Sciencemadness Discussion Board - How do I titrate Sodium Hypochlorite?

How do I titrate Sodium Hypochlorite?

Little_Ghost_again - 19-9-2014 at 02:58

Hi
I have some sodium hypochlorite that is meant to be 10% strength.
I have my doubts on this and would like to test it, it is also a good basic skill to learn.
I have limited acids etc, so I was wondering if I could titrate it with citric acid (lab grade)? If so how do I got about this?
What has me slightly confused is the fact that the solution is meant to be 10%........
So working on the assumption that the solution is in fact 10%, then if I find the MW of sodium hyperchlorite (77.44(Merck 13th ed) )I need to take a figure of 10% of this
(7.74g). So I weigh out 7.74g of the solution and add to a beaker, to this I add a couple of drops of indicator solution.
Now it gets a bit hazy for me.... In the burette I add a solution of citric acid?

My question is how strong is this solution? One part of me says it should be 1 mol, but the other side of me thinks it should be 0.1mol.

I am going to say 1 mol is this correct?

Assuming I am correct then I need to add 192.12g to 1Ltr water Or 19.21g to 100 ml of water.

So I then start to drip it into the hypochlorite solution. If the solution of hypochlorite is 10% strength then I should find I need to add 10ml of the acid to neutralize the hypochlorite.....
Is the above correct? It feels like I am wrong somewhere so I would greatly appreciate a little guidance on this

Many thanks guys
LG
Or LGA

forgottenpassword - 19-9-2014 at 03:28

Google it. This is a common school based practical, and as such there are many detailed procedures.

Amos - 19-9-2014 at 03:49

I'm not going to suggest this, as I don't know how reliable it might be, but could you analyze the density of an exact volume of the solution at a known temperature?

blogfast25 - 19-9-2014 at 04:01

Quote: Originally posted by Little_Ghost_again

Hi
I have some sodium hypochlorite that is meant to be 10% strength.
I have my doubts on this and would like to test it, it is also a good basic skill to learn.
I have limited acids etc, so I was wondering if I could titrate it with citric acid (lab grade)? If so how do I got about this?
What has me slightly confused is the fact that the solution is meant to be 10%........

Your best bet is to take advantage of it being a strong oxidiser.

For instance it will oxidise iodide to iodine very easily. By adding a small excess iodide, the hypochlorite will oxidise an equivalent amount of iodide to iodine (Isub>2</sub>

, which can then be titrated with sodium thiosulphate in a back titration.

Being such a strong oxidiser the hypochlorite may oxidise the iodide to a higher oxidation state like iodate, so that would have to be checked.

But as forgottenpassword said, off-the-shelf methods must be quite easy to find though: don't try and reinvent the wheel.

Zyklon-A - 19-9-2014 at 06:29

Add a known amount of hydrogen peroxide and measure the oxygen.
NaOCl + H₂O₂ = NaCl + H₂O + O₂

[Edited on 19-9-2014 by Zyklon-A]

Little_Ghost_again - 19-9-2014 at 08:02

Thanks for the suggestions. All really good, but I dont have any of the chemicals yet and wanted to do it tonight.
I have googled but was unsure of the maths/method, as I understand it I should get sodium Citrate at the end of it and this can used as a buffer

.

I just wanted to check my method as per above was correct.
Many thanks
LG

AJKOER - 21-9-2014 at 06:45

Quote: Originally posted by Zyklon-A

Add a known amount of hydrogen peroxide and measure the oxygen.
NaOCl + H₂O₂ = NaCl + H₂O + O₂

[Edited on 19-9-2014 by Zyklon-A]

I would suggest a mixture of NaClO with an excess of CaCl2 and CO2. The reactions proceeds as follows:

CaCl2 + 2 NaClO = Ca(ClO)2 + 2 NaCl

which further, in the presence of any added CO2, a bright white precipitate of CaCO3 is formed along with Hypochlorous moving the above reaction to the right:

Ca(ClO)2 + CO2 + H2O ---> CaCO3(s) + 2 HOCl

Then cool and pour out your solution of salt and HOCl (and freeze to store for other applications) and collect and dry the CaCO3.

Calculate the moles produced of CaCO3, and multiply by two to get the real strength of your NaOCl.

The real advantage of this approach over the H2O2 path is that you are not trading the accuracy of one label (on your Hydrogen peroxide bottle) for that on the bleach bottle. Unfortunately, I suspect the H2O2 has a poorer shelf life, especially after openning, than your bottle of bleach. In the chlorine bleach business in the last 20 years, there has emerged a good competition versus substiutes (like Sodium percarbonate) on both price and quality (effective strength in particular). Such a quality concern has not been the case in the H2O2 business, as 2.5% H2O2 looks the same (and feels better) applied to a wound, than 3%. Also, the quicker it stops bubbling at all when applied, the more likely one will buy a fresh bottle, implying that better quality works against sales and profits, in my opinion.

