Spectroscopy Introduction and Applications. A Guide for Use with Ocean Optics OceanView Software

Transcription

1 Spectroscopy Introduction and Applications A Guide for Use with Ocean Optics OceanView Software August 18, 2015

2 Preface Students can now study basic scientific principles on the same world-class equipment used by leading researchers in university and government labs like NASA. Advances in electro-optics, high-speed array detectors, inexpensive optical fibers and powerful computers have made optical spectroscopy the sensing technique of choice for many real-world applications. The development and availability of scientific instrumentation and methods have changed in an equally dramatic way. In the past, cutting-edge scientific instruments were expensive research devices accessible only to well-funded research and development enterprises. Gradually, the technologies filtered into general laboratory use, application-specific instruments and now into the educational setting. Our knowledge of spectroscopy is based on more than 20 years of experimentation in a wide array of disciplines ranging from art to applied physics. All of these fields have their roots in education with educators teaching their students the basics of the field. Thousands of science educators have utilized Ocean Optics spectrometers to create real-world, exciting experiments to teach their students and enrich their lives with a greater understanding and appreciation for science. It is important that today's science and engineering students appreciate the capabilities of optical sensing, the fundamental physics of the measurement process, the design trade-offs inherent in selecting and integrating components and the discipline required to produce quality results. The goal of this lab manual is to provide a vehicle to allow future scientists to study the fundamentals of spectroscopy using modern instrumentation. We would like to offer special thanks to the educators who contributed to this lab manual as part of the ongoing Ocean Optics grant program. Note to Educators: If you would like to contribute to future compilations, please send an to [email protected]. Trademarks: All products and services herein are the trademarks, service marks, registered trademarks or registered service marks of their respective owners. Limit of Liability: Ocean Optics has made every effort to ensure that this manual as complete and as accurate as possible, but no warranty or fitness is implied. The information provided is on an as is basis. Ocean Optics, Inc. shall have neither liability nor responsibility to any person or entity with respect to any loss or damages arising from the information contained in this manual. 2

3 Table of Contents Preface... 2 Table of Contents... 3 Spectroscopy Concepts... 5 Overview... 5 Light... 5 Wavelength and Energy... 6 The Interaction of Light with Matter... 7 Types of Spectroscopy... 8 Color and Wavelength Spectroscopy Instrumentation The Beer-Lambert Law Instrumentation Absorbance of Light vs. Concentration Materials Safety The Absorbance of Commercial Food Dyes Data Analysis Discussion Questions The Absorbance Spectrum of Chlorophyll Materials Safety The Absorbance of Food Coloring The Absorbance of Chlorophyll Discussion Questions Determining the Rate law of a Chemical Reaction: Kinetics of Crystal Violet Bleaching Reaction of Crystal Violet with OH Materials Safety Measuring the Kinetics of Crystal Violet Bleaching Data Analysis Extension Emission Spectra... 41

4 Materials Safety Measuring Light Source Emission with a Spectrometer Data Analysis Spectrophotometric Analysis of Commercial Aspirin Prelab Questions Materials Safety Hydrolyzing the ASA and Aspirin samples Data Sheet Data Sheet Calculations Post Lab Questions Spectrophotometric Characterization of Spice Extracts Materials Safety Procedure Data and Analysis

5 Spectroscopy Concepts Overview Scientific discoveries are based on observations. Scientists look for patterns in what they see, hear, feel, smell and taste to formulate theories and make predictions. Originally, scientists depended solely on their own senses to make observations. But as science has evolved, scientists have developed instruments to extend their observational powers beyond our sensory limits. Telescopes have enabled astronomers to see more of the sky and vastly improve our understanding of the heavens. Likewise, microscopes have enabled biologists to view ever smaller parts of living organisms in their quest to understand living systems. Astronomers are only limited by the size of the telescopes they can build and the distorting effects of the earth's atmosphere. As technological developments have allowed for bigger mirrors and space-based platforms, astronomers have been able to see ever further into space and make more and more discoveries. Unfortunately, the situation is very different going in the other direction. There is a physical limit to the size of objects that can be "seen." This limit is due to the nature of light itself. Light Light is a type of electromagnetic radiation consisting of little packets of energy called photons with both particle and wave-like properties. As shown in the complete electromagnetic spectrum in Figure 1, light in the visible region (~400 to 700 nm) makes up only a small region of the entire spectrum of electromagnetic radiation. Figure 1: Electromagnetic Spectrum 5

6 It is the wave properties of light that limit our ability to use light to create images. For any given wavelength, images can be formed for objects larger than the wavelength of light used to visualize the image. Therefore, for visible light (wavelengths between 4 x 10-7 and 7 x 10-7 m) it is impossible to form images of atoms with sizes on the order of 10-9 m. Very large molecular assemblies such as chromosomes (DNA molecules coated wi th protein molecules) are the smallest objects we can image with visible light. The challenge faced by chemists, biochemists and microbiologists is studying what happens at the atomic and molecular level without actually being able to physically visualize atoms and molecules. Fortunately, even though we cannot capture images of atoms and molecules, we can use light to learn more about them. This is because atoms and molecules interact with light providing detailed information on their structure, composition and interactions. The technique for measuring the interaction of light with matter is referred to as spectroscopy. It is arguably the most powerful tool available to scientists to study the molecular world around them. Spectroscopy techniques are used universally in science at the intersection of the disciplines of chemistry, biology, engineering and physics. Wavelength and Energy Our understanding of the nature of light is a relatively recent development. For a long time the different forms of electromagnetic radiation were thought to be individual phenomena. Thus, we have the collection of common names ending in -wave and ray for the various wavelength ranges. This is because the energy of a photon is inversely related to its wavelength. The shorter the wavelength, the higher the energy of the photon. Where: E = Energy of the photon in joules λ = Wavelength in nanometers h = Planck's constant c = Speed of light The enormous wavelength range for known electromagnetic radiation yields photons with energy covering an equally wide range. It is this energy that determines the effect of the photon when it interacts with matter. For example, radio frequency photons have very small energies enabling us to saturate our atmosphere with them without affecting our environment. The amount of energy they impart to whatever absorbs them is almost negligible (you don t get dents in your car from listening to one of the many available radio stations). Infrared photons have enough energy to heat objects and, as a result, they make great heat lamps. Ultraviolet photons have enough energy to break chemical bonds and can cause molecular 6

7 rearrangements resulting in effects like sunburn and genetic damage. X-rays are very energetic and readily break even the strongest bonds causing significant molecular destruction. For this reason the medical use of X-rays is destructive of living tissues and must be done carefully and only in extremely small doses. The Interaction of Light with Matter It is the electrons in atoms and molecules that typically absorb and emit photons of light. It is worth noting that gamma rays are energetic enough to interact with atomic nuclei and generate photons of light, but we'll leave that to the physicists to pursue. When an electron absorbs low energy photons, like radio frequencies, the spin of the electron is flipped. This effect is used in nuclear magnetic resonance spectroscopy (NMR) and can also be u sed to generate the images from magnetic resonance imaging (MRI). When an electron absorb s infrared, visible and ultraviolet photons they change energy level. All electrons have a series of energy levels they can occupy. The lowest energy level is referred to as the "ground state." The highest level is the "ionization energy" or the energy required to completely remove the electron from the influence of the nucleus. In order for an electron to move from one level to a higher level it must absorb energy equal to the difference in the levels. Likewise, to move to a lower level the electron must give up energy equal to the difference. Because there are a limited number of levels the electron can occupy, there are limited amounts of energy it can absorb or give up. A more detailed discussion can be found in Chemistry and Chemical Reactivity, Chapter 7 by Kotz and Treichel. Their figure 7.13 is a detailed presentation of the electronic transitions possible for the simplest atom, the hydrogen atom. Figure 2 is a diagram of the most common transitions possible for a sodium atom. The 3s to 4p transition is in the ultraviolet range. The 3p to 3d transition is in the infrared range. And the 3s to 3p transition is in the orange region of the visible spectrum. This line is the source of the characteristic color of sodium vapor lamps. Figure 2: Common Electronic Transitions for Sodium There are a number of ways an electron can gain or lose energy. The ones of interest here are the absorption or emission of light. An electron can absorb a photon of light that strikes it only if that photon has the exact energy to change the electron to a higher allowed energy level. An electron already at a higher level can emit a photon of light having exactly enough energy to 7

8 change that electron to one of its lower allowed levels. Note that an electron in the ground state cannot emit any photons as it already has the lowest possible energy. The magnitude of the difference in allowed energy levels determines which kinds of light can be used to study particular atoms and molecules. While spectroscopy is conducted in nearly all regions of the electromagnetic spectrum, practical considerations make the infrared, visible and ultraviolet regions the most useful in chemical laboratories. Infrared spectroscopy is particularly useful for studying the vibration of bonds between carbon, hydrogen, oxygen and nitrogen atoms that predominate in organic compounds. Thus, infrared spectroscopy is a key tool of the organic chemist. Infrared spectra can indicate the presence of particular functional groups in unknown organic compounds by the presence of characteristic features. They can also be used to confirm the identity of compounds by comparison with known spectra. Reference books containing thousands of spectra of known organic compounds are available for this purpose. Visible light spectroscopy is useful for studying some organic compounds and elements that have electrons in d-orbitals, such as transition metals. Ultraviolet spectroscopy is useful for studying some organic compounds and most biological samples. All proteins have useful ultraviolet spectra as do nucleic acids. Furthermore, UV spectroscopy can be used to follow biochemical reactions and this tool is commonly found in biochemical laboratories. In clinical laboratories, ultraviolet spectroscopy is often the means for making quantitative determinations on plasma and urine samples. Types of Spectroscopy Spectroscopy is the study of the interaction of light with matter. There are two distinct aspects of this interaction that can be used to learn about atoms and molecules. One is the identification of the specific wavelengths of light that interact with the atoms and molecules. The other is the measurement of the amount of light absorbed or emitted at specific wavelengths. Both determinations require separating a light source into its component wavelengths. Thus, a critical component of any spectroscopic measurement is breaking up of light into a spectrum showing the interaction of light with the sample at each wavelength. Light interacts with matter in many ways. Two of the most common interactions are light that is absorbed by the atoms and molecules in the sample and light that is emitted after interacting with the atoms and molecules in the sample. 8

