors are absorbed equally. The color circle shown in Figure 2 depicts an organization of colors based on color theory. All shades of various colors can be made from mixtures of the three primary colors, red, yellow, and blue. Complementary colors are shown as opposites on the color circle and can be determined by the color that would result when mixing the other two primary colors. For example the complementary color of red is green, a mixture of yellow Figure 2. Color circle. and blue. Many food dyes are chemically synthesized colorants. A number of these have been approved for use by the FDA, and they are routinely added to cereals, candies, fruits, and even meats to enhance appeal and even out the coloring. This experiment uses three common food colorings, Red 40 (allura red), Yellow 5 (tartrazine), and Blue 1 (brilliant blue), to prepare a rainbow of colored solutions. The chemical structures of these compounds are shown below. Red 40Yellow 5Blue 1 The i-LAB? visible spectrometer will be used to analyze each colored solution. With this spectrometer, the sample is irradiated with white (visible) light. The light that passes through the sample is separated into its component colors and detected. The transmittance, or the fraction of light that passes through the sample, at each wavelength is measured. The measured Pg. 4 of 52 transmittance signal can be converted to an absorbance measurement using Equation 3. A = absorbance A = -log Τ = -log (I/I)(Equation 3)0T = transmittance I = intensity of transmitted light I = intensity of incident light0 Figure 3. The relationship between incident light (I) and transmitted light (I).0 For the sample analysis, you will be collecting an absorbance spectrum, which is a plot of absorbance versus wavelength. An example absorbance spectrum of green food coloring is provided in Figure 4. The lambda max (λ) is the wavelength of light at which the sample most max strongly absorbs, and it is often used to characterize a sample. In the case of green food coloring, there are two major peaks at approximately 420 nm and 630 nm. The λ for this max sample is approximately 630 nm, the more intense peak. 0.25 0.2 0.15 0.1 absorbance 0.05 0 400450500550600650700 wavelength (nm) Figure 4. Sample Absorbance spectrum of a green food coloring solution. Pg. 5 of 52 Objectives ,Create a rainbow of colored solutions using food dyes of the three primary colors. ,Collect visible absorbance spectra for those solutions.,Determine the λ for each solution.max ,Relate the wavelength of the λ and the overall absorbance spectrum to the perceived max color of the solution. Materials VIS i-LAB? Pasteur pipets or droppers Red 40 solution Small beakers, test tubes, or vials Yellow 5 solution 10 ml cuvettes or round vials Blue 1 solution Methods *Asterisk directs you to record information in the Data and Observations section. 1.Collect approximately 20 mL of each dye solution and a beaker full of water. 2.Prepare approximately 10 mL of orange, green, and purple solutions from mixtures of the *primary color solutions. 3.Prepare an approximate 1:4 dilution of the red solution by mixing one dropper full of the red solution with three droppers full of water. 4.Collect an absorbance spectrum for each colored solution using method SPEC 1. Follow method prompts and record the data as directed. *a.Make a sketch of each spectrum and record its λ.max 5.Compare the spectra of the original red solution and the diluted red solution. *a.Record the intensity of each at the λ.max Refer to Basic i-LAB Instructions for general method and sampling directions. Pg. 6 of 52 Introduction to Spectroscopy: A Rainbow of Colors Experiment for the i-LAB Microspectrometer Name ______________________________ Pre-lab Questions 1. How does light interact with an object that is red in color such that we perceive the object as being red? 2. Briefly describe the parts of a spectrometer and what they do.3. In an absorbance spectrum, what is the λ?max Data and Observations Table 1. Preparation of colored solutions. Brief description of how the solution Description of the color and intensity of was preparedthe solution Orange Green Purple Pg. 7 of 52 Table 2. Spectra of the rainbow of solutions (from 400 to 700 nm) = Orange λ = Red λmaxmax 1.0 –1.0 – AbsAbs Wavelength (nm)Wavelength (nm)Yellow λ = Green λ = maxmax 1.0 –1.0 – AbsAbs Wavelength (nm)Wavelength (nm)Blue λ = Purple λ = maxmax 1.0 –1.0 – AbsAbs Wavelength (nm)Wavelength (nm)Table 3. Comparison of red and dilute red solutions.Overlaid sketch of red and dilute red spectra Red Dilute Red 1.0 – ,max Abs Abs at λmax Wavelength (nm) Pg. 8 of 52 Results and Discussion 1. Complete the following data table Color of solution,(nm)Color of light at λmax max Red Yellow Blue 2. Refer to the table in question 1. How does the color of the solution and the color of λmax compare? Explain why those colors are not the same. 3. How do the spectra collected for the red and blue solutions compare to the spectrum collected for purple? Justify any similarities and differences. 4. When you mix red and yellow together to form orange, is a new compound created? Explain. Use the collected spectra to explain why we perceive the solution to be orange in color.5. Compare the spectra collected for the red and dilute red solutions. Justify any similarities and differences. 6. What is the relationship between absorbance and concentration? Justify this relationship. Pg. 9 of 52 Advanced Questions: 7. Is it possible to determine the concentrations of two dye solutions that are added together?Explain how to do that if: a) if there is no overlap in peak Absorbance values, and b) if there is an overlap in spectra. 8. Looking at the chemical structures of the dyes (Red Dye 40, Yellow Dye 5, and Blue Dye 1) there is abundance of pi (unsaturated) electrons. Explain how the pi electrons contribute to color and interact with EMR in general. 9. Looking at the structure of the same dyes, sulfonate salt groups (S03-) are present in each molecule. Explain why the sulfonate groups are needed for dye solubility in the water. Speculate on the dye solubility in water without the sulfonate groups. 