[Edited on 21-9-2014 by AJKOER]

Boffis - 21-9-2014 at 08:56

I do this all the time on my hypochlorite solutions immediately before I use the. I use the excess standard KI method and then titrate the liberated iodine with standard Thiosulphate solution. The procedure is that give in Vogel's "Textbook of Quantitative Analysis". This procedure is actually for bleaching powder but is easily modified for sodium hypochlorite solution. This textbook is available online and can be downloaded for free from bookfi.org but I have attached the relevant pages.

I also ran a google search and got 4 procedure from the first 6 hits, all of them look very basic and specific for hypochlorite bleach.

Attachment: Vogel hypochlorite analysis Textb Quant Anal p396-397.pdf (2.1MB)
This file has been downloaded 660 times

Little_Ghost_again - 21-9-2014 at 09:08

Quote: Originally posted by Boffis

I do this all the time on my hypochlorite solutions immediately before I use the. I use the excess standard KI method and then titrate the liberated iodine with standard Thiosulphate solution. The procedure is that give in Vogel's "Textbook of Quantitative Analysis". This procedure is actually for bleaching powder but is easily modified for sodium hypochlorite solution. This textbook is available online and can be downloaded for free from bookfi.org but I have attached the relevant pages.

I also ran a google search and got 4 procedure from the first 6 hits, all of them look very basic and specific for hypochlorite bleach.

thank you very much guys

Little_Ghost_again - 21-9-2014 at 09:17

I wasnt trying to reinvent the wheel or anything, I wanted a quick way to do it with the chemicals I had on hand.
Looks like I am going to have to get more chemicals!! I had ear marked the rest of my savings for chemicals but something came along................

blogfast25 - 21-9-2014 at 12:29

Quote: Originally posted by AJKOER

This is plain nonsense, as CaCl2 + 2 NaClO = Ca(ClO)2 + 2 NaCl doesn't occur. All species in that 'reaction' are water soluble.

On gassing with CO2 one would simply precipitate all calcium regardless of NaClO strength, PERIOD!

Again, why complicate things when iodometry works fine here?

jock88 - 21-9-2014 at 12:45

There is probably some conversation regarding Hypochlorite titration in the Hdyrazine thread.

Little_Ghost_again - 21-9-2014 at 13:08

Quote: Originally posted by jock88

There is probably some conversation regarding Hypochlorite titration in the Hdyrazine thread.

Ok thanks I will go look.

I have been reading up a fair bit, I found plenty on using Hydroxide to determine citric acid concentration but not much the other way around

.
The other wall I hit was STANDARD SOLUTION !. I had assumed (wrongly by the looks of it), that I could just weigh out some of my anhydrous citric acid add it to a carefully measured amount of distilled/RO water and off I go............
Reading more about it looks a bit more complicated, havnt yet got my head around why you need to add something else to the acid (Assumption is to make it stable).
Then it got into primary standard solutions and secondary standard.
But I guess the crux of it is just how accurate you need the result, in my case I am dealing with scales that at the very best has a resolution .01g, it also needs 50mg minimum before it registers. These scales go to 300 grams.
So straight away it would introduce an error to some degree anyway. But that isnt the point here, I need a reasonable guide on how concentrated some of my solutions are. I missed the chance of some analytical scales, these were from a lab closing down. Great condition very very accurate and only £40 (special price for me).
I turned it down thinking at the time I didnt have use for them and £40 is a huge amount of cash to me.
I still think for now what I have is fine for what I am doing. I have made a small amount of money go a very long way so far

.
Also I have finally opened the last couple of box's from the last job lot I got at auction. Yet more Burettes!!! All kinds in there, some though have no glass stop cock. I need to get some burettes on Ebay (or here) as I must have nearly 70 of them. some are ml some are cm3. Most look like good makes such as Griffin.
The other thing I am swimming in is pipettes! But I doubt they would fetch much on ebay.

Sorry I am waffling lol

aga - 21-9-2014 at 13:49

Something to get your head around early on is Molecular Mass (aka MM or Molecular Weight) and Mols.

MM can be taken directly from the atomic table, e.g. ptable.org and just add up the weights of each element in a compound.

E.g. NaCl

Na = 22.989
Cl = 35.45
so NaCl = 58.439g

That means Grammes per Mole.
A Mol is grammes per 6.022 * 10^23 units of the stuff, also called Avogadro's Number.

No idea why that number is significant (carbon madness) but a number was needed, and that's the number someone chose, so we use that one.

So 58.439 grammes of pure table salt has exactly 1 Mol of NaCl in it.

When you work out a Balanced Equation for a reaction, you can then work out how many Mols of each substance react with each other.

Once you know that, you can calculate (using the MMs) how many Grammes of each substance you need to fully react them together.

