9: Iron Determination by Titration

INTRODUCTION

Quantitative analysis is the measuring of the amount of a substance in a mixture. There are multiple standard quantitative techniques which can be used to measure concentration, and one that we are going to learn about today is titration.

Titration works like this:

a) find a compound (the titrant) that reacts with the compound (the analyte) that you want to measure in a mixture,

b) determine the stoichiometry of this reaction,

c) make up a solution of the titrant whose concentration is accurately known,

d) add just enough titrant to balance the amount of analyte in your unknown sample,

e) from the volume of titrant, concentration of titrant, and reaction stoichiometry, calculate the amount of analyte originally present in the sample.

Two problems have to be overcome for a titration to work. First, the concentration of the titrant needs to be accurately known. Often this solution is not easily made to an accurate concentration. Thus, it needs to be standardized (its concentration needs to be accurately determined). Secondly, how do we know when just enough titrant has been added to react with the amount of analyte (the equivalence point) in the unknown sample? Often we can add an indicator which changes color as soon as there is an excess of titrant in the solution which is defined as the end point. The end point is usually just after (but quite close to) the equivalence point. Thus, we will take the end point to represent the equivalence point.

In this lab, we will measure the amount of analyte, that is iron II (Fe²⁺), in a solution by measuring the amount of an oxidizing agent, KMnO₄, that it takes to oxidize it to iron III. Potassium permanganate easily oxidizes iron II to iron III by taking an electron from each iron II cation. Obviously, the stoichiometry is not one-to-one. The following is the balanced reaction for the reduction of the permanganate anion under acidic conditions:

8 H⁺(aq) + MnO₄^-(aq) + 5 e^- à Mn²⁺(aq) + 4 H₂O

So, one must determine how many iron II atoms need to be oxidized for the reduction of each permanganate anion. You will determine this ratio in your pre-lab assignment.

Now how do we know when we have added an equivalent amount of titrant to just match the iron II in our sample? This is quite easy because the permanganate ion (MnO₄^-) is dark purple, but the reduced manganous ion (Mn²⁺) is colorless. Thus, as you begin to add the titrant, the purple immediately disappears, but as soon as all the iron II is oxidized and a small excess of permanganate has been added, a faint purple persists in the solution. You want to stop adding the titrant just as the faintest purple (looks more like pink) persists in the solution. The amount of titrant added will be very accurately determined using a buret. The amount of titrant in the buret is recorded both before and after the titration. Therefore, the difference is the amount of titrant added to the reaction. Because this reaction produces a color change at the end point, we will not need to add an indicator in this experiment.

Permanganate solutions cannot be made up with an accurately known concentration; thus, each of us will help standardize the titrant at the beginning of the lab. Then each of you will titrate your unknown sample three times. You will use your results to calculate the amount of iron II present in your unknown.

PRE-LAB QUESTIONS

The following is the balanced reaction for the reduction of the permanganate anion under acidic conditions:

8 H⁺(aq) + MnO₄^-(aq) + 5 e^- à Mn²⁺(aq) + 4 H₂O

What is the balanced reaction for the oxidation of iron II by permanganate under acidic conditions? Note this balanced equation will be needed in lab.

What is the ratio of MnO₄ to iron II? _________________________________

In today’s lab, are we going to use the equivalence point or the endpoint? Explain why in your own words.

Why is it so important to look for the faintest color change that persists for 30 seconds? Explain in your own words.

PROCEDURE

Preparing for the titration

1. Obtain a buret and buret clamp and mount them on your stand keeping the buret vertical. Burets should be clean, but you should test them in any case. Close the stop cock, place a beaker under the tip, fill to close to the 0 mark with DI water, and open the stop cock to let the water empty into the beaker. If any water droplets remain on the inside of the buret, it is not clean. Titrant will be left behind during your titration, and the amount you determine for titrant added will be inaccurate. During lab we will discuss proper cleaning before and after use. Also, make sure that your buret does not leak.