10. Dye purity and a narrow peak width are important in many applications, including laser and graphic art dyes. Please comment on why purity and a narrow peak width are desired for these specialty dyes? Pg. 10 of 52 Accuracy and Precision: Quantitative Analysis of Aspirin Experiment for the i-LAB? Visible Spectrophotometer Introduction Aspirin is one of the most widely used medications in the world with approximately 40,000 metric tons being consumed annually. Acetylsalicylic acid (ASA) is the active ingredient in aspirin and can be used to treat aches, pains, and fever and as a preventative medication for heart disease. Like any other pharmaceutical, it is critical that one aspirin pill contains the accurate amount of ASA of the correct purity. Quality control testing ensures that such standards are met. ? Aspirin (ASA) (Acetylsalicylic Acid; 2-(acetyloxy)benzoic acid) In this experiment, you will perform quality control testing of a commercial aspirin sample by quantifying how much ASA is in an aspirin tablet. Aqueous solutions of ASA are colorless and so do not inherently absorb light in the visible range. However, through a colorimetric reaction, a colored complex is formed which does absorb light in the visible range. As a first step, the ASA is hydrolyzed with sodium hydroxide, forming a sodium salt of the salicylic acid (2-hydroxybenzoate). Then the SA salt complexes with iron forming a violet complex as follows: +-NaOOHOO HO+-2HOOOOOHNaHO++ CHCH33 acetic acid2-(acetyloxy)benzoic acid2-hydroxybenzoate +-NaOO2+OOFeHO2Pg. 11 of 5243+HOFe2+OH6OH Yellow Violet The intensity of the violet color is proportional to the concentration of SA in solution (which is proportional to ASA). You will use an i-LAB? spectrometer to measure the transmittance of the colored solutions at 530 nm. In spectroscopy, transmittance refers to the fraction of light at a given wavelength that passes through a sample. In other words, it is the ratio of transmitted light ). Figure 1 and Equation 1 describe this relationship. intensity (I) to the incident intensity (I0 T = I/I(Equation 1)0 T = transmittance I = intensity of transmitted light I = intensity of incident light0 Figure 1. The relationship between incident light (I) and transmitted light (I).0 The measured transmittance signal can be converted to an absorbance measurement using Equation 2. For quantitative purposes, absorbance is the more useful measure. Absorbance is the inverse log of transmittance; its usefulness arises from the fact that absorbance is directly proportional to concentration. A = -log(T) (Equation 2) Quantitative measurements are often assessed based on their accuracy and precision. Accuracy refers to the closeness to the true value, and precision is a measure of reproducibility of repeat measurements. Consider a dart board analogy with the bull’s eye representing the true value. Figures 2a-c demonstrate three possible scenarios regarding accuracy and precision that can result with four replicate measurements symbolized by four darts. Figure 2a. Figure 2b. Figure 2c. Accurate and precise. Precise but not accurate. Neither accurate nor precise. Pg. 12 of 52 Standard deviation, s, (Equation 3) and relative standard deviation, RSD, (Equation 4) are two measures used to assess the precision of measurements. In both cases, a smaller number is indicative of better precision. RSD can be considered a normalized standard deviation, and this makes it possible to compare the precision of various data sets using RSD. In general, measurements with RSDs less than 5% can be considered as having “excellent” precision, RSDs between 5-10% are indicative of “good” precision, and RSDs between 10-15% are indicative of “fair” precision. One way to assess accuracy is by calculating a percent error (Equation 5) of the average experimental result relative to the true value. Similar generalities hold true with percent error as with RSD. In terms of accuracy, a percent error of less than 5% can be considered “excellent,” percent error between 5-10% is “good,” and percent error between 10-15% is “fair.” (Equation 3) RSD = s/average *100%(Equation 4) % error = (theoretical-experimental)/theoretical *100%(Equation 5) N = number of measurements x = one data pointi x = average experimental result ? = summation In this experiment, you will react hydrolyzed ASA with aqueous iron to form a colored complex. The transmittance of this complex will be monitored using the i-LAB microspectrometer. A series of ASA standards and three aspirin tablets will be analyzed. The ASA content in the aspirin tablets will be quantified using a calibration curve, and you will assess the accuracy and precision of your results. Pg. 13 of 52 Objectives ,Explain how colorless compounds can be analyzed using visible spectroscopy.,Use spectroscopic measurements to determine the amount of ASA in a commercial aspirin. ,Assess the accuracy and precision of your data. ,Understand sources of experimental error in the analytical process.Materials VIS i-LAB with round vial adaptorHot plates 0.04 M Fe complex solution10 mL graduated cylinders Solid acetylsalicylic acid 1.0 mL measuring pipette or micropipette Commercial aspirin tablets (325 mg)20.0 mL calibrated vials or 1.0 M Sodium hydroxide solution 25.0 mL volumetric flasks 125 mL Erlenmeyer flasksGlass funnels 250 mL volumetric flasksWeighing paper Analytical BalancePolycarbonate cuvette or 25 mm round vial Methods *Asterisk directs you to record information in the Data and Observations section. Preparing the Standard 6.Weigh approximately 0.4 grams of reagent grade acetylsalicylic acid (ASA) to the nearest milligram (0.001g) on a piece of tared weighing paper. Transfer the sample to a 125 mL Erlenmeyer flask. Add 10 mL of 1.0 M NaOH to the ASA and heat the mixture to boiling on a hot plate to hydrolyze the ASA. Carefully remove flask and let cool. 7.Quantitatively transfer the resulting solution of sodium salicylate salt solution to a 250 mL volumetric flask through a glass funnel and dilute to mark with DI water. Mix thoroughly. This is the stock solution. 8.Prepare dilutions of the stock solution in calibrated 25 ml* vials or volumetric flasks with a measuring pipette or a micropipette according to the table below: StandardStock Soln. 0.04 M Fe Complex DI Water (mL)Soln. (mL) blank0.000.50Dilute to 25 mL mark A0.100.50Dilute to 25 mL mark B0.200.50Dilute to 25 mL mark C0.300.50Dilute to 25 mL mark D0.400.50Dilute to 25 mL mark E0.500.50Dilute to 25 mL mark * Can also use 20 ml vial or volumetric flask for all standards and samples. Pg. 14 of 52 Preparing the Sample 1.Add one commercial aspirin tablet to a 125 mL Erlenmeyer flask. Add 10 mL 1.0 M NaOH solution, heat to boiling, safely remove from hot plate and then cool. 2.Quantitatively transfer the sample to a 250 mL volumetric flask. Dilute and mix. Transfer a 0.50 mL portion into a calibrated vial (or volumetric flask), and add 0.50 mL of the Fe complex solution. Dilute to volume (20.0 or 25.0 mL) with DI water and mix. This is sample ASP1 (Asprin 1). 3.Prepare two additional samples using the same procedure (ASP2 and ASP3). Standard and sample analysis 9.Using method ASPIRIN, first blank and then measure the transmittance of Standards A- E and samples ASP1-3 at the λ, 530 nm. Follow method prompts and record the data max *as directed. *10.Convert transmittance measurements to absorbance measurements. 11.Create plots of transmittance versus concentration and absorbance versus concentration *from the standard data. 12.Using the equation of the line from the calibration curve and the sample absorbance *results, calculate the aspirin concentration in your samples. (Make sure that your sample absorbance measurements are in the calibration range.) *13.Convert the concentration results to mass (mg) aspirin per tablet.Refer to Basic i-LAB Instructions for general method and sampling directions. Pg. 15 of 52 Accuracy and Precision: Quantitative Analysis of Aspirin Experiment for the i-LAB Microspectrometer Name ______________________________ Pre-lab Questions 1. How is it possible to analyze a colorless compound like acetylsalicylic acid using visible spectroscopy? 