PROCEDURE_Crystal Violet Analysis by Spectrophotometry_Fall 2023

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Dec 6, 2023

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CHEM 3051 Fall 2023 CRYSTAL VIOLET COLORIMETRY: A BEER'S LAW INVESTIGATION LEARNING O B J E C T I V E S The objectives of this exercise are to:  Illustrate the basic principles of colorimetry.  Demonstrate the components of a colorimeter and how a colorimeter is interfaced to a computer.  Discover the Beer's Law relationship and apply it to the analysis of an unknown solution. 1.0 BACKGROUND Colorimetry is an instrumental method based on the measurement of light absorption by colored solutions and is widely used for performing chemical analyses. The Micro LAB colorimeter utilizes 10 Light Emitting Diodes (LEDs) ranging from 400 to 644 nanometers, spaced about 35 nanometers apart. They are scanned sequentially and the transmittance or absorbance data is presented as 10 colored bars in a bar graph as shown in Figure 1. Figure 1. MicroLAB spectrum profile for blank illustrating the color bars for each wavelength at 100 % transmittance . These components are arranged such that light from each LED passes through the crystal violet solution and falls on a photocell. The photocell circuit produces a current in microamps (I) which is proportional to the light intensity striking the photocell surface and is converted into Transmittance and Absorbance by the MicroLAB software. These components are all housed in the colorimeter chamber of the MicroLAB interface. Crystal violet (also called methyl violet) is an organic dye (See the reaction equation below). In crystal violet solutions where the molarity of H + is greater than 0.1, the solution color is yellow. If H + molarity less than 0.1, the solution color is purple. Our studies will involve purple solutions of the dye.

In this experiment you will prepare five crystal violet solutions of known concentrations varying from 2.00 x 10 -6 M to 10.0 x 10 -6 M. The Transmittance for each of these standard solutions, for a water blank, and for a crystal violet solution of unknown concentration will be measured. The Transmittance should decrease with increasing solution concentration while the Absorbance should increase with increasing solution concentration. The Spectrum Profile for the absorbance of a CV solution is shown in Figure 2. Note that the maximum absorbance is at the 590 nm wavelength. Figure 2. Absorbance Spectrum Profile for CV. Maximum absorbance occurs at 590 nm. .

One objective in this experiment is to find a mathematical function of Transmittance that is directly proportional to sample concentration. Such a function does exist and forms the basis for what is known as Beer's Law . A graph of this function vs. concentration, which is linear and intersects the origin, is called a Beer's Law plot. To explore this concentration/Transmittance relationship, you will first Hand Enter some "ideal" simulated data into the spreadsheet program. You will compute and graph several current functions (the square, reciprocal and logarithm) until a function is that has a direct proportionality to concentration is found. This function is then used for creating a Beer's Law plot from your actual experimental data. Finally, the concentration of the unknown crystal violet solution is obtained directly from this plot. 1.1 Beer-Lambert Law in terms of Absorbance Beer-Lambert Law, more commonly known as Beer's Law, states that the optical absorbance of a chromophore in a transparent solvent varies linearly with both the sample cell path length and the chromophore concentration. Beer's Law is the simple solution to describing the interaction of light with matter. In practice, Beer's Law is accurate enough for a range of chromophores, solvents and concentrations, and is a widely used relationship in quantitative spectroscopy. Absorbance is measured in a spectrophotometer by passing a collimated beam of light at wavelength λ through a plane parallel slab of material that is normal to the beam. For liquids, the sample is held in an optically flat, transparent container called a cuvette. Absorbance (A λ ) is calculated from the ratio of light energy passing through the sample (I 0 ) to the energy that is incident on the sample (I): A λ = -log(Transmittance) = -log (I/I 0 ) Beer's Law follows: A λ = ε λ bc ε λ = molar absorptivity or extinction coefficient of the chromophore at wavelength λ (the optical density of a 1-cm thick sample of a 1 M solution). ε λ is a property of the material and the solvent. b = sample path length in centimeters c = concentration of the compound in the sample, in molarity (mol L -1 )

