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Titrating for the Concentration of Spent HCl

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[editor appended this entry to this thread which already addresses it in lieu of spawning a duplicative thread]
Q. Sir, how to find out acid content in gm/Ltr in Flux in galvanizing industry.

⇩ Related postings, oldest first ⇩

Q. Dear Gentlemen:

Just to be confident with my analyses of spent HCl concentrations using acid-base titration with phenolphthalein ⇦ this on eBay & Amazon [affil links] as indicator, I'd like to solicit widely accepted procedure or practice in the determination of its end-point. I really cannot help but remember my chemistry teacher's instruction to stop at the "very first persistent color change" that being faint pink for phenolphthalein; however, spent HCl gives a light yellow to greenish black color as the volume of base titrant is increased.

A friend of mine who has worked with galvanizing plants stops at the greenish-black color. I tried it but the disparity in percentage HCl equivalent from Normality/Molarity result is quite high from stopping at the first color change (almost like three-fold). 1% from 4% or vice-versa is quite a large value to consider.

I really would appreciate if you tell me which is which considering your expertise in the field.
Kind regards.

A. Hi Barlow, The problem that you are having is that phenolphthalein is the wrong indicator to use for titration of any acid that has dissolved metals of any kind in it. The reason is that the high pH that it turns color is already precipitating some metals, so you get a bad lie reading. Switch to methyl orange ⇦ this on eBay or Amazon [affil links] or methyl red ⇦ this on eBay & Amazon [affil links] or similar range of indicator color change. You should titrate with with 0.5N sodium hydroxide ⇦ this on eBay or Amazon [affil links] with slow additions or you will be dropping out some of the metals because of the local concentrations. Metal hydroxides do not re-dissolve as fast as you might think in anything except very low pH range. Spin bars help a lot also. James Watts - Navarre, Florida A. Dear Barlow Campano, Sometimes none of the procedures are right for the case at hand. In theory, I would say your teacher is right. In practice there could be a big difference. If you have salts with buffering properties in your solution and if they form nice strongly coloured components as well, it could be that the old fashioned volumetric manual methods are not suited at all. If you look up volumetric HCl analysis in the well known reference book "Analysis of Electroplating and Related Solutions" by - K.E. Langford & J.E. Parker [on AbeBooks or eBay or Amazon affil links] , they describe methyl orange ⇦ this on eBay or Amazon [affil links] as indicator for this type of analysis. Methyl orange is active in the pH-range 2,9 - 4,6, and phenolphtalein in the range 8,3 - 10. Using one or the other could give quite a difference already. In your situation, where obviously other components are present as well, I would investigate the titration curve with a pH meter -- in that case you would be able to define who's right, as finally you will be sure that all HCl is neutralized if you go from approx. 5 to approx pH 9. In your case, you could do your first pH measurement at the first color change and after that continue until you have reached the other colour and see what the differences are. And for your own interest, try the titration with a new, uncontaminated solution as well, this will teach you something about the differences you can have. Of course the readers here are very curious to hear which of you was right. Hope you're not going to lose vast amounts of money in your bet. Best regards, Harry van der Zanden consultant - Tilburg, Netherlands

A. You might consider just using a pH electrode and stop at a pH where your indicator would normally change.

Terry Tomt
- Auburn, Washington

A. Sir:
James Watts is correct.

Regards,

Dr. Thomas H. Cook
Galvanizing Consultant - Hot Springs, South Dakota, USA

Dear Mr. Cook,

for sure Mr. Watts is right, but still the question of Mr. Barlow Campano remained unanswered.
I've always learned too that you have reached the end point of a titration as the first colour change occurs.
Even with the wrong indicator, in an unsuited medium I'm wondering if Mr. Barlow will find a pH above 7 as the first colour change occurs (I guess he will).
In that case, he and his (and my) teacher were right, so still curious to learn what happens when measuring with a pH meter during the titration.
Could be that all other colour changes occur, due to the precipitations that occur.
So still hoping for a follow-up.