Sometimes in a reaction you need some extra H+ flying around, or OH- or other stuff, just to make it work.

That's generally why you have to add an acid or a base (or even a catalyst) to your mixture to get it to work at all, even though the equation with just the Reactants and Products looks fine.

Quick tests :-

1. What is the Molecular Mass of Water H2O ?
2. What is the MM of NaCl ?
3. How many Mols are there in 10g of NaCl ?

4a. NaOH + HCl ==> NaCl + H2O
To get 10g of NaCl from this reaction, how many grammes of NaOH and HCl would you need ?

4b. (slightly tricky) How many ml of water will be made ?

Edit:

Don't actually do that reaction.
It is seriously hot, violent and dangerous at all stages apart from the end bit.

[Edited on 21-9-2014 by aga]

What am i saying ?
If you Do do that reaction, first dissolve the NaOH in distilled water.
It will get very hot. Chill/Let it cool down.
Add the HCl drop by drop.
The mixture will get hot again. Chill/Let it cool down.

Obviously you'll be using all Protective Gear including eye protection and gloves, and observing all the usual Lab safety protocols, or you'd be an idiot.

If done badly (i.e. dump all the HCl in a pot of dry NaOH) it explodes due to the Steam created, and covers everything, including you, with a hot mix of Acid and Lye.

It's a useless reaction really : just turns two really useful reagents into saltwater.

[Edited on 21-9-2014 by aga]

blogfast25 - 22-9-2014 at 04:14

Quote: Originally posted by Little_Ghost_again

The other wall I hit was STANDARD SOLUTION !. I had assumed (wrongly by the looks of it), that I could just weigh out some of my anhydrous citric acid add it to a carefully measured amount of distilled/RO water and off I go............
Reading more about it looks a bit more complicated, havnt yet got my head around why you need to add something else to the acid (Assumption is to make it stable).
Then it got into primary standard solutions and secondary standard.
But I guess the crux of it is just how accurate you need the result, in my case I am dealing with scales that at the very best has a resolution .01g, it also needs 50mg minimum before it registers. These scales go to 300 grams.

Proper primary standard solutions prepared from standard materials with a 0.1 mg scale and calibrated volumetric flasks of course give the most accurate result.

But even with more modest means titrations can give good indications of the strength of a solution. It's just less accurate that way. If you're new to it, just try and work as accurately as you can with what you've got.

Your idea of titrating with citric acid is unlikely to work though because commercial hypochlorite solutions contain excess sodium hydroxide; you'd be titrating the SUM of NaOH and NaClO. Also, titrating with a weak acid (triprotic to boot!) isn't generally recommended.

[Edited on 22-9-2014 by blogfast25]

Little_Ghost_again - 22-9-2014 at 06:36

Quote: Originally posted by blogfast25

Quote: Originally posted by Little_Ghost_again

You have answered many questions in my mind there! Thank you so much. Also weak triprotic (A new term and I have no idea what it means) PLEASE DO NOT TELL ME. That is something I will need to look into.

You have me wondering now, when I was trying to make chloroform from IPA and sodium hypochlorite I got nowhere at first, I had no idea that commercial hypochlorite contained sodium hydroxide. I got no reaction with IPA until out of frustration and general experiment I added sodium hydroxide. So now I wondering if it was actually the hydroxide part in the hypochlorite that had degraded...... Is this possible or likely? I have tried google but its a kind of random search at the moment. So any opinion on this welcome.

Many thanks
LG

blogfast25 - 22-9-2014 at 11:28

Quote: Originally posted by Little_Ghost_again

I had no idea that commercial hypochlorite contained sodium hydroxide. I got no reaction with IPA until out of frustration and general experiment I added sodium hydroxide. So now I wondering if it was actually the hydroxide part in the hypochlorite that had degraded...... Is this possible or likely? I have tried google but its a kind of random search at the moment. So any opinion on this welcome.

Many thanks
LG

NaClO is really only stable (-ish) in strongly alkaline conditions.

For the chloroform reaction, there's plenty of decent threads on it here. Search and ye shall find!

Little_Ghost_again - 22-9-2014 at 12:27

Quote: Originally posted by blogfast25

Quote: Originally posted by Little_Ghost_again

NaClO is really only stable (-ish) in strongly alkaline conditions.

For the chloroform reaction, there's plenty of decent threads on it here. Search and ye shall find!

