UncleJoe1985
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Easy way to measure pH < 0
I'm trying to measure the pH of the solution in my Birkeland Eyde reactor. I've done it by titrating using a red cabbage indicator, resulting in 22
volumes of acid to neutralize 1 volume of 10% ammonia => 0.267 moles HNO3 / L => pH 0.57 (after running 6 hours using a 9000KV @30mA NST). The
total amount of solution is ~1L, so I estimate the production is about 2.8 grams HNO3 / hour. Based on the empirical efficiency of 68g / KwH for the
industrial scale operations, my setup only has 15% efficiency 
Sorry for ranting, but my main question is how to measure the pH of the solution in an easy, non-consuming fashion? I was considering getting a pH
meter, but I don't think they can measure negative pH. I heard I could measure the voltage directly from the pH probe (59mV / pH unit), but I also
heard the response is non-linear when the H+ concentration gets to 1 M. Any ideas?
[Edited on 14-10-2008 by UncleJoe1985]
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jarynth
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Dilute your sample first, thereby lowering the H+ concentration in a controlled way. Measure the pH of the diluted sample, then work backwards and
compute the original pH (which varies with the log of the conc...). As the acid dissociates almost completely (at your concentrations), the
computational error will be small.
Or just titrate it like you did but adding DROPS of ammonia to a few ml of sample, not the other way round.
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UncleJoe1985
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| Quote: | | Measure the pH of the diluted sample, then work backwards |
I was thinking about that, which will be easier using a pH meter, but it still requires precise liquid measurement, which I want to avoid. I currently
only have a 5 mL medicine dropper. I'm going to need a pipette to measure smaller amounts during titration instead of 1 mL ammonium hydroxide and 22
mL acid like I did earlier.
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Magpie
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If you are really just looking for a simple, unobtrusive way to indicate the concentration of HNO3, consider electrical conductivity (or resistivity).
Here's some data found with a google search:
Comparison of Electrical Conductivity Data for Nitric Acid
LITERATURE VALUES.............EXPERIMENTAL DATA
Weight %....Conductivity......Weight %......Conductivity
HNO3...........(mS/cm)............HNO3.............(mS/cm)
2.0...............100...................1.82...............86.8
4.0...............195...................3.39..............167.7
6.2...............312...................4.86..............238
10.0.............440
12.4.............542
24.8.............768
Data from www.smartmeasurement.com (4/23/03) and from Yokagawa Systems conductivity plots (8/29/89)
Nitric Acid Con
I realize you likely don't have a conductivity meter laying around either, but perhaps you do have an ohmmeter that could indicate the solution's
resistivity.
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jarynth
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| Quote: | Originally posted by Magpie
If you are really just looking for a simple, unobtrusive way to indicate the concentration of HNO3, consider electrical conductivity (or resistivity).
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Remember to start with distilled water for that. There would also be some nitrous acid in solution.
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vulture
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Do also mind that nitric acid might eat its way through standard (for pool use) pH/conductimeters.
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UncleJoe1985
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I don't mind as long as I can use it for a few years. I don't see how the acid will dissolve anything - the electrode is glass. Maybe it might weaken
the plastic parts. Plus, the glass bulb is stored in mildly acidic KCl to begin with.
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bfesser
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The glass bulbs are usually surrounded by a protective plastic tube, which may be attacked by nitric acid.
Are you trying to measure the pH, or the total concentration of HNO<sub>3</sub>. If you're looking to do this for a few years, just
invest in a buret, a small volumetric pipet, a volumetric flask, a few erlenmeyer flasks, some decent quality sodium hydroxide, and some
phenolphthalein or other suitable indicator. Titrate small diluted samples of the acid with the NaOH.
[Edited on 10/16/08 by bfesser]
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UncleJoe1985
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After considering the steps needed in using a pH meter for measuring a pH < 0, it seems no labor is saved over titration because you need to
dilute.
