Determination of the Rate Law and Activation Energy for the Dissociation of Ferroin
Kim, Taewoo, Trey L Arnold, Kyle A Leland, Aimee M Morey, Department of Chemistry, USAF Academy, CO 80840
ABSTRACT
Blue wavelength was used to measure absorbance; using initial data and Beer’s law, calculated molar extinction coefficient was 10953 L∙mol^(-1)∙cm^(-1). Variations of concentration over time were graphed to figure out dissociation of ferroin being first-order reaction and sulfuric acid being zero-order. Thus, under the rate law, the solution’s rate was found to be k〖[Ferroin]〗^1 〖[Sulfuric acid]〗^0, where rate constant k = 9.12x〖10〗^(-4) s^(-1). Half-life was 759.87 sec and Ea (activation energy) was 56.82 KJ.
INTRODUCTION
Chemical kinetics, also
…show more content…
known as reaction kinetics, is study of rates of chemical processes. It uses various conditions such as catalyst, light, temperature or concentration to observe any change in reaction rate. It is closely related to our lives as speeding up reaction or down depending on the situation could bring dramatic change into industry. For example, Haber-Bosch process for making ammonia is comprised of concepts like position of the equilibrium or selection of a catalyst. It is also possible to figure out sequences from reactants to products by using kinetics. In the past, three individual experiments were used to determine rate law, activation energy and equilibrium constant. Under this inefficient system, students had difficulty understanding complicated concept of kinetics. Under this circumstance, Ferroin, used for redox indicator which is also called an oxidation-reduction indicator, has proven very effective in determining rate law, activation energy and equilibrium constant through a single experiment that can produce the same results. In this experiment, complementary color wheel diagram is used to select wavelength with best absorbance for the solution. Then, the colorimeter is used to monitor absorbance of ferroin along the time. Beer’s Law is used throughout the experiment to get concentration from measured absorbance by us of ε (Molar attenuation coefficient) determined from initial values. By graphing relationship between variations of concentration with time, one graph will show straight line which is the order of the reaction. Knowing the order, Rate Law of dissociation reaction as well as rate constant k can be easily found. Half-life is the time required for the concentration to fall half of its original value. Different temperature will be used to draw Arrhenius plot or use equation to find out activation energy for the reaction. EXPERIMENTAL Prepare mixture of 5.0mL of 1.0×〖10〗^(-4)M 〖〖[〖Fe(Phen)〗_3]〗^(2+)〗_((aq)) with 5.0mL of 0.40M sulfuric acid and two water baths with 40°C and 45°C.
Use blue LED color for maximum absorbance of Ferroin and scroll down the colorimeter screen to view absorbance at your chosen wavelength. Measure the initial absorbance of the mixture with colorimeter, record it and use this information to determine the molar extinction coefficient for Ferroin. Place the cuvette into 40°C water bath and let it heat up. Remove cuvette from water bath to measure the absorbance of mixture with two minutes of interval in between, putting the cuvette immediately back after the measurement. Be sure to dry the cuvette (ex. with paper towel) before putting into colorimeter. Continue until the absorbance drops below 0.2 or when you have the 10th measurement. For second experiment, repeat procedure using 0.20M sulfuric acid by diluting 0.40M sulfuric acid. For the last experiment, use 0.40M sulfuric acid again but put into 45°C water bath instead of …show more content…
