LI PO CHUN UNITED WORLD COLLEGE SESSION MAY 2016
DETERMINATION OF IRON CONTENT IN IRON TABLET Chemistry Internal Assessment (Standard Level) Wong Long Sang (Isaac) 10/27/2015
Introduction My 87 years old grandmother (Figure 1) has anemia which was caused by an iron deficiency in her body system. As part of her treatment from the public public hospital, the doctor has prescribed prescribed her with a certain type of iron tablet (ferrous sulphate (FeSO4)) as part of her diet. However, my grandmother has not received much improvement in health after such consumption. This situation stimulated two questions in my mind. Since my grandmother takes two iron pills per day, how much iron content is she actually consuming into her body? Is the iron content as stated in the capsule bottle really consistent with the actual iron content? This chemistry internal assessment gave me just the opportunity for me to answer this question in practical terms. I decided to determine the iron content of iron tablets with my knowledge on redox and acid base chemistry. For research purposes, I bought the exact same brand and type of iron tablets (Figure 2), produced by Apt Pharmaceuticals, my grandmother was prescribed. I tried to see if there were any difference in the iron content the doctor thinks there is and what my grandmother is actually taking in.
Figure 1 - I (right) am very close to my grandmother (left). This investigation is dedicated to my beloved grandmother.
Figure 2 - The iron tablets my grandmother consumes
Methodology For a more effective and accurate investigation of the iron content in ferrous sulphate, I used two methods based on chemical knowledge knowledge learnt in class. class. Below, I will describe the the theory and the relevant procedures of each method. g of iron per tablet. From the information stated on the capsule bottle, there is 60mg/0.6 g of
Method 1 – Redox Titration In an acidic solution, potassium permanganate oxidizes iron(II) to iron(III), while itself would be reduced to manganese(II) in the following equation.
-
2+
+
MnO4 + 5Fe + 8H
2+
Mn
3+
+ 5Fe + 4H2O
Potassium permanganate is deep purple in colour while Mn 2+ ions are of light pink. To determine the iron content, I titrated an iron solution with KMnO 4 until the endpoint is reached as from the above chemical equation. Knowing the concentration and the volume of MnO 4- used to reach the endpoint, I would also know the number of moles MnO 4- in the chemical equation. I can then calculate the number of moles of Fe2+ ions with a stoichiometric mole ratio of 1:5. The iron content can then be determined after multiplying the number of moles of Fe 2+ with the molar mass of Fe (55.85). For the iron solution to be titrated, I dissolved the iron tablets into deionized water and added 2M sulphuric acid to form an acidic a cidic medium.
Figure 3 – The dissolving of 5 iron tablets into deionized water with a magnetic stirrer (Step 1)
Experimental procedure:
1.
Using a magnetic stirrer, 5 iron tablets were dissolved in 200 cm 3 of deionized water.
2. Wearing gloves, safety googles and and a lab coat , the the solution was transferred into a 250 250 cm3 volumetric flask and filled up to to the graduation mark mark with 50 cm3 2M sulphuric acid. Prior to reaching the mark, a water sprayer was used around the glass surface near the opening to wash any residue of Fe into the flask. 3. After filling up a burette with 0. 1M of KMnO4, an initial reading was taken while holding a white tile behind behind the marks for a clearer view. This is due due to the very dark purple nature of KMnO4 which would affect affect the reading of the marks engraved engraved in brown. 4. 25 cm3 of iron solution was pipetted into a conical flask. 5. Using a burette, the solution was titrated until it reached an endpoint of turning from colorless to light pink. 6. The reading in the burette after titration was recorded, again while holding a white tile for clearer and more accurate reading.
` Figure 4- Conical flask at endpoint (Pink in colour)
Method 2 – Precipitation method In method one, the iron content was determined through multiplying the stoichiometric mole ratios of Fe2+ ions and multiplying it with the relative atomic mass of Fe. However, there is was underlying assumption that the ferrous sulphate in the tablets consisted of only of Fe 2+ ions which would all be oxidized to Fe3+ with KMnO4. During the manufacturing or storage process, some Fe2+ may have been oxidized (or rusted) into Fe 3+ ions. This would affect the accuracy of the iron content calculated in method 1 as it was based solely on the number of moles of Fe 2+ on the reactant side, neglecting the possible presence of Fe 3+ ions. This leads on to the method of precipitation which addresses this problem. I first dissolved 5 iron tablets into deionized water then added concentrated nitric acid to oxidize all Fe2+ to Fe3+. This would result in an acidic iron solution with only Fe 3+ ions.
3Fe2+ + NO3- + 4H+ 3Fe3+ + NO + 2H2O
I then added excess sodium hydroxide solution to the above resulting solution in order to precipitate iron(III) hydroxide. This is due to the precipitation reaction which occurs between the sodium hydroxide solution with iron(III) nitrate solution. The excess sodium hydroxide was added to neutralize the acidic solution caused by the addition of nitric acid before the reaction could occur.