If the voltage/pH function (which I assume is non-linear) is known, then dilution won't be needed. However, this article says above a certain concentration, strong acids don't ionize completely, so the [H3O+] < [HNO3], which is bad because I want to
measure [HNO3]. If anyone knows that HNO3 ionizes completely between 1M to 10M and the voltage/pH function, I might change my mind.
It seems the only thing the pH meter buys you is accuracy.
Therefore, I plan to use titration by sodium carbonate. I used 10% ammonia earlier, but think the 10% might not be accurate and the concentration will
decrease as the NH3 boils off.
However, according to the titration curve, you can't use red cabbage indicator because the end point is outside its useful range. If you keep adding acid until the bubbling
has stopped, how accurately will that capture the end point? I only plan to use 1 mL of acid, so it might be hard to see the bubbling.
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not_important
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given the solubility of CO2 in water, not very accurate. Picking up a couple of pipettes and a volumetric flask for diluting and measuring the pH of
that would be quicker than titrating, and avoid indicator and end point problems (carbonates are not the best base to use)
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UncleJoe1985
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| Quote: |
diluting and measuring the pH of that would be quicker than titrating. |
I don't see how using a pH meter is easier - only more accurate. I will need to measure 0.5 mL of acid and 10 mL of water to dilute it to pH > 0.
If I were to titrate, I would store a basic solution already mixed with an indicator and add drops until the end point.
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not_important
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One pipette to pull the sample of acid, deposit in volumetric flask and rinse into same, which is then filled to the mark with DW, then stick in the
pH probe to measure. That all takes 15 to 20 seconds, 5 simple steps.
Titration - make the solution of known base strength, which will take about as many steps as diluting the acid. Then titrate, counting drops - slow,
more error prone, or use a proper calibrated burette to quickly run in a measured amount, then switch to drop by drop mode until the end point is
reached.
Alkaline solutions intended to be used as titration reagents can be a bit iffy. Strong alkali absorbs CO2 and attacks glass. Carbonates have softer
and fuzzier end points because of the H2CO3 <=> H2O:CO2 <=> H2O + CO2 slow equilibrium. Weak bases like ammonia also have less clear end
points, solutions of ammonia may suffer loss of ammonia as a gas, as well as CO2 absorption.
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bfesser
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Titrate with aqueous NaOH standardized with dried potassium hydrogen phthalate using a buret and phenolphthalein in ethanol as the indicator. Do not
premix the indicator with the titrant--that's wrong. Titrate to <em>very</em> faint pink which lasts for about a minute with constant
swirling/stirring. Use a volumetric pipette to transfer the acid.
[Edited on 11/2/08 by bfesser]
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UncleJoe1985
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| Quote: | | Do not premix the indicator with the titrant--that's wrong |
How's that bad? Is it because the indicator deteriorates over time? I'm trying to minimize my labor and as long as I get 0.1 M accuracy, I'm
satisfied.
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bfesser
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Yes. Decomposition of the indicator is a main reason. You're also wasting the indicator. And if you're lazy enough that adding 1 drop of something
from a bottle to a flask is too much work, well...
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UncleJoe1985
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Good thing I didn't store my ammonia with indicator premixed. However, I read that red cabbage juice can only be stored for a few days refrigerated.
I've left my red cabbage indicator at room temperature for at least a month and it is dark red probably from sugars converting to acetic acid.
I also noticed its taking longer to for my reactor to produce 1.3 M acid based on titration. Does anybody know if the indicator could deteriorate and
shift its green to red transition to a much lower pH? If it's a shift of +- 3 from 7, then it shouldn't be a problem because it's negligible relative
to the [HNO3] I'm trying to measure.
BTW, according to this, it says sugars destroy anthocyanin (I boiled down my indicator to a fairly concentrated level, so it's probably high in sugar), but I'm not
sure what it means in terms of indicator usability.
[Edited on 10-11-2008 by UncleJoe1985]
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