I did but most dont use IPA, It was all I had on hand, my concern at the time was the bleach, it had been opened and part used over a year ago. It was also stored badly, when I came to use it the normally very strong bleach smell wasnt there.
So in the end I added sodium Hydroxide more to see if it got hot due to the chorite having a lot of water in it.
As it turns out I did get Chloroform (a fair bit) and the reaction never went above 25C in a room with a temp of 15C. I got rid of the chloroform and kept the waste, this has now gone from cloudy to clear with three distinct layers! non of them are chloroform as I had removed all that.
I dont know what the three layers are yet and its been put in a cupboard out the way with a stopper on the flask that has a air bubble trap like for wine making (just in case!) One layer I think is excess IPA and probably very dry due to the amount of Hydroxide I added. I will get around to investigating it but felt now wasnt the time with my lack of knowledge!

JJay - 14-5-2016 at 16:20

I recently found some really cheap bleach and bought several bottles, and I have been thinking about doing some reactions that involve precise hypochlorite measurement. I have all of the chemicals and most of the equipment needed for doing the titration, but I was wondering: is there any reason I can't use sulfuric or hydrochloric acid instead of glacial acetic acid? I am looking in particular at this titration procedure: http://seniorchem.com/chlorine_thiosulfate_titration.pdf Vogel's also indicates glacial acetic acid.

blogfast25 - 14-5-2016 at 17:15

Quote: Originally posted by JJay

I recently found some really cheap bleach and bought several bottles, and I have been thinking about doing some reactions that involve precise hypochlorite measurement. I have all of the chemicals and most of the equipment needed for doing the titration, but I was wondering: is there any reason I can't use sulfuric or hydrochloric acid instead of glacial acetic acid? I am looking in particular at this titration procedure: http://seniorchem.com/chlorine_thiosulfate_titration.pdf Vogel's also indicates glacial acetic acid.

It's a good question.

HCl plus bleach means of course more of:

2 H⁺ + ClO^- + Cl^- ⇌ Cl₂ + H₂O

But I can't see how sulphuric acid could interfere here. And I have read your link, BTW.

JJay - 14-5-2016 at 18:25

Quote: Originally posted by blogfast25

Quote: Originally posted by JJay

It's a good question.

HCl plus bleach means of course more of:

2 H⁺ + ClO^- + Cl^- ⇌ Cl₂ + H₂O

But I can't see how sulphuric acid could interfere here. And I have read your link, BTW.

I would think the chlorine would very quickly oxidize the iodide to iodine, so I don't think that's a problem. Perhaps it would be a good idea to add the bleach dropwise after the iodide... that might not be a good idea with highly concentrated acid... with my planned experimental procedure, I don't think the order will really matter much, but if I see flying acid or smell chlorine gas, I'll know differently. :cool:

I have read that sulfuric acid can actually oxidize iodide, but I think that generally only happens at high concentrations, and I'm just not seeing any advantage to using highly concentrated acid... for that matter, at high concentrations, hydrogen iodide can oxidize iodide as well. I can't buy acetic acid here, so mine is made rather laboriously by distillation from sulfuric acid and anhydrous sodium acetate.... I'd rather use something I can buy at the grocery store.

Boffis - 15-5-2016 at 04:59

The reason for using acetic acid is that mineral acids break down the standard thiosulphate solution liberating sulphur. However, this reaction is fairly slow at the acidity required for the this titration so if you are using the excess iodide and back titration against standard thiosulphate method the liberated iodine rapidly converts the thiosulphate to dithionate with minimal loss. Keeping the solution cold and dilute greatly reduces the problem. I generally use hydrochloric acid but I don't see why you couldn't use sulphuric acid too, the reaction mixture does not need to be strongly acid.

If you are using the thiosulphate + bleach and then back titrating the excess thiosulphate against standard iodine solution it is much more of a problem unless you work very quickly. I recommend the former method, I use it routinely on all hypochlorite solutions immediately before use.

JJay - 19-5-2016 at 23:19

Hmm.. let's see here... 40 mL of GAA in 1L of water... I could probably get away with using commercial vinegar.

AJKOER - 7-6-2016 at 12:19

Quote: Originally posted by Zyklon-A

Add a known amount of hydrogen peroxide and measure the oxygen.
NaOCl + H₂O₂ = NaCl + H₂O + O₂

[Edited on 19-9-2014 by Zyklon-A]

On second thought, the most tangible low cost route to employ the above path is to measure the strength of the H2O2 first by adding a very small amount of iodide and recording the volume of liberated oxygen. Chemistry:

H2O2 (aq) + I- (aq) = H2O (l) + OI- (aq)
H2O2 (aq) + OI- (aq) = H2O (l) + O2 (g) + I- (aq)

Net: 2 H2O2 (aq) = 2 H2O (l) + O2 (g)

Reference: http://cldfacility.rutgers.edu/content/catalytic-decompositi...

The only side reaction that may under estimate the strength of the H2O2 that I can envision is the possible disproportionation of some of the OI- to IO3-. The latter reaction appears to be slow in basic solutions (see, for example, http://www.nrcresearchpress.com/doi/abs/10.1139/v86-375 ).

[Edited on 7-6-2016 by AJKOER]