Laboratory handbook Biochemistry CHML341 Medgar Evers College Department of Chemistry and Environmental Science By Dr

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Laboratory handbook Biochemistry CHML341 Medgar Evers College Department of Chemistry and Environmental Science By Dr. Harsha Rajapakse Content Experiment 1: Use of Pipetmen 1 Experiment 2: Spectroscopy and Dilution 7 Experiment 3: Preparation of Buffers 17 Experiment 4: Titration Curve of Amino Acids 25 Experiment 5: Gel Exclusion Chromatography 43 Experiment 6: Factors affecting enzyme catalyzed reactions 49 Experiment 7: Enzyme Kinetics; Determination of KM and 61 Vmax of the Enzyme Alkaline Phosphatase Experiment 8: Product Inhibition of Alkaline Phosphatase 71 Experiment 9: Lipids 83 Experiment 10: Ripening Bananas: What's Happening and 91 Can We Control It? Experiment 11: Assay of Tabletop Sweeteners Using a 99 Modified Biuret Reagent Experiment 12: Agarose gel electrophoresis 111 Experiment 1: Use of Pipetmen Prelab Questions 1. Your lab partner hands you a P200 Pipetman set as shown in the diagram at right. What volume is it set for? Is this the proper Pipetman for this volume? Why (or why not)? 2. Write the numbers read on each pipette. Indicate the unit Page 1 Page 2 Experiment 1: Use of Pipetmen In biochemistry, the ability to accurately and reproducibly measure and transfer small volumes of liquids is critical for obtaining useful results. For volumes less than 1 ml, the most common method for measuring liquid volumes involves the use of a device known as a pipetman. (Note: “Pipetman” is the brand name of the most commonly used of these types of pipets; however, all of these pipetting devices work on similar principles.) A drawing of a pipetman is shown at right. The devices you use may not look exactly like the one shown. The pipetmen used in this course come in three different types: P1000, P200, and P20. P1000 are useful for volumes from 200 to 1000 µl. P200 are useful for volumes from 20 to 200 µl. P20 are useful for volumes from 0.5 to 20 µl. Make sure that you are using the correct pipetman for the volume you need. Also, make sure that the pipetman is actually set for the volume you need by looking in the “volume window”, and, if necessary, turning the “volume control knob” until the pipetman displays the correct volume (the pipetmen do not read your mind; because several people will use the pipets, they may not always be set as you expect them to be). Do not attempt to set pipetmen for volumes larger than their maximum, or for volumes less than zero; doing so will damage the pipetman. All pipetmen use disposable tips (do not pipet liquids without using the appropriate tip, because this will contaminate the pipetman and may damage it). When attaching the tip, make certain that the tip is the correct type for the pipetman you are using, and that the tip is properly seated on the end of the pipetman. Try depressing the plunger. As the plunger depresses, you will feel a sudden increase in resistance. This is the first “stop”. If you continue pushing, you will find a point where the plunger no longer moves downward (the second “stop”). When using the pipet, depress the plunger to the first “stop”, place the tip into the liquid, and in a slow, controlled manner, allow the plunger to move upwards. (Do not simply let the plunger go; doing so will cause the liquid to splatter within the tip, resulting in inaccurate volumes and in contamination of the pipet.) Now, take the pipetman (carrying the pipetted liquid in the tip) to the container to which you wish to add liquid. Depress the plunger to the first, and then to the second stop. If you watch carefully, you will note that depressing to the second stop expels all of the liquid from the tip. (Actually, this is true for most aqueous solutions. In some cases, however, such as for organic solvents, or for solutions containing large amounts of protein, it is often difficult to get all of the liquid out of the tip. In these cases, it is best to “wet” the tip, by pipetting the original solution once, expelling it, and then taking up the liquid a second time.) Page 3 Although pipetmen are tremendously useful, they have a potential drawback. If used improperly, pipetmen will transfer inaccurate volumes. In addition, pipetmen may lose calibration. If used incautiously, therefore, pipetmen may yield misleading or even totally useless results. Checking the calibration of pipetmen is a simple procedure that can save considerable time, energy, and reagents. In this experiment, you will learn how to use pipetmen of various sizes and measure their accuracy, precision, and calibration. Experimental Procedures Materials Pipetmen Pipet tips Water Balances Weigh Boats Methods For each Pipetman you will want to check the accuracy and precision over the entire range using at least two different volumes (for example: if you are using a P1000 you will want to check 300 µl and 1000 µl). Accuracy is a measure of proximity to the true value or the expected value for a measurement. Precision is a measure of reproducibility. For example, obtaining weights of 0.5, 1.0, and 1.5 grams for a P1000 set for 1000 µl would be accurate, but not precise; obtaining 0.67, 0.68, and 0.67 g for the same setting would be precise, but not accurate. For P100 and P1000 1. Acquire Pipetmen and the correct size tips. 2. Place a weigh boat on the balance and tare the weight to zero. 3. Draw up the designated volume of deionized water into the pipet tip and dispense it onto the weigh boat. Record the weight of the water added. 