9 Absorption Spectroscopy Absorption spectroscopy is the study of light absorbed by molecules. For absorbance measurements, white light is passed through a sample and then through a device (such as a prism) that breaks the light up into its component parts or a spectrum. White light is a mixture of all the wavelengths of visible light. When white light is passed through a sample, under the right conditions, the electrons of the sample absorb some wavelengths of light. This light is absorbed by the electrons so the light coming out of the sample will be missing those wavelengths corresponding to the energy levels of the electrons in the sample. The result is a spectrum with black lines at the wavelengths where the absorbed light would have been if it had not been removed by the sample. Emission Spectroscopy Emission spectroscopy is the opposite of absorption spectroscopy. The electrons of the sample are promoted to very high energy levels by any one of a variety of methods (e.g., electric discharge, heat, laser light, etc.). As these electrons return to lower levels they emit light. By collecting this light and passing it through a prism, it is separated into a spectrum. In this case, we will see a dark field with colored lines that correspond to the electron transitions resulting in light emission. In Figure 3, hydrogen absorption and emission spectrum lines are shown with a continuous spectrum for reference (top spectrum). The absorption (middle spectrum) and emission (bottom spectrum) of a substance share the same wavelength values. In the absorption spectrum, wavelengths of absorbance appear as black lines on a colored field. In the emission spectrum, the wavelengths of emission appear as colored lines on a black field. Figure 3: Hydrogen Emission and Absorption /astro211/lectures/hydrogen_atom.html Qualitative Spectroscopy One of the most useful aspects of spectroscopy derives from the fact that the spectrum of a chemical species is unique to that species. Identical atoms and molecules will always have the 9

10 same spectra. Different species will have different spectra. For this reason, the spectrum of a species can be thought of as a fingerprint for that species. Qualitative spectroscopy is used to identify chemical species by measuring a spectrum and comparing it with spectra for known chemical species to find a match. As an example, consider the discovery of the element helium. It was first observed, not on the earth, but in the sun! In 1868 the French astronomer, Pierre Jules César Janssen, was in India to observe a solar eclipse when he detected new lines in the solar spectrum. No element known at that time would produce these lines and so he concluded that the sun contained a new element. This initiated a search for the new element on planet earth. By the end of that century, the new element had been identified in uranium ores and was named helium after the Greek word for the sun (Helios). Today, spectroscopy finds wide application in the identification of chemical species. Quantitative Spectroscopy Quantitative spectroscopy is one of the quickest and easiest ways to determine how many atoms or molecules are present in a sample. This is because the interaction of light with matter is a stoichiometric interaction. At any given temperature, the same number of photons will always be absorbed or emitted by the same number of atoms or molecules in a given period of time. This makes spectroscopy one of the few techniques that can provide a direct measure of the number of atoms or molecules present in a sample. Quantitative emission spectroscopy requires heating samples to very high temperatures to enable electrons to emit light. Most often, this is done by feeding the sample into a burner flame. As a result, it is not practical for use with most molecular compounds but is frequently employed for elemental analysis. Absorption spectroscopy is performed by passing light of all wavelengths through a sample and measuring how much of the light at each wavelength is absorbed. The statement made above that "the absorption spectrum will appear as black lines on a colored field" is a considerable oversimplification. The interactions of atoms and molecules with water molecules make the absorbance of light in solutions a very complex phenomenon. Nevertheless, the patterns are repeatable and predictable, thus making them useful. By making absorbance measurements at various wavelengths and then plotting the result, one can create what is known as an absorbance spectrum. 10

11 In Figure 4 the absorbance spectra for different heme-containing proteins is shown. Even though the proteins are closely related, the absorbance spectra are distinct enough to enable discrimination of the different proteins. Absorbance spectra are like fingerprints. Each compound has its own unique spectrum. In some cases the spectrum can be used to identify the presence of certain compounds in a sample. More often, it is used to determine the amount of a compound present. Figure 4: UV-Vis Absorbance Spectra for Heme Proteins Due to the nature of the electronic changes that give rise to absorption spectra, the peaks are generally broad. Therefore, absorption spectroscopy is less useful than other molecular techniques for the purpose of identifying compounds. For example, infrared or Raman spectroscopy is much better for identifying the component species when compared to UV- Visible absorption spectroscopy. Color and Wavelength The visible region is a good place to begin a discussion of spectroscopy because color vision is a critical feature of our everyday world. Our perception of color is the eye's response to light of different wavelengths. When photons of a narrow wavelength range interact with our retina, we perceive the effect as color. The apparent color of an object is due to the wavelengths of the photons of light reaching our eyes from that object. This is true whether the object is emitting its own light or reflecting light from another source. In a sense, our eyes operate like a spectrophotometer. White light is a mixture of light of all wavelengths (colors). When white light strikes an object and is completely reflected, we see equal amounts of light of all colors and perceive the object to be white. When all light striking an object is absorbed, no light enters our eyes and we perceive the object to be black. A sheet of paper is white because all light striking it is reflected and none is absorbed. The print on the paper is black because all light striking it is absorbed. None is reflected. We perceive color when some wavelengths of light are reflected (or transmitted in the case of a solution) more than others. There is a rather complex pattern to the absorption of light by colored objects. The statement that "an object appears red because all red light is reflected and all other light is absorbed" is a considerable oversimplification. In fact, varying amounts of light of different wavelengths are 11

12 absorbed in most colored objects and the color we perceive is more closely related to the color that is most absorbed rather than to the color that is reflected. The brain assigns color to an object by a process known as complementary color vision. According to this theory, all colors of light have a complementary color. This is often displayed through the use of a "color wheel" like the one shown in Figure 5. A color and its complementary color are opposite each other on the color wheel. The perception of color occurs when the optic nerve and the brain compare the amount of light of a particular color with the amount of its complementary color. If the two amounts are the same, we see gray or white. If not, we see color. A fire extinguisher appears red in white light because more bluegreen light (the complementary color of red) is being absorbed than any other color. Of course, this also means that more red light is being reflected than its complementary color, blue-green. For all other colors, relatively equal amounts of each color, and its complement, are being reflected. Spectroscopy Instrumentation Figure 5: Color Wheel A large variety of instruments are used to perform spectroscopy measurements. They differ greatly in the information they provide. What they all have in common is the ability to break light up into its component wavelengths. Spectroscopes A spectroscope is the simplest spectroscopic instrument. It functions to take light from any source and disperse it into a spectrum for viewing with the unaided eye. In Figure 6, a diagram of a simple spectroscope is shown. The light from the source passes through the slit and into the prism where it is dispersed into a spectrum. The telescope is used to focus on the light coming out of the prism. The third arm contains a wavelength scale that can be superimposed over the spectrum by shining a white light into it. Spectroscopes are useful for determining 12

13 what wavelengths of light are present in a light source, but they are not very useful for determining the relative amounts of light at different wavelengths. Spectroscopes are most commonly used for qualitative emission spectroscopy. Spectrometers Figure 6: Diagram of a spectroscope A spectrometer is a spectroscope with a meter or detector so it can measure the amount of light (number of photons) at specific wavelengths. It is designed to provide a quantitative measure of the amount of light emitted or absorbed at a particular wavelength. Some spectrometers are constructed so that the wavelength can be varied by the operator and the amount of radiation absorbed or transmitted by the sample determined for each wavelength individually. Others have a fixed light dispersing element (e.g., diffraction grating) that disperses multiple wavelengths of light onto a multi-element detector. Using a spectrometer, it is possible to measure which wavelengths of light are present and in what relative amounts. Spectrometers are common in astronomy where they are used to evaluate light collected by telescopes. They are the only source of information we have about the chemical composition of the universe outside our own solar system. The diagram of a simple spectrometer is shown in Figure 7. Light enters the spectrometer via the entrance slit and then passes through several parts: an objective lens, a grating, and an exit slit. This combination of parts functions as a monochromator, a device that selects only one color (actually, a narrow band of wavelengths) from all of the wavelengths/colors present in the source. A particular wavelength is selected, using the wavelength control, by adjusting the angle of the grating. This works because different wavelengths of light reflect off the grating at different angles. The net result is the separation of white light into a "rainbow," much like the effect of light transmitted through a prism of glass. The selected wavelength is at the center of the narrow band of wavelengths passing through the slit. The light then strikes a detector that 13