2. What is the difference between accuracy and precision?3. What is quality control testing and why is it important for pharmaceuticals? Data and Observations Table 1. ASA standard results. StandardConc. of Transmittance Absorbance at *****ASA (M)at 530 nm530 nm A B C D E **Calculated based on MV = MV. Show sample calculation below:1122 *** Calculated based on A = -log(T). Show sample calculation below: Figure 1. Using the standard data from Table 1, plot the absorbance versus concentration. Add a linear trend line and include the equation of the line on the graph. This is your calibration curve. (Attach graph.) Pg. 16 of 52 Equation of the line: ________________________Table 2. Aspirin sample results. SampleTransmittancAbsorbance Conc. of Mass (mg) of *********ASA (M) ASA in tablete at 530 nmat 530 nm ASP1 ASP2 ASP3 ** Calculated based on A = -log(T). Show sample calculation below:***Calculated based on the equation of the line of the calibration curve. Show sample calculation below: ****Calculated in two steps: (1) MV = MV and (2) based on molarity, total volume, and formula weight (180.2 1122 g/mol). Show sample calculations below: Table 3. Data analysis. Show sample calculations below:True value **(mg ASA from label)% error = (theoretical-experimental)/theoretical *100% Experimental average (mg ASA)**% Error*** ***Standard Deviation (s) ********RSD (%)RSD = s/average *100% Pg. 17 of 52 Results and Discussion 1. Are your results accurate? Explain. 2. Are your results precise? Explain. 3. Is there evidence of systematic error in your data? In other words, are all your results higher than the label claim, or are all your results lower than the label claim? Give a potential reason for any bias that is demonstrated in your data. 4. Which results do you think are more correct, your experimental data or the label claims? Explain. 5. Why did you analyze three different aspirin tablets rather than analyzing the same aspirin tablet three times? Pg. 18 of 52 Advanced Questions: 6. List possible sources of error and “questimate” where the largest source of error was; explain your reasoning? 7. Explain why- the SA solution was colorless? – the 0.04 M iron complex was yellow? - the iron-SA complex turned violet? 8. Preparation of aspirin (ASA) usually results in very high purity (>99%). However, other more expensive drugs are occasionally subject to the addition of cheap fillers. Explain why good QC is needed in the chemical and pharmaceutical industries. Pg. 19 of 52 How an Acid-Base Indicator Works & Determination of its Equilibrium Constant Experiment for the i-LAB? Visible Spectrometer Introduction Acid-base indicator solutions change color as the pH changes. The indicator solutions are typically prepared using a compound that is itself a weak acid or base. Weak acids and bases are defined as acids and bases that do not completely ionize. They only partially ionize and thus reach a state known as chemical equilibrium. Chemical equilibrium occurs in a reversible reaction when the concentrations of the reactants and products in an aqueous solution (or pressures of gases) become constant. When a reaction reaches equilibrium, the reaction has not stopped. Rather, the rate of the forward reaction becomes equal to the rate of the reverse reaction. This means that there is no net change in the concentrations of both the reactants and products. Since the reaction is still proceeding (just at the same rate in both directions), this state is known as dynamic equilibrium. Any reversible reaction can be characterized by its equilibrium constant or K. For a general eq chemical reaction: aA + bB ? cC + dD (aq)(aq)(aq)(aq) its equilibrium constant is:cdabK = [C][D]/[A][B]eq Pure solids and liquids do not affect equilibrium and therefore are not included in the equilibrium constant expression. Bromothymol blue (BTB) is an example of an acid-base indicator. Like other acid-base indicators, BTB is a weak acid whose protonated and deprotonated forms are two different colors. BTB is a complex molecule, and so the acid-base reaction is simplified by using HBTB -to represent the protonated form and BTB to represent the deprotonated form of BTB. The acid- base reaction of BTB can be represented by the following equilibrium: -+HBTB + HO ? BTB + HO (aq)2(l)(aq)3(aq) yellow colorless blue colorless Protonated Bromothymol blue (HBTB) Deprotonated Bromothymol blue (BTB-) yellow blue Pg. 20 of 52 Accordingly, its equilibrium constant can be written: -+ = K = [BTB][HO]/[HBTB] Keqa3 Since BTB is a weak acid, its equilibrium constant also represents the acid dissociation constant or K.a +Because the K value is a constant, at high concentrations of HO (low pH), the concentration eq3-of HBTB must be high (and the concentration of BTB must be low). This effect is visible; at low pH, bromothymol blue solutions are yellow in color. Similarly, at low concentration of +-HO (high pH), the concentration of BTB must be high (and the concentration of HBTB must 3 be low). As expected, bromothymol blue solutions are blue at high pH. In this lab, you are going to experimentally determine the equilibrium constant for bromothymol blue by measuring the equilibrium concentrations of the hydronium ion, the protonated form of BTB, and the deprotonated form of BTB. To determine the hydronium concentration, you will +measure the pH of the solution since pH = -log[HO]. You will use the i-LAB spectrometer and 3-Beer’s Law to determine the equilibrium concentrations of HBTB and BTB. Beer’s Law relates absorbance (A) to concentration (c) using the following equation: A = εbc -1-1The molar absorptivity or extinction coefficient, ε (typically in units of Mcm), is a measure of how strongly a particular compound absorbs light at a specified wavelength. The symbol b represents the path length that the light travels through and is typically 1.0 cm for standard spectrometer designs. By measuring the absorbance of a standard at a known concentration, you can calculate the compound’s extinction coefficient. The calculated extinction coefficient can thus be used with subsequent absorbance measurements to determine the concentration in unknowns. In the case of bromothymol blue, each form (protonated and deprononated) has a unique λmax (wavelength of maximum absorbance) and extinction coefficient. Standards of BTB will be -prepared such that only one form predominates in each standard (i.e., [HBTB] >> [BTB] or -[BTB] >> [HBTB]). Pg. 21 of 52 Objectives ,Use Beer’s Law to calculate extinction coefficients for both the protonated and deprotonated forms of bromothymol blue. ,Use pH and absorbance measurements to determine equilibrium concentrations of the three relevant species. ,Calculate an experimental K value for bromothymol blue and compare it to the eq literature value. ,Explain how an acid-base indicator works. Materials VIS i-LAB with cuvette or round vial pH meter adaptoracidic solution (0.10 M HCl) Stock bromothymol blue (BTB) standard basic solution (0.10 M NaOH)-4 solution (0.04% (w/v) = 6.4*10M)Pasteur pipets or droppers pH 4 buffer solution 5.0 mL graduated pipets (0.010 M acetate buffer)25 mL volumetric flasks pH 7 buffer solution 25 or 50 mL beakers (0.010 M phosphate buffer)10 mL and 25 or 50 mL graduated cylinderspH 10 buffer solution Cuvettes or 25 mm glass vials (0.010 M carbonate buffer) Methods *Asterisk directs you to record information in the Data and Observations section. Standard Preparation and Analysis 14.Prepare HBTB standard by measuring 2.0 mL stock BTB solution into the 25 mL volumetric flask and diluting to mark with pH 4 buffer. Cap and invert several times to ensure thorough mixing. -15.Prepare BTB standard by measuring 1.0 mL stock BTB solution into the 25 mL volumetric flask and diluting to mark with pH 10 buffer. Cap and invert several times to ensure thorough mixing. -16.Using method BTB, collect an absorbance spectrum of the HBTB and BTB standards. Record the λand absorbance at the λfor each. Follow method prompts and record max max *the data as directed. 