3202 Chemistry Laboratory Investigation to Determine the Percentage of H2O2 in a Commercial

3202 Chemistry Laboratory Investigation to Determine the Percentage of H2O2 in a Commercial Sample.

Objective: To standardize a solution to be used in a redox titration and then to determine the concentration of an analyte in a solution

This lab has two major tasks. The first task is to standardize the concentration of a KMnO4 solution. This task is necessary in order to complete the second task, which is to evaluate how close commercial H2O2 solutions are to their labeled concentrations. The permanganate ion is commonly used as a titrant (in the BURET), it changes from a purple color to colorless (or pinkish). When the titration is just past the endpoint, the solution is just barely pink (the ENDPOINT, stop the titration). The point just before that was the equalization point. You will first standardize a permanganate ion solution, then you will use the standardized solution in a titration to determine the concentration of an analyte (component in a solution)

Standardization: Titration 1: MnO4 - (purple) + Fe 2+ + 8 H + → Mn 2+ + Fe 3+ + H2O

The unbalanced equation for TITRATION 2: MnO4 -(aq) + H2O2(aq) + H +(aq) → Mn 2+ (aq)(pink)+ O2 (g) + H2O(l)

Store bought peroxide claims to have a concentration of 3%. That value means that 3g of hydrogen peroxide is in 100ml of solution (which is 100grams of solution and is 97% water) (or 0.030g/ml) and this is equivalent to 0.88M.

Materials:

~0.020 M KMnO41.0 ml graduated pipet100 ml Volumetric flask

0.250g (Fe(NH4)2(SO4)2·6H2O)Mohr’s salt4 Erlenmeyer flasks3% hydrogen peroxide solution (5 ml)

50ml buret3M H2SO4graduated cylinder

Distilled waterBuret clamp and Ringstand

Procedure: Here are some guidelines to follow:

1. On a balance, weigh about 0.35 g KMnO4 crystals on a piece of weighing paper. Add the crystals to a beaker to warm and dissolve in about 50 to 60 mL of water to the flask and heat the solution with occasional swirling to dissolve the KMnO4 crystals. Do not boil the solution. This may take about 30 minutes. WRITE DOWN the KMnO4 used, you don’t have to get 0.35 g exactly, just close. hen transfer to a 100 ml VOLUMETRIC flask and get the solution to 100ml exactly. Now calculate the purple permanganate concentration. Stopper solution, stuff paper towel in or get a stopper or wax sheet. You will need this solution for day 2. PUT NAME ON IT or on a piece of paper under it. UNLESS the solution is already made, then just go up and get 50 ml of stock solution at front AND the acid

Titration 1 : Standardization with Fe +2 Complex or Mohr’s salt

1. You will make, or be given a 0.0500 M Fe +2 solution, to which you can add 10 ml of 3M acid. This will be in the FLASK, and the PERMANGANATE (purple) is in the buret. If you have to make the solution yourself, then weigh out close to 0.25 g to 0.30g of Iron +2 , in weighing boat or small beaker. Don’t be exact and take up all your time at the balance, but whatever mass you get RECORD it exactly, as you did above with the permanganate.

2. To the Iron add 10 or 20 ml water to dissolve, transfer it ALL to an Erlymeryer flask. WRITE down the exact mass you weighed.

3. Titrate, which means open the stopcock and slowly add the purple “juice” till pink.

Titration 2: H2O2

1. Using a pipet, transfer 1.00 mL of the commercial hydrogen peroxide solution into a 125-mL Erlenmeyer flask.

2. Add about 20 mL of distilled or deionized water to the flask.

3. Measure 15 mL of 3 M sulfuric acid into a graduated cylinder and carefully add the acid to the solution in the Erlenmeyer flask. Gently swirl the flask to mix the solution. CAREFUL this is NASTY acid. NEVER take off googles.

4. Titrate 3 times to pink again.

Note: the H2O2 is approximately 44 times as concentrated as the permanganate (0.88M H2O2 vs 0.020 M for permanganate).

Data: Create a chart for the values you should be recording for each titration. First for the permanganate standardization:

1.Calculate the moles of iron(II) ions in the sample

2.Calculate the moles of KMnO4 used in the titration (refer to the equation in the information section for the mole ratio to use)

3.Calculate the molarity of the KMnO4 solution

4.Record the molarity that represents the average for the whole class

Now for the peroxide titration:

5. Calculate the moles of KMnO4

6. Calculate the moles of H2O2

7. Calculate the molarity of the H2O2

8. What is the average molarity of H2O2

Conclusion Questions:

1.Was your molarity different than the molarity claimed (0.88M)? If so, was it higher or lower?

2.How would your calculated concentration of the H2O2 solution been affected (i.e. would it be too high or too low) if when you stopped the titration, the solution in the flask was dark purple instead of very light purple? Explain your answer.