Analysis of Iron Oxidation of Metals If you’ve made it this far in your chemistr

Question

Analysis of Iron

Oxidation of Metals
If you’ve made it this far in your chemistry career, you know that metals in one oxidation state can be oxidized to another oxidation state. An atom of copper I (Cu+), for instance, can become copper II (Cu2+) if it encounters an oxidizing agent strong enough to pluck one more electron off it. As this happens, the oxidizing agent is itself reduced.

Calculations Example
Let’s say 13.03 mL of an 0.06451 M solution of peroxide is needed to titrate 25.00 mL of a mixed iron solution (one with Fe2+ and Fe3+).
First we need to know how many moles of permanganate were used, which found as follows. Note conversion of mol/L to mol/mL is not depicted:

And finally, calculate the concentration of Fe2+ in the 25.00 mL of it titrated (again, conversion of mL to L not depicted):

Determining the total [Fe]
In the second phase of this experiment, you will determine the total concentration of iron in solution; that is, the concentration of Fe2+ plus the concentration of Fe3+. This will involve two reactions. For the first of these, you will use some zinc metal to “convert” all the Fe3+ in solution to Fe2+. For the second, you’ll use another permanganate titration to determine how much Fe2+ is in this solution.

Determining the concentration of iron in this solution is done the same way as the previous solution. The only difference is what concentration exactly you’re calculating. When we calculated the concentration of iron in the solution that didn’t have zinc added, we determined the concentration of Fe2+ in the original solution. When we calculate the concentration of iron in the solution that did have zinc added, we’re calculating the total concentration of iron (Fe2+ and Fe3+) in the original solution.

PROCEDURE
Part A: Set Up
1. Obtain about 50 mL of KMnO4 solution and ready a burette to dispense it.

You’ll probably want to use a burette that has white markings, but that’s up to you.

Note that a portion of the credit you receive for your burette data will depend on reading it correctly and making appropriate initial and final readings, with the correct number of significant figures.

2. Your TA will either assign or let you choose an unknown iron solution to analyze. Acquire approximately 100 mL of whatever you’re given.

NOTE: You’re going to want to exclude air from your iron sample as much as possible (Why? Well, air has oxygen in it. What might this do to your sample?). Keep all sample containers covered with foil and try to avoid agitating the sample unless you have to.

3. Pipette TWO aliquots of your iron solution into each of two clean 250 mL Erlenmeyer flasks.

That’s two pipette’s worth in each flask.

4. Add 9-10 mL of 3 M sulfuric acid to each flask

Part B: Reduction of Iron
5. Add one scoop full of zinc metal to one of the flasks. Gently heat this flask until it is just a little too warm to comfortably hold in your hand, then swirl it for a minute to mix the contents. Heat it back up (if it needs it), and swirl it again, then let it sit on the plate with the dial set to low while you do the next part of the experiment. You can take it off the plate after it’s been there for about 15 minutes, even if you’re not yet done with the next part.

You’ll titrate the other, non-zinc flask while the one with zinc warms. Periodically give the “zinc” flask a swirl and don’t let it get too warm.

Part C: Titration of Iron Solutions
6. Add 9-10 mL of 8.5% phosphoric acid to the flask you’re about to titrate.

7. Titrate the solution with KMnO4 until the titration reaches its endpoint. Gently swirl the flask as you do this. You’ll titrate the “non-zinc” flask while the “zinc” flask warms. When you’re done, titrate the “zinc” flask.

Part D: More Data
8. No proper scientific experiment relies on one data point and neither will this one. Wash your flasks out and perform another set of titrations with them.

Part E: Clean-up
9. Dispose of all remaining solutions in the waste container in the fume hood.
10. Wash all glassware with detergent, rinse thoroughly with water, and place on your towel to dry.
11. Clean-up all spills and throw all garbage in the garbage cans.

One More Thing: Your Equipment

Every piece of glassware in your drawer must be clean and your drawer organized before you leave the lab room; basically it has to look like the picture below, with most of your pieces lined up in their proper place on the placemat inside.  Some of the points you receive for this experiment will be riding on how well you do this.

QUESTIONS:

Explanation / Answer

1) The purpose of the experiment is to demonstrate redox reactions. A more important find out from the experiment is to quantitatively determine the mass percent of Fe(II) and Fe(III) present in a sample of iron by titrating the Fe(II) with permanganate (this is a titration) and then converting the Fe(III) to Fe(II) by reduction and carrying out the permanganate titration.

2) Permanganate (MnO4-) oxidizes Fe(II) to Fe(III) and itself gets reduced to Mn2+. The end point of a redox titration will be indicated by a purple color (due to excess MnO4-) of the solution. This happens because, as long as there is Fe(II), MnO4- will oxidize Fe(II) to Fe(III). When all the Fe(II) is exhausted, MnO4- will no longer have a reducing agent to act upon and will be in excess. The end point of the titration is the volume of MnO4- which quantitatively converts all Fe(II) to Fe(III).

Fe(III) is yellow in color and can interfere with the end point detection. This problem is mitigated by adding phosphoric acid to the solution. Phosphoric acid forms a colorless complex with Fe(III) and hence, the end point detection is not compromised.

3) The total Fe concentration/amount can be determined by adding a reducing agent, zinc (Zn) metal to the solution of Fe(II) and Fe(III). Zn reduces Fe(III) to Fe(II) and Fe(II) can be estimated quantitatively by titrating with MnO4- in presence of phosphoric acid.

4) The end point of the titration will be purple when the titrant (MnO4-) is in excess. The end point is reached when all the Fe(II) is quantitatively oxidized to Fe(III) and the solution contains a slight excess of MnO4-.

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