Fe3+ (aq) +3OH-(aq)
Fe(OH)3(s)
The precipitated Fe(OH) 3 was then be filtered using filter paper with a vacuum flask. The iron content could then be determined after drying the filter paper in an oven.
Figure 5 - Iron solution with reddish brown iron precipitation after step 3. The bottle on the far left is concentrated nitric acid.
Experimental procedure:
1.
Using a magnetic stirrer and a beaker, dissolve 5 iron tablets into 200 mL of 2M sulphuric acid (H2SO4) to fully dissolve the iron tablets.
2. Wearing gloves, safety googles and and a lab coat, add 100 mL of concentrated concentrated nitric acid (16M) into the iron solution in a fume cupboard. 3. Add 300 mL of 5M sodium hydroxide solution solution (NaOH) to neutralize the excess acid used and precipitate the iron in the solution. 4. Weigh a filter paper paper and put it on top of a vacuum vacuum flask. 5. Filter the precipitate ( Fe(OH)3) with the vacuum flask 6. Repeat step 5 until all visible Fe(OH)3 have been filtered. 7. Take out the filter paper and put in oven to dry and set to 60 degrees. 8. Weight the filter paper every 30 minutes minutes until the weight is constant. constant.
Figure 1 - The solution s olution with iron precipitate. Note the dark blue pH paper which shows the solution to be alkaline (~10-12 pH). An alkaline solution was needed in order for the reactino to occur.
Figure 7 - A vacuum flask was used for quicker filtering. The process was repeated several times with the “filtered” solution.
Analysis For each method, I have performed five trials and 3 trials respectively. I will then mean out the values for my calculations. calculations. Uncertainties Uncertainties will also also be taken into into account.
Results of Method 1 (Redox Titration) Volume of MnO4- to reach reach endpoint (cm 3) 7.9 ± 0.1 8.2 ± 0.1 8.6 ± 0.1 7.7 ± 0.1 8.4 ± 0.1
Trial Number 1 2 3 4 5 Mean of all trials + standard deviation uncertainty
8.2 ± 0.4
Table 1 - Results of Method 1
The uncertainty of each volume calculated is ±0.1 since the uncertainties of two readings, each of ±0.05, had to be added after their subtraction. I will then calculate the iron content there is per iron tablet. No. of moles of MnO 4- in the 0.1 M solution:
.
× 0.1 = 0.00082
No. of moles in 25 cm 3 of Fe2+ from 1:5 mole ratio from MnO 4- : 0.00082 x 5 = 0.0041 No. of moles in 250 cm 3 solution: 0.0041 x 10 = 0.041 Mass of Fe2+: 0.041 x 55.85 (molar mass of FeSO 4) = 2.28985 g Mass of Fe2+ per tablet: 2.28985 ÷ 5 = 0.45797 g Uncertainty calculation: Percentage uncertainty for the volume of MnO 4- needed to reach endpoint: . .
× 100% = 4.89% = ± 5%
Percentage uncertainty for volumetric flask for 25 mL of iron solution transferred (± 0.24): .
× 100% 100% = ± 0.96% = ± 1%
Absolute uncertainty uncertainty for iron content: content: 0.45797 x (5%+1%) = ± 0.03 Therefore the mass of Fe2+ iron in one tablet is 0.46 ± 0.03 g.
Results of Method 2 (Precipitation method) Trial Number 1 2 3
Mass of Fe(OH) 3 per pill (g) 0.973± 0.002 0.865 ± 0.002 1.038 ± 0.002
Mean of all trials + standard deviation uncertainty
0.959 ± 0.09
Table 2 - Results of Method 2
All the values in Table 2 were divided by five since five pills were dissolved in sulphuric sulphuric acid. The uncertainty of the Mass of Fe(OH)3 is ± 0.002 since it combines the uncertainties of ± 0.001 the digital balance has when measuring the weight of the filter paper and the measurement of the mass of filter paper plus Fe(OH)3 . After drying the filter filter paper in the oven, oven, the filter paper paper was weighted every thirty minutes until the weight was constant. Below are two pictures of the filter paper after 30 minutes and when the filter paper was completely dried. The value from Table 2 is a subtraction of the mass of the initial filter paper from the mass of the filter paper plus Fe(OH) 3 .
