Sciencemadness Discussion Board - Titrating a HCl in acetic acid solution

Titrating a HCl in acetic acid solution

CrimpJiggler - 13-3-2012 at 08:49

If you wanted to determine the amount of HCl present in a solution of HCl in acetic acid, obviously you couldn't titrate with a strong base like NaOH because the base would react with the acetic acid. If you were to use sodium bicarbonate as the titrant instead, would I be right in assuming that it will only neutralize the HCl, not the CH3COOH?

ScienceSquirrel - 13-3-2012 at 08:51

I suspect the only way to do it is with a pH meter.

bbartlog - 13-3-2012 at 09:42

Sodium bicarbonate will react with (aqueous) acetic acid, so that wouldn't work unless you're talking HCl in glacial acetic acid. More generally, unless you know how much acetic acid is present (do you?), how will you know when you've reached the endpoint?
An old school way to do this might be to add lead(II) acetate solution in excess and measure the weight of the precipitated PbCl2. Of course that would involve filtering, drying and so on. It has the advantage that you don't need to know how much acetic acid you've got.
If you do know how much acetic acid is present in your solution and you have a decent way of measuring pH, then using NaOH or Na2CO3 (not NaHCO3) and titrating until you reach the (calculated) pH that just the acetic acid should give seems like a good approach.

UnintentionalChaos - 13-3-2012 at 09:42

Quote: Originally posted by CrimpJiggler

If you were to use sodium bicarbonate as the titrant instead, would I be right in assuming that it will only neutralize the HCl, not the CH3COOH?

Ever seen a baking soda and vinegar volcano?

Dilute a sample with water. Titrate with sodium acetate and a pH meter. The pH curve will flatten dramatically after you consume the HCl and have nothing but a solution of NaCl and AcOH/NaOAc buffer.

blogfast25 - 13-3-2012 at 10:44

Assuming the HCl to be present in significant quantity, it's perfectly possible to titrate both (weak and strong) acids separately with a strong base like NaOH, in one single titration. The presence of a strong acid in a solution of weak acid almost completely suppresses the already weak dissociation of the weaker acid and thus two distinct end points are observed.

For HCl:

HCl(aq) + H2O(l) === > H3O+ (aq) + Cl- (aq) . Dissociation is almost 100 %

For acetic acid (HAc):

HAc(aq) + H2O(l) < === > H3O+ (aq) + Ac- (aq). Equilibrium pushed almost completely to the left due to presence of H3O+ 'from the HCl'. Nearly 100 % of the acetic acid present as undissociated HAc.

The first endpoint will be about at the pH of the HAc solution (as if no HCl was present).

This also explains why multiprotic acids like H3PO4 can be titrated for three different end-points, sufficiently distant from each other to be discerned easily.

A combination of a strong and a weak acid can thus be titrated in one single titration with a strong alkali, obtaining two distict end points. It's best to measure a pH v. volume titrated master curve, that way the most suitable indicators can then be chosen, based on the observed end points.

[Edited on 13-3-2012 by blogfast25]

Nicodem - 13-3-2012 at 13:00

I think aqueous dilution of the sample, followed by titration with sodium or potassium acetate would give the end point with least interference. However, this or any other titration would not be the direct evidence of the HCl concentration in your acetic acid samples. It would only give you the combined concentrations of all comparably strong acids. You would need to couple acid titration with the chloride determination (e.g., gravimetric measurement via AgCl or similar) to obtain the HCl concentration.
The concentration of HCl can then be declared to be the measurement which gives the lower value, either the strong acid titration or the chloride concentration. This is because you can have two situations, either c(Cl-

> c(H+

or c(H+

> c(Cl-

(where H+ represents all strong acids indistinguishable by titration). In any of these situations, the concentration of HCl equals the lower value.
You would also need to do a validation of such a method on standard samples or else it would be of poor value.

blogfast25 - 13-3-2012 at 17:57

Quote: Originally posted by Nicodem

However, this or any other titration would not be the direct evidence of the HCl concentration in your acetic acid samples. It would only give you the combined concentrations of all comparably strong acids.

Hmmm... the problem was stated in terms of HCl + HAc, no mention of other strong acids was made.

Using Na or K acetate as titrant solution allows only to determine the titer of the strong acid. Where both acids are sufficiently different in Ka both acids can be titrated in one single operation, sufficiently accurately for most 'industrial' purposes. Using standard samples of course always makes sense, whatever analysis you're setting up. And to be fair, the problem was posed only in terms of determining HCl, not HAc.

[Edited on 14-3-2012 by blogfast25]

Nicodem - 14-3-2012 at 11:23

Provided that the samples are composed only of HCl and acetic acid, then yes, I agree with you. A titration would be enough. I looked at it as an analytical problem in which case a titration would not give the answer in regard to the concentration of HCl. I'm sure the original poster knows his problem better than I do, so I leave it to him.

PS: blogfast25, that (stupid) organic shorthand "Ac" stands for acetyl and should be used as such (or better yet, not used at all). You use it for acetoxy which is non-standard. Any organic chemists looking at "AcH" would be a little bit angry because that stands for acetaldehyde while it is obvious you talk about acetic acid. I know I'm being annoying, also because I myself find the idea of using the symbol of actinium for acetyl kind of organically stupid.

Lambda-Eyde - 14-3-2012 at 11:40

Quote: Originally posted by Nicodem

PS: blogfast25, that (stupid) organic shorthand "Ac" stands for acetyl and should be used as such (or better yet, not used at all). You use it for acetoxy which is non-standard. Any organic chemists looking at "AcH" would be a little bit angry because that stands for acetaldehyde while it is obvious you talk about acetic acid. I know I'm being annoying, also because I myself find the idea of using the symbol of actinium for acetyl kind of organically stupid.

The compendium I used in the general chemistry labs last semester used "Ac" instead of "AcO" all the time... I find it hard to believe that professional chemists (with PhDs) wrote it.

Also, few chemists work with actinium, and I doubt anyone but them are seriously annoyed by the convention we organic people use.

blogfast25 - 14-3-2012 at 12:43

Quote: Originally posted by Nicodem

Oh dear. Non-standard, am I? Good to hear, in a sense...

Scientists that don't recognise context (it's abundantly clear we are talking about acetic acid, ooops, ethanoic acid, here) should get a life and/or a different career.

My Uni textbooks (admittedly 30 years old by now) all use HAc in that context. I guess they were all stupid, these authors. And that my old habits die hard...

Still, to avoid future petty disputes, HOAc it is from now on (have I got it right? I didn't 'peek')

[Edited on 14-3-2012 by blogfast25]

Titrating HCl in acetic acid

mycotheologist - 12-4-2012 at 14:39

How would you determine the concentration of HCl in a HCl in acetic acid solution? All I can think of is using a very weak base that only HCl can protonate but I don't know of any weak bases like that. The weakest base I can think of is bicarbonate but that reacts with vinegar.

barley81 - 12-4-2012 at 15:10

Neutralize with NaOH until neutral, and then titrate with standardized silver nitrate using potassium/sodium chromate as an indicator.

Paddywhacker - 12-4-2012 at 17:06

That depends on whether or not there is water present. Strong acids in anhydrous acetic acid are used to titrate weak bases such as aniline. 0.1 M perchloric acid in acetic acid is a common reagent in analytical laboratories.

If there is water present then titration is more difficult. You could construct a graph of pH vs volume of NaOH added and deduce the neutralisation points from inflections in the curve.

Or you could determine the chloride gravimetrically, as, say, silver chloride.