Best regards,

A. Harry, If you use a pH meter you will find an inflection (large pH change over a small amount of added base) near a pH of 4.5 to 5. This inflection is due to the end point of strong base/strong "free" acid HCl. The next end point(s) are due to the titration of metal cations precipitating with OH-1 groups and generally occur near "neutral" pH's (e.g. 7 and even above). I like bromophenol blue ⇦ this on eBay or Amazon [affil links] sodium salt best for these titrations because of the sharp color change (I can see it nicely.) It you titrate to pH 7 then you are precipitating ALL the iron Fe+2 and or Zn+2 that was ever dissolved in the acid. To go to this high of pH also is complicated by the various colors of the metal hydroxides. At any rate going to pH 7 is just plain wrong as Mr. Watts has said. Regards, Dr. Thomas H. Cook Galvanizing Consultant - Hot Springs, South Dakota, USA A. The actual pH or the depth of indicator color at the point that you end your titration isn't significantly important because if you plot a titration curve in mL titrated on the x-axis vs. pH on the y-axis, you will see that the slope of the curve is very steep near the equivalence point. That is to say that the pH changes rapidly over a very small titrant volume increment. So, if the HCl to NaOH equivalence point occurs at an approximate pH of 7.0, by the time you get the first permanent phenolphthalein pink (about pH 8.0) you are already past the equivalence point. If you went to a darker pink, the pH would be about 9.0. But, it doesn't really matter to your end result because the curve is so steep in that region. I hope this helps. Jon Barrows, MSF, EHSSC GOAD Company Independence, Missouri Thank you for your kind responses gentlemen. I noticed the same is true when using methyl orange ⇦ this on eBay or Amazon [affil links] as indicator - you're tempted to pursue the blackish green color after the orange has turned to yellow. I never had tried reading a pH meter though while performing the test but I will next time to be sure. Anyway, I believe I have to end up to the first persistent color change rather than going for the blackish green (we could always check the excess through back titration). I also find methyl orange ⇦ this on eBay or Amazon [affil links] good for pickle acids with high HCl concentrations but I prefer phenolphthalein ⇦ this on eBay & Amazon [affil links] (with NaOH) for the quench water and flux. Best regards. Barlow Campano galvanizing chemist - Jeddah, KSA

A. Barlow,
Phenolphthalein has no application in the three applications that you mention. There are many mistakes by the people who have responded in this thread. Mr. Watts is still correct. I am finished with my comments.

Regards,

Dr. Thomas H. Cook
Galvanizing Consultant - Hot Springs, South Dakota, USA

A. It looks like I need to amend my earlier comment to say that I certainly agree with Dr. Cook and others that phenolphthalein is an incorrect indicator when there are dissolved metals. I was answering, in a general case, a portion of the thread that seemed to be asking a question about phenolphthalein ⇦ this on eBay & Amazon [affil links] in an acid-base titration. The dissolved metals in solution change what occurs around the equivalence point because there are multiple reactions taking place as you increase the pH beyond about 4.0. That is why this particular titration should be ended at a low pH where you are avoiding the interferences from metals.

Jon Barrows, MSF, EHSSC
GOAD Company

Independence, Missouri

Q. Kind Sirs:

It seems that having you around in this site makes me lazy. However, granting that this is the reason why this site was created, I am hoping I'd be excused. Nevertheless, I did my homework on comparing both phenolphthalein ⇦ this on eBay & Amazon [affil links] and methyl orange ⇦ this on eBay or Amazon [affil links] indicators (since I have the luxury of having both at the lab) on various acidic plant solutions and the results are as follows:

NOTES:
1. New acid preparation is about 15 % w/w (40% - 29% pure acid; 60% - used rinse water)
2. Acid No. 2 is relatively new
3. Flux formulation is double salt

Hereunder also are my remarks:

1. Conventionally, colorless acidic solution becomes red (actually ranges from light pink to fuchsia) when phenolphthalein is used as acid solution indicator with the change occurring at around pH 8.00. However in the above case, the acid contains appreciable amounts of metals reactive to base and the expected color change did not turn out to be. It may not follow the expectation but the next color that it changes to (say from colorless to yellow orange as in the acid no. 2 sample), have exactly about the same pH as with using a methyl orange as indicator. Endpoint then can still be detected vividly using phenolphthalein as indicator - one just needs to have an idea when to stop with.
2. Pushing further with the greenish black color as endpoint from yellow (in spent acid) when using phenolphthalein has the tendency to magnify the actual free acidity by consuming more caustic at a relatively lower pH rise (below equivalence pH).
3. Methyl orange does not give a clearly detectible endpoint (color change) when used as indicator in the determination of flux free acidity both with sodium hydroxide and sodium carbonate as base titrant. The % HCl given by phenolphthalein, though, should be confirmed or correlated with Dr. Cook's findings (if I may request) on HCl strengths formed from various flux formulations such as a double salt flux. I'd like to note also that I happen to pull out an electronic copy of a booklet on pipe galvanizing published in the internet by a well known source describing methyl orange as indicator for detecting free acidity in flux. From the above result, it would show that the author/s have made an error on the type of indicator to be used or I just do not know how they did it.
4. Cresol red ⇦ this on eBay or Amazon [affil links] is identified as an indicator which changes color from yellow to red at about a pH of 7.00 (Quantitative Chemical Analysis, Hamilton and Simpson ⇦[this on eBay or Amazon or AbeBooks affil links] , 12th Ed.,page 138). Wonder why it is not commonly used.

