Practical investigations
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I am trying to find the exact concentration of a saturated aqueous solution of chlorine by back titration, and it has been suggested to use an Iron(II) compound as the reducing agent. However I am unsure as which suitable iron(II) compound to use (can I use any?), and how to determine its excess in only knowing the appoximate concentration of the chlorine (7g dm-3).
Jenkin replies:
The most convenient iron(II) compound is probably ammonium iron(II) sulphate, (NH4)2SO4.FeSO4.6H2O. You will need to add dilute sulphuric acid when you make up a solution to prevent hydrolysis and air-oxidation. You could titrate with standard potassium manganate(VII) (0.02 mol dm-3 is probably a convenient concentration) to determine the excess.
As to the concentration to use, your chlorine solution is about 0.1 mol dm-3, and 1 mole of chlorine reacts with 2 moles of Fe2+. so I would suggest you use a solution of about 0.3 to 0.4 mol dm-3 of the iron compound. But, of course, you must make it up accurately.
If I were doing this work, I would be inclined to use potassium iodide rather than iron(II), and titrate the liberated iodine with standard sodium thiosulphate solution.
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Risk assessment
Before attempting any practical work based on the advice and suggestions on this website, you must do the following. Identify any hazards, assess the risks from these hazards, and then decide appropriate control measures to reduce the risks. You must have these approved by those in authority in your school or college laboratory. Do not rely on what is said on this website.
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updated: 25 January 2006