4. Repeat the procedure twice for each volume (yielding a total of three weights for each volume). For P20 The P20 uses very small volumes, which have very small weights. In order to obtain accurate readings with the relatively low precision balances available, you may need to pipet the volume of water several times (5 or 10 times is recommended) for each volume being tested. For example, pipet 10 µl 10 times in succession, and record the weight of the 100 µl total volume as one measurement. Analyze the data you collect as described in the section on data reduction (below). Page 4 Experiment 1: Use of Pipetmen Data Page Name: P1000 Trial Weight of 300µl of water 1 2 3 Average P100 Trial Weight of 50µl of water 1 2 3 Average Weight of 1000µl of water Weight of 100µl of water P10 Trial Weight of 50µl of water (5ul *10) Weight of 100µl of water (10ul*10) 1 2 3 Average Answer the following six questions for each volume measured in the table given: 1. Record the weight you measured for the three trials. 2. Average the three weights. 3. Calculate standard deviation for your average.1 4. What is the standard deviation as a percent of the average value? 5. Is the Pipetman accurate at this volume? 6. Is the Pipetman precise at this volume? P1000 P100 P10 Average Weight Std Std as a percentage of the average Is the pipetman accurate at this volume Is the Pipetman precise at this volume Do the accuracy and precision vary over the range of volumes for the Pipetmen? Page 5 S= Standard deviation, X= measurement, X bar = mean, n=the number of measurements. Most calculators and all spreadsheets will calculate standard deviation. Note, however, that some spreadsheets also have algorithms for “population standard deviation” (in which the denominator is “n” rather than “n–1”). The use of population standard deviation is inappropriate for these data. In Excel, the correct function is “STDEV(cell range)” Citation/Reference http://www.rose-hulman.edu/~brandt/publications/422_Manual_3rd_Ed.pdf Page 6 Experiment 2: Spectroscopy and Dilutions Prelab Questions 1. What is the reliable range of absorbance measurements? 2. What is the unit of absorbance? 3. Can you directly compare absorbance readings of the same sample measured by different people using different instruments? 4. You have a 0.5 M stock solution of Tris base. How would you make 100 ml of 0.03 M Tris base? 5. If you perform a 1:4 dilution on 50 mM Tris base, what is the final concentration? 6. A student diluted 1ml of a stock solution (1M) of a compound with 9 ml of water. Then took 1ml from that solution and diluted with another 9ml of water. Then took 10ml of the resultant solution and added 90ml of water. What is the final concentration of the solution? 7. Is the extinction coefficient for a molecule the same for all wavelengths? 8. After obtaining calibration curve for your known solution to have slope of 2.5 and intercept of 0.01, you determine that your unknown has Absorbance of 0.2. Is it enough to determine the concentration of your unknown and how? 4. You prepare several dilutions of an unknown compound. You measure the absorbance of each solution at 340 nm using a 1 cm cuvette (your results are listed in the table below). What is the extinction coefficient (in (M•cm)-1) of the compound? (Hint: assume that each of the individual values contains some degree of experimental error.) Concentration (µM) Absorbance at 340 nm 2 4 8 16 32 0.009 0.020 0.061 0.111 0.189 Page 7 If an unknown concentration of the same compound gave an absorbance of 0.08, find its concentration using the above standard curve you prepared. Page 8 Experiment 2: Spectroscopy and Dilutions A spectrophotometer is an instrument for measuring the absorbance of a solution. Absorbance is a useful quantity. The Beer-Lambert law states that: A = ecl where A is the absorbance of the sample at a particular wavelength, is the extinction coefficient for the compound at that wavelength in (M•cm)-1, c is the molar concentration of the absorbing species, and l is the path length of the solution in cm. Thus, if the extinction coefficient of an absorbing species is known, the absorbance of the solution can be used to calculate the concentration of the absorbing species in solution. (This assumes that the species of interest is the only material that absorbs at the wavelength being measured.) The