In this lab, 0. 010 M purple-colored potassium permanganate solution was standardized by redox titration with iron (II) ammonium sulfate hexahydrate (FAS).
The average mass of the three flasks of FAS was 0. 483 grams. Once the concentration of the standard solution of KMnO4 (aq) was determined, it was used to determine the concentration of Fe2+ in iron pills. On average, there was 0. 01813 L of solution used.
With this information and the balanced net-ionic equation 10Fe(NH4)2(SO4)2*6H2O+2KMnO4+H2SO4 –> 5Fe2(SO4)3+(NH4)2SO4+K2SO4+2MnSO4+H2O, the average percentage by mass composition of Fe2+ in iron pill samples was 16. 99%. (103) INTRODUCTION: The scientific concept of this lab was to titrate potassium permanganate with iron (II) ammonium sulfate hexahydrate in order to oxidize the Fe2+ to Fe3+. Then the concentration of the standard solution of KMnO4 (aq) will be used to determine the concentration of Fe2+ in iron pills. The purpose of the lab is to first standardize a stock KMnO4 (aq), and then determine the percentage Fe2+ in iron pills.
The expected outcome of this lab was for the percentage of Fe2+ in the iron pills to be 17%. This is what was on the bottle of iron pills, and the point of the lab was to use titration to retrieve as much if the iron as possible. MATERIALS AND METHODS: The procedure for the standardization of KMnO4 (aq) was to first fill a clean 50 mL buret with 0. 0100 M KMnO4 (aq).
Three clean Erlenmeyer flasks needed to be labeled, and a piece of FAS needed to be weighed to 0. 5g on a piece of weighing paper. Flask 1 was tarred and tapped into the FAS. The mass was then recorded. This was then repeated with flask 2 and 3. 0 mL of water and 5 mL of 3M H2SO4 was added to the three flasks. 50 mL of water was put in a beaker and 1 drop of permanganate solution was added. The color intensity of the mixture matched the standard and remained for 5 seconds or more. The initial volume was recorded to 0. 01 mL. Permanganate was then added to the FAS solution in flask 1 until the equivalent point was reached. The final volume was recorded and the permanganate solution used was determined. The titration process was repeated using FAS in flask 2 and 3. The flasks were then washed down the drain and rinsed with distilled water.
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Moles of Fe2+ ions present was calculated from the mass of FAS. The balanced equation was used to find the moles of KMnO4 needed to reach the same point. The three molarities of the permanganate solutions were then calculated. The molarities were added to the class data which would calculate the overall molarity for permanganate solutions. The procedure for the analysis of the iron pill was to first grind 2 iron pills in a mortar and pistol. There was 63 mg of iron per pill. 0. 3 grams was weighed on weighing paper, and an Erlenmeyer flask was tarred. The powder was then put in the flask and massed to 0. 01 g. This was then repeated with the second sample. 25 mL water, 15 mL 3M H2SO4, and a few drops of H3PO4(aq) was added to each flask and swirled until the iron pill was dissolved. The initial volume of the reading in the buret was recorded, and permanganate was added to the iron pill solution in flask 1 until the same point is reached. The final volume reading was recorded, and the exact volume of the permanganate used was found. The titration process was repeated using the iron pill in flask 2. The mixtures in flasks were washed down the drain, and the flasks were rinsed with water.
Laboratory rules 1) Prepare a lesson carnestly before the experiment. Clarify the related basic principle, sequence of the operation, and safety measures in the experiment. 2) Put on the working clothes before entering the laboratory. Record the experimental phenomena and data. 3) Keep the floor neat. Don’t throw anything into the water trough so as to avoid stops up. 4) After every experiment, ...
The buret was rinsed with tap water and the stopcock was open for storage. RESULTS: Standardization: FAS mass Flask #1: 0. 500 gtarred: 0. 485 g FAS mass Flask #2: 0. 501 gtarred: 0. 496 g FAS mass Flask #3: 0. 497 gtarred: 0. 467 g KMnO4 |Initial V (mL) |Final V (mL) |Solution used (mL) | |Titration #1 |0. 01 |25. 45 |25. 44 | |Titration #2 |0. 00 |26. 36 |26. 36 | |Titration #3 |0. 00 |24. 31 |24. 31 | |Iron Pills: Iron Pill mass Flask #1: 0. 286 g Iron Pill mass Flask #2: 0. 293 g KMnO4 |Initial V (mL) |Final V (mL) |Solution used (mL) | |Titration #1 |0. 00 |18. 09 |18. 09 | |Titration #2 |0. 00 |18. 17 |18. 7 | |The data found in this lab allowed for the discovery of the percentage by composition of Fe2+ in the iron pills. For the standardization data, the different titrations is the independent variable, and the volumes of KMnO4 is the dependent variable. In order to find the percentage by composition of Fe2+ in the iron pills, the KMnO4 needed to be titrated and the starting, final, and solution used needed to be recorded. The solution used will later be used in liters to find the molarity of each flask of KMnO4. The data above the table is how much FAS was in each flask when it was tarred, and with the flask itself.
