Sciencemadness Discussion Board

Testing for the molarity of a solution.

CHRIS25 - 8-5-2012 at 07:07

In this case - copper chloride which I have prepared. The only way I can tell when the copper chloride has reached a saturation point is when it turns to mud brown, but I am wondering if there is a way of chemically/scientifically testing in order to maintain a constant molarity strength, by regenerating with either the oxidising agent or HCl as required.

I tried on google but maybe my search terms are too braod, if you know a place where I can read about it then I am just as happy. Thankyou.

smaerd - 8-5-2012 at 07:34

I'm not sure I completely understand what you're trying to do here. It might help to give more explicit details. A redox titration?

Copper(I) Chloride or Copper(II) Chloride?

I do think it's funny that, "The only way I can tell when the copper chloride has reached a saturation point is when it turns to mud brown", while your signature says "Colour is not an inherent property of matter or substance, its existence is purely subjective..." hehehe.

CHRIS25 - 8-5-2012 at 07:51

Quote: Originally posted by smaerd  
I'm not sure I completely understand what you're trying to do here. It might help to give more explicit details. A redox titration?

Copper(I) Chloride or Copper(II) Chloride?

I do think it's funny that, "The only way I can tell when the copper chloride has reached a saturation point is when it turns to mud brown", while your signature says "Colour is not an inherent property of matter or substance, its existence is purely subjective..." hehehe.


That's easy, it's muddy brown to me but it might appear reddish brown to you depnding on what angle you are looking at it from, plus the colour teperature at that time of day - let's just use the word "Dark" :)
The following is the whole process it's a cyclic event self perpetuating as follows:
Cu + H2O2 + 2HCl = 2H2O + CuCl2
Then
Cu + CuCl2 = 2CuCl
Then
2CuCl + H2O2 + 2HCL = 2H2O + 2CuCl2

It starts off very pale blue slowly going to pale green then dark green and then darker and then much darker - this is the saturation point at which time it must be re-generated in order to keep the reaction going.... I just thought (erroneously) that there might be a way of determining the strength/molarity of the solution by some method other than relying on colour changes, no matter how subjective:)

[Edited on 8-5-2012 by CHRIS25]

Hexavalent - 8-5-2012 at 09:27

You could try and work out molarity in a few potential ways; looking for a reaction where CuCl2 in solution is a reactant, write a balanced equation for the reaction, use a known molar amount of the other reactant and then continually add the solution to it in 0.1ml increments until the reaction is complete. . .then work out the concentration by looking at the total volume dispensed, and thus the amount of CuCl2 you must have had in your aliquot for the reaction to be complete.

Another way would be to evaporate the crystals down to the dihydrate, then prepare a standard solution using these crystals. Then you will know the precise concentration of your solution.

A third way, albeit much less practical for the amateur, would be to use an analytical technique such as spectrophotometry or something to find the concentration.

Fourthly, although I can't be sure about CuCl2, many salts when dissolved in a solvent will have a particular pH at a specific concentration. You can then measure the actual pH, ideally with a lab pH meter, and then compare it against these published values for particular concentrations and work out your molarity from that. It would be the work of a few minutes on the internet to find if this is can be done for CuCl2, but those should be your few minutes, not mine. However, I wouldn't expect brilliant results from this method, although if you can try the technique out on something at some point it can be a good experience and the knowledge of the calculations could be very useful for you at some point in your chemical career.

Hex

[Edited on 8-5-2012 by Hexavalent]

Magpie - 8-5-2012 at 10:23

Finding a solution to one's problems that is easy, simple, and cheap is not always possible. Here's a method that is easy and simple, but not cheap.

http://www.outreach.canterbury.ac.nz/chemistry/chloride_mohr...

watson.fawkes - 8-5-2012 at 10:27

There's a very good site on this very system called "Etching with Air Regenerated Acid Cupric Chloride", by Adam Seychell. The original site appears to be down, but it's been copied extensively. There are titration instructions there for determining concentrations, as well as a handy photograph of the various colors, sampled at different concentrations.

CHRIS25 - 8-5-2012 at 13:14

Hi Magpie - Mmm...I need my silver for the silver nitrate....

Hexavalent - Plenty to get a log with there, thankyou for the pointers, it's often the pointers in the right direction that I am not so good at finding when it gets a little more complicated, so Thankyou.

Watson.Fawkes, yes a good site, I found it, confirming my suspicions too, but plenty of info there.

Thankyou guys for your excellent help.