EXPERIMENT III. Determination of Iron in Iron Oxide, (Fe 2 O 3 ), Using Dichromate Method. Chemical Overview

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1 EXPERIMENT III Determination of Iron in Iron Oxide, (Fe 2 O 3 ), Using Dichromate Method Chemical Overview This is a direct titration using K 2 Cr 2 O 7, a primary standard, as the titrant. As such the analysis is especially sensitive to the determination of the end point. Reactions: (1) Fe 2 O H+ 2 Fe H 2 O (2) Sn Fe3+ Sn Fe2+ (3) Sn+2 + 2Hg2+ + 2Cl- Hg 2 Cl 2 (s) + Sn+4 (4) Sn+2 + Hg2+ Hg0(s) + Sn+4 (5) K 2 Cr 2 O 7 + 6Fe H+ 2Cr3+ + 6Fe3+ + 7H 2 O HCl reacts with Fe 2 O 3 to generate Fe+3 (iron III). Stannous chloride (SnCl 2 ) is added, to a slight excess, converting iron III, Fe+3 to iron II, Fe+2. The reaction mixture now contains Fe+2, Sn4+, and Sn2+. (The iron II is relatively unstable and subject to air oxidation). Since Sn2+ is subject to oxidation by K 2 Cr 2 O 7, the Sn+2 must first be converted to Sn+4 state before the titration is started. Mercuric chloride (HgCl 2 ), a mild oxidizing agent, is added to the solution which converts the excess Sn2+ to the Sn4+ state while leaving iron II unreacted and forms Hg 2 Cl 2, a milky white precipitate as a product. This step also tests if: 1) Too large an excess of Sn II was added (reaction (4) takes place) and Hg0 (black precipitate) forms. This sample must be discarded as Hg0 can be oxidized by K 2 Cr 2 O 7. 2) Insufficient Sn II has been added, which results in no Hg 2 Cl 2 (a white precipitate) forming. This indicates that all of the Fe3+ was not converted to the Fe2+ state. This sample must also be discarded. Purpose 1. Test of procedual skills learned in Experiments I and II. Since this experiment uses a primary standard solution directly, both your weighing techniques and how precisely you detect your end point will effect your results. 2. To learn the subtleties of redox reaction which allow selective reaction in a solution containing several different ions. 39

2 Terms: Direct titration, oxidation, reduction, primary standard. Procedure Overview Procedure I. Prepare Potassium Dichromate II. Dissolution of iron ore You may stop analysis at this point. Cover samples with plastic wrap and store. Each sample is treated individually from this point on as Iron (II) is easily oxidized by air. III. Prereduction of Iron (III) IV. Titration of Iron (II) I. Preparation of Potassium Dichromate (K 2 Cr 2 O 7 ) Equipment: unribbed watch glass 600 ml beaker 500 ml volumetric flask 1) On a clean, dry, preweighed unribbed watch glass, weigh 2.452X grams ± 0.1 gram of primary standard potassium dichromate (That is, weigh between 2.352X and 2.552X grams.) Record the weight to four decimal places. 2) Quantitatively rinse the watch glass into the 500 ml beaker. 3) Transfer the diluted dichromate to a 500 ml volumetric flask. Quantitatively rinse the beaker. Dilute to mark. 4) Calculate the molarity of solution: Note that if the weight is not exactly grams, the normality will not be exactly You are calculating molarity. Formula: Molarity K 2 Cr 2 O 7 = II. Dissolution of iron ore (Fe 2 O 3 2 Fe3+) grams K 2 Cr 2 O 7 1 mole K 2 Cr 2 O 7 M.Wt. K 2 Cr 2 O 7 Purity 40 Volume in liters Equipment: ml conical flasks 4 ribbed watch glasses (Speedyvaps) 1) Weigh 0.35XX gram (± 0.1 gram) samples of the iron oxide unknown into 4 flasks. You may have to discard the other three samples and adjust the grams of unknown used if the first titration uses more or less than 35 ml ± 5 ml. This is a time saver gamble. If your first reduction of iron III fails you have back up unknown samples prepared. However because the amount of iron varies in different unknowns, the milliliters of titrant, potassium dichromate, used per unknown sample may fall out