Hope you'll bear with me with the interpretation of the data.

Best regards,
Barlow

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A. Dear Mr. Barlow,

Thank you for your homework, this type of feedback makes the forum more valuable to everyone.

As expected, no big difference in the concentrations found, using the two different indicators. So for practical purposes, you could use either of them now, however I understand it is still difficult to exactly determine the end-point, due to the precipitation of salts in your solution.

As Mr. Watts already indicated, in principle the titration should be stopped before the salts in the solution precipitate.
As the pH of spent acid rises above 2 , iron loses it's solubility and precipitates. Ferrous hydroxide precipitates at pH 7 to 9, ferric hydroxide at pH 2 - 4.
This suggest that the use of an acid-base indicator which changes colour in the pH 2 -3 range should be the preferred choice.
Thymol blue changes in the pH range 1,2 - 2,8 from red - yellow, and as the Fe-precipitates in your solution are reddish-brown, this will make it already difficult.
Like Mr. Cook also already earlier suggested, bromophenol blue ⇦ this on eBay or Amazon [affil links] (changes at pH 2,8 - 4,6 from yellow to blue), may be your best bet for having a clear difference and a repeatable result.
The cresol red ⇦ this on eBay or Amazon [affil links] is already too high in pH, like the others. I'm afraid it will give you the same problem.

If no indicator serves your needs, I would go for the titration, using the pH meter and stop at pH 4 - 4.5.
For the flux, I have no comments, I hope others will help you out.

Thank you for your feedback and work done!

Kind regards,

Harry van der Zanden
consultant - Tilburg, Netherlands

A. Sirs:

James Watts is still correct.

What do you think the "correct" pH's are for: 1) Rinse, 2) Flux and 3) Quench?

Flux is a natural buffer and it certainly does not contain 13% HCl.

I see you have not taken into account the dissolved zinc in your acids.

Regards,

Dr. Thomas H. Cook
Galvanizing Consultant - Hot Springs, South Dakota, USA

A. Dear Mr. BARLOW
I know that you ask the experts here. I am not one of them but I will tell you what I know.

When you make up a new pickling tank the HCl content will be high and the FERROUS CHLORIDE content will be low, at this stage you will not face a problems with the titration, the problem appear at the end life of the pickling tank and it is important to get the true results to decide if you will through the solution and what is the amount of HCl you will add to the new one.

At the end life the FERROUS CHLORIDE content will be high and the HCl content will be low; by the normal titration you will get an inaccurate result -- that is because you will consume NaOH not to neutralize the HCl but for converting FERROUS CHLORIDE to FERROUS HYDROXIDE (dirty green precipitations); at this stage I use the chart that compares the HCl content, FERROUS CHLORIDE and specific gravity, it helps to know the actual HCl content. Note that all the FERROUS CHLORIDE content will convert to FERROUS HYDROXIDE at PH=4.5 to 5.0, it means that is before the HCl neutralization; secondly try to estimate the pickling tank behavior that is by KLEINGARN chart -- it is useful in concentration estimation as will as pickling optimization. Last thing, please measure the FERROUS content before measuring HCl content, it will help you to know what you will face during titration process.

M. Saad
- JEDDAH , KSA
September 19, 2010

Q. Dear kind sirs:

I've tried titrating a diluted reagent grade ferrous chloride ^ ^{ferric chloride} which is inherently acidic to know if I'll get a free acidity measurement and I did (as expected). This will then tell us that the H+ ion from water in the solution did actually dissociate to form the acid HCl right? To be more than sure, I spiked the sample with pure hydrochloric acid and I also did get the spike over the previous reading.

