<ul><ul><li>First we had to put 8.00 mL of stock solution containing .400 grams of FeCl 3 /mL or .4 M. Next we had to obtain 12.0-12.5 grams of K 2 C 2 O 4 *H 2 O and add 20 mL of distilled water to dissolve the white, powdery substance. We heated the solution and stirred until all of the K 2 C 2 O 4 *H 2 O was dissolved. We then poured the hot solution into the FeCl 3 and stirred until a green color appeared. We let the solution cool and then stored it in a refrigerator over night so the crystals could form. Next, we decanted the crystals carefully, making sure no crystals were removed. We then used vacuum filtration to filter the crystals. Finally, we washed the crystals twice with about 5 mL ice water and twice with 5 mL portions of acetone. We then spread the crystals out on the bottom of a clean dry filter paper and let them try overnight. </li></ul></ul>
This is the green color that formed once we poured the hot solution of K 2 C 2 O 4 *H 2 O into the beaker containing FeCl 3 .
Brett couldn’t resist from eating the watermelon rock candy...I mean…green crystals!
On Day 2 of Experiment 1… Mrs. Males demonstrated how to use vacuum filtration. The ‘vacuum’ sucks up some of the water that surrounds and is absorbed in the green crystals.
Brett and I mimicked Mrs. Males instructions on vacuum filtration and applied it to our self-made green crystals. This is what our crystals looked like after vacuum filtration.
After drying out our crystals in an oven, this was our final product. Each time an experiment was conducted using green crystals, these are the crystals that were used.
<ul><li>1) Why should the crystals be cooled overnight in the refrigerator? </li></ul><ul><li>Because this will cause the substance to cool, crystallize, and form into a solid. A good amount of the water and acetone mixture will evaporate in this process but some will also get absorbed into the formation of each salt crystal. </li></ul><ul><li>2) When rinsing the crystals with water and acetone, why should only small amounts of those solvents be used and why should they be cold? </li></ul><ul><li>Using small amounts of water and acetone will help prevent excess dilution. A lower temperature will hinder the dissolving of the crystals and keep them solid. </li></ul>
<ul><li>Purpose : to prepare and standardize KMnO 4 to use in another experiment in order to determine the percentage of Oxalate in the crystal. The standardization of KMnO 4 is necessary to determine exactly how much MnO 4 - is in our stock solution because the KMnO 4 is not 100% pure to begin with. </li></ul>
<ul><li>In order to prepare the KMnO 4 , first we had to calculate the mass of KMnO 4 required to make 250 mL of a .010 M solution. We then had to weigh out the KMnO 4 . We transferred the KMnO 4 crystals into a 250 mL volumetric flask and added distilled water to bring the solution up to the calibration line. Next we mixed the contents by flipping the flask up and down about 10 times. Then we transferred the solution into a dark bottle to be stored. Before transferring, we cleaned this bottle with 1-2 mL portions of KMnO 4 . </li></ul><ul><li>To standardize the KMnO 4 , first we weighed between .12 and .13 grams of Na 2 C 2 O 4 . We then transferred the Na 2 C 2 O 4 to a clean beaker and added 50 mL of distilled water. Then we added 6 mL of 6M sulfuric acid. Next we heated the solution to just below its boiling point and we began the titration at this temperature. We repeated this procedure for 2 more trials. </li></ul>
Brett heating the beaker and its contents just below boiling point. Titration followed this.
<ul><li>Balanced equation of an acidic solution: </li></ul><ul><li>MnO 4 - + C 2 O 4 -2 Mn +2 +4H 2 O </li></ul><ul><li>2(5e - + 8H + + MnO 4 - Mn +2 + 4H 2 O) </li></ul><ul><li> 5(C 2 O 4 -2 2CO 2 + 2e - ) </li></ul><ul><li>16H + + 2MnO 4 - + 5C 2 O 4 -2 2Mn +2 + 10CO 2 + 8H 2 O </li></ul><ul><li>The oxidizing agent is: Manganese (Mn) </li></ul><ul><li>The reducing agent is: Carbon (C) </li></ul>
Brett is still hungry! He’s still going after that green rock candy!
<ul><li>Purpose : The purpose of this experiment is to determine mass % of C 2 O 4 -2 by titrating a solution of known mass of the green crystals synthesized with the 0.010 M KMnO 4 prepared and standardized in experiment #2. </li></ul>
<ul><ul><li>- First we weighed our green crystals between .12 and .13 grams and transfer the crystals to a beaker to titrate. We added 6 mL of 6M sulfuric acid to the beaker. We then added 1 mL of 85% phosphoric acid. Next we heated to solution to just below the boiling point just like in Experiment #2. We then removed the flask from the heating plate and titrated. These steps were repeated using another green crystal sample. </li></ul></ul>
These were our two green crystal samples used in this experiment’s titration.