3 side the range of 30 to 40 milliliter. If this occurs, you will need to scrap the remaining weighed unknowns and adjust your sample weights according to the ml/wt. ratio of your first titration. 2) Add ml of 12 M. Hydrochloric acid (HCl) to each flask. Swirl to mix, cover with ribbed watch glasses. 3) On the heaters in hood, evaporate solution to about 5 ml. There may be precipitation of iron salts due to concentration. 4) Dilute each solution to 15 ml with distilled water. This will usually dissolve any precipitated salts. Samples may be covered with plastic wrap and stored at this point. From here on do each solution individually to completion. III. Prereduction of Iron (III) to Iron (II) Equipment: 2 small beakers eye dropper 100 ml graduated cylinder 10 ml graduated cylinder 1) In the hood, heat flask #1 with the Iron (III) solution to boiling. While the solution is heating, prepare the chemicals for the prereduction step. 2) Preparation of chemicals: put small amounts of 0.5 M. Stannous chloride and 0.25 M. Mercuric chloride into beakers. Measure 10 ml of Mercuric chloride into 10 ml graduated cylinder Measure 50 ml of distilled water into 100 ml graduated cylinder 3) Remove the flask (boiling) from heat. 4) DROPWISE, add 0.5 M. Stannous chloride until the yellow iron solution is reduced to light green ( Iron (II) ). You may not be able to detect the light green color of the iron (II) unless your initial solution was very concentrated. The solution may simply go from yellow to colorless or very faint green. Do not add more than 2 drops excess. 5) Cool to room temperature. 6) Add 50 ml of distilled water, stir rapidly (or swirl flask). 7) Using a 10 ml graduated cylinder, add all at once 10 ml 0.25 M Mercuric chloride. Mercuric chloride is added all at once to avoid reducing the mercuric ion to mercury metal. 41

4 8) Wait 3 minutes. If your reduction was carried out properly, the precipitate in the flask will appear to be "silky white" at this point. The actual appearance depends on the amount of excess stannous chloride added. Too much causes greying of the precipitate and the loss of the silky appearance. A white precipitate should form, but may appear cloudy. If no precipitate at all forms, there was no excess stannous chloride, discard the sample. If the precipitate is dark grey or black, mercury metal is present, discard the sample. While waiting the 3 minutes, clean and fill buret with K 2 Cr 2 O 7. IV. Titration of Iron (II) 1) Prepare: a) Using 100 ml graduated cylinder to measure,.prepare 200 ml of distilled water in a beaker. b) Wear Safety Goggles.1:5 Sulfuric acid (H 2 SO 4 ) is prepared by first adding 8 ml distilled water to a 10 ml graduated cylinder followed by then, adding 2 ml concentrated H 2 SO 4. c) Add 5 ml of 85% Phosphoric acid to a separate graduated cylinder. bottle. d) Sodium or barium diphenylamine sulfonate is all prepared in a dropper 2) Immediately after 3 minutes if a white precipitate forms: Add in rapid sequence: 200 ml of distilled water 10 ml of 1:5 H 2 SO 4 5 ml of 85% Phosphoric acid 8 drops of Na or Ba diphenylamine sulfonate Swirl flask to mix 3) Place the flask on a magnetic stirrer and put a clean stirring bar in the solution. Titrate slowly with potassium dichromate solution. The end point is a color change from blue-green to grey to purple. At the grey point proceed titrating dropwise to purple. After the endpoint is reached the color may revert back to grey or blue-green if the solution stands for a while. 4) Remember, if the amount of potassium dichromate used to titrate is not about 35 ml, discard the 3 remaining iron sample solutions and adjust the unknown sample weight so that you will be using about 35 ml of potassium dichromate in the titration. 42

5 5) Repeat steps from III. Prereduction of Iron (III) with the 3 remaining samples, if titration of first was about 35 ml. Repeat from the beginning if you are adjusting weight of unknown. 6) Calculations: a) Percentage of iron in sample: Formula: % Fe = Vol.(L) K 2 Cr 2 O 7 M K 2 Cr 2 O 7 6 mole Fe 1 mole K 2 Cr 2 O 7 At. Wt. Fe 100% gms unknown b) Find the mean %Fe. c) Find the mean deviation. d) Fine the RMD in ppt. Notes 43

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