40°C RESULT/DISCUSSION First, molar extinction coefficient (ε) was determined as 10560 L∙mol^(-1)∙cm^(-1) using the initial absorbance (0.528) and concentration (5 x〖10〗^(-5)M) using Beer’s Law (A=εlC). Then, as the absorbance decreased over time, concentration was found using Beer’s Law, dividing absorbance by molar extinction coefficient (ε). Getting all concentration respect to recorded time, three graphs were plotted. To determine the order of sulfuric acid, 0.40M solution was compared to diluted 0.20M solution. First four records were taken to compare, excluding initial point at room temperature. When concentration is doubled, rate would stay same, double or quadruple depending on its order, zero- first- second- respectively. Rate Law = k〖[Ferroin]〗^m 〖[H^+]〗^n, using two graph, (2.541x〖10〗^(-8))/(2.538x〖10〗^(-8) )=〖(0.2/0.1)〗^n, n= 1.70x〖10〗^(-3), which is almost zero. Thus, the order of the sulfuric acid was found to be zero-order. Then, looking at three graphs of ferroin, ‘ln(concentration) VS time’ (figure 2) showed almost straight line (R^2=99.5% ) meaning the order of Ferroin is first-order. Reference was used to confirm this order. Therefore, using Rate Law, Rate=k[Ferroin] was found where from slope = -k, rate constant k= 9.12x〖10〗^(-4) s^(-1) . Figuring out the order of reaction and rate constant k, Half-life (t_(1⁄2)) = 0.693/k = 759.87s was calculated. To get a Ea (activation energy), Arrhenius Plot was made. When temperature changed from 40°C (313K) to 45°C (318K), rate constants changed from k_1=6.47x〖10〗^(-4) to k_2=9.12x〖10〗^(-4). Slope is -E_a/R in this case. -E_a=-6833.9K(8.314∙J/(mol∙K)) = 56.82KJ Using equation [ ln k_2/k_1 =E_a/R(1/T_1 -1/T_2 ) ] where, k_1=6.47x〖10〗^(-4) s^(-1), T_1=313K, T_2=318K and k_2=9.12×〖10〗^(-4) give same result k_2/k_1 = 1.41, ln k_2/k_1 =0.343, R= 8.314 J〖∙K〗^(-1)∙mol^(-1) (1/T_1 -1/T_2 )=5.02x〖10〗^(-5) K , therefore, E_a= 56.82KJ Molecule Molar Extinction Coefficient (units) Rate Constant (units) Half-life (units) Activation Energy (units) [Fe(phen)3]2+ 10953 L∙mol^(-1)∙cm^(-1) 9.12x〖10〗^(-4) s^(-1) 759.87 s 56.82 KJ Molar extinction coefficient (ε) was found out to be 1/3(10940+10560+11360) = 10953 L∙mol^(-1)∙cm^(-1).
ε should be all same for three experiments so error occurred here. Source of this error might be little water or moisture inside colorimeter before starting new experiment. This could affect absorbance of wavelength by colorimeter. Also, I used 45°C graph to determine the order of ferroin because when I plotted all three graphs for 40°C, zero-order graph had higher percentage of R^2, than first-order graph. It must have been human mistake as my lab partner forgot to measure for every 120 seconds but randomly recorded four points instead. (120, 240, 430, 540)
Therefore, rate constant is little higher than it is supposed to be. At 40 °C, k would be -6.47 x〖10〗^(-4) s^(-1) and half-time would be 1071 seconds. It is same for activation energy. I think it is less than it should be because of human mistake while taking the measurement. I am guessing it might be around 100KJ based on references.
REFERENCE
John O. Edwards; Kathleen Edwards; Jorge Palma; The reactions of ferroin complexes. A color-to-colorless freshman kinetic experiment 1975, 52, 408
Simeen Sattar; Unified Kinetics and Equilibrium Experiment: Rate Law, Activation Energy, and Equilibrium Constant for the Dissociation of Ferroin 2011, 88, 457-460
T. S. Lee; I. M. Kolthoff; D. L. Leussing; Reaction of Ferrous and Ferric Iron with 1,10-Phenanthroline. I. Dissociation Constants of Ferrous and Ferric Phenanthroline
1948, 70, 2348-2352 DOCUMENTATION STATEMENT I visited Mr. Dave Summer in QRC to get help to determine the order of the reaction because two graphs for ferroin both looked like a straight line. Also, I got help writing abstract; it is not explaining with words but writing numbers of results which will let reader know very condensed result of my paper. I also visited writing center right after QRC not for specific reason but to have them look at my introduction for structure and grammar.
The analyzed yellow#5 wavelength was determined to 395nm because the actual wavelength 427nm was restricted in the Micro lab. The R2 value of the graph is 0.9827, and the level of data accuracy it indicated extremely weak data correlation. The first one dilution data points excluded from the standard curve because the point is not in the linear curve. The first concentration and absorbance value are the highest point in the graph that cannot connect as linear with another data point. After removing the first data point, the standard curve is clear and make
2. Cooper, M. M., Cooperative Chemistry Laboratory Manual, McGraw-Hill: New York, NY, 2009, p. 60.
The purpose of this experiment is to use our knowledge from previous experiments to determine the theoretical, actual, and percent yields of the lead chromate from the reaction of solutions of potassium chromate and lead nitrate (Lab Guide pg. 83).
