Lab 14 Report - Grade: B PDF

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Redox Titration Ulyana Volkoff CHEM 1A Lab, Department of Chemistry, California State University, Fresno, CA 93740 [email protected] October 30, 2017

Abstract

This experiment will depict a redox titration of KMnO4, Potassium Permanganate, into a solution containing H2C2O4•2H2O and H2SO4. The percent concentration of the aqueous solution, H2C2O4•2H2O, is unknown and the purpose is to figure it out through the redox titration of this reaction. After titration where all the molecules were allowed to react with one another, it was discovered the percent of Oxalic Acid Dihydrate was an average if 6.445% from the two trials performed. Titration of this reaction was effective in finding out the concentrations of everything in the mixture. The method for doing such, as well as any errors, observations, or further clarifications, is outlined further in this experiment.

Introduction Upon starting this experiment, it is important to note that the purpose was to perform an oxidationreduction, or redox reaction, which is a process where one reactant loses elections while another one gains it in exchange. This reaction results in a change in oxidation amount, and the reactants that are to gain electrons will have a reduced charge, thus it is the oxidizing agent. However, the substance that loses its electrons is called the reducing agent, therefore it is oxidized. Titration is the method used in this experiment, and titration is used primarily to find an unknown concentration in a sample. For

example, titration is often used in the food processing industry in order to determine the known base amount that is to be added to a waste vegetable oil batch.6 Acid-Base titration in particular is used in the situation of an unknown free fatty acid content through acid-base titration when using a color indicator such as phenolphthalein.6 In a titration, the purpose is to find an unknown concentration using the relationship between the moles and volume used. The formula for the particular titration of this experiment is: Balanced Chemical Reaction: 2MnO4+5H2C2O4•2H2O+3H2O+3H2SO4 —> 10CO2+2MnSO4+K2SO4+18H2O In this experiment, KMnO4 is the oxidizing agent, or election acceptor, since KMnO4 the Mn ions gain a charge through the loss of its electrons: MnO4 - is reduced to Mn+2 In counter of this, oxalic acid is the reducing agent since it gains elections in the equation. The direct reaction is a slow process throughout titration; for example, the first few drops of Potassium Permanganate that is added to the acidic solution containing oxalic acid, and it remains that dark purple color for a couple seconds before the color is lost. This reaction produces Mn+2 ions, and they in turn act as a catalyst as they react with permanganate to form an oxidation state of manganese, however, these states of matter react swiftly to yield the products. In this regards, the KMnO4 is the catalyst, as Khan Academy states3 in that it speeds up the reaction time without getting consumed in the process. Therefore, Potassium Permanganate is its own indicator since the pink color it emits is distinguishable even in a diluted solution, meaning due to KMnO4’s reduction by the oxalic acid solution, after complete consumption of the oxalate ions, the end point is a light pink color. An end-point, as said by Xavier4 of a titration indicates that an equivalence point, which is the ideal point for the completion of titration, has been met. This is shown when some form of indicator, such as the solution changing color as demonstrated from KMnO4, is being done. The use of an indicator is the true key in any successful titration reaction since it shows when enough of the standard solution has

been added to fully react with the unknown concentration.4 A good endpoint and a bad endpoint is also shown through the indicator:

The image on the left indicates a good endpoint while the image on the right indicates an overly titrated bad endpoint. To over titrate something is the action of, for example, dropping too much KMnO4 into the solution. The endpoint is shown by some kind of indicator towards the end of the titration, while in contrast the equivalence point is when the moles of the titrant equal the moles of the unknown concentration.4 The murky brown color of the titrated solution on the left is a result of the KMnO4 is due to it no being fully oxidized, and the errors that come along with this particular attempt will be fully explained in the discussion section.