17.Using Beer’s Law, calculate the molar extinction coefficient for both the protonated and *deprotonated forms of BTB. Pg. 22 of 52 K Determinationeq 1.Prepare a BTB solution by combining approximately 2 mL of BTB and 18 mL of pH 7 buffer (exact volumes not necessary). 2.Record the pH of this solution and then using method KEQ measure the absorbance -values at both λ’s(for both HBTB and BTB). Follow method prompts and record the max *data as directed. 3.Change the pH of this solution by adding 1-10 drops of acidic or basic solution, maintaining a pH between 6-8. 4.Record the new pH of this solution and then measure the absorbance values at both λ’s max-*(for both HBTB and BTB) 5.Repeat pH change and absorbance measurements two more times for a total of four *trials. *6.Calculate the experimental K based on the collected data.eq Refer to Basic i-LAB Instructions for general method and sampling directions. Pg. 23 of 52 How an Acid-Base Indicator Works & Determination of its Equilibrium Constant Experiment for the i-LAB Microspectrometer Name ______________________________ Pre-lab Questions 1. Why is a chemical system at equilibrium described as “dynamic”?2. What is an extinction coefficient? How can it be used to determine concentration?3. What color would you expect bromothymol blue solutions to be near neutral pH? Explain.Data and Observations Table 1. Bromothymol blue standard results. *****SolutionColor of Conc.,Absorbance εmax -1-1 Solution(M)(nm)(Mcm)at λmax BTB in pH 4 buffer (HBTB form) BTB in pH 10 buffer -(BTB form) **Calculated based on MV = MV. Show sample calculation below:1122 ***Calculated based on A = εbc. Show sample calculation below: Pg. 24 of 52 Table 2. Bromothymol blue pH and absorbance results for equilibrium measurements. Color of pHAbsorbance Absorbance -Solutionfor HBTBfor BTB Trial 1 Trial 2 Trial 3 Trial 4 Table 3. Equilibrium concentrations and experimental K values.eq +*****-*******[HO][HBTB][BTB]K3eq Trial 1 Trial 2 Trial 3 Trial 4 Average **+ Calculated based on pH = - log[HO]. Show sample calculation below:3 ***Calculated based on A = εbc. Show sample calculation below: ****-+Calculated based on K = [BTB][HO]/[HBTB]. Show sample calculation below:eq3 Pg. 25 of 52 Results and Discussion 1. What is the color of the solution buffered at pH 4? What is the color of light at the λ? max What is the relationship between these colors? Explain. 2. To determine the extinction coefficient for the protonated form of bromothymol blue (HBTB), the solution is buffered at a pH of 4. Why? 3. The equilibrium data is collected at a pH between 6 and 8. Why? (Hint: consider the color of the solution.) -84. The literature value for the K of bromothymol blue is 7.9*10. Are your results in good a agreement with the literature value? Explain. 5. How does the color of a bromothymol blue solution relate to its pH? Use your knowledge of the equilibrium reaction of bromothymol blue to explain how it works as an acid base indicator. Pg. 26 of 52 Advanced Questions: 6. In the experiment there were two forms of BHB that depended on pH. If the pH were decreased to <1.0, another form of BHB is made that happens to be red. Knowing that the differences in BHB forms depends on the oxygen groups and an acid adds H+, draw the new form of BHB at pH <1.0 below: 7. Approximately 4% of the general population is color-blind, meaning that they cannot differentiate between certain colors. Discuss why a spectrometer would be good to use for people that are color-blind and for those that are not. 8. If you had a blue solution of BHB at pH 4.0 and starting adding water to it how would the color change and why? (Hint: look at the equilibrium equation and use common sense) Pg. 27 of 52 Kinetics of the Reaction of Methylene Blue with Ascorbic Acid Experiment for the i-LAB? Spectrophotometer Introduction Kinetics is the study of chemical reaction rates. Through kinetics, one can study how experimental conditions affect the rates of chemical processes. Kinetics can also provide information about the mechanism of a reaction (how the reaction occurs on the molecular scale). In this experiment, you will determine the rate law of the reaction between methylene blue and ascorbic acid (vitamin C). Methylene Blue Ascorbic Acid The rate law for a chemical reaction is a relationship that describes how the concentrations (or pressures for gases) of reactants affect the overall reaction rate. For a general reaction:aA + bB? cC (aq) (aq) (aq) its simple rate law is: xyrate = k[A][B] where k represents the rate constant, a term for all factors (except concentration) that affect the rate. [A] and [B] represent the concentrations of the reactants, and x and y represent the order of the reaction with respect to reactants A and B, respectively. For instance, if x = 0 and y = 1, we would say that the reaction is zero order in reactant A, first order in reactant B, and first order (the sum of all orders) overall. (Note: the order of the reaction is independent of the stoichiometry of the reaction!) The order of the reaction gives us information about how the reaction proceeds and provides insight into the mechanism of the reaction. In a zero order reaction, the rate is independent of concentration and so the rate does not change as the reaction proceeds. In a first order reaction, the rate is directly proportional to concentration. In this case, the rate of the reaction will decrease as the reaction proceeds because the reactant is being consumed. In the case of a second order reaction, the rate proceeds relative to the squared concentration. With second order reactions, the reactant is consumed more rapidly as the reaction proceeds, and the rate slows relative to the squared concentration of the reactant. A plot of rate versus concentration is not linear. Pg. 28 of 52 In a zero order reaction, the reaction rate is independent of reactant concentration. The mechanism for this type of reaction may involve a surface reaction or a thermal decomposition. First order reactions are unimolecular reactions. In other words, the rate-determining step involves only one molecule. Two molecules are involved in the rate-determining step of bimolecular reactions. Kinetic characteristics of zero, first, and second order reactions are summarized in Table 1. Table 1. Summary of the various kinetic characteristics of zero, first, and second order reactions. stnd Order2 OrderZero Order1 Conc.Conc.Conc. AAA A TimeTimeTime RateRateRate Conc. AConc. AConc. A slope = -kslope = -kslope = k Conc. ln [A]1/[A] A A TimeTimeTime concentration