The mean for the temperatures is 0.116 and the solvents is 20. We predicted the 37 Celsius would be the most absorbed, but it was the -20 Celsius which can be seen in the graph above.
barbier reaction: In a 50 mL round bottom flask that had a reflux condenser attachment, saturated ammonium chloride (5 mL), THF (1 mL), zinc powder (0.4 g), benzaldehyde (0.500 mL, 0.5225 g, 4.92 mmol), and allyl bromide (0.470 mL, 0.6533 g, 5.40 mmol) were charged with stir bar and stirred at room temperature for 45 minutes. Diethyl ether (10 mL) was added to the reaction mixture and stirred. The mixture was gravity filtered into a beaker that was topped with a watchglass. The filtrate was transferred to a separatory funnel and the organic layer was extracted with deionized water (10 mL) and diethyl ether (15 mL). The organic layer was placed into an Erlenmeyer flask and the aqueous layer was placed into a beaker, which was extracted with
This question refers to the example data given below. Using the rate law and the experimental values given below, calculate k.
Felder, M. Richard, Elementary Principles of Chemical Processes, 3rd ed.; Wiley: New Jersey, 2000; p 631.
For this solution, 5 mL of the solution with 2.5 mL of AMV was placed in the cuvette. The cuvette was placed inside of spectrophotometer and the amount of absorbance was recorded. This procedure that involves a solution with a known concentration was repeated for the concentrations:1.0x10-4 M,5.0x10-5 M,2.0x10-5M, and1.0x10-5M.A unknown solution absorbance was measured by putting 5 mL of unknown solution with 2.5 mL AMV in a cuvette. The cuvette was placed in the spectrophotometer and the amount of absorbance was recorded. The procedure that deals with the unknown solution was repeated 2 more times with the same solution and the same amount of solution and AMV.
...eases, including temperature. It is determined from the data that the reaction is more likely to have a step wise mechanism than a concerted due to the small – ΔS and a relatively large value of ΔH from the tables. Due to some errors, it is best to perform another experiment for future protocols. In addition with the variance the 35°C where at one point the absorbance levels off and then increases. In comparison to the rate constant against temperatures, at 25°C it is higher than 35 and 45. More test is required to ensure proper determination of the rate constant at those temperatures.
...Coauthor, ChemBioChem 2006, 7, 1-10; b) A. Author, B. Coauthor, Angew. Chem. 2006, 118, 1-5; Angew. Chem. Int. Ed. 2006, 45, 1-5.))
the water baths I think were accurate enough but having two thermometers in each bath maybe would have helped to hold the temperature readings more accurately. We were not given any instructions either to shake or not to shake the test tubes with the coloured solutions before inserting them in the spectrophotometer to read the absorbance. By shaking each test tube a certain number of times before putting it in the spectrophotometer could have improved the accuracy of the absorbance of the solutions.
A cuvette was filled 3/ 4ths of the way and the absorbance measured in a spectrophotometer. The data was compiled as a class and recorded. The Spectrophotometer was blanked using a test tube of distilled water.
To examine the interaction between two molecules in solution without isolating the compound Jobfs method is used. Although unstable compounds tend to be generated, this is not a reflection of weak interactions. In some cases, the transition metal species cannot be crystallized from the solution and separated from the other species present. Without Jobfs Method this composition can be very difficult to deduce.
Many factors influence rates of chemical reactions. Some of these factors include: the nature of reactants, for example the formation of salts, acid-base reactions, and exchange of ions are fast reactions, while in reactions where bigger molecules are formed or break apart are typically slow; temperature, frequently, the higher the temperature, the faster the reaction; concentration effect, the reliance of reaction rates on concentrations are called rate laws. Rate laws are expressions of rates in terms of the concentra...
... point was subjective; however it would have been a systematic error because I consistently judged the end point of my experiment. To eliminate this inaccuracy I should have used a colorimeter to judge the end point of my experiment.