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Note the subscript in the absorbance factor, A in the Beer’s law equation. The absorbance of a solution depends on the wavelength or color of the incident light. Substances do not absorb the same amount of light at all wavelengths. At specific wavelengths absorbance reaches a maximum and this particular wavelength is max . All measurements of concentrations must be made at this wavelength. In an absorbance experiment, light is attenuated not only by the chromophore, but also by reflections from the interface between air and the sample, the sample and the cuvette, and absorbance by the solvent. These factors can be quantified separately, but are often removed by defining I 0 as the light passing through a sample "blank" or "baseline" or reference sample (for example, a cuvette filled with solvent but zero concentration of the chromophore is used as the blank). Many factors can affect the validity of Beer's Law. It is usual to check for the linearity of Beer's Law for a chromophore by measuring the absorbance of a series of standards. This "calibration" can also remove errors in the experiment, the equipment, and the batch of reagents (such as cuvettes of unknown path length). 2.0 SAFETY PRECAUTIONS Crystal violet solutions may cause skin and eye irritation. Wear safety goggles at all times. Wear gloves whenever handling chemicals including water. As usual, wash hands with soap and water before leaving the laboratory. Crystal violet stains, so appropriate care should be taken in its handling. 3.0 EXPERIMENTAL PROCEDURE 3.1 Simulated concentration/current data treatment 1. Open Microsoft EXCEL and enter the following "ideal" simulated data into two columns. Use Column A for the concentration values and label as Molarity. Use Column B for the current values and label as Current (I). Concentration (M) Current (I) 0.00 x 10 -6 600.0 2.00 x 10 -6 476.4 4.00 x 10 -6 378.6 6.00 x 10 -6 300.6 2. After entering the data, generate an x-y scatter plot with transmittance on the y-axis and concentration on the x-axis . Right-click on any data point and select Add Trendline . Choose a Linear Curve fit as the type of plot desired, and EXCEL will fit the data with the best possible straight line. Select the boxes to “Display Equation on chart” and “Display R - squared value on chart”. Note in this case, the data points would be better fit by a curved line, and hence, the correlation (Corr.) is poor. An acceptable Correlation value for a linear plot would be at least 0.999. Print this graph for your report.

3. In separate columns on the spreadsheet, transform the transmittance data to I 2 , 1/I, and Log(I). Make a separate plot using each of these values on the y-axis, and concentration on the x-axis. When you discover the linear plot, leave the graph on the screen and wait until all groups arrive at the same point. Your instructor will then discuss the modifications needed to give a straight line with a positive slope which also passes through the origin . 4. Make notes on the development of Beer's Law in your notebook. 3.2 Solution Preparation A stock 10.0 x 10 -6 M crystal violet (CV) solution is provided. Perform dilution calculations to complete the following table (see your text if needed) and have your work checked for correctness by your lab instructor. Dilution Table for Standards CV Std #1 CV Std #2 CV Std #3 CV Std #4 CV Std #5 Desired CV Molarity 2.00 x 10 -6 4.00 x 10 -6 6.00 x 10 -6 8.00 x 10 -6 10.00 x 10 -6 mL CV Total mL 50.00 50.00 50.00 50.00 50.00 Absorbance Transmittance Use graduated cylinders, pipets, or burets to transfer the CV solution volumes calculated in the above table into a series of dry 50 mL volumetric flasks. Use great care in your measurements, as careless work will be rewarded with unacceptable Beer's Law plots and the necessity of repeating the experiment. Add deionized water to the mark. When all five solutions have been prepared, stopper or parafilm the flasks and mix by gently inverting them several times to avoid bubbles. Label each solution according to the above numbering. 3.3 Measurements 5. Select Spectrophotometer Experiment from the Main MicroLAB Menu to carry out the experiment. Your data will come from the MicroLAB Interface. 6. Your first sample should be the water blank (0.00 M ). Fill the colorimeter vial about three- fourths full, wipe the outside faces with a KimWipe, insert the vial into the colorimeter and place the light cover over the vial . The program will automatically shift to Read Knowns after the Blank is read. 7. For each sample of CV, rinse the vial three times with small amounts of the solution, then fill the vial three-fourths full, insert it into the colorimeter, place the cap over the vial and click the ADD button.