This then proves my initial idea that for a double salt flux, the ZnCl2 will still dissociate with massive amounts of water to form HCl and Zn(OH)2 - and you will have a free acidity reading for that; more so, that zinc chloride is normally acidic. For the ammonium chloride however, I was not able to get a free acidity result same as what I expected based on stoichiometry. I expected both salts to give about 10% HCl upon dissociation with water in the flux tank however, I was able to get only around 3% for ammonium chloride alone. This was fortified by my results in titrating a double salt flux (with dosage in accordance with supplier recommendation) solution which gave a free acidity of about 13% w/w as HCl.

By all means I would believe this would be the case because we will be having best of both worlds in fluxing, that having an acidity (as HCl) that will take care of newly formed metal surface oxides AND the tight crystalline coating upon withdrawal and drying that will protect the bare metal from further oxygen attack as Zn(OH)2:Cl2 and the amine (I would appreciate much if someone can help me out on this).

The 13% percentage HCl is quite reasonable as we kind of have an extended pickling in the flux tank PLUS the other benefit. Dr. Cook sir, I re-analyzed our flux solution using phenolphthalein ⇦ this on eBay & Amazon [affil links] as indicator, however, ending at the equivalence point (pH 7.00). I believe the excess NaOH that precipitates the soluble metal hydroxides is not that excessive at this point. Since I cannot clearly distinguish the endpoint unless it becomes pinkish, I made a correction factor that is to divide the volume of titrant used upon reaching the pinkish endpoint with 1.5 - that will give us a volume close to the equivalence point. My result was half of the previously reported in this forum and I believe is logical.

Regarding the presented Kleingarn Curve that some of us are using, I cannot correlate it with my results being that the iron readings that I got have 20 g/L the most at about 3% free HCl left and the chart has as high as 80 g/L with the same free acidity and density/specific gravity values.

The method I used for iron test is just visual colorimetry using Hach's ferrover iron reagent for 5 ml sample but the equipment's precision is high. Aside from that I learned that adding fresh acid to increase the spent acid's free HCl is just a waste of time and money. We're just holding the line of acidity at a shorter usable period which is tantamount to making a new acid solution at extended free operator man-hours - explaining the latter's cost-effectiveness.

Will appreciate your feedback on this concern kind sirs.

Regards,

A. Barlow, The proper pH for flux, and hex chromate quench is 4.5. The proper pH for acid rinse is 4.5 or lower. Regards, Dr. Thomas H. Cook Galvanizing Consultant - Hot Springs, South Dakota, USA Dear Sirs: Please accept my apologies. I tested a dilute sample of 40% w/w reagent grade ferric chloride ⇦ this on eBay or Amazon [affil links] and not ferrous chloride as initially reported. For your perusal. Kind regards, Barlow Campano galvanizing chemist - Jeddah, KSA September 24, 2010

Q. Barlow,

I have a Galvanizing customer testing the virgin HCl (20 Be') using Bromo Phenol Blue ⇦ this on eBay or Amazon [affil links] in the method for determination of acid strength.
What method would you recommend for field use not laboratory.

Thanks!

OS

Q. Dear sir, one of our customers asked us to create lab procedure to test HCl, flux concentration and iron content of process tank.

I am datta shinde, working in hot dip galvanizing plant last 4 years now. Please give me following information details:
1) in galvanizing flux, free HCl & iron content: how to analyze, details, all methods.
2) in galvanizing flux, free HCl & iron content how many minimum & maximum limit
3) in galvanizing flux bath, iron content minimum with what effect on fluxing job & zinc bath
4) in galvanizing flux bath, iron content maximum when what effect on fluxing job & zinc bath.

Now sir, please give me detailed information .

Thanks & regards,

 Hi, cousin Datta.

The above conversation is long and somewhat confusing, and not everyone was in agreement on all points. Still, I think it does address your questions about HCl content.

Other threads here cover other aspects of your questions; for example, thread 28861, "Iron content in HCl" addresses that subject. And we have at least half a dozen threads just about galvanizing flux, including thread 43365, "Hot dip galvanizing flux Q & A's, Problems and Solutions".

Please try to absorb the points presented on this thread, search the site for in-depth discussions about galvanizing fluxes, tell us what kind of flux you use, and request clarification of whatever wasn't sufficiently detailed and clear. If indeed you truly want full details of all methods, you certainly need to consult books, not a forum where answers can't be expected to exceed a couple of paragraphs:-)
Thanks!

Ted Mooney, P.E.
Striving to live Aloha
finishing.com - Pine Beach, New Jersey
July 31, 2011