- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 1 (52.86 KiB) Viewed 23 times
- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 2 (62.69 KiB) Viewed 23 times
- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 3 (70.66 KiB) Viewed 23 times
- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 4 (60.52 KiB) Viewed 23 times
- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 5 (72.64 KiB) Viewed 23 times
- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 6 (72.56 KiB) Viewed 23 times
- The Absorbance Is Defined As A Log 2 Log T 2 The Absorbance A Of An Ideal Solution Is Directly Proportional To 7 (41.56 KiB) Viewed 23 times
The absorbance is defined as, A=-log--2-log (%T) (2) The absorbance, A, of an ideal solution is directly proportional to the concentration of absorbing ions or molecules in that solution. The equation that relates absorbance to concentration is known as Beer's Law: Where: A=cle (3) e is called the molar extinction coefficient (aka molar absorptivity). It is an attribute that is specific to the absorbing species. It is a measurement of how strongly a chemical species absorbs light at a given wavelength (at a given wavelength, 7., c is a constant) (We will use 2-447 nm.) I is the path length (width of the cuvette) (another constant) which contains the solution that is being analyzed and . . c is the concentration of the absorbing species expressed as molarity, M. (The absorbing species in this experiment is FeSCN. Hence, the concentration, c, of FeSCN is directly proportional to the absorbance of that solution with a y-intercept equal to zero.) A = (constant) c A graph of absorbance vs. concentration is known as a Beer's Law Plot. The constant is the slope of the calibration curve made from the data in Procedure 1A. Once the constant is determined, you can measure the absorbance of any solution and then know its concentration of FeSCN. Let me restate this with direct reference to the procedure for this lab: once the calibration curve has been constructed using Procedure 1B, then any absorbance you measure in Procedure 2 can be used to calculate the concentration. of FeSCN³. It is the concentration of FeSCN³ that is needed to then determine K. for reaction (5). In this experiment, each group will use different cuvettes. It is imperative that the path length I is constant for all of these cuvettes. Therefore, each of these cuvettes is made with a width that is precisely 100 cm. Each cuvette is marked with a white line near the lip. This white line must be aligned precisely with the black groove in the sample holder to ensure that the path length is 1.00 cm. Operating instructions for the Spectronic 20-D are available at the end of this experiment. B. Determining the Equilibrium Constant The main purpose of this lab is to determine the equilibrium constant for the complex ion formed by the reaction of Fe" with SCN"," Fe (aq) + SCN (aq) = FeSCN²³(aq) (5) In this experiment, we will be calculating K. for reaction (5) five different times. Each of those five times, we will get a different number for K.. While K, for this reaction is constant, our attempts to experimentally measure and calculate K. will not be constant. There is experimental error, both procedural and human-made, in this experiment. Therefore, we will take the average of the five experimentally determined values of K. Rev: 2019-2020 As noted in lecture and the text, when K. is>> 1, a reaction is product-favored, goes essentially to completion and the concentrations of products at equilibrium are much larger than the concentration of reactants (essentially, there are no reactants left). When K, is <<1, a reaction is reactant-favored, does not go at all and the concentrations of reactants at equilibrium are much larger than the concentration of products (essentially, no products form).
