Question

Below is the background info for the lab assignment. The 4 blank boxes are the questions I would like answers too. The end point for the fine titration was 35.90 mL in the burette when the solution turned bright green again. And the coarse titration I got 35.36mL as the end point. 35.9 mL is exact in case you need that info. Thanks!

Background

In this lab, we will determine the amount of alcohol (ethanol), C₂H₅OH, in a commercial vodka sample using a method called a redox titration. The technique of back-titration will also be used. In a back-titration, a known excess of a standard reagent is added to the unknown solution, and then the excess is titrated with another standard.

Most chemical methods of analysis involve oxidation of the ethanol with a strong oxidizing agent and measurement of the amount of oxidant consumed in the reaction. The products of the oxidation of ethanol with most oxidants depend greatly on the experimental conditions used. Therefore, it is necessary to adhere closely to those conditions when the stoichiometry of the reaction is known. In this experiment, the oxidizing agent is the dichromate ion (Cr₂O₇^-2) in the reagent potassium dichromate (K₂Cr₂O₇^-2).

Using an acid and dichromate ion (Cr₂O₇^-2), small amounts of ethanol will be oxidized to acetic acid (CH₃COOH) at room temperature according to the equation:

2Cr₂O₇^-2(aq)+16H⁺(aq) + 3C₂H₅OH(l) → 4Cr⁺³(aq)+11H₂O(l)+ 3CH₃COOH(aq)

The proton (H⁺) is supplied by addition of dilute sulfuric acid (H₂SO₄). Note the 2:3 ratio in which the dichromate ions react with the ethanol molecules. To ensure that this reaction goes to completion, an excess of Cr₂O₇^-2 must be used.

To determine the moles of ethanol (C₂H₅OH) in solution, we must determine the amount of dichromate ion (Cr₂O₇^-2) that reacted. To do so, the unreacted dichromate is back-titrated with a standard iron(II) ion (Fe⁺²) solution using the reagent iron(II) ammonium sulfate hexahydrate, Fe(NH₄)₂(SO₄)₂^. 6H₂O. The back-titration allows us to determine the amount of excess dichromate from the first reaction. The balanced equation for this reaction is

6Fe⁺²(aq) + Cr₂O₇^-2(aq) + 14H⁺(aq) → 6Fe⁺³(aq)+ 2Cr⁺³(aq) + 7H₂O(l)

An iron(II) ion solution is added to the solution from the first reaction that contains the excess dichromate ion (Cr₂O₇^-2) until all of the dichromate ions are completely reduced. A problem is that it is difficult to determine the end point of such a titration because dichromate solutions are orange while the Cr(III) solutions are green. To address this issue, a redox indicator sodium diphenylamine sulfonate is used. The indicator diphenylamine sulfonate ion can be reversibly changed between its oxidized state (deep purple) and its reduced state (colorless). In the presence of dichromate ions it is a deep purple; it is colorless in their absence.

In this lab, once the dichromate ions have been converted to chromium(III) ions, the color of the solution will change suddenly from the purple color of the redox indicator to the bright green color of the chromium(III) solution. (Because all of the dichromate ions have been used, the diphenylamine sulfonate solution becomes colorless, allowing the green color from the Cr⁺³ ions to be seen.

Note the 6:1 ratio in which the iron(II) ions react with the dichromate ion. You will need to take this into account when calculating the moles of excess dichromate ion.

Procedures

Experiment 1: Preparing the Materials

PLEASE NOTE: The procedures described in this lab assume that you have already done the Titration Tutorial and are familiar with the technique.

1. Take four volumetric flasks from the Containers shelf and place them on the workbench.

2. In one flask, prepare a standard solution of potassium dichromate (K₂Cr₂O₇):

a. From the Materials shelf, add 4.00 g of dry potassium dichromate to the volumetric flask.

b. Add 30 mL water to dissolve the dichromate compound.

c. Complete the solution by filling the volumetric flask to the 100.00 mL mark with water (do this by checking the "Fill To Mark".)

LABEL: Be sure to label this flask and other containers in this experiment by double-clicking on the container and renaming it.

Record in Lab Notes: Be sure to record all measurements for all solutions, results, and observations in your lab notes.

3. In two of the empty flasks, prepare standard solutions of iron(II) ammonium sulfate hexahydrate (Fe(NH₄)₂(SO₄)₂^.6H₂O). (These solutions will be used for the titrations in Parts 2 and 3 of Experiment 2.)

a. From the Materials shelf, add 4.00 g of iron(II) ammonium sulfate hexahydrate to each empty volumetric flask.

b. Add 30.00 mL water to each to dissolve the compound and release the water of hydration.

c. Complete both solutions by filling the volumetric flask to the 100.00 mL mark with water (do this by checking the "Fill To Mark".)

4. In the last empty flask, add 2.00 mL of the commercial Grey Moose vodka from the Materials shelf and fill with water to the Fill Mark to make a 100.00 mL solution.
The vodka has now been diluted to 1/50th, or 2.00%, of its original ethanol concentration.

