Chemical Forums

Please login or register.

Login with username, password and session length

Sponsored links

Pages: [1]   Go Down

Author Topic: Sodium Thiosulphate/Iodine Titration Molarity Issue  (Read 16837 times)

0 Members and 1 Guest are viewing this topic.

josharoon

  • New Member
  • **
  • Mole Snacks: +1/-0
  • Offline Offline
  • Posts: 8
Sodium Thiosulphate/Iodine Titration Molarity Issue
« on: October 27, 2010, 11:11:35 PM »

Hi all,
I'm currently in year 11 chemistry, and I've recently been doing a back titration method of determining the ethanol levels in wine, and have ran into some issues with standardizing a solution of Sodium Thiosulphate (Na2S2O3) against an Iodide (I-) solution, so problematic in fact, that not even my chemistry teacher has any idea on what may be occurring and is considering calling up the local university for help etc. However, he said that if I were able to research, hear of, find out (etc etc) an answer to this problem that he would forget about the rest of the titration and would grade me an A no questions asked.
From my calculations, which should be right, I have made a thiosulphate solution 0.1M by adding 6.20g to 250ml of water (the thiosulphate's molar mass is 248.18g/mol)
248.18/10 = 24.818
24.818/4 = 6.2045 ~ 6.20g

So my calculations of molarity should be fine, however when we ran into some errors during the wine titration, my chem teacher decided to check the molarity of the thiosulphate against a standardised solution of iodine which had been derived from the reaction between Potassium Iodate and Iodide and found that my first batch of thiosulphate was 0.2M and my second batch was 0.176M. I was not with him when he did this standardizing as he did this after that lesson, however he did give me the method of which he standardized the solution:
KIO3-+5I- +6H+  :rarrow: 3I2 + 3H2O
I2 + 2S2O32- :rarrow: 2I- + S4O62-
The solution of the KIO3 was 5*10-3M with 10ml used, a total of 5*10-5 moles of Iodate used and 2.5*10-4 moles of Iodide, creating 1.5*10-4 moles of Iodine and 1.5*10-4 of water.
He then diluted a sample of the sodium thiosulphate solution I hade made up by a factor of 1:5, so it should now have a concentration of 0.02M and used 8.4ml of it with the Iodine until the endpoint was found.

So from here I'm somewhat confused...
I know that I made my solution 0.1M, however according to his calculations my solution is out by 50% on the most leanient percentage errors and taking into account of water of crystalisation, but if theres anything that is wrong with those calculations, does have any suggestions on why my solution is out?
Many Thanks to whoever is able to help me with this problem, as it is an absolute pain in the backside.  :)
Josharoon
Logged

Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • **
  • Mole Snacks: +1333/-364
  • Offline Offline
  • Gender: Male
  • Posts: 20458
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #1 on: October 28, 2010, 02:10:39 AM »

From your description it is not clear what was the concentration of iodide solution prepared by your teacher, so it is hard to tell what may have happen. No idea why he used exact number of moles of iodide, correct procedure is to use excess, this ensures iodate is the limiting reagent and increases solubility of iodine.

Compare:

http://www.titrations.info/iodometric-titration-standardization

(scroll down, standardization with use of iodate is listed as second method).

Common problem with thiosulfate solutions is that their concentration is lower than expected, as thiosulfate easily decomposes.
Logged
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

MRO

  • Very New Member
  • *
  • Mole Snacks: +0/-0
  • Offline Offline
  • Posts: 1
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #2 on: October 28, 2010, 12:10:07 PM »

If you look at the moles of iodide ions added you will see that they are in excess. Josh was merely demonstrating his understanding of the stoichiometric relationship to the IO3- ions of the standard present.
The issue is all the calculations, weighings and volumes recorded were correct, I even checked the calibration of the balance and ensured that the iodate primary standard used was dry anhydrous, so why does the caslculated concentration come in at 50% stronger?
Weaker I can understand and explain due to possible deterioration of the thiosulfate, although a fresh batch was used and freshly made up.
So I too would like to know where the discrepency could lie.
Logged

josharoon

  • New Member
  • **
  • Mole Snacks: +1/-0
  • Offline Offline
  • Posts: 8
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #3 on: October 28, 2010, 08:18:58 PM »

Does anyone think it could be possible that the Sodium Thiosulphate, which is pentahydrate, could have dehydrated so that it returned to its non-hydrated form with a molar mass of 158.108g/mol, therefore increasing the moles of actual thiosulphate per gram, hence creating a stronger solution.
That means that a required 3.95g would be placed in 250ml to make a 0.1M solution.
So assuming the original amount of thiosulphate being 6.20g was added when its actual molar mass was 158.108g/mol, a 0.157M solution would be created, this being ~10% from that of my teachers calculations from his titration against the iodine.
Josharoon
Logged

Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • **
  • Mole Snacks: +1333/-364
  • Offline Offline
  • Gender: Male
  • Posts: 20458
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #4 on: October 28, 2010, 09:14:07 PM »

If you look at the moles of iodide ions added you will see that they are in excess. Josh was merely demonstrating his understanding of the stoichiometric relationship to the IO3- ions of the standard present.

No, if I look at the moles of iodide

a total of 5*10-5 moles of Iodate used and 2.5*10-4 moles of Iodide

they are given as stoichiometric. I understand that it is not necessarily procedure that was followed, but we have no choice but to rely on the information supplied.