14 generates a voltage in proportion to the intensity of the light hitting it. That voltage is then used to drive a read-out device that is designed to provide data in a format such as intensity. Figure 7: Diagram of a simple spectrometer As with all electronic devices, the design and operation of spectrometers has been greatly impacted by the developments of the latter half of the 20th century. Perhaps the most crucial was the development in the early 1970s of the Charged Coupled Device (CCD). Originally conceived as a new mode of data storage, it was soon discovered that CCDs held great promise as imaging devices. An imaging device is something that electronically mimics what photographic film does. CCDs consist of a number of elements between which charge can be shifted. In an image sensor, light falling on the array of elements produces a pattern of charges corresponding to the image. This image can then be electronically transported to some other location, such as a monitor, and reconstructed. CCDs were first employed to replace photographic plates in telescopes. The first such device was installed on the I-meter telescope at Kitt Peak National Observatory in Today, CCDs are the detectors that make digital cameras not only possible, but affordable. Soon after its successful application to astronomical problems, it was determined that CCDs could greatly enhance the performance of spectrometers. This was achieved by replacing both the exit slit and detector with a CCD array. Now, it was no longer necessary to measure light intensity one wavelength at a time. The number of wavelengths that can be monitored simultaneously is determined by the number of elements in the CCD array. Figure 8 is a schematic of a spectrometer outfitted with a CCD array. The array generates an output that can be used to reconstruct the intensity of light striking each of the elements in the array. This output can be sent to a monitor or a printer for display. The output is instantaneous across the spectrum. No longer is it necessary to "scan" back and forth across the spectrum to identify light intensity at individual wavelengths. 14

15 Figure 8: Diagram of a CCD Spectrometer Spectrophotometers An instrument that includes a light source is known as a spectrophotometer. It is constructed so that the sample to be studied can be irradiated with light. The wavelength of light incident on the sample can be varied and the amount of light absorbed or transmitted by the sample determined at each wavelength. From this information, an absorption spectrum for a species can be obtained and used for both qualitative and quantitative determinations. Spectrophotometers measure the amount of light transmitted through a sample. Once the transmission for a sample is measured, it can be converted into other values. Percent transmittance (%T) is the ratio of the transmitted light (I) to the incident light (I o ) expressed as a percent. T I X100% I o The %T calculation is easy to design into a spectrophotometer and was a common output in primitive early spectrometers. Percent transmittance can also be converted to absorbance, A or Abs, which is directly related to the molar concentration of the chemical species in the sample. Absorbance values are calculated from % T values using the following equation. A = Log 1 % T Io log I 15

16 There is an assumption inherent in the calculation of either %T or absorbance. The assumption is that all light not transmitted to the detector is absorbed by the chemical compounds in the solution. Two other possibilities exist. One is that the light is being scattered by the solution. Light interacting with any particle that is larger than its wavelength can scatter light. Because of this, samples containing solid material or that are cloudy or turbid in nature are difficult to analyze using a spectrophotometer. Samples encountered in the commercial world (biological fluids, soil solutions, etc.) are often cloudy and extra steps must be employed before analysis by absorption spectrophotometry can begin. Another consideration is that light can be scattered or absorbed by the container used to hold the solution. Care must be taken to ensure that the sample cells are clean and free from fingerprints so they do not affect the measurement. The cells must also be constructed of absolutely transparent material free of defect. If measurements are to be made below 350 nm, the cells must be made of quartz or other materials that readily transmit UV light. The Beer-Lambert Law The relationship between absorbance and concentration is known as the Beer-Lambert Law, or sometimes simply Beer's Law, Where: A = ε l c A = Measured absorbance, c = Concentration of the absorbing species, l = Pathlength of the sample (width of the cuvette) ε = Proportionality constant known as the molar absorptivity with units of (M -1 cm -1 ). The molar absorptivity (ε) is constant for a specific chemical compound at a specific wavelength. For most compounds, there is typically at least one wavelength where ε reaches a maximum. This wavelength is often chosen to carry out absorption spectrophotometry of that compound. For example, consider the absorbance spectra for hemoglobin shown in Figure 4. There are three wavelengths in the visible range that would be suitable for characterizing the absorbance of hemoglobin: 412, 541 and 576 nm. If the molar absorptivity is known at a particular wavelength, the concentration of a chemical compound present in a transparent sample can be calculated from the measured absorbance using Beer's Law. The simplest way to determine ε is to take a solution of known concentration, select the wavelength for which the value of ε is desired (usually the wavelength where the 16

17 absorbance has its greatest value), measure the absorbance there and measure the pathlength. The above equation can be rearranged to solve for ε (ε = A /lc) and the value computed from the experimental measurements. It is important to remember that spectrometers are limited in their ability to measure absorbance accurately; therefore, the results for very concentrated samples (with high absorbance values) may not be reliable. For example, an inexpensive spectrometer may produce reliable results only in the absorbance range of 0.01 to 1.5. Absorbance values outside this range are not reliable due to instrument limitations. For the best results, multiple measurements of a number of samples under a variety of conditions are required to provide an accurate answer. A more accurate method to determine ε is to measure the absorbance of a number of solutions of different concentrations and construct a calibration plot or standard curve. Beer's Law is a linear equation of the form A = ε l c A = ε l c y = mx + b (b, the y intercept, is zero and therefore does not appear in the Beer's Law equation.) A plot of absorbance versus concentration should produce a straight line with a slope equal to ε l. A sample standard curve is shown in Figure 9. With a zero intercept, a concentration value of zero should produce a zero absorbance and the origin of the plot (0, 0) should be included as a point on the plot. Figure 9: Absorbance Standard Curve Determining ε requires calculating the slope of the best-fit line through the data points. Consider the data graphed in Figure 9. By selecting two points on the line and reading their coordinates, the slope can be calculated. To avoid biasing the readings, the points selected for this determination should not be the same as any of the data points. εl = 17

18 As long as the pathlength (l) through the sample can be measured, ε can be calculated from the slope. A quick measurement of pathlength can be made with a ruler. A more rigorous method is to measure the absorbance of a standard solution having a known concentration and molar absorptivity and then calculate the pathlength from Beer's Law. Note that the value of l may vary from cuvette to cuvette. It will also vary with the orientation of the cuvette in the sample holder if the cuvette does not have a uniform pathlength in all directions. To maintain optimal accuracy, one should always use the same cuvette or at the very least make sure cuvettes are oriented the same way every time they are placed in the spectrophotometer. It is possible to read the concentration of an unknown sample directly from a calibration plot by interpolation using the measured absorbance of the unknown sample. In the example shown in Figure 10, an absorbance reading of 0.45 produces a concentration of 2.7 x 10-5 M. Instrumentation Figure 10: Interpolating Concentrations from a Beer s Law Plot Ocean Optics Spectrometers The Ocean Optics spectrometer is a quick and easy to use instrument for generating spectra in the UV, visible and near-infrared (near -IR) regions of the electromagnetic spectrum. The spectrometer (mirrors, grating, slit and detector) are housed in an optical bench that is small enough to fit into the palm of your hand. 18

19 The spectrometer accepts light energy either transmitted through an optical fiber or in free space -- and disperses it via a fixed grating across the linear CCD detector. The detector is designed to provide an electronic output for its wavelength range. The output from the detector is then fed into the computer to software, processed, and then displayed in the appropriate units. The spectrum you see is the result of multiple detector elements being fed into the computer and processed. This happens fast enough for you to be looking at the spectra generated by the instrument in real time. Ocean Optics Spectrophotometers Ocean Optics spectrometers can also be used to make absorbance measurements when used in conjunction with a light source and cell holder as shown in Figure 11. To measure absorbance, the instrument must be calibrated. This is done by first measuring the number of counts at each of the detector elements from the light source as it passes through a reference solution (solvent without the analyte of interest). Next, the background or number of counts when the light source is blocked (no light is entering the spectrometer) is measured. Both of these operations are automatically performed by the Figure 11: Ocean Optics Flame Spectrometer with Vis-NIR Integrated Sampling System software when the appropriate buttons are pushed. Once the instrument has been calibrated and a sample is inserted into the holder, the computer calculates the ratio of the counts hitting the detector to the stored reference counts for each of the detector elements and converts these to absorbance values displayed as an absorbance spectrum. 19

20 Absorbance of Light vs. Concentration When light is passed through a sample of material, the compounds in the sample can absorb some of the light. When this occurs, the intensity of the light beam that exits the sample will be less than the intensity of the light beam that entered the sample. A spectrophotometer is an instrument that measures the intensity of the light entering a sample and the intensity of light exiting a sample. The ratio of the intensity of the exiting light to the intensity of the entering light is expressed as transmittance, % transmittance (transmittance x 100%) or absorbance. Different materials absorb different wavelengths of light. Therefore, the wavelength of maximum absorption (λ max ) of a material is one of the characteristic properties of that material. In this experiment, the λ max of several food dyes will be determined. It is also possible to relate absorbance at a given wavelength to the concentration of the absorbing material present in a solution. This direct relationship between absorbance and concentration for a solution is known as the Beer-Lambert Law, or more commonly Beer s Law. You can use Beer s Law to test several samples of a solution, of known molar concentrations, and calculate a best-fit line equation to relate the absorbance of the solutions to their concentrations. Where: A = ε l c A = Absorbance of the solution ε = Molar absorptivity (L/mol cm, specific to the chemical species and wavelength of light used) l = Cell path length (cm) c = Concentration of absorbing species (mol/l) The more concentrated the light-absorbing material in the solution, the more light is absorbed at the λ max and the less light is transmitted. This behavior is investigated in this exercise and used to determine the concentration of a dye in a commercial drink. In this experiment, you will measure the absorbance spectra for several food dyes and determine the λ max of each dye. You will then prepare and measure absorbance spectra for several samples of dye with a known concentration and use Beer s Law to determine the concentration of an unknown dye in a commercial drink. 20