Experiment 2: Titrating the Vodka Sample

Part 1: Oxidation of Ethanol in Vodka

1. Take a 150 mL Erlenmeyer flask from the Containers shelf and place it on the workbench.

2. Add 5.00 mL of diluted vodka solution from the volumetric flask to the Erlenmeyer flask.

3. Add 35.00 mL of water from the Materials shelf to the Erlenmeyer flask.
Note that this further dilutes the vodka sample by a factor of eight; the ethanol concentration is now 1/8th of 2.00%, or 0.250% of the original ethanol concentration of the bottled vodka.

4. Acidify the vodka solution in the Erlenmeyer flask by adding 5.00 mL of sulfuric acid (H₂SO₄) solution.

5. Add 5.00 mL of the standard potassium dichromate solution to the Erlenmeyer flask. This is enough to reduce all of the ethanol in the vodka and leave an excess of dichromate ions. Note that the solution has turned bright green; this is the color of the reduced Cr⁺³ ions. Remember to record data and observations in your Lab Notes.

Part 2: Coarse Titration

1. To the green solution in the Erlenmeyer flask, add 0.500 g of the redox indicator, sodium diphenylamine sulfonate. In the presence of the excess dichromate ions it turns the solution a deep purple because it is in its oxidized state.

2. Take a burette from the Containers shelf and place it on the workbench. Fill the burette with 50.00 mL of the standard iron(II) solution. Record the initial burette reading.

3. Place the Erlenmeyer flask on the lower half of the burette – this will connect them.

4. Perform a coarse titration: Adding large increments of the standard iron(II) solution from the burette by pressing and holding the black knob at the bottom of the burette until the solution turns suddenly from intense dark purple to green. Each time you add the standard iron(II) solution to the flask, record the volume remaining in the burette. Be sure to record all the significant figures in volume reading.

As the iron(II) is added, the dichromate ions (Cr₂O₇^-2) are reduced to Cr⁺³ ions. At the end point of the titration, there are no dichromate ions left and the redox indicator becomes colorless and the deep purple color suddenly disappears leaving the solution bright green again, the color of the Cr⁺³ ions.

5. In your Lab Notes, identify both the last burette volume that the solution was dark purple, along with the burette volume at which the solution first turned green. This gives the range where the endpoint occurs.

6. Discard just the Erlenmeyer flask in the recycling bin underneath the workbench.

Part 3: Fine Titration

1. Set up the titration as before:

a. Prepare a new solution to be titrated by following the steps in Part 1. Then add 0.500 g of the redox indicator, sodium diphenylamine sulfonate, to the Erlenmeyer flask, as you did in Step 1 of Part 2. The solution should turn a deep purple color.

b. Connect the Erlenmeyer flask to the lower half of the burette.

c. Note the current volume of standard iron(III) solution in the burette. To make sure you have enough titrant, you should add more iron(III) solution to the burette from your volumetric flask on the workbench. Record the initial burette reading.

2. Click and hold the black knob of the burette to quickly add enough standard iron(II) solution to just get into the range of the coarse titration (the first number you recorded), but still have the solution in the flask appear dark purple. This is near, but not yet at, the titration's end point.

3. Add standard iron(II) solution from the burette in small increments, down to one drop at a time, until the addition of just one more drop causes the solution in the flask to turn green. Record the final burette reading.

4. Clean your workstation by placing the Erlenmeyer flask in the recycling bin underneath the workbench.

5. Repeat the fine titration once more, and record the results. If the results from the two fine titrations do not closely agree to within 0.20 mL of titrant added, do at least one more fine titration to obtain precise results.

LNL7 Alcohol Content of Vodka by Dichromate Titration

Assignment

Introduction Questions:

Refer to the Background and Procedure information to this lab to answer these questions.

Experiment 1: Preparing Solutions

Show work and include units and the correct number of significant figures on all measurements and calculated results.

6. Calculate the molarity concentration of potassium dichromate, K₂Cr₂O₇, in the first volumetric flask. This is also the concentration of dichromate ion.

           

7. Calculate the molarity concentration of Fe(NH₄)₂(SO₄)₂·6H₂O in the second two volumetric flasks. This is also the concentration of Fe²⁺ ion.

Experiment 2: Titrating the Vodka Sample

8. For the Course Titration, in what volume range did the solution turn from purple to green?

Answer 1

6) molarity of potassium dichromate will be

Molarity = Mass of dichromate dissolved / Molecular weight of dichromate X volume of water

Molecular weight of potassium dichromate = 294

Molarity = 4 / 294 X 0.1 = 0.136 M

7) molarity of Fe(NH₄)₂(SO₄)₂·6H₂O

Molarity = Mass / Molecular weight X volume in litres

Molecular weight of Fe(NH₄)₂(SO₄)₂·6H₂O = 392

Molarity = 4 / 392 X 0.1 = 0.102 M

8) The solution will change colour when 36mL of solution will remain in burette

so the volume of solution used = 14mL