4

More details on the titration process is displayed in the image above, emphasizes the importance of understanding how to perform the process of titration itself for the acid-base redox reaction. It is also important to note that oxalic acid is a weak acid, thus it isn’t very effective by itself in titration so it is necessary to dilute it with sulfuric acid since by itself it won’t provide a strong acidic medium. However, despite it being a weak acid, it is still important to use it to prevent hydrolysis and provide enough H+ ions in the solution so the reaction continues, also sulfuric acid at an oxidized state is very stable. Since oxalic acid is weak, it is not able to enough H+ ions to keep the reaction continuing while still being a medium that won’t interfere with the oxidizing agent. If a strong acid were to be used instead, such as HCI (aq), the Mn+2 ions would oxidize the Cl- ions and form the gas, Cl2, thus making it not available to react with C2H2O4. KMnO4, as a result of the sulfuric acid and oxalic acid, becomes much stronger under acidic conditions as show between the two chemical equations down below. 𝑀𝑛𝑂!! + 8𝐻 ! + 5𝑒 ! ⇔ 𝑀𝑛!! + 4𝐻! 𝑂 This reaction depicts the dark purple MnO4 Permanganate into Mn2+, the colorless ion which can only occur under acidic conditions. However, under more neutral conditions, such as when too much Permanganate is added to the acidic solution, it would turn into a brown Manganese(IV) oxide as depicted in the equation below:

𝑀𝑛𝑂! ! + 4𝐻 ! + 3𝑒 ! ⟺ 𝑀𝑛𝑂! + 2𝐻! 𝑂 Heating the oxalic acid is also a vital step in the titration process as all reaction require activation energy and correct orientation in order to actually react.1 If the mixture is not heated, then in the short time of the reaction, the reactants will not have adequate time in order to have enough activation every to start the reaction, regardless of whether or not they are in the correct orientation. However, if the solution is heated, that heat energy gained by the particles will be converted into kinetic energy, meaning that they will now move with higher speed than before.1 This results in a higher frequency of collisions between the molecules, allowing someone to have an increased rate of reaction and the ability to obtain the desired results in a much shorter timeframe. Now, if the temperature was to be too low, the reaction between oxalic acid and potassium permanganate will move much too slowly and in a lab experiment, it will not display the results in an appropriate amount of time. In this experiment, titration will be used to determine the unknown concentration of the solution and to find the end point where the number of equivalents of oxidizing agent will equal the number of equivalents of reducing agent. Experimental General Methods and Materials. Materials used in this procedure were provided by the Fresno State Freshman Chemistry Stockroom. Equipment provided in this lab includes a buret, 600 mL beaker, 250 mL beaker, 50 mL beaker, graduated cylinder, Erlenmeyer flask, ring stand, thermometer, hot plate, and funnel. The chemicals and substances used in this experiment are 25 mL Potassium Permanganate, DI water, 1.0 mL unknown Oxalic Acid Dihydrate, and 25 mL of 3M sulfuric acid. Procedures. In this experiment, it is important to start with diluting the Potassium Permanganate. Using the set calculation for diluting it, as show in calculations 1, pour 25 mL KMnO4 and mix with 225 mL of DI water into a beaker. Transfer the solution to the 500 mL plastic bottle and cap it. Then rinse a clean

buret with 5 mL portions of the diluted KMnO4. Pre weigh a beaker and record the amount in grams. Take the unknown sample of H2C2O4•2H2O and transfer 1.0 mL of the liquid to a small pre-weighed beaker, then determine the mass of the liquid to the nearest 0.0001 g using the analytical balance. Transfer it to a to a 250 mL flask and add 20 mL of Di water. Next, swirl the mixture and slowly add 25 mL of 3 M sulfuric acid, then heat it to 70ºC to 80ºC. Titrate the standardized KMnO4 solution very slowly at first until that one drop causes light pink coloration. Record the volume from the upper edge of the KMnO4 meniscus. Results In this titration, it was discovered that by adding 0.76 g of a 1.0 mL Oxalic Acid Dihydrate, mixed with 20 mL DI water and 25 mL 3 M Sulfuric acid and 7 mL Potassium Permanganate, as shown by table 1 down in supporting information, the total of the solution would be 53 mL for trial 1. However, since the solution for trial B was only 0.711 g, which is a 0.049 g difference from the first trial, the amount of Potassium Permanganate was 8 mL and the final volume of the solution was 54 mL. With this information, known, it became possible to calculate the percent of Oxalic Acid dihydrate in trial’s 1 and 2. As show in calculations 3 and 4, the percent concentration of Oxalic Acid dihydrate in trial 1 was 5.8% and in trial 2 it was 7.09% Oxalic Acid dihydrate. After that information was available, standard deviation of the two trials aided in determining the overall average of the two trials to be 6.445% Oxalic Acid dihydrate, as show in calculations 5 down below. The end point, as show by a simple titrations curve for an acid-base reaction5, is show by graph