21 Materials Ocean Optics Flame Spectrometer and direct attach light source-cuvette holder (or standalone light source, cuvette holder and optical fibers) Computer with OceanView software installed Commercial food dyes Distilled water for diluting food dyes Commercial drinks containing food dyes 10 ml volumetric flasks Transfer/Beral pipettes Cuvettes Safety Wear safety goggles and other appropriate safety gear (gloves or apron) during the experiment. Take all proper safety precautions when handling the solutions. The Absorbance of Commercial Food Dyes Several samples of commercial food dyes are available for this lab. Work in groups of two to three to determine the λ max of the food dyes. The Ocean Optics spectrometer will be used to test the absorbance of the dye over a selected wavelength range. 1. Connect the spectrometer to your computer using the USB cable and turn on your light source. If you are using a direct attach light source, start the OceanView software and click the Strobe/Lamp Enable box in the Acquisition Group Window to turn the lamp on. 2. Allow the spectrometer and light source to warm up for at least 15 minutes before proceeding. 3. If you have not already done so, start OceanView. Click the Start button, then select All Programs Ocean Optics OceanView when you installed the software. OceanView or use the Desktop shortcut created 4. Select the Spectroscopy Application Wizards option on the Welcome Screen. 21

22 5. Select the Absorbance (Concentration) option in the Spectroscopy Wizards window to start the Absorbance Wizard. 6. Select the Absorbance only option and follow the steps through the Wizard to optimize the acquisition parameters for the spectrometer and acquire background and reference spectra for your absorbance measurement. 7. Place a cuvette filled with the solvent (water) in the cuvette holder and set your acquisition parameters in the Set Acquisition Parameters window. 8. Click the Automatic button to automatically adjust the Integration Time to the optimum value. 22

23 9. Set Scans to Average to 10 and Boxcar Width to 5 to reduce measurement noise and improve the measurement. You can set the Scans to Average to a higher value but this will slow down your acquisitions (total measurement time equals number of averages multiplied by the integration time). 10. Place a checkmark in the Nonlinearity Correction box if this feature is available for your spectrometer. 11. Click Next. 12. Click the Store Reference (yellow light bulb) button to store a reference spectrum. Click Next. 13. For direct attach light sources, click the Strobe/Lamp Enable box to close the shutter of the light source for your Background measurement. For other light sources, block the light from entering the spectrometer (do not turn your light source off). Click the Store Background (gray light bulb) button to store a background spectrum. 14. If you are using a direct attach light source, click the Strobe/Lamp Enable box to open the shutter. Click Finish. You are now ready to measure absorbance spectra. 15. The Absorbance spectra will be displayed in the AbsorbanceView window. You can adjust the graph appearance using the graph tools above the absorbance graph. For this experiment, click the Manually Set Numeric Ranges button to reset the wavelength scale to your region of interest from 400 to 700 nm. Click Apply to accept your changes and Exit to close the dialog box. Click the Scale Graph Height to Fill Window button to zoom in on the absorbance spectra. 16. Rinse the cuvette with the solution to be tested and fill it with the solution. Insert the cuvette into the cuvette holder and acquire the absorbance spectrum. 17. Save your absorbance spectrum by clicking the Configure Graph Saving button the File Writer. to configure a. Select the directory where you want to save the file. b. Enter a filename for your spectrum. c. Click Apply and then Exit to close the dialog box. 23

24 18. Click the Save Graph to Files button to save your absorbance spectrum. 19. Click the Convert Active Spectrum to Overlay button to create an overlay of the spectrum. 20. To find the wavelength of maximum absorbance, click anywhere on the graph to activate the cursor. A panel below the graph identifies the wavelength and intensity at the cursor position. Move the cursor on the graph until you find the wavelength where the highest absorbance occurs for each of the dyes. Record the λ max of the dye. 21. Repeat Steps 12 through 16 until all the dyes have been tested. 22. If a printer is available, print your data by clicking the Print graph button. Attach the printout of your absorbance spectra or draw the graph in your notebook. Determine the Relationship of Absorbance to Concentration at λ max 1. Obtain a 20 ml sample of one of the food dyes. Each student needs to investigate one dye. 2. Label the undiluted sample Solution A and record its concentration. 3. Make a series of dilutions of the sample using the dye and distilled water. Record the volume used for each dilution. When diluting a sample of solution A to make a new solution, the amount of solute does not change. Therefore, the concentrations of the new solutions can be determined by the following relationship: M A V A =M B V B where M is the molarity of the solution (or concentration in this case), V is the volume of solution, A is the initial solution and B is the final solution. Solution B: Pipette 5.00 ml of Solution A into a ml volumetric flask. Dilute to the mark with water. Mix thoroughly and label as Solution B and its concentration. Solution C: Pipette 2.00 ml of Solution A into another ml volumetric flask. Dilute, mix as before and label. Solution D: Pipette 1.00mL of Solution A into a third ml volumetric flask. Dilute, mix as before and label. 4. Place the cursor in your graph window on the λ max for your dye. 5. Record the absorbance for solutions D, C, B, and A at λ max. It is advisable to begin with the least concentrated solution to avoid sample carry over in the cuvette. 24

25 6. Obtain a sample of one of the commercial drinks that contains the dye you are investigating. Take the absorbance spectrum of the sample and record the absorbance at λ max. Data Analysis 1. Use a Microsoft Excel spreadsheet to graph the absorbance versus concentration. 2. By linear regression using Excel or a graphing calculator, determine the equation for the "best fit" line through your data points and record the value of R 2. An R 2 value close to 1.00 indicates that the line fits the data points well. 3. Determine the equation for the best fit line of the graph of absorbance versus concentration: Where: y = Absorbance x = Concentration m = Slope of the line b = y intercept y = mx + b 4. Report the concentration of the dye in the commercial drink using your Beer's Law standard curve (the best fit line equation) to relate the commercial dye absorbance to concentration. Discussion Questions 1. Discuss the relationship between absorbance and concentration using the data from this experiment. This relationship is called Beer's Law. 2. Student Problem: A soft drink company found that stores in their market area were not buying as many cases of their product as usual. Investigation showed that the stores had purchased the same number of cases but not from the company representative. The company was sure someone was counterfeiting their labels and selling imitations of their products. The chemical analysis lab for which you work has been asked to submit a bid for the job of analyzing the suspected merchandise. Your job as a chemist is to work with samples of a commercial soft drink containing two food dyes and to develop a method to: Determine the food dyes in a drink Determine the concentration of each food dye in the drink Separate the food dyes in the drink samples so that further analysis by an organic laboratory can be done on each dye separately (NOTE: A nonpolar column must be used for this problem.) 25

26 The Absorbance Spectrum of Chlorophyll The mixture of two chlorophyll molecules (chlorophyll a and b) from green, leafy plants absorbs several wavelengths of visible light, with five distinct absorbance peaks: three in the blue range (413, 454, and 482 nm) and two yellows (631 and 669 nm). The combination of these wavelengths is green to the human eye, but different ratios of these chlorophylls can create many shades of green. In this experiment you will extract chlorophyll from spinach (o r some other fresh green leafy sample) and measure its absorbance spectrum. While you wait for the extract to develop, you will measure the absorbance of blue and yellow food color samples so you will know what to expect when you measure the absorbance of your chlorophyll extract. Materials Ocean Optics Flame Spectrometer and direct attach light source-cuvette holder (or standalone light source, cuvette holder and optical fibers) Computer with OceanView software installed Spinach or fresh green leaves Isopropanol or ethanol Yellow and blue food color solutions Three small beakers Mortar and pestle Two 10 ml graduated cylinders Funnel and filter paper Ring stand and ring Plastic Transfer/Beral pipettes Cuvettes Safety Wear safety goggles and other appropriate safety gear (gloves or apron) during the experiment. Take all proper safety precautions when handling the solutions. 26

27 The Absorbance of Food Coloring 1. Obtain small amounts of blue and yellow food coloring solutions. 2. Connect the spectrometer to your computer using the USB cable and turn on your light source. If you are using a direct attach light source, start the OceanView software and click the Strobe/Lamp Enable box in the Acquisition Group Window to turn the lamp on. 3. Allow the spectrometer and light source to warm up for at least 15 minutes before proceeding. 4. If you have not already done so, start the OceanView software. Click the Start button, then select All Programs Ocean Optics OceanView OceanView. 5. Select the Spectroscopy Application Wizards option on the Welcome Screen. 6. Select the Absorbance (Concentration) option in the Spectroscopy Wizards window to start the Absorbance Wizard. 27

28 7. Select the Absorbance only option and follow the steps through the Wizard to optimize the acquisition parameters for the spectrometer and acquire background and reference spectra for your absorbance measurement. 8. Fill a cuvette ~⅔ full with dis lled water to serve as a blank/reference. Place the cuvette in the cuvette holder and set your acquisition parameters in the Set Acquisition Parameters window. 9. Click the Automatic button to automatically adjust the Integration Time to the optimum value. 10. Set Scans to Average to 10 and Boxcar Width to 5 to reduce measurement noise and improve the measurement. You can set the Scans to Average to a higher value but this will slow down your acquisitions (total measurement time equals number of averages multiplied by the integration time). 11. Place a checkmark in the Nonlinearity Correction box if this feature is available for your spectrometer. 12. Click Next. 13. Click the Store Reference (yellow light bulb) to store a reference spectrum. Click Next. 14. For direct attach light sources, click the Strobe/Lamp Enable box to close the shutter of the light source for your Background measurement. For other light sources, block the light from entering the spectrometer (do not turn your light source off). Click the Store Background (gray light bulb) button to store a background spectrum. 28