8. The program will direct you to enter the concentration. It will then display the transmittance for that sample. Follow this procedure for each of the standard solutions in order of increasing concentration. Be sure to wipe the vial with a KimWipe each time. 9. Although the software records the transmittance and absorbance, note these quantities and concentrations in your notebook also. Make a table of concentration vs. absorbance and transmittance. 10. For your unknown solution, click on Read Unknowns and measure the transmittance in the same manner as was done for the standards. Note the Unknown letter in your notebook . 11. Using the digital calipers, determine the path length of one of your sample vials. Record this value in your notebook . 4.0 CLEAN UP 12. Pour all crystal violet solutions into the designated waste container. DO NOT POUR ANYTHING DOWN THE SINK! 13. Discard empty vials in the broken glass box. 14. Wash all glassware with detergent and place above sink to dry. 15. Exit MicroLAB software, and shut down the computer. 5.0 DATA ANALYSIS 16. Obtain printouts of: • Each of your Hand Entered data graphs, as modified by EXCEL for each mathematical transformation. • The dilution table used to make your CV solutions. • Your table of absorbance/trasmittance values for each CV and unknown solution (from MicroLAB). • An example spectrum profile (from MicroLAB). • Your Beer's Law plot made in EXCEL with your data. 17. Use your Beer's Law plot regression equation to calculate the concentration of the unknown, and then compare it against the value that was given in MicroLAB. Show your calculations for the concentration of your unknown. As discussed by your instructor, Beer's Law is expressed by the equation A =  bc. In this equation, A is the absorbance of the solution, b is the cell path length (measured with calipers), c is the molar concentration of the light-absorbing solute (CV), and, is a proportionality constant called the molar absorptivity. 18. Calculate the molar absorptivity for each of the five standard CV solutions and determine the average and standard deviation of the five values. Watch significant figures and be sure to include proper units. Note how well your values agree with one another. After all, it is supposed to be a constant!

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19. Compare the numerical value of your average molar absorptivity with the slope of the best straight line given on your Beer's Law plot. 6.0 REPORT Prepare a title page and abstract according to the format detailed in the syllabus. Complete and include the calculations and data tables as instructed in Section 5.0 above. Also include in the report answers to the questions/problems listed below in Section 7.0. 7.0 Questions and Problems 1. When iron (III) chloride is added to a solution of salicylic acid a purple solution forms. The photocell current read with the colorimeter for a series of standard solutions is given below. Convert these values to absorbencies. (The blank current is 100.0.) FeCl3 Transmittance Absorbance 40.0 x 10 -5 M 17.9 32.0 x 10 -5 M 25.0 24.0 x 10 -5 M 35.7 16.0 x 10 -5 M 50.2 8.0 x 10 -5 M 70.8 2. What would be the effect of the following on the Beer's Law plot? a. At the higher concentrations used to construct the plot, some solute molecules react with one another to form a species that doesn't absorb visible light. b. A fingerprint is present on the solution cell where the light beam passes through. c. Some solute molecules react with one another to form a species that doesn't absorb visible light.

Additional Reading Dye Structure and Color What dyes are: In short, dyes are colored, ionizing, aromatic organic compounds. It must be appreciated that they are individual chemicals, and like all chemicals, they are similar in their reactions to some other chemicals, and distinctly different from others. It may seem that this is a statement of the obvious, but we sometimes appear to view dyes as something other than ordinary chemicals, and I want to stress that the same rules that apply to sodium chloride, acetic acid, benzidine and a host of others also apply to dyes. This includes the possibility that they are toxic. They may be carcinogenic or mutagenic, or harmful to your health in some other way. Just because we call something by an appealing name and because it has an appealing color, does not make it any the less harmful. Handle dyes with care! Put your own safety first! What color is: Dyes are aromatic organic compounds, and as such are based fundamentally on the structure of benzene. To us, benzene appears to be a colorless fluid. In fact it absorbs electromagnetic radiation just as dyes do, but it does so at about 200 nm so that we do not see it. The perception of color is an ability of some animals, including humans, to detect some wavelengths of electromagnetic radiation (light) differently from other wavelengths. Normal daylight, or white light, is a mixture of all the wavelengths to which we can respond and some to which we cannot, in particular the infra-red and ultra-violet rays. We respond to wavelengths between about 400-700 nm. When an object absorbs some of the radiation from within that range we see the waves that are left over, and the object appears colored. In reality this range we see makes up only a very small fraction of the electromagnetic spectrum. In scientific terms there is nothing special about the wavelengths in the visible range, other than being the major components of sunlight which are not removed by the earth's atmosphere. Their special importance is based exclusively on the ability of human retinas to respond to them, and to discriminate between them to a significant degree. These discriminations are what we call color. Wavelengths just outside the visible range are considered colorless, even though there is no substantive difference between them and the limiting wavelengths inside the range. Some animals (bees, for example) can see these other wavelengths but, because humans do not, we consider them colorless. The point is that color is a subjective phenomenon, and thinking of color as something objective is misleading. For that reason, we should refer to the wavelengths involved rather than describe the human response to them.