independentunimolecular mechanismbimolecular mechanism (surface phenomenon or +decomposition) AACAC Pg. 29 of 52 Ascorbic acid, also known as vitamin C, is a molecule that has antioxidant properties. Antioxidants prevent the damaging effects of free radicals and other oxidizing species by participating in reduction-oxidation reactions with these species. The reaction between ascorbic acid and methylene blue mimics the type of reduction-oxidation reaction that occurs between an antioxidant (ascorbic acid) and an oxidizing compound (methylene blue). (Methylene blue, however, is not a free radical species. It is just an oxidizing agent.) Methylene blue solutions are bright blue in color. However, when methylene blue solutions are exposed to antioxidants like ascorbic acid, these solutions lose their color as the methylene blue is reduced to leucomethylene blue, a colorless compound. This reaction and its corresponding rate law are simplified below: Ascorbic acidAscorbic acid (reduced form)(oxidized form)xy MB + AA-H ? MB-H + AArate = k[MB][AA-H]methylene blueleucomethylene blue (blue)(colorless) ? HHHH ++++HHHH Leucomethylene Blue Ascorbic Acid (oxidized) With spectroscopy, we can monitor the loss of blue color over time to obtain information about the kinetics of this reaction. In reactions where [AA-H] >> [MB], the concentration of ascorbic acid remains essentially constant throughout the course of the reaction. In this case, the rate law can be simplified to: xyrate = k[MB]where k= k[AA-H]11 If, under these circumstances, x = 1, we say that the reaction is pseudo first order in methylene blue. In this experiment, you will use the i-LAB? spectrometer to monitor the rate of reduction of methylene blue by ascorbic acid. By monitoring the rate at various concentrations of ascorbic acid, you can determine the overall rate law for this reaction. Pg. 30 of 52 Objectives ,Monitor the kinetics of the reaction between methlyene blue and ascorbic acid with spectroscopy. ,Determine the order of the reaction with respect to methylene blue and ascorbic acid.,Calculate the rate constant and establish the overall rate law for this reaction.Materials VIS i-LAB with cuvette or round vial Deionized water adaptor1.0 mL transfer pipettes Methylene blue (MB) stock solution 5.0 mL measuring pipettes-4 (4.0*10M)10.0 mL volumetric flasks or calibrated vials Ascorbic acid (AA) stock solution Small beakers (0.050 M) Methods *Asterisk directs you to record information in the Data and Observations section. 18.Collect approximately 10 mL of MB stock solution and 20 mL of AA stock solution.19.Scan a blank (deionized water) as directed by method KINETICS. Follow the method prompts and record the data as directed. 20.Prepare Trial 1 solution by combining 1.0 mL MB and 5.0 mL AA and diluting to 10.0 mL. a.Quickly mix Trial 1 solution, aspirate into sampler, and begin collecting absorbance readings at 665 nm as prompted by method KINETICS. b.This method will collect six absorbance readings at approximately 30 sec intervals. *c.Return to MENU and view LOG to retrieve data. 21.Repeat this procedure for three additional trials based on the following table: TriaVol. MB Vol. AA With Water, lSolution (mL)Solution (mL)Dilute to (mL) 21.04.010.0 31.03.010.0 41.02.010.0 Pg. 31 of 52 Kinetics of the Reaction of Methylene Blue with Ascorbic Acid Experiment for the i-LAB Microspectrometer Name ______________________________ Pre-lab Questions 1. What do we mean when we say “order” of a reaction?2. How does the concentration of a reactant affect the rate of a chemical reaction? 3. Describe briefly how a spectrometer can be used to monitor the reaction between methylene blue and ascorbic acid. Data and Observations Table 1. Kinetic data for Trial 1. Table 2. Kinetic data for Trial 2.TimeAbsorbance ln (A)1/ATimeAbsorbance ln (A) (sec)(A)(sec)(A) 00 3040 6080 90120 120160 150200 Pg. 32 of 52 Table 3. Kinetic data for Trial 3. Table 4. Kinetic data for Trial 4. TimeAbsorbancln (A)TimeAbsorbancln (A) (sec)e (A)(sec)e (A) 00 5060 100120 150180 200240 250300 Figures 1-3. Use the data in Table 1 to make a series of plots: absorbance versus time, ln(A) versus time, and 1/A versus time (attach graphs).Figures 4-7. Make a plot of ln(A) versus time for each trial. Add a linear trendline and include the equation of the line on the graph. (Attach graphs.)Table 5. Data used to determine the order of the reaction with respect to ascorbic acid. Trial[Ascorbic Acid] kk if zero k if first order k if second order 1** ***12(M)(-slope)order (k=k/(k=k/[AA])(k=k/[AA])1110[AA]) 1 2 3 4 **Calculated based on MV = MV. Show sample calculation below:1122 ***Obtained from the equation of the line in Figures 4-7. Pg. 33 of 52 Results and Discussion 1. Refer to Figures 1-3 generated from your experimental data. Is the reaction between ascorbic acid and methylene blue zero, first, or second order in methylene blue? Explain. 2. What is the relationship between ascorbic acid concentration and k? Justify this relationship.1 3. Refer to the calculated rate constant data in Table 5. In which data set does the calculated k value remain essentially constant? Based on this finding, is the reaction between ascorbic acid and methylene blue zero, first, or second order in ascorbic acid? 4. Write out the complete rate law, including the rate constant, for the reaction between ascorbic acid and methylene blue. 5. Use the order of the reaction to hypothesize a general mechanism for how this reaction proceeds. Pg. 34 of 52 Advanced Questions: xy6.The rate = k[A][B] has x and y as the rate orders of the kinetic reaction. In the experiment and in the examples there have been zero, first and second order reactions. Can a rate order be a non-integer number (e.g., 0.89) as opposed to 1? Explain why or why not? 7.Kinetics may also be influenced by temperature and pressure. Please comment on these two factors in general as to how they may affect the reaction kinetics.8.Catalysts are materials that assist reactions without wearing themselves out; often by lowering activation energies. Comment, in general, on how a catalyst may affect reaction kinetics. Pg. 35 of 52 LeChatlier’s Principle: Shifting Equilibrium of an Acid-Base Indicator Experiment for the i-LAB? Spectrophotometer Introduction Chemical equilibrium is defined as a state in a reversible reaction where there is no net change in the concentration of either products or reactants. For a particular reaction, this state is characterized by its equilibrium constant or K. Whether the reaction starts off as mostly eq reactants or as mostly products, it will reach the same “balance point” as defined by its K. This eq balance point or equilibrium occurs when the rate of the forward reaction is equal to the rate of the reverse reaction. When these rates become equal, the concentrations of reactants and products remain constant. LeChatlier’s Principle is a key concept in chemistry that helps us understand reversible chemical systems. LeChatlier’s Principle states that when a stress is applied to a system at equilibrium, the system shifts its equilibrium to partially counteract the change. In other words, a new balance point or equilibrium state is reached that minimizes the stress on the