above is an explanation of why we measure absorbance: absorbance allows us to calculate the concentration of compounds in solution. However, it does not explain what absorbance is. Another definition of absorbance is: A = log æ I0 ö è Iø where I0 is the amount of light entering the sample, and I is the amount of light leaving the sample. Absorbance is therefore a measure of the portion of the light leaving the lamp that actually makes it to the detector. A little thought will reveal that when absorbance = 1, only 10% of the light is reaching the detector; when absorbance = 2, only 1% of the light is reaching the detector. The typical internal arrangement of a Spectrophotometer Absorbance values greater than 1 are unreliable, because too little light is reaching the detector to allow accurate measurements. When measuring absorbance, note the values; if the reading is greater than 2, dilute the sample and repeat the measurement. Spectrophotometers measure the decrease in the amount of light reaching the detector. A spectrophotometer will interpret fingerprints on the optical face of the cuvette, or air bubbles, or objects floating in your solution as absorbance; you therefore need to look carefully at your cuvette before putting it into the spectrophotometer to make sure that your reading is not subject to these types of artifacts. Cuvettes are usually square objects 1 cm across (as shown in the above figure). In some cases, the liquid reservoir is not square; in those cases, make sure that the 1 cm dimension is aligned with the light path (note the orientation in the diagram above.) Page 9 Some cuvettes are designed for visible light only. When the spectrophotometer is set for ultraviolet wavelengths (wavelengths of 340 nm or less) make sure that your cuvette does not have a large absorbance when it contains only water. The term “spectroscopy” comes from the word “spectrum” which originally referred to the multiple colors of light apparent in an analysis of white light using a prism. “Spectroscopy” therefore implies the use of multiple wavelengths of light. Spectrophotometers have the ability to specifically measure absorbance at specific wavelengths. The most commonly used method to allow this involves a “monochromator”, a device (either a prism, or more commonly, a diffraction grating) that splits the incident light into its component wavelengths, and allows only light of the desired wavelength to reach the sample. The ability to measure absorbance at different wavelengths is very useful, because the extinction coefficient of a compound varies with wavelength. In addition, the absorbance spectrum of a compound can vary dramatically depending on the chemical composition of the compound, and depending on the environment (such as the solvent) around the compound. The graph at right shows the absorbance spectrum of a protein. The protein has a strong absorbance peak near 280 nm, but exhibits very little absorbance at longer wavelengths. For this protein, the only chromophores (chemical groups within a compound that absorb light) are the aromatic amino acids tryptophan and tyrosine. For many proteins, these two residues are the only chromophores; because tryptophan and tyrosine only absorb in the ultraviolet portion of the spectrum, such proteins are colorless molecules. Colored proteins, such as hemoglobin, exhibit their color due to chromophores (heme, in the case of hemoglobin) that absorb in the visible portion of the spectrum. The extinction coefficient of a molecule at a given wavelength can be calculated using the BeerLambert equation from absorbance measurements for solutions of known concentration. Dilutions Many solutions used in biochemistry are prepared by the dilution of a more concentrated stock solution. In preparing to make a dilution (or series of dilutions), you need to consider the goal of the procedure. This means that you need to consider Page 10 both the desired final concentration and required volume of the diluted material. A simple equation allows the dilution to be calculated readily: C1V1 = C2V2 where C1 is the concentration of the initial solution; V1 is the volume of the initial solution available to be used for dilution (this may not be the total volume of the initial solution, and instead may be a small fraction of the initial solution), C2 is the desired final concentration, and V2 is the desired final volume. In most cases, the initial concentration and the final concentration are either known or are chosen in order to work correctly in the experiment being planned. The final volume is usually an amount that is chosen based on the amount required for a given experiment. This means that at least three of the required terms are either known or can be chosen by the experimenter. Let us consider an example. You are setting up a standard curve. You have a stock solution of 1000 µg/ml BSA, and for one of the points on the curve, you want 200 µl of 20 µg/ml. In this