5

The equivalence point, as depicted on the graph, is met when reactant A has nothing to react with reactant B in order to make product C. Based on the graph, it is possible to find the concentration of the unknown through the equivalence point itself. As such, when all the Oxalic Acid dihydrate has been oxidized and the solution is a very light pink, then the equivalence point has been met. The starting solution was a clean liquid and the ending was a light pink once fully oxidized as depicted from these images.

The starting image on the left shows titration as the pink Potassium Permanganate reacts with the acid solution and disperses, while the one on the left shows the complete reaction as a light pink. Discussion The experiment underwent a couple sets of errors in the beginning trials. The first error, as show in the image above in the introduction, resulted in a dark brown titration. This was caused by method error and lack of experience. It is important to drip the Potassium Permanganate slowly, not all at once since this unbalance is caused by the oxalic acid being overpowered by the KMnO4. The brown color is the result of the reduction of permanganate (Mn2+) and the oxidation of the oxalate to CO2 and H2O. When the Mn+2, which is normally the self indicator due to the faint pink color it emits to the solution, reacts

with permanganate to form murky brown MnO2, which leads to the haziness of the fluid as shown by the equation below: 2𝑀𝑛𝑂!! + 3𝑀𝑛!! + 2𝐻! 𝑂! ⟺ 5!𝑀𝑛𝑂!↓ + 4𝐻 ! Comments and Conclusions Despite the errors occurring in method, the last two trials were an overall success and the titration yielded the desired products. The end point was met in the form of all the Oxalic Acid dihydrate having been oxidized, proving that Potassium Permanganate was the oxidizing agent while Oxalic Acid was the reducing agent. Acknowledgment. The author thanks lab instructor Kumar. Supporting Information Available: Table 1. Trials of Titration Measurements

Trial 1

Trial 2

Small Flask Weight (Initial 51.9 g beaker) Unknown, H2C2O4•2H2O weight 52.660 g

51.9 g

Unknown, H2C2O4•2H2O weight 0.76 g subtracted from initial beaker DI H2O 20 mL

0.711 g

3 M H2SO4

25 mL

25 mL

KMnO4

7 mL

8 mL

Unknown solution in mL

53 mL

54 mL

52.611 g

20 mL

Equation 1. Formula for determining Dilution 𝑀! 𝑉! ! = ! 𝑀! 𝑉! ! M1 = concentration of the first solution

V1 = volume of the first solution M2 = concentration of the second solution V2 = volume of the second solution Equation 2. Formula for determining Sample and Precision (Standard Deviation)

𝑆=

Σ(𝑥 − 𝑥)! 𝑁

S = the standard deviation of a sample Σ = sum of X = each value in the data set X– = all values in the data set N = number of values in the data set Calculations 1. Diluting the KMnO4 Stock Solution Preparing Titrating strength KMnO4 from stock, more concentrated KMnO4 Prepare 250 mL of 0.020 M KMnO4 that will be used for titrations in this experiment 0.02!𝑀!×!250!𝑚𝐿 = ! 𝑉! !×0.2!𝑀 𝑉! = !25!𝑚𝐿!𝐾𝑀𝑛𝑂! ! 250!𝑚𝐿!– !25!𝑚𝐿!𝐾𝑀𝑛𝑂! ! = !25!𝑚𝐿!𝐻! 𝑂! Calculations 2. Standardizing the diluted KMnO4 Solution Volume of H2C2O4•2H2O Standard that would use about 15 mL of KMnO4

15!𝑚𝐿!×!