Brett measures 6mL of 6M sulfuric acid in a graduated cylinder.
<ul><li>What is the purpose of the addition of phosphoric acid before heating the sample? </li></ul><ul><li>Due to Le Châtelier's Principle, the reaction will be driven to completion if we can remove water (thus inhibiting the reverse reaction) and phosphoric acid serves as a dehydrating agent. It is necessary to remove some of the water that is absorbed into the green crystals. </li></ul>
<ul><li>Purpose : The purpose is to of this experiment is to standardize NaOH, which will then be used in experiment #5 to determine %K and Fe in the crystals. We must first prepare a solution of approximately the desired concentration, and then find its exact concentration by titrating against a standard substance, or KHP in this case. </li></ul>
<ul><ul><li>First we made a solution of approximately .10 M NaOH by measuring about 2 grams of NaOH and diluted it with about .5 liters of distilled water. We then obtained 0.4 to 0.6 grams of previously oven-dried sample of KHP and placed it in a beaker with 40 mL of distilled water. We swirled until completely dissolved. We then added 3 drops of phenolphthalein acid to the solution before beginning the titration. This will allow us to see a color change when an equivalence point has been reached. While titrating with NaOH, we constantly stirred the contents in the beaker. We repeated this process in order to complete 3 trials. </li></ul></ul>
Ashley rinsing the buret with 7 mL of NaOH before titration.
After our equivalence point was reached in our first trial, this is the color change that occurred. Adding 40mL of distilled water to the beaker.
<ul><li>Purpose : to determine both the %K and %Fe in a single titration after passing a solution containing a known mass of the complex salt down an ion exchange column. </li></ul>
<ul><li>To prepare the column for this experiment, the procedure was as follows. First we weighed out .16 grams of our green crystals and placed them in a beaker. We put 4 mL of distilled water into the beaker and swirled until the green crystals were all dissolved. We then placed a large beaker under the ion exchange column and transferred the dissolved green salt into the column. We then let the water run until it was just above the resin. We then placed 4 mL of distilled water into the ion exchange column and let the water run until it was just above the resin. We repeated this process 2 more times. Once we were ready to titrate, we placed a beaker on a magnetic stirrer and dropped a stirring bar into our beaker. We then set up our pH sensor using a graph on Logger Pro . We then set up for our titration by obtaining a buret and cleaning it with .10 M NaOH solution a few times. We filled the buret so it had 50 mL of NaOH. Once the buret was filled and ready to titrate, we started our data collection. We added the NaOH solution to the beaker so that the pH would increase by .15mL. At each addition, we recorded our results. Once we noticed a visibly large increase in pH upon the addition of 1 drop of NaOH solution, we knew we had reached an equivalence point. In this lab, we reached two equivalence points: one was more clearly defined than the other. Once we were finished with the collection of data, we disposed of the beaker contents and cleaned up our lab station. </li></ul>
Ashley allowing distilled water to pass through the ion exchange column and into the beaker of our green crystals. This picture was taken while we were titrating and taking pH to collect data to find equivalence points.
After both equivalence points were found, this is what our beaker looked like.
<ul><li>1) How does this titration differ from any others completed thus far in this lab? </li></ul><ul><li>This titration is different because we used pH to help determine when the equivalence point was reached. It is also different because we use a method known as column separation. This allows us exchange any K + ions with H + . We titrate to find the quantified amount of H + and Fe 3+ in an acid-base titration. </li></ul><ul><li>2) What one precaution should you take when opening up the column? </li></ul><ul><li>When opening the column, we must make sure that the water level does not drain below the resin level. </li></ul><ul><li>3) What does the 2 nd equivalence point represent? </li></ul><ul><ul><li>It represents the completion of the precipitation of iron (III) hydroxide in the acid-base titration. </li></ul></ul>
<ul><li>Purpose: The purpose of this lab was determine the %water in our green crystals since these crystals are hydrated with absorbed water. </li></ul>
<ul><ul><li>First we weighed 2 evaporating dishes. Then we added 1 gram of our crystals to each of the dishes. We then placed both the dishes in the oven overnight. The next day, once the dishes were removed from the oven and cool, we weighed them once again to record the final weight. </li></ul></ul>
<ul><li>Why should you label the dishes before weighing them? </li></ul><ul><li>Labeling the dishes prior to weighing them is very important because if labeling tape is added after, then the recorded mass will be wrong. This will result in inaccurate data. </li></ul>
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