Titrimetric Redox Determination of Iron in an Ore Using Dichromate

CHM 2290
Experiment #5
Titrimetric Redox Determination of Iron in an Ore Using Dichromate

A.  Introduction

Many chemical compounds can be oxidized or reduced completely by reacting them with suitable reagents.  Oxidation-reduction reactions can be utilized for quantitative determinations if a suitable means of detecting the equivalence point can be established.  If more than one compound in a mixture is capable of oxidation or reduction, we will either need to (i) run a prior separation or (ii) utilize a reagent with an electrochemical potential which will make it selective for one component in a mixture relative to all others.  In the current experiment, potassium dichromate, K2Cr2O7, is used as the titrant to determine the amount of iron in an ore sample.   Since the potassium dichromate is ionized in solution, the redox reaction taking place during the titration can be written as:
6 Fe2+  +  Cr2O72-  +  14 H+       6 Fe3+  +  2 Cr3+  +  7 H2O
In this reaction, each Fe(II) ion loses one electron (so that it is oxidized) while each Cr gains 3
electrons (so that it is reduced).  Since there are two chromiums in each dichromate ion, a total of 6 electrons is gained by Cr2O72- so that 6 irons are reduced for each dichromate added.

The major advantage of utilizing dichromate as the titrant is that it can be obtained in a primary standard grade so that a solution prepared from a weighed K2Cr2O7 sample has a known concentration.  Moreover, dichromate solutions are stable over a reasonable length of time whereas potassium permanganate (KMnO4) solutions (which are also often used) must be continuously re-standardized.  The main advantage of permanganate is that its solutions are a deep violet color so that the titrant serves as its own indicator.  Dichromate solutions are orange-yellow but, at the concentration levels generally utilized, the color is too faint to be used for reliable endpoint detection.  Thus, a separate indicator is usually added to aid in the endpoint detection.

As with all oxidation-reduction titrations, it is first necessary to make certain that all of the material to be titrated is in the desired oxidation state.  This is often one of the more difficult steps in the determination and great care should be taken in this experiment to reduce all of the iron to Fe2+ without leaving any other oxidizable material in the solution.

B.  Preparation of Potassium Dichromate Solution

The primary standard grade potassium dichromate in the laboratory should already be dried and cooled and ready for your use.  Weigh out approximately 2.5 g of this dried material, transfer it to a 500-mL volumetric flask, and dilute to volume with distilled water.  This will result in a solution which is approximately 0.017 M in K2Cr2O7.  Based on the exact amount you weighed out, you will need to calculate the exact molarity of your standard solution using a molecular weight of 294.19 for K2Cr2O7.  This solution is then ready for use.


C.  Dissolution of Iron Ore Sample and Pre-reduction of Iron(III)

1. Obtain a sample of unknown iron ore from your laboratory instructor.  Weigh out three samples of about 0.5 g into 500 mL conical flasks.  Add 25 mL of 6 M HCl to each flask; cover with a watch glass and, in the fume hood, heat below boiling until the iron ore
dissolves (about 20 minutes).  The solution will turn yellow as the iron dissolves and Fe(III) is complexed with chloride ion.  A white flocculent residue of silica will also be observed.  A heavy black residue may be observed if insoluble sulfides or silicates are present.
RXN: 2 Fe3O4  +  18 HCl  +  1/2 O2       6 FeIIICl3  +  9 H2O

2. Since all of the iron must be in the Fe(II) oxidation state before the start of the titration and since Fe(II) is easily oxidized to Fe(III) upon standing in air, COMPLETE THE FOLLOWING STEPS AND THE TITRATION FOR ONE SAMPLE AT A TIME.  Do not allow the sample to sit around after the iron has been reduced to Fe(II).

3. Heat one solution to boiling, remove from heat and immediately add 0.5 M stannous chloride, SnCl2,  dropwise until the yellow Fe(III) is completely reduced to light green Fe(II).  Most samples will be so dilute that the Fe(II) will appear virtually colorless.  DO NOT ADD MORE THAN TWO DROPS IN EXCESS (but you must have an excess).
RXN: 2 FeIIICl3  +  SnIICl2       2 FeIICl2  +  SnIVCl4

4. Cool to room temperature, add 50 mL of water, stir rapidly, and add 10 mL of 0.25 M mercuric chloride (HgCl2) ALL AT ONCE (that is, DUMP IT IN FASt!).  {If you add the Hg(II) slowly, some of the Hg(II) may be reduced all the way to elemental mercury which will cause problems during the dichromate titration.}  This step oxidizes any excess Sn(II) to Sn(IV) so that it does not interfere in the subsequent titration (since Sn(II) would also react with Cr2O72- .  
RXN: SnCl2 (excess)  +  2 HgIICl2       SnIVCl4  +  Hg2Cl2 (s)
The mercurous chloride (Hg2Cl2) which is formed will precipitate immediately so that it also
does not interfere in the subsequent titration.  If no precipitate is seen, you failed to add enough SnIICl2 in step 3.  DISCARD THE SAMPLE!  Also, if the precipitate appears to be gray or black, mercury metal is present.  DISCARD THE SAMPLE!

5. If a white precipitate is observed following the preceding step, immediately add 200 mL of water, 10 mL of 1:5 sulfuric acid, 5 mL of 85% phosphoric acid, and 8-10 drops of barium diphenylamine sulfonate indicator to the flask containing Fe(II).  Titrate slowly with the standard potassium dichromate solution until you reach the end point.   As the dichromate reacts with the Fe(II), it is reduced to chromic ion, Cr(III), which appears blue-green.  The combination of this color with the indicator color results in a color change of blue-green to a grayish tinge to a purple.  The titration should be conducted dropwise when the gray tinge is noted because the indicator oxidizes somewhat slowly—that is, you must allow time for the indicator to react after every addition of titrant.
RXN: 6 Fe2+  +  Cr2O72-  +  14 H+       6 Fe3+  +  2 Cr3+  +  7 H2O

6. Calculate the percentage of iron in the ore sample using the atomic weight of iron (55.847).  Remember that the stoichiometric ratio of the titration reaction is 6:1.
my questions require only these calculations listed below:
1-PLEASE CALCULATE THE MOLARITY IN PART B IF I GET A WEIGHT OF 2.5077g
2-CALCULATE THEBPERCENTAGE OF IRON IN THE ORE SAMPLE AS MENTIONED IN # 6 PART C.
THE THREE TITRATION RESULT I GOT IN THE LAB ARE AS FOLLOW:
THE TOTAL WEIGHT OF MY UNKNOWN WAS 1.2690g
titration 1  m1=0.3051  buret reading V1=22.1 ML
TITRATION 2  m2=0.3013  buret reading  V2=22.3 ml
titration 3  m3=0.3061  v3=22.1 ml