Quote
The issue is all the calculations, weighings and volumes recorded were correct

Agreed when it comes to calculations, I have checked them in several ways (mostly using my programs - concentration and stoichiometry calculator, see signature - they really save a lot of time). That's not to suggest there is something wrong with the procedure, but as you have seen above information we deal with is not always perfect.

Does anyone think it could be possible that the Sodium Thiosulphate, which is pentahydrate, could have dehydrated

I have checked the idea that it is about anhydrous thiosulfate yesterday (in CASC checking it is a one click), but it doesn't explain the difference.

According to my handbook thiosulfate exists either as anhydrous or pentahydrate, so depending on the way it was stored exact composition can change (that's one of the reasons why hydrates are poor candidates for standard substances). But first - it is very unlikely that it dehydrated completely, and second - even if, you would be still far from being able to prepare 0.2M solution.

I will forward the question to CHEMEDL.

Edit: have you tried to repeat the solution preparation procedure? This will help eliminate human error during weighting/dissolving.
Logged
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

josharoon

  • New Member
  • **
  • Mole Snacks: +1/-0
  • Offline Offline
  • Posts: 8
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #5 on: October 28, 2010, 10:27:19 PM »

Edit: have you tried to repeat the solution preparation procedure? This will help eliminate human error during weighting/dissolving.
Yep, Mr.O's first thought was that my measurements must have been incorrect, as his first calculated value for the molarity was 0.2M, however when we repeated it the molarity was assumed to be 0.176M, so on both occasions they were close to double the desired molarity.

Also my brother who is a research scientist at a uni was talking to a chemist they have there, and apparently theres a possibility that the acid added to the iodate and iodide, would be remaining and can react with thiosulphate, however I'm not so sure how this would work, as I would have assumed this would be in an acid/base titration.
Logged

Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • **
  • Mole Snacks: +1333/-364
  • Offline Offline
  • Gender: Male
  • Posts: 20458
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #6 on: October 30, 2010, 11:57:03 PM »

Any side reactions I could think off will consume thiosulfate, that would make its apparent concentration lower, not higher.

As told earlier, I have forwarded the question to CHEMEDL, several points were raised, but after all nobody had a better idea than those already discussed.

How many times have you prepared the solution - are we talking all the time about one, single sample, or have you tried to repeat whole procedure? I am asking because after all other problems have been eliminated, there is always a human factor left - and stupid mistakes happen, like weighting 23g instead of 32g, or using wrong bottle.

Many years ago during organic synthesis lab I was not able to do some reaction, even if I tried several times. It turned out the lab assistant each time gave me cyclohexanol instead of cyclohexanone (in Polish the names look even more similar, cykloheksanol and cykloheksanon). When asked by lab supervisor what she gave she said "cyclohexanone" and pointed to cyclohexanol bottle. She was absolutely sure she was OK. And she was not some greenhorn, she was there for many years and she was quite good at what she was doing.
Logged
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

josharoon

  • New Member
  • **
  • Mole Snacks: +1/-0
  • Offline Offline
  • Posts: 8
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #7 on: October 31, 2010, 12:36:59 AM »

The solution has been prepared twice, the original which was tested and gave the result of 0.2mol/L.
So I made up a new batch being extremely careful on the measurements that were added (6.20g in 250ml water)
I guess there could always be the possibility that the supposedly pentahydrate thiosulphate which was in the container had been topped up with anhydrous thiosulphate and the lab assistant may not have relabelled it, as I am sure that my measurements were correct.
I will speak with Mr.O about this tomorrow and hopefully will discover a solution to this issue.
Logged

josharoon

  • New Member
  • **
  • Mole Snacks: +1/-0
  • Offline Offline
  • Posts: 8
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #8 on: November 07, 2010, 01:21:40 AM »

Ok so after talking to Mr.O the following day I standardised the thiosulphate against the iodide myself and found that the solution was actually 0.089M, so there must have been some error within his calculations which led to the large differences. We assumed that the 11% difference in molarity would be due to the thiosulphate going bad over time since Mr.O's standardisation.
Logged

nimbus

  • New Member
  • **
  • Mole Snacks: +0/-1
  • Offline Offline
  • Posts: 5
Re: Sodium Thiosulphate/Iodine Titration Molarity Issue
« Reply #9 on: November 17, 2010, 07:50:47 AM »

Thiosulfate readily decomposes when exposed to light.
You can slow this process by storing it in a dark place, and wrapping your volumetric flask with tin foil when you have it on the lab bench. Proper analytical procedure is to standardize all solutions when titrating anyways, so the discrepancy shouldn't matter if you are standardizing it with an iodometric method. If that's not part of your lab procedure, then I'm not sure what you can do. You could explain to your teacher the mechanisms with which thiosulfate degrades and prove to him that you are not messing up the solution prep yourself.

When you prepare the thiosulfate solution it should be free of metal ions, and the water you use should have been boiled prior to making the solution, as CO2 and metal ions degrade thiosulfate. The pH must be near neutral to increase stability as well. You can add a tiny amount of sodium carbonate to achieve this. We do a backtitration to standardize the solution, by using KIO3 solution and KI solid, with 1M HCL. When the solution goes to pale yellow, we add starch indicator and titrate till the solution is clear (this is more accurate).

Hope this helps!
Logged
Pages: [1]   Go Up
 

Mitch Andre Garcia's Chemical Forums 2003-Present.

Page created in 0.096 seconds with 24 queries.