2. Using 6 M H₂SO₄, prepare 300 mL of 1 M H₂SO₄. Download a local copy of the linked Excel Spreadsheet. In the Spreadsheet, calculate the amount of 6M sulfuric acid needed. First add the DI water that you determined was needed to make the 1 M H₂SO₄ to your 400 mL beaker, then add the amount of 6 M H₂SO₄ you determined was needed. (This will be accurate enough for this experiment.). Mix with your stirring rod.

3. Obtain an unknown iron II sample and record its label in your report sheet.

4. Obtain about 100 mL of titrant (KMnO₄), rinse your buret with this titrant as described at the beginning of lab, and fill your buret to close to the 0 mark, but do not go over. Make certain that all air/air bubbles from the tip of the buret have been eliminated.

Standardization of Potassium Permanganate

1. Using weigh paper, obtain about 0.6 g of ferrous ammonium sulfate hexahydrate (Fe(NH₄)₂(SO₄)₂·6H₂O) and record the mass as accurately as possible. Record this and the following titration data in your local copy of the Excel Spreadsheet. This is the iron II standard with a molar mass of 392.2 g/mol. Add this solid to a 250 mL Erlenmeyer flask (washing all of it in with a little DI water from a wash bottle), add 40 mL of the 1 M sulfuric acid, and swirl to completely dissolve the solid. To prevent yellowing of the solution add about 3 mL of 85% H₃PO₄. This yellowing can mask the initial pink color at the end point.

2. Record the initial volume of titrant in the buret to hundredths of a mL. With swirling, add titrant to the sample until the faintest pink color possible persists for at least 30 seconds. Record the final volume of titrant in the buret.

3. Using the volume of titrant, moles of iron II standard, and stoichiometry, calculate the concentration of the titrant. Input this value into the cloud version of the linked Excel spreadsheet under sheet “Class data”.

4. After everyone has entered their values, copy the data into your local copy of the Excel spreadsheet and calculate the average and standard deviation for the concentration. You can begin your unknown titrations before you know this value.

Titration of Your Unknown Iron II Sample

1. Using weigh paper, obtain about 0.35 g of your unknown sample and record the mass as accurately as possible. Add this solid to a 250 mL Erlenmeyer flask (washing all of it in with a little DI water from a wash bottle). Add 40 mL of the 1 M sulfuric acid and swirl to completely dissolve the solid. Add about 3 mL of 85% H₃PO₄.

2. Record the initial volume of titrant in the buret, with swirling add titrant to the faintest pink endpoint and record the final titrant volume. See Excel sheet.

3. Repeat steps 1 & 2 with two more 0.5 g samples of your unknown.

4. Waste disposal: All purple colored solutions will go into the waste container labeled “permanganate waste”. All pink solutions can be poured down the drain.

5. For each of your three titrations, use the volume of titrant, standardized concentration of titrant, mass of unknown, and reaction stoichiometry to calculate the percent mass of iron II in your sample. See Excel sheet.

6. Calculate the mean and standard deviation for the percent mass of iron II from your three replicate values. Report these values on your Excel sheet.

Final Questions

1. How close was your standardized KMnO₄ concentration to the average of the whole class? Calculate the % difference between your value and that of the class (this is much like a % error calculation). %diff = (your molarity – class avg molarity)/class avg molarity * 100%

2. How would your results turn out if an additional unknown metal was present in the sample as a contaminant that is also oxidized by KMnO₄? Would your calculated result be lower, higher, or the same as the actual amount of Fe²⁺ present? Why?

3. What are the standard deviations for the whole class standardization of the KMnO₄ and for your three titrations of your unknown? Show your math for calculating the standard deviation for your three replicates. What do these standard deviation values tell you about the precision or accuracy or both of the reported values for the KMnO₄ concentration and the % iron in your unknown sample?

4. This type of experiment can be quite accurate with proper care. List three steps that you think have the most potential for introducing error into your final result. Explain how each would add error to your results. Order them from most severe to least severe.