29 15. If you are using a direct attach light source, click the Strobe/Lamp Enable box to open the shutter. Click Finish. You are now ready to measure absorbance spectra. 16. The Absorbance spectra will be displayed in the AbsorbanceView window. You can adjust the graph appearance using the graph tools absorbance graph. Click the Scale Graph Height to Fill Window button on the absorbance spectra. above the to zoom in 17. Pour out the water from the blank cuvette, rinse and fill it ~⅔ full with the yellow food color solution. Place the cuvette in the cuvette holder. The absorbance spectrum for the yellow solution is displayed. 18. Save your absorbance spectrum by clicking the Configure Graph Saving button to configure the File Writer. a. Select the directory where you want to save the file. b. Enter a filename for your spectrum. c. Click Apply and then Exit to close the dialog box. 19. Click the Save Graph to Files button to save your absorbance spectrum. 20. Pour out the yellow solution from the cuvette, rinse and fill it ~⅔ full with the blue food color solution. Place the cuvette in the cuvette holder. The absorbance plot for the blue solution is displayed. Repeat Steps 14 and 15 to save the absorbance spectrum for the blue solution. 21. Mix equal amounts of the blue and yellow solutions in a small beaker. Pour out the blue solution from the cuvette, rinse and fill it ~⅔ full with the mixture. Place the cuvette in the cuvette holder. Repeat Steps 14 and 15 to save the absorbance spectrum for the mixture. 29

30 The Absorbance of Chlorophyll 1. Tear up a small sample of spinach into tiny pieces and grind them with a mortar and pestle. Add ml of 70% isopropanol (IPA) and transfer the mixture to a small beaker. Allow the mixture to sit for 30 minutes. 2. Set up an apparatus using the ring stand, ring and filter paper to filter the IPA/chlorophyll extract into a clean beaker. 3. After the IPA/chlorophyll extract has been soaking for 30 minutes, filter the extract into a clean beaker. 4. Measure the absorbance spectrum of the chlorophyll extract. You will need to run the Absorbance Wizard again because the solvent in the chlorophyll extract is IPA, not water. 5. Click the Create a new spectroscopy application button to open the Spectroscopy Application Wizards window. 6. Select the Absorbance (Concentration) option in the Spectroscopy Wizards window to start the Absorbance Wizard. 7. Select New in the Existing Absorbance window to display the chlorophyll extract data in a new graph window. 8. Fill a cuvette ~⅔ full with IPA to serve as a blank/reference. Place the cuve e in the cuvette holder and set your acquisition parameters in the Set Acquisition Parameters window. 9. Click the Automatic button to automatically adjust the Integration Time to the optimum value. 10. Set Scans to Average to 10 and Boxcar Width to 5 to reduce measurement noise and improve the measurement. You can set the Scans to Average to a higher value but this will slow down your acquisitions (total measurement time equals number of averages multiplied by the integration time). 11. Place a checkmark in the Nonlinearity Correction box if this feature is available for your spectrometer. 12. Click Next. 13. Click the Store Reference (yellow light bulb) to store a reference spectrum. Click Next. 14. For direct attach light sources, click the Strobe/Lamp Enable box to close the shutter of the light source for your Background measurement. For other light sources, block 30

31 the light from entering the spectrometer (do not turn your light source off). Click the Store Background (gray light bulb) button to store a background spectrum. 15. If you are using a direct attach light source, click the Strobe/Lamp Enable box to open the shutter. Click Finish. You are now ready to measure absorbance spectra. 16. The Absorbance spectra will be displayed in the AbsorbanceView window. You can adjust the graph appearance using the graph tools absorbance graph. Click the Scale Graph Height to Fill Window button on the absorbance spectra. above the to zoom in 17. Measure the absorbance of the chlorophyll extract. Pour out the IPA from the blank cuvette, rinse and fill it ~⅔ full with the chlorophyll extract. Place the cuve e in the cuvette holder. The chlorophyll absorbance spectrum is displayed. 18. Save the chlorophyll absorbance spectrum by clicking the Configure Graph Saving button to configure the File Writer. 19. Select the directory where you want to save the file. 20. Enter a filename for your spectrum. 21. Click Apply and then Exit to close the dialog box. 22. Click the Save Graph to Files button to save your absorbance spectrum. 31

32 Discussion Questions 1. Consult a reliable resource to identify the major absorbance peaks of chlorophyll a and chlorophyll b. Examine the absorbance spectra for chlorophyll. Does your graph clearly show these absorbance peaks? Are there other, unidentified peaks on your graph? Identify them and speculate about what caused these peaks. 2. How did your tests of the absorbance of the blue and yellow food coloring solutions compare with the tests of the chlorophyll extract? 3. If distilled water had been used as the calibration blank for the chlorophyll test, would it have affected the absorbance measurements? 4. Were the distinguishing features of chlorophyll a and b evident in the absorbance spectrum graph of the chlorophyll extract? Explain. 32

33 Determining the Rate law of a Chemical Reaction: Kinetics of Crystal Violet Bleaching Chemists are always interested in whether a chemical reaction can occur and exactly how it occurs. The first question is answered through thermodynamics while the second is the domain of kinetics. In a kinetics experiment, a chemist attempts to understand the step-by-step transformation of reactants to products. Taken together these elementary steps give us the mechanism by which the reaction proceeds. Note that a reaction's kinetics are very much tied to the pathway the reactants take to the products (i.e., the mechanism), which is very different from the reaction's thermodynamic properties (i.e., H (Enthalpy), S (Entropy) and G (Gibbs Free Energy) that do not depend on the path. While the thermodynamics and kinetics of a reaction may at times seem complementary, and at other times seem contradictory, it is always important to have a detailed understanding of both. In this experiment, you will determine the rate law for a chemical reaction. The rate law is a mathematical expression relating to chemical reactions. The amount of time it takes a reaction to occur correlates to the concentrations of the starting materials. The disappearance of reactant over time depends on the rate constant and the concentration of each reactant raised to some power. This power is known as the order of reaction with respect to that reactant. The sum of the individual orders is the overall order of the reaction. The order of reaction with respect to each reactant, as well as the rate law itself, cannot be determined from the balanced chemical equation; it must be found experimentally. The rate law is the basic equation of kinetics and it will be the standard against which possible mechanisms are judged. Reaction of Crystal Violet with OH - In this experiment, you will determine the rate law and order for the reaction of a dye, crystal violet (CV) with sodium hydroxide ( OH - ) in aqueous solution according to the balanced net ionic equation shown below. 33

34 We will define the rate of reaction as the disappearance of the colored CV over time, which can be expressed in differential form as d[cv]/dt so the rate law for this reaction can be written as shown in the equation below in terms of the concentration of CV and OH - and the rate constant for the reaction, k. In writing this equation we assume that both CV and OH - are involved in the reaction (that is x and y are both not zero and are likely integers). Only the experiment will tell us whether these assumptions are valid. rate d[cv ] dt k[cv ] x [OH ] y The point of any kinetics experiment is to determine the order with respect to each reactant (i.e., find x and y) and to find the value of k. This is a challenge if we have more than one reactant, as is the case here. When there is more than one reactant, the isolation method is often used, which entails making the concentration of all but one of the reactants very high so their concentrations do not change appreciably over the course of the reaction. The order with respect to the isolated reactant is then determined. The process is then repeated, isolating each of the other reactants until all of the orders have been determined. In this experiment we will use a vast excess of OH - (make the OH - concentration ([OH - ]) very large) and therefore, essentially constant. In this case, the rate and order of the reaction will be in crystal violet only. We will define the rate of the reaction as the disappearance of crystal violet over time, which can be expressed in differential form as d[cv]/dt We can then simplify the rate law to the equation below, where we have defined a new rate constant, k obs, which is the observed rate constant at some specific [OH - ]. 34

35 rate k obs [CV ] x The relationship between k obs and the intrinsic rate constant, k, for this reaction is given by the equation below. k obs k[oh ] y Under conditions of high, constant [OH - ], the order with respect to CV can be determined by graphically applying the integrated rate laws. Since the absorbance of a CV solution is directly proportional to the concentration of CV ([CV]), according to Beers' Law, the actual [CV] can be replaced by A max, the solution's maximum absorbance (somewhere around 600 nm). A graph of A max as a function of time will give a straight line if the reaction is zero-order in CV (x = 0). If the reaction is first-order in CV (x = 1), then a graph of ln(a max ) as a function of time is linear. Finally, if a graph of 1/A max as a function of time is linear, it indicates that the reaction is second-order with respect to CV (x = 2). In each case, if a particular relationship is linear, then the slope of that graph can be used to determine k obs. Note that only one of these three graphs will be linear! In some instances it is not possible to isolate one of the reactants, because the concentration of that reactant must remain high for the system to behave predictably. In this reaction the [OH - ] must remain high, but the order of the reaction with respect to OH - and k, can still be found. First we need to change the rate law into an easily graphed form by taking the natural logarithm of both sides to give ln(k obs ) y ln[oh ] ln(k) To determine the order with respect to OH - and k, we first perform the kinetics experiment at different, albeit still high, OH - concentrations and then graph ln(k obs ) for these reactions as a function of ln[oh - ]. The slope of this graph is y, the order with respect to OH -, and the intercept is ln(k). In this experiment, you will measure the absorbance of the reaction over time, at a specific wavelength of visible light. You will first measure the absorbance spectrum of a crystal violet solution and select one wavelength to examine during the reaction. As the reaction proceeds, the bright purple color of the crystal violet solution will fade and the absorbance will decrease. 35