0.02!𝑀 1 5!𝑚𝑜𝑙 !×! !×! = 37.5!𝑚𝐿 1 2!𝑚𝑜𝑙 0.02!𝑀

Calculations 3. Percent Oxalic Acid Dihydration in Sample Trial 1 Known:

Molar Mass of H2C2O4•2H2O =126.07 g/mol Concentration of MnO4 = 0.02 M Volume of MnO4 solution = 7.0 mL 5 Moles of oxalic acid reacts with the 2 moles of MnO4 consumed Find Grams H2C2O4•2H2O 7.0!×!

1!𝐿 0.02!𝑚𝑜𝑙!𝑀𝑛𝑂! 7.0!×!0.02!𝑚𝑜𝑙!𝑀𝑛𝑂! ! 5!𝑚𝑜𝑙!𝐻! 𝐶! 𝑂! • 2𝐻! 𝑂 = 1.4!×!10!! !×! = !×! 1000!𝑚𝐿 2!𝑚𝑜𝑙!𝑀𝑛𝑂! 1!𝐿 1000!𝑚𝐿 𝑔 = 3.5!×!10!! !𝐻! 𝐶! 𝑂! • 2𝐻! 𝑂!×126.07 𝐻 𝐶 𝑂 • 2𝐻! 𝑂 𝑚𝑜𝑙 ! ! ! = 0.0442!𝑔!𝐻! 𝐶! 𝑂! • 2𝐻! 𝑂

Find % H2C2O4•2H2O !.!""#! !.!"!

!×!100 = 5.8%!H2C2O4•2H2O

Calculations 4. Percent Oxalic Acid Dihydration in Sample Trial 2 Find Grams H2C2O4•2H2O 8.0!𝑚𝐿!×!

1!𝐿 0.02!𝑚𝑜𝑙!𝑀𝑛𝑂! 5!𝑚𝑜𝑙!𝐻! 𝐶! 𝑂! • 2𝐻! 𝑂! = 1.6!×10!! !×! !×! 1!𝐿 2!𝑚𝑜𝑙!𝑀𝑛𝑂! 1000!𝑚𝐿 𝑔 = 0.0505𝑔!𝐻! 𝐶! 𝑂! • 2𝐻! 𝑂! = 4.0!×!10!! !𝑚𝑜𝑙!×!126.07 𝑚𝑜𝑙

Find % H2C2O4•2H2O 0.0505!g !×!100 = 7.09!%!𝐻! 𝐶! 𝑂! • 2𝐻! 𝑂! 0.711!g Calculations 5. Calculation Average % Oxalic Acid Dihydrate & Standard Deviation of 2 Trials 𝑇𝑟𝑖𝑎𝑙!1!5.8% + 𝑇𝑟𝑖𝑎𝑙!2!7.09% = 12.89 ÷ 2 = 6.445%!𝑎𝑣𝑒𝑟𝑎𝑔𝑒

𝑆=

Σ(5.8 − 7.09)! = 0.83205!𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑!𝑑𝑒𝑣𝑖𝑎𝑡𝑖𝑜𝑛!𝑣𝑎𝑟𝑖𝑎𝑛𝑐𝑒 2

References 1. Tro, Nivaldo J. Chemistry: A Molecular Approach Edition 4 2. Department of Chemistry Chemistry 1A Laboratory Manual, Fall 2017 Edition, Star Publishing: Fresno, CA 2017. 3. Types of catalysts (article) | Kinetics https://www.khanacademy.org/science/chemistry/chemkinetics/arrhenius-equation/a/types-of-catalysts (accessed Oct 29, 2017). 4. Xavier, Lauren, Titration Fundamentals. Libretexts. [Online] 2016 https://chem.libretexts.org/Core/Analytical_Chemistry/Lab_Techniques/Titration/Titration_Fundamenta ls (accessed October 26, 2017). 5. Acid-base titration curves https://www.khanacademy.org/test-prep/mcat/chemicalprocesses/titrations-and-solubility-equilibria/a/acid-base-titration-curves (accessed Oct 29, 2017). 6. Use of Titration https://sciencing.com/use-titration-5855864.html (accessed Oct 29, 2017)...