camura
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pH strips for amateur strong acid-strong base titration
Hello. I don't have access to normal indicators like phenolphthalein or methyl orange, but I do have pH test strips, and was wondering if they could
do in a pinch.
To test this idea:
I made a 1M solution of NaOH from some store bought granules assumed to be monohydrate for my titrant. I pipetted 2ml of store bought HCl into a
glass, and diluted it with hard tap water as the analyte. I placed a pH strip in the analyte, and performed the titration.
The pH strip started red, then as the titration progressed, suddenly turned blue. This sudden change was expected from the titration curve of a strong
acid-strong base titration. The blue colour then detached from the paper, becoming a suspension.
This was performed twice. The original HCl solution was calculated to be 22% and 23.5% concentrated by weight from the two tests. This fits within the
10-35% range in the supplier's SDS. (what an oddly wide range)
So, what are your thoughts? Could pH papers work for strong acid-strong base titrations? I have only a very cursory education in chemistry, so I'm
sure there are things I'm missing. Reactions between the paper and the solutions, for example. I think I could nail down a better precision if I was
paying more attention with the titration. Also, this is my first post to the forum, so I apologise if I am making an error in my etiquette.
[Edited on 11-7-2025 by camura]
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Sulaiman
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welcome
although you can titrate a strong concentrated acid vs. a strong concentrated base
it is better to dilute at least one or the other, (the one that is added dropwise), or preferably both.
This improves accuracy by using larger volumes so finer control over the neutralusation end point,
hence a more accurate measurement.
It uses much less quantities of (usually) not-cheap ingredients.
very strong acids and bases can destroy some indicators.
....................
eg add 1.00 ml concentrated acid to 99 ml water.
CAUTION : Hobby Chemist, not Professional or even Amateur
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camura
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Quote: Originally posted by Sulaiman  | although you can titrate a strong concentrated acid vs. a strong concentrated base
it is better to dilute at least one or the other, (the one that is added dropwise), or preferably both.
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I certainly would have gotten a better result if I had diluted the NaOH more, I'll do this next time. Although even at 1M, I was using half the
burette's capacity for the titration.
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Radiums Lab
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NaOH in monohydrate form, how? Did you assume that 40gm of it is containing 18gm H2O.
Water is dangerous if you don't know how to handle it, elemental fluorine (F₂) on the other hand is pretty tame if you know what you are doing.
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j_sum1
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Fundamental to any titration process is to calibrate your setup. NaOH is not reliable since it so readily absorbs moisture and CO2 from the air.
A good base standard to use is home made sodium carbonate. Simply heat sodium bicarbonate to drive off CO2. Heat until no mass reduction is
observed. Or you could use borax.
A suitable acid standard to use is "cream of tartar", potassium bitartrate.
It helps for accuracy in your process that your primary standards are relatively high molecular mass. This gives more precision when weighing out a
specific number of moles.
Other things to look for in a standard are long-term stability, lack of unwanted side reactions, and known stoichiometry. For acid base analysis, it
is better to have at least one of your species being a strong acid or base. This makes the titration curve steeper and the transition more sudden.
However, there is an advantage to a weak acid/weak base combination when you are setting up your system – read on.
Now, I know your question is more about indicators. There are options other than commercial indicators which you said are hard for you to obtain. A
classic, which is surprisingly good is red cabbage juice. (Other purple vegetables work too but cabbage is the best in my experience.) Red cabbage
does present with a range of colours as an indicator, and these can vary a bit from cabbage to cabbage. So, what you need to do is determine the
colour at a neutral pH. here is how you can do it.
Mix up some standard solutions using of your acid and base. Calculate how much is needed to exactly neutralise. Then mix together. Effectively you
are creating a buffer solution with a pH of 7. You can check with your pH strips if you like. Then add some of your cabbage indicator. (I would use a
measured amount.) Then record the colour. A camera and reproducible lighting is helpful. Now you know what to look for as you do your titration.
Other options:
Go to a pool shop and buy some phenol red. It has a transition from yellow to red/pink at a pH around 7.2. It is quite inexpensive.
Or, do what I do with HCl. I distil it and obtain something close to the azeotrope (20.2%). From experience I know my of-the-shelf acid needs to be
diluted with water in a volume ratio of 5:2 and then it nearly all comes over as the azeotrope – the distillation temperature remains very constant.
This has the advantage of removing unwanted junk. It also means that my stock is non-fuming and does not rust my lab. When I need more concentrated
acid I make it from NaCl and H2SO4.
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camura
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Thank you, this is excellent advice. There is so much here I had not even considered! Combining cabbage juice with a calibration is brilliant.
A brief story of science foolishness: my original idea was to do a NaCO3 titration, using the production of gas as the indicator of neutralisation. I
abandoned the idea when I realised part of the NaCO3 could become NaHCO3 without giving off gas. That, and the gas production isn't always obvious,
and part of the gas could remain dissolved in the water. A worthwhile misadventure, because I learned about producing anhydrous NaCO3 from store
bought NaHCO3.
Who knew phenol red was so widely available! I had only found pH strips in pool stores previously, and written them off.
I had assumed it was in monohydrate form because that seemed the most likely form it would be sold in, and I could not figure out any way to verify it
in any case, so I just went with it.
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Texium
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Perhaps it would nominally be NaOH•H2O at the time of manufacture, but it absorbs moisture from the air so rapidly that you can’t
safely assume, especially when using it for analytical purposes.
Also, sodium carbonate is Na2CO3 not NaCO3. I’m not just saying that to be pedantic. It’s
important to make sure you get your formulas right, especially if you’re interested in titrations, since with the correct formula in mind, it
becomes clear that 2 equivalents of acid are required to fully neutralize a carbonate.
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camura
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Whoops, that's embarrassing. Thanks for the correction.
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