As a hint to the answer for this lab, the value of K, for reaction (5) is somewhere close to 1 (0.001 K.<1,000), meaning that the reaction gets stuck somewhere in the middle. What we will see is that we can manipulate the concentrations of reactants to make this reaction either go approximately 50% of the way to completion or as far as 99.5% to completion. This is the basis for this experiment. In Procedure 1B, the concentration of Fe" initially, after accounting for dilution, will be the equilibrium or final concentration of FeSCN because the above reaction will go essentially to completion. We will use LeChatelier's Principle (the subject of experiment 5-we're just using it here) to drive the reaction to the right. We will add a known concentration of Fe". We will then add more than 100 times higher concentration of SCN than Fe" to the reaction mixture. Using LeChatelier's Principle. the excess SCN will push the position of equilibrium far to the right and essentially to completion. The Fe" is the limiting reactant and all of the Fe" will be reacted to produce FeSCN³ FeSCN absorbs light at 2=447 nm. We will construct a "calibration" or "standard" curve (a.k.a. Beer's law Plot) that plots the known [FeSCN³] on the x-axis and the absorbance A of FeSCN² on the y-axis to establish a linear relationship between these two variables. The power of the calibration curve is that for any solution containing FeSCN" as the only absorbing species, we can measure the absorbance and use the linear relationship to calculate the [FeSCN³]. This is exactly why we are preparing the calibration curve. For Procedure 2, we add approximately equal concentrations of Fe" and SCN. Under these conditions. the reaction will not go to completion. In fact, we will not know how far the reaction went to completion because we don't know the equilibrium constant K. However, if we measure the absorbance of the unknown solutions in Procedure 3, then we can use the calibration curve prepared from Procedure 1 to calculate what the equilibrium (FeSCNJ is in that solution. Once we have the equilibrium [FeSCN³], we can fill in the rest of an ICE table to determine K.. To summarize, in this experiment we will: 1. Create a calibration curve by measuring the absorbance of a series of solutions of known [FeSCN³]. For these solutions, reaction (5) will be essentially complete. The calibration should yield a straight-line relationship between [FeSCN³] and absorbance. 2. Measure the absorbance of a series of solutions for which [FeSCN] is unknown. For these solutions, reaction (1) will not be complete, but there is still the same relationship between absorbance and concentration. Use this relationship (equation of the line) and the absorbance to calculate [FeSCN³"]. 3. Determine the value of the equilibrium constant K, for reaction (5). II. Experimental A. Equipment Needed: From stock room: 1 spectrophotometer cuvette, 10mL. volumetric flask(s) Equipment in lab: Spectronic 20-D spectrophotometer, digital micropipettes, aluminum foil 0.200 M KSCN in 0.500 M HNO₂. 2.00x10" M Fe(NO) in 0.500 M HNO, 0.500 M HNO, Chemicals in lab: B. Disposal: Combine and collect all used chemicals from this experiment in a large beaker. At the end of the lab period, transfer waste chemicals to the provided waste container. Rev: 2019-2020
C. Experimental considerations 1. The product formed in this experiment is light sensitive. Direct sunlight will slowly decompose the product, but room light will not. Use the aluminum foil to wrap around any flasks with FeSCN solution if you will be waiting for more than an hour to use the spectrophotometer (which will most likely NOT happen). 2. All of the ions in this experiment are dissolved in 0.500 M H HNO, to prevent the side reactions that produce other complex ions. Henceforth, we will not mention this fact. 3. Please do not take any more of the standardized solutions than necessary. D. Before Starting Experimental Work (Before Class) 1. In your notebook, enter the experiment title, date, your name and name of partner. 2. Write the purposes of the lab. 3. Draw and label the major internal components of a spectrophotometer. 4 Explain why the reaction goes to completion in Procedure I but does not go to completion in Procedure 2 even though the value of K. is constant in both procedures. 5. Write an executive summary of the procedures described in this lab. E. Procedure 1A. Preparing Standard Solutions for Absorbance Measurements. The data for this procedure will be pooled together as a class. 