case, C1 = 1000 µg/ml; C2 = 20 µg/ml, and V2 = 200 µl. This leaves V1 as the unknown value (i.e. how much of the stock solution must be diluted to 200 µl final volume to yield the desired concentration). Rearranging the dilution equation gives: V1 = V2 * C2/C1 4µl = 200 µl * 20 µl/ml 1000 µl/ml Thus, you need to dilute 4 µl of the stock solution to a final volume of 200 µl (i.e. by adding 196 µl). If, in the example, you wished to make a solution of 1 µg/ml, the same equation would indicate that you need 0.2 µl of the 1000 µg/ml stock solution for 200 µl of the final diluted sample. This is a problem: 0.2 µl is very difficult to measure accurately. You have two choices: change the final volume (i.e. if V2 is larger, then V1 must also increase), or perform serial dilutions (i.e. instead of diluting the stock solution by a factor of 1000 in one step, dilute the stock solution, and then make a further dilution of the diluted stock). In many cases, while the final concentration is important, the final volume is not (as in the previous paragraph). In these cases, do what was explained in this example: use a convenient dilution: a dilution that involves volumes that are easily pipetted. Pipetting 1.3333 µl is usually less accurate than pipetting 4 µl, both because 4 µl is a larger volume, and because it is difficult to set the pipet for 1.3333 µl. In this case, 4 µl is a convenient volume, while 1.3333 µl is not. In some cases, you may not know the actual starting concentration. If, for example, you need to measure the enzyme activity in a sample, and you find that the activity is too high to measure accurately, you will need to dilute the starting material. Since you don’t know the actual starting concentration, all you know is the concentration ratio between starting and final solutions. As long as you keep track of the concentration ratio in all of your dilutions, you can easily determine the enzyme activity in the initial solution, even though you cannot measure it directly. Page 11 Concentration ratios are frequently of considerable value. For example, you have a stock solution of buffer that contains 450 mM Tris-HCl, 10 mM EDTA, and 500 mM NaCl. You actually wish to use a final concentration of 45 mM Tris-HCl, 1 mM EDTA, and 50 mM NaCl. In each case the concentration of the final buffer is one-tenth that of the original. Simply performing a 1:10 dilution of the stock solution then gives the appropriate final concentration of each component. The stock solution of buffer is typically called a 10x stock, because it is ten-times more concentrated than the final, useful buffer. Note, in the previous paragraph, the “1:10” dilution. The description uses the chemistry convention for this term, which will be used throughout this course. The 1:10 dilution mentioned is performed by taking one part of the initial solution, and adding nine parts of solvent (usually water). This results in a final concentration that is ten-fold lower than the original. In this experiment, you will learn how to prepare solutions using dilutions, and learn how to use a spectrophotometer. Experimental Procedures Materials Pipetmen Pipet tips Parafilm Water 1.0 M CuSO4 solution UV cuvettes Vis cuvettes Quartz cuvettes Unknown CuSO4 solution Simple Dilutions-1 1. Your instructor has set the spectrophotometer to wavelength 700 nm 2. Prepare 1 ml of the following dilutions of the CuSO4 solution using water: 1:2, 1:5, 1:10, 1:50, and 1:100. 3. Determine the A700 for each dilution using water as the blank. Simple Dilutions-2 4. Assuming the CuSO4 solution is a 5X stock, prepare the following solutions: 0.5X, 1X, 2X. 5. Determine the A700 for each dilution. Serial Dilutions 6. Prepare the following solutions of the CuSO4 solution using serial dilutions: 1:5, 1:25, 1:125, and 1:625 7. Determine the A700 for each dilution. Page 12 Experiment 2: Spectroscopy and Dilutions Data Page Name: Simple Dilutions-1 1. Give the volume of water and CuSO4 solution used for each sample dilution. Solution dilution Volume of CuSO4 (ml) 1:2 1:5 1:10 1: 50 1:100 Volume of Water absorbance (ml) 4. Plot the results on an A700 vs. Concentration of CuSO4 solution graph. 5. Calculate the extinction coefficient for aqueous CuSO4 solution. Page 13 Simple Dilutions-2 1. Give the volume of water and CuSO4 solution used for each simple dilution. Absorbance(A700 Solution dilution Volume of Volume of Water ) CuSO4 (ml) (ml) 0.5X 1X 2X Page 14 Serial Dilutions 1. Describe how each serial dilution was prepared. 1:5 1:25 1:125 1: 625 Page 15 Absorbance Results 1. Based on your calculated extinction coefficient, estimate the concentration of an unknown CuSO4 solution with an absorbance of 0.65? Page 16 Experiment 3: Preparation of Buffers Prelab Questions Calculate the grams/ml of the reagents needed to prepare following buffers. MW 1. Na3C6H5O7.2H2O 2. C6H6Na2O7 3. C6H7NaO7 4. K3PO4 5. K2HPO4 6. KH2PO4 7. H3PO4 8. CH3COONa 9. Ch3COOH 10. NaHCO3 11. Na2CO3 - 294.10 gmol-...