system. However, that balance point is still defined by the K for the reaction. The K is a constant (for eqeq a given temperature and pressure); it does not change. Consider the following general reaction at equilibrium. aA + bB? cC (aq) (aq) (aq) Now, if we apply a stress to the system by adding more of reactant B, how does that affect the equilibrium? We’ll use a seesaw analogy as depicted in Figure 1 to explain LeChatlier’s Principle. Figure 1a. Chemical system at Figure 1b. Stress applied to system. More Figure 1c. New equilibrium state achieved. equilibrium.of reactant B has been added. System is Reaction proceeds in the forward direction initially not at equilibrium.to counteract stress. To counteract the addition of reactant B, the reaction proceeds in the forward direction (shifts to the right). Some of reactants A and B have been consumed (as depicted by a decrease in size), and more of product C has been generated (as depicted by an increase in size) to establish a new balance point. In other words, equilibrium has been re-established as the concentrations have adjusted back to the K value.eq In an aqueous system, a stress that induces a shift in equilibrium can be anything that affects the concentration of a reactant or product. Temperature does not affect concentration. However, Keq values are constant only for a particular temperature (generally measured at 25?C). So Pg. 36 of 52 temperature can cause a shift, but that shift is a result of a new K rather than a response to a eq change in concentration. Bromothymol blue (BTB) is a weak acid whose color depends on its form. The “protonated” -state, simplified as HBTB, is yellow in color, and the “deprotonated” form, simplified as BTB, is blue in color. Its reaction in water is characterized by the following equilibrium: -+-+HBTB + HO ? BTB + HO K = K = [BTB] [HO]/[HBTB] (aq)2(l)(aq)3(aq)eqa3 yellow colorless blue colorless Figure 2 shows the absorbance spectra of BTB at three different pH values. At low pH (Fig. 2a), we see mostly the protonated form with a λ(wavelength of maximum absorbance) at 431 nm. max At high pH (Figure 2c), we see mostly the deprotonated form with a λat 616 nm. Around max neutral pH (Figure 2b), both the protonated and deprotonated forms are evident. 0.70.90.60.80.60.50.70.50.40.60.40.50.30.4 0.3absorbanceabsorbanceabsorbance0.20.30.20.20.10.10.1000400450500550600650700400450500550600650700400450500550600650700wavelength (nm)wavelength (nm)wavelength (nm)- Figure 2a. BTB solution at pH 4. Figure 2b. BTB solution at pH 7. Both Figure 2c. BTB solution at pH 10. BTB-HBTB form predominates.HBTB and BTB forms are evident.form predominates. Because absorbance is directly proportional to concentration (as governed by Beer’s Law), we can use spectroscopy to monitor equilibrium shifts. For instance, if a stress induces the BTB reaction to proceed in the reverse direction, more reactant is generated and product is consumed. This change is verified through spectroscopy by a measured increase in the HBTB peak and - decrease in the BTBpeak. In this experiment, you will use spectroscopy to monitor equilibrium shifts for a bromothymol blue solution. You will use your knowledge of LeChatlier’s principle to justify the observed shifts. In addition, you will calculate the K of bromothylmol blue at two different temperature eq extremes. Objectives ,Use spectroscopy to detect equilibrium shifts in an acid-base indicator solution. ,Explain how components participating in a chemical reaction can affect the equilibrium. ,Explain how components not directly participating in a chemical reaction can affect the equilibrium. ,Explain how temperature affects chemical equilibrium. Pg. 37 of 52 Materials VIS i-LAB with cuvette or round vial Salt (solid NaCl) adaptor150 mL beakers Stock bromothymol blue (BTB) solution 10 mL and 100 mL graduated cylinders-4(0.04% (w/v) = 6.4*10M)Small test tubes and racks pH 7 buffer solution Hot water bath (hot plate or bunsen burner) (0.01 M phosphate buffer)Ice water bath acidic solution (0.10 M HCl)Polycarbonate cuvettes or 25 mm round vialsbasic solution (0.10 M NaOH) Methods *Asterisk directs you to record information in the Data and Observations section. 22.Using 150 mL beakers, prepare a hot water bath and an ice bath.23.Prepare 100 mL of BTB solution at equilibrium by mixing approximately 5 mL of BTB solution with approximately 95 mL of pH 7 buffer. *24.Make predictions about how each of the following changes would affect the equilibrium: Add acid, add base, add salt, add buffer, heat solution, cool solution.25.Label 7 small test tubes: control, acid, base, salt, buffer, hot, cool.26.For the test tube labeled buffer, add 4 mL of your prepared BTB solution and 4 mL of pH 7 buffer. Cap and invert several times to mix. 27.Fill the remaining 6 test tubes about ? full with your prepared BTB solution. 28.Alter each BTB solution as directed by the following table: Test tubeAdjustment controlNo change acidAdd approximately 5 drops of 0.10 M HCl and mix baseAdd approximately 5 drops of 0.10 M NaOH and mix saltAdd approximately .20 g NaCl (small scoop) and mix hotPlace in boiling water bath at least 10 min coolPlace in ice water bath at least 10 min 29.Using method LECHAT, measure the absorbance of each solution at 431 nm (λfor the max protonated form of BTB) and 616 nm (λfor the deprotonated form of BTB). Follow max *method prompts and record the data as directed. *30.Calculate the K for the hot and cool samples.eq Pg. 38 of 52 Add samples to cuvette or round vial. Place in adaptor and measure referring to Basic i-LAB Instructions for general method and sampling directions. Pg. 39 of 52 LeChatlier’s Principle: Shifting Equilibrium of an Acid-Base Indicator Experiment for the i-LAB Microspectrometer Name ______________________________ Pre-lab Questions 1. Re-state Le Chatlier’s Principle in your own words. 2. What do we mean when we say “shift” in equilibrium? 3. If a BTB solution at equilibrium shifts to the right, how would the color of the solution change? How would this equilibrium shift affect the absorbance spectrum of BTB?Data and Observations Table 1. Predictions on how various changes might affect the equilibrium of BTB.ChangeJustify your predictionPrediction (shifts left , , shift right , , or no change) Addition of acid Addition of base Addition of salt Addition of buffer Heating of solution Cooling of solution Pg. 40 of 52 Table 2. Absorbance measurements for both the protonated and deprotonated forms of BTB solutions at equilibrium. Test TubeObserved color AbsorbancAbsorbancShift relative to control (left , , right , , or no change)of solutione at 431 e at 616 nmnm ControlNA Acid Base Salt Buffer Hot Cool Table 3. K data for the hot and cool samples.eq EquatioHotCool n+ *-7 -7 [HO]11.0*10M1.0*10M3 [HBTB]2-[BTB]3 K4eq*Assuming a pH = 7.0. Show sample calculations below:+(1) pH = log[HO]34 -1-1(2) A = εbc (where ε = 2.25*10Mcm for HBTB, A from Table 2)4 -1-1-(3) A = εbc (where ε = 1.82*10Mcm for BTB, A from Table 2)-+(4) K = [BTB] [HO]/[HBTB]eq3 Pg. 41 of 52 Results and Discussion 1. What is the purpose of the control in this experiment? 2. Why is a spectrometer used for this lab rather than just using our eyes to observe the color change? 