1. Turn on the spectrophotometer. It must warm up for 10-15 minutes. 2. Insert Table 1 in your notebook if you have not already done so. 3. Fill your smallest (10mL.) beaker 2/3 full of 2,00x10" M Fe. Use the micropipettes to deliver the volumes listed in Table 1 of 2.00x10³ M Fe" into each of five 10 ml. volumetric flasks. (You may need to use a combination of micropipettes to obtain some volumes.) 4. Add 2.0 ml of 0.200 M SCN to each volumetric flask, then dilute with 0.500 M HNO, exactly to the 10.00 mL mark. Use a dropper to deliver the last few drops of HNO,. 5. Mix by alternately inverting and shaking for at least 30 seconds. F. Procedure IB: Measuring the Absorbance of the Standard Solutions. 1. To standardize the Spectronic 20-D meter: a) Let the spectrophotometer warm up for at least 10 minutes. Set the wavelength dial (on top) to 447 nm. With the cell holder empty, adjust the left knob to zero transmittance. (When the cell holder is empty, the light path is blocked, so no light gets to the detector. We want the meter to recognize this electronic signal as 0%T.) b) Fill a cuvette with 0.500 M HNO). Make sure you fill the cuvette more than halfway. Wipe any fingerprints from the cuvette with a Kim-Wipe (and not a paper towel-the paper towels scratch the cuvettes), and then place it in the cell holder. Ensure that the vertical white mark on the cuvette is aligned with the mark in the spectrophotometer. Close the cover for the sample holder. Adjust the lower right knob until the meter reads 100% transmittance. (This solution is considered a "blank" solution. It does not contain any of the species of interest, FeSCN. We want the meter to recognize this electronic signal as no light absorbed from the species of interest, or 100%T.) c) Switch from Transmittance to Absorbance mode. To scroll between Transmittance and absorbance, press the "Mode" button. Rev: 2019-2020 6-4
2. To gather the data for the Beer's Law Calibration Curve: a) One at a time, starting with the most dilute solution, pour a small amount of each solution to be measured into the cuvette. Condition the cuvette by rinsing the walls of the cuvette with the solution by tilting and rotating much like it was a buret or pipette. Dispose of the liquid in a waste beaker. This is necessary if the cuvette is not dry so that the solution you are putting into the cuvette is not diluted by water or the previous solution left in the cuvette. b) Wipe off fingerprints from the cuvette with a Kim-Wipe. Place the cuvette into the sample holder, replace the cap, and read and record in Table 1 (which should be created in your notebook) the ABSORBANCE, A. (% transmittance may also be recorded but, for our purposes. A has a direct relationship to the quantity of interest, namely, concentration of iron thiocyanate). c) Measure the absorbance A for each of the standard solutions. d) Make a hand graph of your group's data by graphing the diluted concentration of FeSCN² on the x-axis and the measured absorbance on the y-axis. (Note, the calculations for the diluted concentration of FeSCN are described in section III A 1 on page 6.) Each student in the group must make their own graph. Check the linearity of your data by drawing a best-fit line through your data. Calculate the slope of your calibration line using two points on your line and dividing the rise over the run. e) Show your hand graph to your instructor. If any one or 2 point(s) are substantially off of the line, your instructor may require your group to re-prepare the solution and measure its absorbance. After approval of your data, discard the solutions into the container in the hood and record your data on the board to be part of the class data set. f) Record the class data set into your notebook. Perform any Q-tests requested by your instructor for suspect data points and then compute the average absorbance for each solution. G. Procedure 2: Absorbance of Unknown Solutions 1. Insert Table 2 in your notebook. 2. Each group should prepare 100 mL of 4.00×10 M SCN by adding 2.00mL of 0.200M KSCN (using a digital micropipette) to a 100mL volumetric flask, dilute with 0.500M HNO, and mix well. 3. Prepare five solutions with the volumes of 2.00×10 M Fe" and the above prepared 4.00×10 M SCN solution as indicated in Table 2 using the 10.0 mL volumetric flasks. Fill to 10.00 ml with 0.500 M HNO, and mix well. 4. Determine the absorbance of each solution and record the value in Table 2 and on the board to be part of the class data set. This will be used to create an Excel graph. 