When some of the wavelengths found in white light are absorbed, then we see what is left over as colored light. The color that we see is referred to as the complementary color of the color that was removed. For instance, if the red rays are removed from white light, the color we detect is blue-green. Blue-green is complementary to red, and red is complementary to blue-green. Complementary Colors Removed Observed Removed Observed Yellow-green Violet Violet Yellow-green Yellow Blue Blue Yellow Orange Cyan Cyan Orange Red Blue-green Blue-green Red Purple Green Green Purple Absorption of light leaving the complementary color The perception of color is merely a human evolutionary adaptation to the absence of some wavelengths in white light. Suppose however, that the same thing happens outside the range to which our eyes respond. Suppose a chemical removes radiation which has a wavelength about 200 nm, as benzene does. Is this colored? Well, obviously it is not. There has been no impact on the wavelengths of white light, so it is not colored. However, what has happened to the ultra-violet radiation that was absorbed by the benzene is no different to what happens to the blue radiation that is removed when we look through a solution of acid fuchsin. The difference is solely whether we can detect it visually. Color is not an objective phenomenon; it is the human detection and perception of electromagnetic radiation. Why dyes are colored Color in dyes is invariably explained as a consequence of the presence of a Chromophore . Since, by definition, dyes are aromatic compounds their structure includes aryl rings which have delocalized electron systems. These are responsible for the absorption of electromagnetic radiation of varying wavelengths, depending on the energy of the electron clouds. For this reason, chromophores do not make dyes colored in the sense that they confer on them the ability to absorb radiation. Rather, chromophores function by altering the energy in the delocalized electron cloud of the dye, and this

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alteration results in the compound absorbing radiation from within the visible range instead of outside it. Our eyes detect that absorption, and respond to the lack of a complete range of wavelengths by seeing color. Chromophores are atomic configurations which can alter the energy in delocalized systems. They are composed of atoms joined in a sequence composed of alternating single and double bonds. Double bonds in organic compounds can be of two types. If the atoms with double bonds are not adjacent, they are termed isolated double bonds, and exist independently of other double bonds in the same molecule. If adjacent atoms have double bonds they are termed conjugated double bonds and the bonds interact with each other. Chromophore configurations often exist as multiple units, having conjugated double bonds, and are more effective when they do so. This is due to the interaction between the double bonds, which causes partial delocalization of the electrons involved in the bonds. In this case, although specific atoms are involved in the bonds, the electrons are distributed over a larger area than the specific atoms and also involve adjacent atoms that have double bonds. The point of this is that conjugated systems have partially delocalized electrons, and the energy in these delocalized electrons can impact on the energy of the delocalized electrons of the parent aromatic compound by extending the number of electrons involved in the system and the energy needed to keep the whole system in place. Conjugated double bonds (red) Another common chromophore is the nitro group. This chromophore is a nitrogen with two oxygen atoms attached. One oxygen is shown attached with a single bond, the other with a double bond. In fact, like the carbon atoms in benzene, these two oxygen atoms are attached to the nitrogen with bonds of equal strength. The extra electrons are delocalized between the three atoms. Nitro group The quinoid ring is found in many dyes. It is a ring structure with two points at which chromophores can attach. It should be thought of as a closed system of conjugated double bonds. The attachment of configurations which add delocalized electrons to the system at one point, and the attachment of configurations which extend the atoms involved in the delocalization at the other, causes very dramatic shifts in the wavelengths which these compounds absorb. For this reason, quinoid ring configurations are considered to be extremely powerful chromophores, producing very intensely colored compounds. Ortho- and para- quinoid ring chromophores To sum this up, chromophores are atomic configurations that contain delocalized electrons. Usually they are represented as nitrogen, carbon, oxygen and sulfur that have alternate single and double bonds. By incorporating the delocalized electrons in these configurations into the delocalized electrons in the aryl rings of aromatic compounds, the energy contained in the electron cloud can be modified. If the energy incorporated into the electron cloud is changed, then the wavelength of the radiation it