3. How does the addition of base affect the BTB equilibrium? Explain.4. How does the addition of buffer affect the BTB equilibrium? Explain.5. How does the addition of salt affect the BTB equilibrium? Explain.6. How does temperature affect the K value? Explain.eq 7. Make a note of how many predictions you got right!) Pg. 42 of 52 Advanced Questions: 8. Why is a pH neutral solution of bromothymol blue (BTB) green in color?9. Consider the following general reaction at equilibrium. aA + bB? cC (aq) (aq) (aq) If more of cC (aq) was added to an equilibrium mixture, what would happen to the concentration of aA (aq)? Of bB(aq)? To the K? eq 10. Comment on the structure and color of the forms of bromothylmol blue below; meaning what causes the color in general, and what changes in the molecule’s structure to cause a color shift? ?? BHB- (blue) HBHB (yellow) Pg. 43 of 52 Quantification of Tartrazine in Mt. Dew: A Beer’s Law Experiment Experiment for the i-LAB? Spectrophotometer Introduction Mt. Dew provides a refreshing pick-me-up to many students. The yellow dye tartrazine (Yellow Dye Number 5) gives this soda its special color. This lab experiment uses spectroscopy and the Beer-Lambert Law to determine the amount of tartrazine in Mt. Dew. Tartrazine (Yellow Dye Number 5) Quantitative spectroscopy is based on the fairly straightforward principles of transmittance and absorbance, shown in figures 1a and b. In one sample (Fig. 1a), a concentrated colored dye is poured into water to form a transparent, but colored solution. In the other sample (Fig. 1b), another colored solution is prepared, but with twice the amount of water. The solution on the right is lighter in color because it is less concentrated and more light can pass through. In the solution on the left, the dye is more concentrated in the sample causing less light to pass through. Actually, the solution that is twice as concentrated has two times as many dye molecules that absorb more light, or transmit (pass through) less light, and thus appears darker. One could assume that the solution depicted in Fig. 1b would be half as dark as that in Fig. 1a, but that is not always a valid assumption. Fig. 1a.Fig. 1b Colored and Transparent Dye SolutionColored and Transparent Dye Solution at Half the Concentration as that in Fig 2a. In spectroscopy, Transmittance refers to the fraction of light at a given wavelength that passes through a sample. In other words, it is the ratio of transmitted light intensity (I) to the incident intensity (I).0 Pg. 44 of 52 T = I/I0 For quantitative purposes, Absorbance is often a more useful measure. Absorbance is the inverse log of Transmittance; its usefulness arises from the fact that Absorbance is directly proportional to concentration. A = -log(T) A α concentration The Beer-Lambert Law or Beer’s Law characterizes this relationship: A = εbc -1-1The molar absorptivity or extinction coefficient (Mcm), ε, is a constant value for a particular compound at a specified wavelength. The path length (b) that the light travels through the liquid is also a constant, typically a 1.0 cm cuvette for a standard spectrometer. (Special cuvettes of 0.10 cm have been used for more concentrated solutions. Most cuvettes are made of clear polycarbonate and are disposable. High precision cuvettes are made of quartz.) Since ε and b are constants, Absorbance is directly proportional to the concentration of the dye or colored molecules increases. A calibration curve (Beer’s Law plot) can be constructed of Absorbance versus concentration (Fig. 2) showing a direct linear relationship. To generate such a curve, a series of standard solutions must first be made. The standards will contain a known and specific amount of our analyte (compound or species being tested). The absorbance of each standard is measured at a particular wavelength. Typically the wavelength of maximum peak height (λ) is used from max spectra such as Fig. 3, since it is the most sensitive wavelength and provides the most precise results. 0.9 0.8f(x) = 40.54x + 0.01 R? = 0.990.7 0.6Sample absorbance 0.5 0.4 0.3Absorbance 0.2 0.1 0 0.0000.0050.0100.0150.0200.025 Concentration (M) Fig. 2. Example calibration curve. In this case, a sample Fig. 3. Example spectra showing increasingabsorbance of 0.56 corresponds to a conc. of Absorbance with concentration, and maximumapproximately 0.014 M. peak height. Pg. 45 of 52 ese results are used to plot the calibration curve. Then the absorbance of the sample (having Th an unknown concentration of analyte) is measured. This absorbance result (y-value) is used in the equation of the line from the calibration curve to calculate the concentration of the analyte in the sample (x-value). Figure 2 shows the relationship between a sample result and a calibration curve. There are three important factors to keep in mind when using Beer’s Law to perform quantitative spectroscopic analysis: (1)The standard absorbance measurements should bracket the sample absorbance. (In other words, standard concentrations should be chosen such that the standard Absorbance values fall both above and below the sample absorbance.) (2)Beer’s Law becomes non-linear at higher concentrations (concentrations > 0.01 M). As a rule of thumb, quantitative work is typically done using Absorbance measurements less than 1.0. (3)The standard matrix (everything in the standard except the analyte) should match the sample matrix as closely as possible to minimize matrix effects. Matrix effects occur when the matrix falsely increases or decreases the analytical signal. For instance, a compound in the sample matrix may absorb light at the analytical wavelength. This would result in a falsely elevated absorbance reading for the sample which would suggest that there was more analyte in the sample than there actually was. To offset matrix effects, an alternative calibration technique known as standard addition can be used. With standard addition, the sample is spiked with a known amount of standard. The premise with this technique is that the sample matrix affects the standard (in the spike) to the same degree that it affects the analyte in the sample. A simple approach to standard addition is to perform a single spike. The concentration of the analyte in the sample can be calculated using the following relationship: A = absorbance from initial (unspiked) solutionx A = absorbance from final (spiked) solutionS+X [X]A[X] = concentration of analyte in initial solutionii x =[X] = concentration of analyte in final solution f[S]+ [X]A[S] = concentration of standard in final solution ff f S+X In this lab, you will prepare a series of tartrazine standards and construct a calibration curve from Absorbance measurements collected using the i-LAB? spectrometer. Mt. Dew samples will be prepared and will also be spiked with standard for standard addition analysis. The tartrazine content in the Mt. Dew sample will be quantified using both the calibration curve and standard addition. Pg. 46 of 52 Objectives ,Prepare a series of tartrazine standards and measure their absorbance at the λ.max ,Construct a calibration curve of Absorbance versus concentration for the tartrazine standards. ,Use the calibration curve and sample absorbance to quantify