5. Record the class data set into your notebook. Perform any Q-tests requested by your instructor for suspect data points and then compute the average absorbance for each solution. III. Further Instructions A. Required Calculations and Graphs Approach to calculations for the calibration curve. For the standard solutions, the number of moles of SCN is much greater than the number of moles of Fe". Using LeChatelier's Principle, the excess SCN drives reaction (5) to completion (or >99% of the way to completion). Therefore, for the standard solutions used to prepare the calibration curve, Fe is the limiting reactant and the initial concentration (accounting for dilution to 10 mL) of Fe" is converted entirely to FeSCN Rev: 2019-2020 6-5
Calculating the concentration of a diluted solution. This calculation will be used any time you mix solutions together to calculate the concentration of ions in the new solution. Ex: If 0.100 mL of 2.00×10 M SCN is diluted to a volume of 10.00 ml., what is the final concentration of SCN ion? M₂V₁M₂V₂ (2.00×10 M) (0.100 mL)- M: (10.00 ml.) M: 2.00×10 M Units on concentration and volume can be any units, as long as they cancel out. 1. Calculate the diluted [Fe] in each of the standard solutions prepared during Procedure 1A using MV = M₂V₂. Set this equal to [FeSCN³") in the data table 1. Show an example calculation for the first solution and enter all values obtained into Table 1. 2. Do any required Q-tests first. Any values of absorbance that are considerably different than thei average for a given solution should be tested using the Q test (see Statistical Functions in the introductory lab materials) Show your calculations even if they show that no data were required to be excluded from the data set. Be sure to include a comparison to the critical value of Q in the table and a conclusion to keep or discard the suspect data. 3. Determine the average of the absorbance measurements for each solution of FeSCN³* in Procedure 1A. Only include values that were not discarded by the Q-test. Approach to calculations for Determination of the Equilibrium Constant. For the solutions in Procedure 2, the number of moles of Fe" is approximately equal to the number of moles of SCN". In this case, we will observe that the reaction does not go to completion. When a reaction does not go to completion, we can count on setting up some kind of ICE table. The ICE tables and the equilibrium values of the concentrations will allow us to determine the value for K.. the equilibrium constant. Ideally, all of our calculations will lead to similar values of K... 4. Create Table 3 in your notebook. a.) Calculating the concentration of each diluted solution. • Calculate the initial concentration of Fe" and SCN, accounting for dilution using CIV₁-C₂V₂, for each solution in Table 2. Show an example calculation for solution A for both the iron(III) and thiocyanate ions. Enter all calculated values into Table 3. b.) Creating the graph. • Create a Beer's Law Plot (a.k.a. calibration curve) using the class average data from Table 1. Use Excel to create this graph. The concentration should be plotted on the x-axis and the absorbance on the y-axis. • Plot the best-fit line and report the equation of the line (y=mx+b) that relates conc. (x) to absorbance (y). This can be done by selecting Chart/Add Trendline from the menu. On the "Type" tab of the dialog box, select "linear". On the "options" tab, be sure to include the equation of the line and the R² value on the graph. Include all other items required for a good graph. c.) Determining "Final (FeSCN³] M" for Table 3 from the calibration curve • Use the equation from your Excel graph to solve for concentration of FeSCN at equilibrium for each solution in Table 2 and enter the results in Table 3. (Remember that it is impossible for the conc. of FeSCN to be greater then the initial conc. of either reacting species.) Show a sample calculation for the (FeSCN1 from the equation of the best-fit line for Solution A. Rev: 2019-2020 6-6
Soln. 1 2 3 mL Fe³+ soln. (2.00×10 M) 0.100 0.200 0.400 0.600 Table 1: Absorbance of Standard Solutions Diluted [Fe³] M = [FeSCN²] M 0.800 0.103 413 0.098 ttt 0.169 Absorbance Values at 447 nm 0.0350 0.0590 10.121 0.097 0.331 0.44 0.092 0.070 075 0.106 0.150 0.151 0.172 0.273 0.185 0.251 0.219 5 0.290 0.217 0.366 0.194 0.285 Note: All solutions contain 2.0 mL 0.200 M SCN & final volumes were made up to 25 0.500 M HNO3. Average Absorb. ml with