absorbs will also change. If this change in the wavelength to be absorbed is sufficient to cause any absorption at all within the visible range, then the compound will be colored. Below are the usual chromophores seen in histological dyes: – – C=C – – C=N – – C=O – – N=N – – NO 2 Quinoid rings Other effects There is more than one effect when chemical groups are attached to aryl rings. Any delocalized electrons and their energy can simply be added to that already present, thus increasing it. Also, the delocalized electrons may be shared by more atoms than those in the original structure, by adding in to the delocalized system the atoms in the chromophore and any modifiers that may be present. Whether these effects occur jointly or separately, the final impact is an alteration in the overall energy of the electron cloud with a subsequent effect on the wavelength of radiation to which the whole molecular system reacts. Another possibility is that electrons may be removed from the electron cloud, and this may result in loss of color. Removal of electrons may cause the remaining electrons to revert to local orbits. A good example would be Schiff's reagent. When sulfurous acid reacts with pararosanilin, a sulfonic group attaches to the central carbon atom of the compound. This disrupts the conjugated double bond system of the quinoid ring, causing the electrons to become localized and the ring to cease being a chromophore. Consequently, the dye becomes colorless. Auxochrome Below are the usual auxochromes found in histological dyes: – – NH 3 – COOH – HSO 3 – OH Auxochromes are groups which attach to non-ionizing compounds yet retain their ability to ionize. While this definition is largely correct, it is also inadequate. This is because it restricts the definition of the auxochrome to ionization, and does not comment on the effect of auxochromes on the absorbance of the resulting compound. The word auxochrome is derived from two roots. The prefix auxo is from auxein , and means increased. The second part, chrome means color, so the basic meaning of the word auxochrome is color increaser. This word was coined because it was noted originally that the addition of ionizing groups resulted in a deepening and intensifying of the color of compounds.

Color enhancing by an auxochrome To the left is naphthalene , a colorless compound. The addition of a single hydroxyl group to naphthalene produces 1-naphthol which is also a colorless compound, but one which can ionize. If instead of a hydroxyl group we add the nitro group, which is a chromophore, we get the compound 2,4- dinitronaphthalene . The addition of this chromophore has caused it to become pale yellow. If instead of a hydroxyl or nitro groups, both a hydroxyl and nitro groups are added, we get the deep yellow dye, martius yellow . The addition of both an auxochrome and a chromophore results in a much stronger alteration of the absorption maximum of the compound. The hydroxyl group must have deepened the color, showing that auxochromes are also chromophores. Sometimes the term auxochromophoric is used to denote the action of an auxochrome that modifies the color as well as permitting ionization. This term infers that the color modifying effects of auxochromes are rare, but this is not the case. The effect on absorption should not be considered an incidental aspect of the auxochrome's action, but an integral and fundamental part of it. The word auxochromophoric is redundant. Edward Gurr proposed the terms colligator and non-colligator to distinguish between ionizing auxochromes and color modifying effects. Auxochromes are of two types. They may have a positive charge as the amino group and its substituted variants. Or they may be negatively charged as the carboxyl and hydroxyl groups, and the sulfonic group. This last is commonly used to convert basic dyes to acid dyes. Both negatively charged and positively charged auxochromes may be present on a single molecule. Resonance The process by which electrons are stimulated by radiation is resonance . It should be made clear at the outset that resonance is not the same as vibration. Resonance is the induction of a response in one energy system from another energy system in close proximity, which is operating at the same energy level (frequency). In aromatic organic compounds,

