posted by kelly .
I'm not suppose to react Iron Chloride with Silver Nitrate in order to get Iron nitrate. If i'm suppose to keep one of the above ingredients and obtain a pure dry sample, what should i do?
I don't know the question. Why are you not supposed to react the two? and which ingredient do you want to keep? what pure dry product do you want to obtain? The silver in silver nitrate will precipitate silver chloride when reacted with iron nitrate. If you had a solution of the two, the white silver chloride, AgCl, will ppt and you can filter it out. The iron nitrate will stay in solution. Evaporation of the water will give you the iron nitrate crystals and it will be pure IF YOU ADDED JUST THE right AMOUNT OF AgNO3 to ppt all of the chloride with no excess AgNO3. You must add enough AgNO3 to ppt ALL of the chloride but not add excess AgNO3. I hope this helps. If I have guessed wrong about what you are trying to do, please repost and clarify exactly what you are looking for in an answer.
Sorry if i sounded confusing.
I have to obtain a pure dry product of Iron (II) Nitrate by modifying the experiment below:
Mixing Iron (II) Chloride with Silver Nitrate and filtering the AgCl and crystallising the remaining solution.
At the same time, I have to identify failures in the above experiment and keep either Iron (II) Chloride or Silver Nitrate and redesign this experiment to get Iron (II) Nitrate.
As both Iron (II)Nitrate and Iron (II) Chloride are soluble in water, we might obtain an impure sample if FeCl and AgNO3 do not fully react...
FeCl2 + AgNO3 ==> AgCl + Fe(NO3)2
I see. If you had too little AgNO3, then some FeCl2 will contaminate the Fe(NO3)2 and if you add too much AgNO3, then AgNO3 will contaminate the Fe(NO)3. In fact, that statement would be a good one to use for the failures. The ONLY way you can get pure Fe(NO3)2 this way is by adding just enough AgNO3 to react fully and completely with the last atom (more or less) of chloride from the FeCl2 without adding an excess of AgNO3. One way you can do that is to modify the experiment in such a way that you know when that exact amount of AgNO3 has been added. There are at least a couple of ways you might proceed. One way is to perform several "mini" experiments on a predetermined amount of FeCl2; for example, say 5 mL of your FeCl2 solution. Then very carefully add AgNO3, say 1 mL at a time, stir, let the ppt settle, then add another drop of AgNO3 and see if a ppt forms. If it does, you know you don't have enough AgNO3. Add another mL of AgNO3. Continue until you DON'T get a ppt when adding one more drop of AgNO3. That way you will bracket whether the FeCl2 requires 1, 2, 3, 4,or so mL. THEN take a new sample of FeCl2, add the required amount of AgNO3 (let's say you found 3 mL was not enough but 4 mL was too much--in this case add 3 mL). Then add a drop of AgNO3, stir, add another drop AgNO3 and observe if a ppt forms. Continue until the added drop does not produce AgCl. That way you will have no more than one drop too much AgNO3. I don't know if this will be pure enough for your experiment since I don't know all the rules but I suspect it will do. A second procedure is to add an indicator that will tell you when the right amount of AgNO3 has been added. One problem with that method is that then you have contamination from the indicator. I hope this is enough to get you started in the right direction.
You're quite welcome. Please come again.
vbpmqrhcy kedxf vfxa jyigzwl gwhftn oferaxbc hnsxwgj