36 Materials Ocean Optics Flame Spectrometer and direct attach light source-cuvette holder (or standalone light source, cuvette holder and optical fibers) Computer with OceanView software installed M crystal violet solution 0.05 M NaOH solution Two 10 ml graduated cylinders or pipettes Three small beakers Plastic Transfer/Beral pipettes Cuvettes Safety Wear safety goggles and other appropriate safety gear (gloves or apron) during the experiment. Take all proper safety precautions when handling the solutions. Measuring the Kinetics of Crystal Violet Bleaching 1. Obtain ml of the crystal violet (CV) and sodium hydroxide (OH - ) solutions. 2. Connect the spectrometer to your computer using the USB cable and turn on your light source. If you are using a direct attach light source, start the OceanView software and click the Strobe/Lamp Enable box in the Acquisition Group Window to turn the lamp on. 3. Allow the spectrometer and light source to warm up for at least 15 minutes before proceeding. 4. If you have not already done so, start OceanView. Click the Start button, then select All Programs Ocean Optics OceanView OceanView or use the Desktop shortcut created when you installed the software. 5. Select the Spectroscopy Application Wizards option on the Welcome Screen. 36

37 6. Select the Absorbance (Concentration) option in the Spectroscopy Wizards window to start the Absorbance Wizard. 7. Select the Absorbance only option and follow the steps through the Wizard to optimize the acquisition parameters for the spectrometer and acquire background and reference spectra for your absorbance measurement. 8. Fill a cuvette ~⅔ full with the 0.05 M NaOH solu on to serve as a blank. Place the cuvette in the cuvette holder and set your acquisition parameters in the Set Acquisition Parameters window. 9. Click the Automatic button to automatically adjust the Integration Time to the optimum value. 37

38 10. Set Scans to Average to 10 and Boxcar Width to 5 to reduce measurement noise and improve the measurement. You can set the Scans to Average to a higher value but this will slow down your acquisitions (total measurement time equals number of averages times the integration time). 11. Place a checkmark in the Nonlinearity Correction box if this feature is available for your spectrometer. 12. Click Next. 13. Click the Store Reference (yellow light bulb) to store a reference spectrum. Click Next. 14. For direct attach light sources, click the Strobe/Lamp Enable box to close the shutter of the light source for your Background measurement. For other light sources, block the light from entering the spectrometer (do not turn your light source off). Click the Store Background (gray light bulb) button to store a background spectrum. 15. If you are using a direct attach light source, click the Strobe/Lamp Enable box to open the shutter. Click Finish. You are now ready to measure absorbance spectra. 16. The Absorbance spectra will be displayed in the AbsorbanceView window. You can adjust the graph appearance using the graph tools absorbance graph. Click the Scale Graph Height to Fill Window button on the absorbance spectra. above the to zoom in 17. Click the Create Strip Chart button. a. Select the data source for the strip chart. Make sure to highlight the data source with Spectrum Type Absorbance. b. In the Update Rate panel, check the box next to Stop after this amount of time and change the time to 10 minutes. c. In the Wavelength Selection panel, select One Wavelength and set it to 585 nm. d. DO NOT click the Finish button until you have prepared your reaction as described in the steps below. 38

39 18. To prepare for the reaction, pour out the NaOH solution from the blank cuvette and rinse the cuvette with distilled water. 19. Mix 10 ml of the OH - and CV solutions in a third beaker. Swirl the beaker gently, and then use a plastic Beral pipette to transfer 2-3 ml of the reaction mixture to the cuvette. 20. Place the cuvette in the cuvette holder. Click the button. A graph of absorbance versus time will be displayed in the Trend window. 21. If the strip chart is not visible in the window, adjust the graph scale using the graph tools above the absorbance graph. 22. If the strip chart is not scrolling (trend runs off the screen), click the Automatically Scroll the Graph button to scroll the strip chart during the acquisition. 23. After the strip chart run ends, copy the data to the clipboard for additional analysis in a program like Microsoft Excel. a. Click the Copy Data to Clipboard button b. Open Excel and paste the data into the worksheet. Data Analysis 1. Prepare three plots: absorbance versus time, ln absorbance versus time and 1/absorbance versus time. Calculate the best-fit line equation for the plot that is the most linear and write down the equation. 2. Based on the information provided in the introductory remarks, what is the order of the reaction with respect to crystal violet? 3. Based on the information provided in the introductory remarks, write the rate law for this reaction. 39

40 Extension In some instances it is not possible to isolate a reactant because the concentration of the reactant must remain high for the reaction to behave predictably. This is the case with the reaction between crystal violet and sodium hydroxide; the [OH - ] must remain high. However, the order of the reaction with respect to OH - and the subsequent rate law constant can still be determined. To determine the order of the reaction in OH -, you must conduct the reaction using different concentrations (albeit in vast excess) of OH -. From each data collection run you will calculate a value of k. Using the values of k and the [OH - ], you will prepare a plot of ln [OH - ] (the Y-values) versus ln k (the X-values). The best-fit line for this plot takes the form: ln (k abs ) = m ln [OH - ] + ln k. The slope, m, is the order of the reaction in OH - and the Y-intercept is the natural logarithm of the rate constant. 40

41 Emission Spectra A fascinating feature of spectroscopy is how one can make use of light to learn about atomic and molecular structure. Under certain conditions, an atom or molecule will absorb or emit light. By examining and measuring the light that is absorbed or emitted by a substance, certain physical properties are revealed. The electrons of atoms and molecules exist in specific energy states. The energy emitted by the excitation of electrons is limited to differences between these states, thus specific energies of light are emitted. The color of a glowing LED, for example, is a result of the energy of the emitted light. The energy and wavelength of the light is described by the equation E = hc/λ where λ is the wavelength, h is Planck's constant ( J sec), and c is the speed of light ( m/sec). If you are measuring the emission spectrum of a gas trapped in a discharge tube, only certain wavelengths of light are emitted by the gas and the pattern that is produced is unique for that substance. In this experiment you will use the spectrometer outfitted with a fiber optic cable to measure the emission spectra of various sources of light. Materials Ocean Optics Flame Spectrometer Fiber optic cable Computer with OceanView software installed Light sources: LEDs, Lamps, Discharge Tubes Safety Wear safety goggles and other appropriate safety gear during the experiment. Take all proper safety precautions when working with light sources including eye protection. 41

42 Measuring Light Source Emission with a Spectrometer 1. Connect a fiber optic cable to the threaded SMA 905 connector on the housing of the spectrometer. 2. Connect the spectrometer to your computer using the USB cable. 3. Start the OceanView software. Click the Start button, then select All Programs Ocean Optics OceanView OceanView. 4. Select the Quick View option on the Welcome Screen. 5. Allow the spectrometer to warm up for at least 15 minutes before proceeding. 6. Turn on the light source. Aim the tip of the fiber optic cable at the light source. 7. There are two methods for optimizing the spectrum showing intensity versus wavelength. Note that one method may be better than the other depending on the light source you are measuring. Option 1: Adjust the distance between the light source and fiber. Set the distance between the light source and the tip of the fiber optic cable so that the peak intensity on the graph is ~55,000 counts for the Flame spectrometer. Option 2: Adjust the integration time. Click the Automatic button the Acquisition Group Window to automatically adjust the Integration Time to the optimum value. in 42

43 8. When you are satisfied with your emission graph, save the spectrum by clicking the Configure Graph Saving button a. Select the directory where you want to save the file. to configure the File Writer. b. Enter a filename for your spectrum. c. Click Apply and then Exit to close the dialog box. 9. Click the Save Graph to Files button to save your emission spectrum. 10. To overlay your emission spectrum on the graph, click the Convert Active Spectrum to Overlay button to create an overlay of the spectrum. 11. To analyze your emission spectrum graph: a. Click anywhere on the graph to activate the cursor. Note the vertical line marking a given wavelength on the graph. b. Click on each peak of the emission spectrum. A legend below the graph displays the counts at the wavelength of the cursor location. Write down the wavelength for the peak or peaks in your spectrum, as well as any other distinguishing characteristics of the graph, in your lab book. 12. Repeat Steps 6-11 to plot and capture the emission spectrum of a second light source. 13. Use the Convert Active Spectrum to Overlay option to show more than one emission spectrum on the same graph. If you wish to remove a set of graphed data, click the Delete Overlay Spectrum button. Highlight the spectrum you want deleted (all of the spectra are color coded) and click OK. 14. To export the graphed emission measurements to Microsoft Excel: a. Click the Copy Data to Clipboard button b. Open Excel and paste the data into the worksheet. 43

44 Data Analysis 1. Examine your first graph of emission. Identify the peak or peaks. Describe the distinguishing characteristics of the graph. 2. Examine your second graph of emission. As before, identify the peak or peaks and describe the distinguishing characteristics of the graph. 3. Are there any features of either graph that stand out as being unusual or unexpected? 4. Below are two emission graphs. The graph on the left is the emission from standard fluorescent office lighting. The graph on the right is the emission from a mercury discharge tube. Using these graphs, make a case either for or against the presence of mercury in the office lighting. The X- and Y-axis ranges for the graphs are identical. 44