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including dyes, the two energy systems are electromagnetic radiation , and the delocalized electron cloud . Do not confuse resonance with the imaginary resonance hybrids used in an older explanation for the structure of benzene. Modifiers Absorption of benzene compared to some dyes Color modifiers such as methyl or ethyl groups alter the color of dyes by altering the energy in the delocalized electrons. By themselves they cannot do this enough to cause absorption in the visible range, but they can affect the shade significantly when absorption is already in that range. Adding more of a particular modifier results in a progressive alteration of color. Compounds that differ from each other in this kind of regular fashion are called homologues . A very good example is seen with the Methyl violet series. Alteration of color by modifiers Without any methyl groups the parent dye is called pararosanilin and is red. When four methyl groups are added we get the reddish purple dye methyl violet . As more methyl groups are added we get the purple blue dye crystal violet which has six such groups.

The Color The color of the dye is caused by the absorbance of electromagnetic radiation. We have constantly referred to the wavelength that is absorbed in the singular, but a simple scan of a dye solution with a spectrophotometer shows that dyes do not remove a single wavelength. Rather they absorb radiation on either side of the wavelength most completely removed (the absorption maximum). Plotting the wavelength absorbed against the degree of absorbance usually results in a display resembling a bell curve. If any part of this curve is in the visible range, the dye will appear colored. White light is a mixture of wavelengths. Some of these have a relationship to the energy in the delocalized electron cloud of the dye molecule. By the process of resonance, previously described, the electron cloud will respond to the energy contained in that radiation by absorbing it, and removing it from the spectrum. As a consequence the white light will cease to be white and will display the colors of the wavelengths left over. The transmitted light will have the complementary color to the wavelengths removed. The Effect When light illuminates a dye, some of it is absorbed as energy. Since energy is not destroyed, something must then happen. We could use an analogy of heating water - the water's temperature rises, molecular vibration increases. However, we do not see anything else, as the water just sits there being water. The same can happen with dyes, we may not observe anything particular as the effect may be at the atomic level. There are several possibilities, however. A. The energy level in the electrons in an unaffected dye is called the ground state. When electromagnetic radiation, as light energy, is absorbed the electrons become more energized. B. With most dyes, there is then a gradual decay and the electrons return to the ground state. We do not see anything. Nevertheless, something may happen that we do not see. Possibly there is an increase in temperature, or some chemical changes occur which disrupt the dye's structure and cause it to lose color - fading . C. Another possibility is that the return to the ground state is not gradual, but sudden. If this is accompanied by emission of any residual energy in the form of light, we observe the dye glowing - fluorescence . Since the emitted light must always contain less energy than the absorbed light, as some was used to energize the electrons, the emitted radiation is always at longer wavelengths than the absorbed radiation. By manipulating the light available, we can cause ultra-violet light to be absorbed and visible light to be emitted. If a seventh methyl group is added, the resulting dye is methyl green .

D. A third possibility is that the electrons stabilize in their newly energized state. After a passage of time they then return to the ground state. If this happens gradually, we may observe nothing, with the same possibilities regarding fading and temperature increase as before. E. If return to the ground state happens suddenly, and the residual energy is emitted as light, we once again see the dye glowing - phosphorescence . As with fluorescence, the light emitted is always a longer wavelength than the light absorbed, but the disparity is greater with phosphorescence due to the greater energy consumed in keeping the electrons in the excited state. Note that the difference between fluorescence and phosphorescence is in whether the electrons stabilize in the excited state before returning to the ground state. With any stability, no matter how long (or short), it is considered phosphorescence. Conclusion The explanation of the relationship between structure and color depends on the basic atomic structure of the aryl ring, and the shared or delocalized electrons that this atomic arrangement has. The ability to absorb radiation is inherent in this structure. The effect of other atomic configurations is to modify the energy contained in the delocalized electron cloud so that the compound absorbs electromagnetic radiation at a wavelength in the visible range. Some also ionize, enabling the compound to chemically react with ionizing tissue groups. Color, fading, fluorescence and phosphorescence are all seen to be different effects of the same fundamental process.

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