Figure 8 - Filter paper, 30 minutes in drying in Trial 3
Figure 9 - Filter paper, completely dried in Trial 3
Averaging the values values from Trial 1 to 3, the amount of Fe(OH)3 per pill is 0.959 ± 0.002 g. To determine the iron content, this valued is multiplied with the percentage by mass of iron in Fe(OH)3 . Percentage by mass of Fe in Fe(OH)3 :
55.845 = 55.845+(15.9994+1.00794)×3
52.2565 %
Mass of Fe in Fe(OH) 3 : 0.959 x 52.2565% = 0.501 g Uncertainty calculation: 0.009 0.959
× 100% = ± 9%
Absolute uncertainty: uncertainty: 0.501 x 0.2% 0.2% = 0.004509
Therefore the mass of Fe in one iron tablet using the precipitation precipitation method is 0.501 ± 0.005 g
Conclusion and evaluation The initial aims of the investigation have been achieved. I have attempted to determine the iron content in one iron tablet of ferrous sulphate, produced by Apt pharmaceuticals, as prescribed to my grandmother. Due to the consideration of the presence of Fe 3+ ions, as stated in the methodology section, a complementary experiment using precipitation was performed in addition to the redox titration. My calculated results show that the iron content by Method 1 and Method 2 are 0.46 ± 0.03 and 0.501 ± 0.005 g respectively. The calculated iron content determined by method 2 is larger than method 1, thus justifying my assumption that Fe3+ ions were present as Method 2 measures the presence of all Fe 2+ and Fe3+ ions while Method 1 only takes into the account of Fe 2+ ions only. However, conservatively, I would still conclude conclude the range of my calculated iron content as between 0.43 to to 0.506 mg. My calculated value is between 9~17 mg off from the theoretical value of 60mg as stated on the label of the capsule bottle. Yet, I would not be confident to conclude that the pharmaceutical company is providing false information to my grandmother’s doctor as I am aware of the many potential errors and limitations my investigation has.
Evaluation and limitations I will list the potential errors specific to each experiment (Method 1 & 2) followed by overall limitations related to both experiments.
Me th od 1 1.
Potassium permanganate solution has a colour of deep purple which would affect the reading taken. The burette has marks that are in dark brown and not being able to clearly take readings may affect the accuracy of my stoichiometric calculations. To reduce this systematic error, I put a white tile behind the burette glass when I take my readings. This was to create a plain and clear background to facilitate a more accurate reading. Despite this measure, there may still be an effect of the accuracy of my resulting calculations. Ultimately, this random error can be reduced by repeating more trials and taking an average.
2. The tablets I bought have a thin foil of pink on the outside. When dissolved into water and sulphuric acid, there is a light pink suspension on the iron solution. On the other hand, the end point of the redox titration with potassium permanganate is pink also. Despite the very distinct difference between these two types of pink and the fact that the pink would fade a little as time went on, this might be a contribution towards potential inaccuracies in interpreting the endpoint reached. To improve, again, more trials could
have been done. Had it not been my personal initiative to investigate in my grandmother’s pill, I could have also changed the type of iron tablet investigated.
Me th od 2: 1.
The small size of the Fe(OH)3 being filtered and the wet medium resulted in some of it sticking on the sides of the filtering funnel as seen in Photo 10. I had to wash the funnel in order to retrieve the residue, a process I repeated many times. In addition, the small size of Fe(OH)3 also meant I had to repeatedly filter the precipitate solution in order to get as much residue possible. However, there may still be traces of Fe(OH) 3 unfiltered or stuck in the filtering funnel which would affect my calculations. To improve, I could have used finer filter paper in order to get more iron residue as more Fe(OH) 3 would be filtered. filtered.
Figure 10 - Filtering funnel with traces of Fe(OH)3
2. There also may have been other impurities and substances that could have adhered to the filtered Fe(OH)3, possibly making my result larger than what it should be. To improve, I could have used ice cold water to wash the filtered residue so that other substances would dissolve in the water and thus be washed away. The insoluble Fe(OH)3 would stay due to its insolubility. This would increase my accuracy in terms of the Fe(OH)3 measured. 3. Despite taking a reading of the weight of filter paper only when the weight has been constant after continuously drying it in an oven, there may still be inaccuracies in my measurement. Since I could not put the filter paper on the digital balance while it is still hot straight from the oven, the process of leaving it to cold down may have increased the weight it has. Water vapor from air may have added on to the weight of the filter paper during the wait. This would make my measurement larger due to an increased weight. To improve, the paper could have been put in a desiccator to stay dry while it cools down.
Me th od s 1 an d 2 Overall the presence of other chemicals may also have affected my experimental results. From photo 7, the precipitated solution with Fe(OH) 3 is in yellow despite despite being in an alkaline medium medium of around 10 to 12 pH, meaning meaning excess NaOH has been added to to precipitate the solution. As a matter of interest, I kept on adding NaOH yet could see no change. Therefore, I assume the presence of other chemicals may have affected the precipitation of Fe3+, a potential explanation of the yellowish colour. This may, again, affect my calculation results as Fe 3+ in all pills are not fully precipitated and may have adhered with other impurities. To improve, a thorough investigation regarding the chemical composition of the iron tablets could have been done followed by a manipulation experimental design in accordance to the other substances present.
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Bibliography "Experiment 16 Help!!!" Web. 25 Dec. 2015. . 6.htm>. "Iron (III) Hydroxide Precipitate." Web. 25 Dec. 2015. .
Illustrations All photos and illustrations illustrations were taken taken and produced by myself.