the tartrazine content in a Mt. Dew sample. ,Analyze the Mt. Dew sample using standard addition methodology.,Compare the results obtained using the calibration curve to those obtained using standard addition. Materials VIS i-LAB? with cuvette or vial adaptor25ml volumetric flasks Stock tartrazine standard 1.0ml graduated pipets (or micropipets)-4 (3.0*10 M in water)10.0ml transfer pipets pH 4 buffer solution Large test tubes and test tube racks (0.01 M acetate buffer)Small and large beakers Degassed Mt. Dewcuvettes or vials Methods *Asterisk directs you to record information in the Data and Observations section. Standard preparation and analysis 31.Collect approximately 5mL of stock tartrazine standard and approximately 500 ml buffer32.Label a series of test tubes: A – E 33.Prepare Standard A by measuring 0.20 ml stock standard into the 25 ml volumetric flask and diluting to mark with buffer. Cap and invert several times to ensure thorough mixing. Pour into the test tube labeled “A.” 34.Repeat procedure for the remaining standards based on the following table: StandardVol. of stock tartrazine With buffer, standard (ml)dilute to (ml) B0.4025.0 C0.6025.0 D0.8025.0 E1.0025.0 35.Using method BEERS LAW, measure the absorbance of Standards A-E at the λ, 428 nm.max* Follow method prompts and record the data as directed. (Use the buffer as your blank.) Pg. 47 of 52 Plot a calibration curve of absorbance versus concentration for the standards. Add a linear trend 36.*line to the graph. Sample preparation and analysis 37.Collect approximately 50ml of Mt. Dew 38.Label a series of test tubes: MD1, MD2, MD3. 39.Prepare MD1 by measuring 10.0 mL Mt. Dew into the 25 mL volumetric flask and diluting to mark with buffer. Cap and invert several times to ensure thorough mixing. Pour into the labeled test tube. 40.Repeat procedure two additional times to prepare samples MD2 and MD3. *41.Using method TARTRAZINE, measure the absorbance of the three samples at 428 nm. Follow method prompts and record the data as directed. 42.Using the equation of the line from the calibration curve and the sample absorbance results, *calculate the tartrazine content in the Mt. Dew sample. Standard addition and analysis 43.Label two test tubes SA1 and SA2. 44.Prepare SA1 by measuring 10.0ml Mt. Dew into the 25ml volumetric flask and diluting to mark with buffer. Cap and invert several times to ensure thorough mixing. Pour into the labeled test tube. 45.Prepare SA2 by measuring 10.0 ml Mt. Dew and 0.4 ml stock tartrazine standard into the 25ml volumetric flask and diluting to mark with buffer. Cap and invert several times to ensure thorough mixing. Pour into the labeled test tube. *46.Using method MD STD ADD, measure the Absorbance of SA1 and SA2 at 428 nm. Follow method prompts and record the data as directed. *47.Using the standard addition equation, calculate the tartrazine content in the Mt. Dew sample. Refer to Basic i-LAB Instructions for general method and sampling directions. Pg. 48 of 52 Quantification of Tartrazine in Mt. Dew: A Beer’s Law Experiment Experiment for the i-LAB? Spectrometer Name ______________________________ Pre-lab Questions 1. Why do darker solutions absorb more light than lighter solutions?2. What is a calibration curve and how is it used?3. How does the method of standard addition offset matrix effects?Data and Observations Table 1. Tartrazine standard results. ** Calculated based on MV = MV. 1122StandardConc. of Absorbance at Show sample calculation below:**Tartrazine (M)428 nm A B C D E Figure 1. Calibration curve (attach graph). 2Equation of the line: _____________________________R = ____________ Pg. 49 of 52 Table 2. Mt. Dew sample results. SampleAbsorbance Conc. of Conc. of at 428 nmTartrazine (M) Tartrazine (M) *****in dilute samplein original sample MD1 MD2 MD3 Average tartrazine concentration in original Mt. Dew sample = _________**Calculated based on the equation of the line of the calibration curve. Show sample calculation below: ***Calculated based on MV = MV. (Hint: M was just calculated.) Show sample calculation below:11222 Table 3. Standard addition results. SampleAbsorbanceConc. of ** at 428 nmTartrazine (M) SA1A = [X] =xi SA2A =[S]=S+Xf ***Tartrazine concentration in original Mt. Dew sample = _________**SA2 tartrazine concentration same as Standard B. **SA1 tartrazine concentration calculated using standard addition equation. Show calculation below: ***Calculated based on MV = MV. (Hint: M was just calculated.) Show calculation below:11222 Pg. 50 of 52 Results and Discussion 1. Do your standard absorbance values bracket your sample absorbance values? If not, how could you fix this problem? Why is it important that your sample absorbance measurements fall in your calibration range? 2. Are all your standard and sample absorbance values less than 1.0? If not, how could you fix this problem? Why should quantitative absorbance measurements be less than 1.0? 23. The correlation coefficient, or R value, for the equation of the line of your calibration curve is a measure of goodness of fit. In other words, it is a measure how closely the measured data 2points fit to the equation of the line. Do your standard data represent a good fit (R > 0.90)? Explain. Justify any bad standard data points. 4. Experimental results suggest that the actual tartrazine content in Mt. Dew is 0.47 mg per can (355 ml). Convert this content to a molarity number and calculate the percent error for each quantification method. (The formula weight of tartrazine is 534.4 g/mol.) Actual content = _____________M Quantification MethodTartrazine Conc. in % error from **Mt. Dew (M)true value Calibration curve (average) Standard addition **Calculated based on % error = (theoretical – experimental)/theoretical *100%. Show sample calculation below:5.Which method, calibration curve or standard addition, do you think is better for the quantification of tartrazine in Mt. Dew? Justify your choice. Pg. 51 of 52 Advanced Questions: 6.The Beer’s Law experiment that was done was performed at specific concentrations of tartrazine, using a bracketing method. If the concentration of the colored dye continued to increase or decrease, how would the calibration curve look? Please draw it in the chart below: ? ? ABSORBANCE ? Tartrazine?Concentration 438?nm?(M/L) 7.Tartrazine is an “azo” dye, meaning that it has an N=N bond. Azo dyes are fairly common dyes, since they are relatively easy to make compared to other dyes. The dye is made from two molecules having primary amines, R-NH, and combining them under mild acidic conditions and heat to produce an azo bond. Please circle the azo bond in the tartrazine (product) and add the reactants for making tartrazine below: + Tartrazine (Product) 8.Tartrazine is yellow because of the interactions of light with the molecule’s pi (π) bonds, represented generally as R=R, where R can be C, N, or O. The more pi bonding, or “conjugation”, the longer the wavelength of light absorption. With that in mind, discuss if you think the tartrazine reactants (7. above) should absorb in the visible region, and explain your reasoning. Also, discuss what would happen to the wavelength of absorbance if the tartrazine molecule increased in conjugation. Pg. 52 of 52