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  • Level: GCSE
  • Subject: Science
  • Document length: 20005 words

How much Iron (II) in 100 grams of Spinach Oleracea?

Extracts from this document...

Introduction

How much Iron (II) in 100 grams of Spinach Oleracea? Spinach Oleracea Name : Jade Taylor Candidate Number : 3689 Centre Number : 58203 Year of Entry : 2004 Contents Plan Aim To found out how much Iron (II) is present in 100 grams of Spinach Oleracea. The factors that I am going to investigated in this experiment include finding the best method to determine the concentration of an Iron (II) Ammonium Sulphate (aq) by trying colorimetry, an electrochemical cells experiment and a redox titration with Potassium Manganate (aq). After this I will extract Iron (II) from Spinach Oleracea using various methods (i.e. boiling the Spinach Oleracea for a range of times in different solutions) and use this spinach extract solution to determine the volume of iron extracted. I will take into account the presences of Oxalate ions and change my experiment accordingly (i.e. heating the spinach extract solution before titration's). Introduction It is important to know the Iron (II) content in 100 grams of Spinach Oleracea, as this allows people to calculate how much Spinach Oleracea needs to be eaten in order to obtain the Recommended Daily Allowance (RDA) of Iron (II). As shown in the table below the Recommended Daily Allowance for Iron (II) varies with age and sex (2) The Recommended Daily Allowance of Iron (II) (2) Age Amount (mg) Youth 1-3 6 4-12 8 Males 13-18 10 19+ 8 Females 13-49 14 50+ 7 Pregnant Females Second Trimester 18 Third Trimester 23 As the table shows, the volume of Iron (II) needed by the human body increases with age; woman generally require more Iron (II) than men as this is needed to replenish the Iron (II) that is lost during menstruation. As different age groups and sexes require different volumes of Iron (II) they will need to consume different volumes Spinach Oleracea (2). It is important to obtain the correct amount of Iron (II) ...read more.

Middle

Finish (cm3) Titre (cm3) 1.25 21.10 19.85 21.10 40.80 19.70 10.20 30.00 19.80 30.00 49.95 19.95 7.50 28.10 20.60 28.10 48.30 20.20 26.00 45.90 19.90 Average titre = 19.85 + 19.70 +19.80 + 19.95 +19.90 5 = 19.84 cm3 Experiment Two - How accurate is a Calibration Graph in determining the concentration of Iron (II) Ammonium Sulphate (aq) Experiment A - I was unable to produce data for this experiment, as the concentrations I had proposed to use were too concentrated to be able to dissolve the solid fully in a standard volume of Sulphuric Acid (aq). I was able to dissolve them by heating the Iron (II) Ammonium Sulphate (s) with Sulphuric Acid (aq) in a beaker over a Bunsen burner but as the solution cooled, the Iron (II) Ammonium solidified and separated from the acid. Experiment B - Concentration (mol dm-3) 1.0 0.8 0.6 0.4 0.2 Absorbance at 440 ?/nm -0.3 -0.3 -0.3 -0.3 -0.4 Absorbance at 470 ?/nm 0.1 0.01 0.0 0.0 0.0 Absorbance at 490 ? /nm -0.22 -0.04 -0.03 -0.03 -0.05 Absorbance at 520 ? /nm -0.07 0.0 -0.02 0.0 -0.01 Absorbance at 550 ? /nm -0.15 0.01 -0.04 -0.03 -0.04 Absorbance at 580 ? /nm -0.37 0.1 0.0 0.0 0.0 Absorbance at 680 ? /nm -0.08 0.03 0.1 0.3 -0.3 Experiment Three - How accurate is an Electrochemical Cell in determining the concentration of Iron (II) Ammonium Sulphate (aq) Unfortunately I was unable to gain data for this experiment as I ran out of time before I was able to complete it. The results should have shown that as the concentration of Iron (II) Ammonium Sulphate (aq) increased the voltage being passed between the solutions would have increased. This would have created a graph showing that an increase in concentration produces an increase in the voltage produced by the Electrochemical Cell. From this I would have been able to work out the concentration of Iron (II) ...read more.

Conclusion

remained in the beaker. This experiment required me to heat the spinach extract solution to 70oc before using it in the titration. Once the solution was being used in the titration, its temperature quickly dropped. For this reason errors may have occurred. It was difficult to judge the end point of the reaction even though a dramatic colour change took place. This was because a single extra drop of Potassium Manganate (VII) (aq) turned the solution a deeper shade of purple. Also the colour faded very quickly after the Potassium Manganate (VII) (aq) had been added. Even after the end-point had been reached (i.e. the solution remained coloured for 30 seconds) the colour continued to fade so that there was nothing to compare the solutions colour too, to help with determining the end point of the reaction. If the end point had been over shot, the average titre point would be higher; therefore the concentration of Oxalic Acid would appear higher suggesting that the solution was more concentrated. If the end point had not been reached, (i.e. the solution did not appear as dark in colour as the others at the end point) the average titre point would be lower; therefore the concentration of Oxalic Acid would appear lower suggesting that the solution was less concentrated. If either of these had occurred the redox titration would have appear less accurate. The whole solid may not have dissolved fully in the acid/distilled water needed to create the correct concentration of solutions in both the Oxalic Acid (aq) and the Potassium Manganate (VII) (aq). If this had occurred the solutions produced would have been less concentrated resulting in any equations done using their concentration being inaccurate. It was difficult to get all the solute and washings into the volumetric flask without it over shooting the graduation mark, this occurred as the solutions were being made into too concentrated solutions. This over shooting took up much time as every time this occurred a new solution had to be made up. Improvements Jade Taylor 3689 1 Jade Taylor 3689 51 Jade Taylor 87 ...read more.

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