45 Spectrophotometric Analysis of Commercial Aspirin The concentration of acetylsalicylic acid (ASA) is determined spectrophotometrically by the percent transmittance (%T) of visible light at a given wavelength %T = (I t /I o ) x 100% where I t = Intensity of the beam transmitted by the solution, I o = Intensity of the light beam before it is passes through the sample solution. The absorbance of the analyte in solution is related to the transmittance (T) as follows: Beer s Law says that where: A = -log (I t /I o ) A = -log T A = -log (%T/100) A = -log (%T) + 2 A = 2 - log (%T) A = l c = Molar absorptivity of the particular absorbing species in (L /moles cm) l = Pathlength of the light through the solution in cm c = Concentration of the absorbing species in moles/l From these equations, we can determine the concentration c, in mole/l, for the absorbing species if the percent transmission (%T) can be determined from the spectrometer. To determine the concentration of the absorbing species in solution, you need to construct a standard Beer s Law Plot. This is done by preparing a series of known concentrations of the analyte and reading the percent transmission of each standard solution at the maximum absorbing wavelength ( max ) in the visible region. The resulting plot of absorbance at max versus concentration is a straight line plot known as the Beer s Law plot. 45

46 A = x x x slope = k 0.20 x x [C] unk = A/k = / k concentration (M x 10 4 ) The plot of absorbance versus concentration gives a straight line plot passing through the origin (0,0). The slope (k) of the line can be obtained from a linear regression or the least-squares fit method and is equal to the k = A / c, where k (slope of the Beer s Law plot) = x l = molar absoptivity (L /moles cm) l = pathlength of the light (cm) If the absorbance of the unknown solution is measured, then the concentration of the ASA complex can be calculated by the formula c = A/k The analyte of interest in this experiment, acetylsalicylic acid (ASA) complex ion, is formed by hydrolyzing the aspirin sample (ASA) in NaOH solution and complexing it with Fe 3+ ion in acid solution to bring out the color as shown in the chemical reaction. The complex displays a maximum absorption at a wavelength () of 530 nm, and has a crimson red color. 46

47 In this experiment, we will determine how much active ingredient, acetylsalicylic acid (ASA), in mass %, is contained in commercially available aspirin tablets using a visible spectrometer. Prelab Questions Construct a Beer s Law plot for the following experimental data using computer software like Excel or by plotting the points on precision graph paper. Concentration of Cobalt (III) Complex (Standard solution, M) Percent Transmission Absorbance at 530 nm 3.00 x x x x x Unknown Calculate the absorbance of each solution using: A = -log (I t /I o ) = -log T = -log (%T/100) = -log (%T) + 2 = 2 - log (%T) 2. Plot concentration on the x-axis versus absorbance on the y-axis for the five standard solutions. Use computer software or precision graph paper. 3. Calculate the slope (k) from the Beer s Law plot generated in Step 2. Use linear regression to find the best slope, if you are using computer software. The straight line should pass through the origin (0,0), otherwise a large error may occur. 4. Determine the molar concentration of the unknown cobalt (III) solution, using: c = A/k 47

48 Materials Ocean Optics Flame Spectrometer and direct attach light source-cuvette holder (or standalone light source, cuvette holder and optical fibers) Computer with OceanView software installed Reagent grade ASA (acetylsalicylic acid) Aspirin 1.0 M NaOH solution M FeCl 3 /KCl/HCl solution (ph=1.6) DI water Weigh paper or weigh boats Two 125 ml Erlenmeyer flasks Two 250 ml volumetric flasks with stoppers Seven 50 ml volumetric flasks with stoppers 10 ml graduated cylinder Pipettes with suction bulbs Large test tubes with stoppers Glass funnel Hot plate Cuvettes Safety Wear safety goggles and other appropriate safety gear (gloves or apron) during the experiment. Take all proper safety precautions when handling the solutions. Hydrolyzing the ASA and Aspirin Samples Work in pairs to prepare the standard acetylsalicylic acid (ASA) solution and aspirin sample solution simultaneously. 1. Weigh to the nearest mg (0.001 g) on a piece of weigh paper or a weigh boat approximately 0.4 g of reagent grade ASA. Transfer the sample to a 125 ml Erlenmeyer flask labeled Flask #1. Record the exact mass of ASA and weigh boat or weigh paper on your Data Sheet Record your unknown aspirin brand name and weigh each aspirin tablet to the nearest milligram (0.001 g). Record this mass on Data Sheet 2. Repeat the above procedure with the aspirin sample and label the Erlenmeyer flask Flask #2. 48

49 3. Measure 10 ml of 1.0 M NaOH solution in a clean, dry, graduated cylinder. Add the NaOH to the ASA in the 125 ml Erlenmeyer Flask #1. 4. Repeat for Flask #2. 5. Heat the mixtures (Flask #1 and Flask #2) to a mild boil for five minutes on a hot plate to hydrolyze the ASA. Be careful to avoid splattering and do not let the solution dry up to prevent loss of contents. Rinse the inside of walls of the flasks with a small amount of DI water to ensure complete chemical reaction of ASA. 6. Quantitatively transfer the solution of sodium salicylate in Flask #1 to a 250 ml volumetric flask through a glass funnel. Thoroughly rinse the flask and funnel with DI water so that the rinse water flows into the volumetric flask. Add DI water to the solution in the flask until the bottom of the meniscus touches the index mark of the flask neck. Stopper the flask. While firmly holding the stopper with your forefinger, invert the flask 10 times to thoroughly mix the solution. Repeat the process for Flask #2 containing the aspirin sample. 7. Transfer the solutions into 250 ml Erlenmeyer flasks with rubber stoppers and label Standard for the sample from Flask #1 and Aspirin for the sample from Flask #2. Store these flasks in your drawer until needed. NOTE: Aspirin solutions may have milky appearance due to starch fillers. Some buffering agents such as aluminum hydroxide will not dissolve completely in base. In such cases, allow the undissolved material (precipitate) to settle to the bottom of the flask. If a precipitate is present, use your pipette to remove solution from the top portion of the liquid so that you will not draw any precipitate into your pipette. 8. Check with your instructor whether you may proceed beyond this point. Preparing Standard and Aspirin Solutions 1. Clean your pipette with small portions of the solution you are trying to measure by drawing in solution with the suction bulb and discarding it. 2. Pipette a 2.40 ml portion of the standard solution into a clean 50 ml volumetric flask. Allow the solution to drain and gently allow the tip of the pipette to touch the side of the flask. Add M FeCl 3 /KCl/HCl solution (ph=1.6) to the 50 ml volumetric flask until the bottom of the meniscus touches the index mark on the flask. NOTE: You should see the solution turn red at this point. If no color appears after adding FeCl 3 solution, notify the instructor. 49

50 3. Stopper the flask. While firmly holding the stopper, with your forefinger, invert the flask 10 times to thoroughly mix the solution. Label this flask Standard Solution A. 4. Repeat Steps 2 and 3 to prepare Standard Solutions B, C, D and E by diluting 2.00, 1.60, 1.00, 0.40 ml portions of the sodium salicylate standard solution with the FeCl 3 /KCl/HCl solution. 5. Transfer the Standard Solutions to large test tubes with stoppers and label them clearly. 6. Prepare two unknown aspirin solutions. Use a clean, well-rinsed pipette, transfer 1.20 ml of the aspirin stock solution into a 50 ml volumetric flask and dilute it to the mark with M FeCl 3 /KCl/HCl solution. Mix thoroughly as described in Step 3. Make a second unknown aspirin sample using 1.60 ml of the aspirin stock solution. Transfer the unknown aspirin solutions to large test tubes with stoppers and label them clearly. Measuring the % Transmittance using a Spectrometer 1. Connect the spectrometer to your computer using the USB cable and turn on your light source. If you are using a direct attach light source, start the OceanView software and click the Strobe/Lamp Enable box in the Acquisition Group Window to turn the lamp on. 2. Allow the spectrometer and light source to warm up for at least 15 minutes before proceeding. 3. If you have not already done so, start OceanView. Click the Start button, then select All Programs Ocean Optics OceanView OceanView or use the Desktop shortcut created when you installed the software. 4. Select the Spectroscopy Application Wizards option on the Welcome Screen. 50

51 5. Select the Transmission option in the Spectroscopy Wizards window to start the Transmission Wizard. 6. Select the Active Acquisition (Recommended) option and follow the steps through the Wizard to optimize the acquisition parameters for the spectrometer and acquire background and reference spectra for your transmission measurement. 7. Fill a cuvette ~⅔ full with the solvent (FeCl 3 /KCl/HCl solution) to serve as a blank. Place the cuvette in the cuvette holder and set your acquisition parameters in the Set Acquisition Parameters window. a. Click the Automatic button to automatically adjust the Integration Time to the optimum value. 51