For the iron pills data, the different titrations is the dependent variable, and the volumes of KMnO4 is the dependent variable. For each titration, the initial, final, and solution used is recorded on this data table also. The solution used will later be used in liters to determine the grams of Fe2+. The data above the table is how much iron was in each flask measured in grams. CALCULATIONS: Flask 1: 0. 485g (1 mol FAS/392. 16g) (2 mol KMnO4/10 mol FAS) = 2. 473*10-4 mol KMnO4 Flask 2: 0. 496g (1 mol FAS/392. 16g) (2 mol KMnO4/10 mol FAS) = 2. 530*10-4 mol KMnO4 Flask 3: 0. 467g (1 mol FAS/392. 6g) (2 mol KMnO4/10 mol FAS) = 2. 340*10-4 mol KMnO4 Flask 1: (2. 473*10-4 mol KMnO4/0. 02544L) = 0. 00972M Flask 2: (2. 530*10-4 mol KMnO4/0. 02636L) = 0. 00960M Flask 3: (2. 340*10-4 mol KMnO4/0. 02431L) = 0. 00980M Average M = 0. 00971M Flask 1: 0. 01809L (0. 00971 mol/1L) = 1. 7565*10-4 mol MnO4- 1. 7565*10-4 mol MnO4- (0. 00971 mol/1L) (5 mol/1 mol) = 8. 78270*10-4 mol Fe2+8. 78270*10-4 mol Fe2+ (55. 85g Fe2+/1 mol Fe2+) = 0. 04905g Fe2+ Flask 2: 0. 018171L (0. 00971 mol/1L) = 1. 7643*10-4 mol MnO4- 1. 7643*10-4 mol MnO4- (5 mol/1 mol) (0. 00971mol/1L) = 8. 8215 mol Fe2+ 8. 8215 mol Fe2+ (55.
Aim The point of this investigation is to find out the concentration of the sulphuric acid in my experiment. I will do this by titrating the sulphuric acid with sodium carbonate. H2SO4 + Na2CO3 à Na2SO4 + H2O + CO2 Acid + Alkali à Salt + Water + Carbon dioxide Ratio of sulphuric acid to sodium carbonate: H2SO4 : Na2CO3 1 : 1 Concentration of sulphuric acid: H2SO4 is approximately 0.05 – 0.15 ...
85g/1 mol) = 0. 492g Fe2+ (0. 04905g Fe2+/0. 286g Fe2+) * 100 = 17. 15% (0. 0492g Fe2+/0. 293g Fe2+) * 100 = 16. 82% 17. 15% + 16. 82% = 33. 97% (33. 97% / 2) = 16. 99% DISCUSSION: The results from the lab procedure fully support the lab results expected. If the validity of the lab was not reliable, then the percentage of iron in the iron pills would not be 17%. The results from the lab was 16. 99% which is extremly close to the actual percentage of iron. The results compare with the manufacturer’s stated vcalue of %Fe because the results were over the percentage depicted by the company by . 31%. There were 389. 5mg in each pill, and 16. 9% of this would be 66. 18mg. Compared to 65mg which was determined by taking 10 pills, weighing them, and dividing by 10, the desired yield was reached. Any mistake in discrepancy greater than 5% would be a mistake in the titration process or misweighed mass. Titration was used in this lab to oxidize the Fe2+ to Fe 3+ in the FAS, using potassium permanganate solution. It was then used to determine the concentration of Fe2+ in iron pills. The results supported the scientific concept of titration because the process was done right and carefully in order to get the closest results to 17% as possible.
There was one validity error made in the lab though while the titration process was being done. With flask #2, the stopcock on the buret was not turned parallel to the ground in time and too much KMnO4 was added to the flask. This caused the color to be more of a light magenta color, instead of a peachy or salmony color like flask #1 and #3 were. As explained before, this did not interfere with the results of the percentage of iron in the iron pills, as it was . 01% away from being exactly correct. It is still a validity error, and needs to be taken into account.
For the next time this lab is performed, the person doing the titrating should be extremly careful to do it right and take their time, in order to get the best results possible. (Picture: To the left is flask #1, the middle is flask #2 and shows the validity error, and to the right is flask #3. Flask #1 had a translucent peachy color, flask #2 had a translucent light magenta color, and flask #3 had a translucent salmon color (the lightest).
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) CONCLUSION: This lab definitely explains how to do a titration, but also gives a hands on experience and visual of the chemical reaction of iron oxidizing from Fe2+ to Fe3+.
It also gives a hands experience and visual of determining the concentration of Fe2+ in iron pills. This lab was quantitative because it refers to how much the amount of the present element or compound there is. The theoritical data compares to the experimental data because the theoretical data gave the conclusion that there was 17% iron in the iron pills. The experimental results gave a very similar result at 16. 99%. With those results, this lab was a success and brought to the conclusion that even with a small error with the titration, the validity of this lab is very reliable.