52 8. Set Scans to Average to 10 and Boxcar Width to 5 to reduce measurement noise and improve the measurement. You can set the Scans to Average to a higher value but this will slow down your acquisitions (total measurement time equals number of averages times the integration time). 9. Place a checkmark in the Nonlinearity Correction box if this feature is available for your spectrometer. 10. Click Next. 11. Click the Store Reference (yellow light bulb) to store a reference spectrum. Click Next. 12. For direct attach light sources, click the Strobe/Lamp Enable box to close the shutter of the light source for your Background measurement. For other light sources, block the light from entering the spectrometer (do not turn your light source off). Click the Store Background (gray light bulb) button to store a background spectrum. 13. If you are using a direct attach light source, click the Strobe/Lamp Enable box to open the shutter. Click Finish. You are now ready to generate transmission spectra. 14. The Transmission spectra will be displayed in the TransmissionView window. You can adjust the graph appearance using the graph tools transmission graph. Click the Scale Graph Height to Fill Window button on the absorbance spectra. above the to zoom in 15. Click anywhere on the graph to place a cursor on the graph. Position the cursor at 530 nm and record the percent transmittance (%T) of the standard solutions A, B, C, D and E and of the two unknown aspirin solutions. Record the %T values on Data Sheets 1 and 2. NOTE: Your aspirin sample solutions should have % Transmittance between 30% and 70%. If not, make appropriate dilutions or additions so that they fall within the range. Remember to record the dilution or addition factor used to prepare your aspirin sample if you decide to adjust your concentration. 52

53 16. You can save your absorbance spectrum by clicking the Configure Graph Saving button to configure the File Writer. a. Select the directory where you want to save the file. b. Enter a filename for your spectrum. c. Click Apply and then Exit to close the dialog box. 17. Click the Save Graph to Files button to save your absorbance spectrum. 18. You can overlay your spectrum by clicking the Convert Active Spectrum to Overlay button to create an overlay of the spectrum in your graph view. 53

54 Data Sheet 1 Preparing Standard Solutions: Weighing boat or paper, g Mass of ASA + weighing boat or paper, g Mass of ASA, g Concentration of ASA complex in solution, M Solution Concentration, M % Transmittance Absorbance A B C D E Slope of the Beer s Law plot = 54

55 Data Sheet 2 Analyzing Commercial Aspirin Tablets: Sample ID Mass of tablet + weighing boat or paper, g Mass of weighing boat or paper, g Mass of tablet, g Percent Transmittance, %T Absorbance of solution, A Concentration of aspirin in solution, M Mass of ASA in tablet, g Percent ASA in tablet, % Mean percent of ASA, % Trial 1 Trial 2 55

56 Calculations 1. Find the slope (k) from your Beer s Law plot. 2. Convert the percent transmittance (%T) to the equivalent absorbance (A) using: A = -log (I t /I o ) = -log T = -log (%T/100) = -log (%T) + 2 = 2 - log (%T) 3. From the calculated absorbance and the Beer s Law plot slope (k), determine the concentration of ASA in your aspirin solution using: c = A/k 4. Calculate the mass of ASA in each tablet, using: Mass of ASA in grams = [ASA]*(180.2 g /mol ASA)( L)(250 ml/1.60 ml) *... Use your experimental values of [ASA] in mole/l and the volume of the original aspirin solution in ml. 5. Find the percent ASA in each tablet, using : % ASA in tablet = (mass ASA in tablet, g/ mass of tablet, g) x 100% 6. Calculate the mean percent ASA in your commercial brand of aspirin, using: Mean % ASA per tablet ={[ ASA (trial #1)] + [% ASA (trial#2)] }/2 56

57 Post-Lab Questions 1. Explain why the Fe 3+ /H + solution was used as a reference solution. Suggest a procedure you could follow to determine whether it was necessary to use the solution as a reference or whether de-ionized water would have been satisfactory. 2. Most commercial aspirins claim to contain 5.0 grains of ASA per tablet, where a grain is an old apothecary unit of measurement for mass, equal to 65 mg. Compare your calculated value of ASA per tablet with respect to the advertised value (i.e., 5.0 grains) and determine the percent error. 3. Compare the results of the class data, if available, for different commercial brands of aspirin. Is there any difference between the different brands of aspirin? Support your answer using simple statistics. 57

58 Spectrophotometric Characterization of Spice Extracts Spices have a long and rich history around the world. In some senses (pun intended), the cuisine of a particular region identifies itself with specific local spices. For centuries, spices have been used to modify the flavor of prepared foods, help preserve foods and act as herbal remedies for a wide variety of real and imagined ills. Along with the important role of contributing to the flavors of foods, spices are also employed to modify the color of foods. Turmeric adds a strong yellow tinge to curries while saffron offers a softer yellow to paellas. Paprika contributes a deep, vibrant red to chicken paprikash as well as presenting a vivid, flavorful spice garnish to simple cottage cheese. In this experiment, you will use isopropanol, also known as rubbing alcohol, to prepare extracts of a selection of spices. You will then use an Ocean Optics spectrometer and light source to measure the visible light absorbance spectrum of several known spice extracts. You will then measure unknown spice samples to determine whether or not a specific spice is present. Materials Ocean Optics Flame Spectrometer and direct attach light source-cuvette holder (or standalone light source, cuvette holder and optical fibers) Computer with OceanView software installed Spices such as paprika, turmeric, saffron, parsley, jalapeno, cumin and curry 70% isopropanol (rubbing alcohol) Unknown spice solution prepared in 70% isopropanol (rubbing alcohol) 25 ml graduated cylinders 50 ml beakers Transfer/Beral pipettes Cuvettes Safety Wear safety goggles and other appropriate safety gear (gloves or apron) during the experiment. Take all proper safety precautions when handling the solutions. 58

59 Procedure 1. Place a very small amount of each spice in a 50 ml beaker. Add 20 ml of rubbing alcohol (70% isopropanol). Swirl each beaker to mix the spices. Allow the mixtures to sit for 30 minutes so the colorants are well dissolved and any particulate matter settles to the bottom of the beakers. 2. Connect the spectrometer to your computer using the USB cable and turn on your light source. If you are using a direct attach light source, start the OceanView software and click the Strobe/Lamp Enable box in the Acquisition Group Window to turn the lamp on. 3. Allow the spectrometer and light source to warm up for at least 15 minutes before proceeding. 4. Start the OceanView software. Click the Start button, then select All Programs Ocean Optics OceanView OceanView or use the Desktop shortcut created when you installed the software. 5. Select the Spectroscopy Application Wizards option on the Welcome Screen. 6. Select the Absorbance (Concentration) option in the Spectroscopy Wizards window to start the Absorbance Wizard. 59

60 7. Select the Absorbance only option and follow the steps through the Wizard to optimize the acquisition parameters for the spectrometer and acquire background and reference spectra for your absorbance measurement. 8. Fill a cuvette ~⅔ full with 70% rubbing alcohol to serve as a blank. Place the cuvette in the cuvette holder and set your acquisition parameters in the Set Acquisition Parameters window. a. Click the Automatic button to automatically adjust the Integration Time to the optimum value. b. Set Scans to Average to 10 and Boxcar Width to 5 to reduce measurement noise and improve the measurement. You can set the Scans to Average to a higher value but this will slow down your acquisitions (measurement time equals number of averages times the integration time). c. Place a checkmark in the Nonlinearity Correction box if this feature is available for your spectrometer. d. Click Next. 9. Click the Store Reference (yellow light bulb) to store a reference spectrum. Click Next. 10. For direct attach light sources, click the Strobe/Lamp Enable box to close the shutter of the light source for your Background measurement. For other light sources, block the light from entering the spectrometer (do not turn your light source off). Click the Store Background (gray light bulb) button to store a background spectrum. 60

61 11. If you are using a direct attach light source, click the Strobe/Lamp Enable box to open the shutter. Click Finish. You are now ready to measure absorbance spectra. 12. The Absorbance spectra will be displayed in the AbsorbanceView window. You can adjust the graph appearance using the graph tools absorbance graph. Click the Scale Graph Height to Fill Window button on the absorbance spectra. above the to zoom in 13. Pour out the rubbing alcohol from the blank cuvette, rinse and fill it ~⅔ full with the one of the spice extracts. Place the cuvette in the cuvette holder. The absorbance spectrum for the extract is displayed. 14. Save your absorbance spectrum by clicking the Configure Graph Saving button the File Writer. to configure a. Select the directory where you want to save the file. b. Enter a filename for your spectrum. c. Click Apply and then Exit to close the dialog box. 15. Click the Save Graph to Files button to save your absorbance spectrum. 16. Click the Convert Active Spectrum to Overlay option to overlay the spectrum so you can compare it to the next absorbance spectrum on the same graph. If you wish to remove a spectrum, click the Delete Overlay Spectrum button. Highlight the spectrum you want deleted (all of the spectra are color coded) and click OK. 17. Pour out the extract from the cuvette, rinse and fill it ~⅔ full with another spice extract. Place the cuvette in the cuvette holder. The absorbance plot for the solution is displayed. Repeat Steps 14 and 15 to save the absorbance spectrum. 18. Repeat Step 17 with your remaining spice extracts. 19. Obtain an unknown spice extract mixture. Repeat Step 17 to measure its absorbance spectrum. 61

62 20. Before you exit OceanView, print copies of the absorbance spectrum graphs for each of your spice extracts and your unknown spice mixture by clicking the Print button. Alternately, keep OceanView running and use it to compare your unknown solutions to your known spice spectra. Data and Analysis 1. Print a copy of the absorbance spectrum for each of the spice extracts that you tested, as well as the unknown spice mixture. Note the distinguishing characteristics of the spectra that will help you identify the unknown. 2. Which spice, or spices, is contained in your unknown? Explain. 3. If distilled water had been used as the calibration blank for the chlorophyll test, would it have affected the absorbance measurements? 62