aga
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Titration cockups
I have been obsessively doing titrations since the burette arrived, along with Methyl Orange, Yellow and some Phenolphthalein, plus some sodium
carbonate 99.97%.
Titrating a known 6.8M nitric acid solution, i get reasonably close (seeing as i'm a noob) using phenolphthalein, and surprisingly, red cabbage juice,
but an error of at least 50% using the Methyl Orange and Yellow.
The Methyl orange says that it is Orange below pH 3.1, and Yellow above pH 4.4.
Can i assume that i see a rapid colour change at pH 4.4, but that does NOT indicate total neutralisation of the acid using the Methyl compounds as
indicators ?
I'm putting the titrant in the burette, and a measured volume/concentration of the titrator + indicator in the beaker, into which the titrant solution
is slowly added.
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blogfast25
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MO and MY are much harder to use than phenolphthalein, which has a very sharp end point. But MO's end point is orange, so it goes from red to orange
or from yellow to orange depending what you are titrating. It can be hard to see.
I use MO only to titrate standardised Na2CO3 with HCl, it then goes from yellow to orange. If you go to red you've missed the end point. It takes
careful observation.
Not so easy, these unassuming little titrations 
[Edited on 20-4-2014 by blogfast25]
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aga
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Definitely Not So Easy (depending on how much red cabbages can be trusted).
I guess 'getting it right' teaches a young beginner the essentials, as in Planning, Patience, Being Meticulous, and Questioning all of the Assumptions
made, and basically Thinking about it.
I'm not so young, so i think i'll be buying an autotitrator, or a digital pH meter at the very least.
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smaerd
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Digital pH meter is a good idea, but an autotitrator is probably unnecessary for most tasks. Edit building one would be more interesting. A good trick
when you are approaching the endpoint (pH meter or not) is as follows; spin the stop-cock literally as fast as you can twist your middle and thumb
fingers 180 degrees (1/2 tau or 1pi for you radian freaks  ). This allows for a
so called 'fraction of a drop'. Then wait briefly. Adequate stirring is required. Repeat 'fraction of drop' additions with brief periods of waiting.
This is one of those tricks that got me through some general chemistry and analytical experiments with consistent results where many other class-mates
suffered from 'poor' technique.
To be honest, I've had more significantly more issues with 'cheap' auto-titrator devices then doing it with the good old burette. Especially of the
'drop counter' variety. Your mileage may vary.
Of course if your analyte is cheap it's a good idea to take a little bit of it and do a quick titration with a more dilute titrant to calculate
approximately where your end point 'should' be. Of course for more involved analytical work, don't forget things like ionic strength, and activity
coefficients. The obligatory mention of standardizing titrants should also be stated.
Best wishes to you.
[Edited on 21-4-2014 by smaerd]
sorry last edit,
This image may help answer your question about indicator changes that exhibit a continuum of color changes.
http://chemistry.beloit.edu/classes/Chem220/indicator/indica...
[Edited on 21-4-2014 by smaerd]
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aga
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Cool ! That is very useful chart. Many thanks.
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blogfast25
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@smaerd:
Your fraction of a drop method only works if your burette allows very precise readings. Even a whole drop is only about 0.05 ml. Poor results with
titration can be caused by a whole raft of problems, not just fractionated drops or 'technique'.
@aga:
Titrating against a pH meter is what I used to do but it isn't easy either. Beginners often think they need to titrate to pH = 7.00 but in most cases
that simply isn't true.
When titrating against the meter, the analyst is looking for the largest jump in pH (call it ΔpH), near the end point. A sequence of pH readings
near the end point might go as follows:
8.0; 7.9; 7.7; 7.4; 6.1; 5.9, 5.8 etc. The biggest ΔpH is 7.4 - 6.1 = 1.3 and that's your end point. In the end I resorted back to good old
indicators.
Robotitrators are only useful in environments that do loads and loads of the same titrations, for instance for QC or process control. Robotitrators
basically determine that biggest ΔpH mathematically.
[Edited on 21-4-2014 by blogfast25]
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smaerd
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Yea that is why I put 'poor' in quotes. The idea is it usually still takes 5-20 of those 'fractions of drops' to get to a desired end-point anyways,
but it provides a miniscule amount of extra control. For phenolphthalein its the difference between a faint pink that just barely sustains verses a
well you know what kind of pink I mean. With other indicators sometimes I hit the point that I believe is the end point. Write down the volume. Then
drive it a bit further to be sure. Especially for a first run with an indicator color change that I am unfamiliar with.
I'm not seasoned with analysis like you probably are blogfast but my experience with pH meter's hasn't been too great either. Albeit it's probably due
to different reasons. Some of the ones at my university are probably about the same size as an old CRT computer monitor with maybe 1/3 the depth
(circa 1980). Two carefully calibrated pH meters read the same solution with a difference of I think it was 0.30 pH 'units'. Needless to say, I'm a
bit more trusting of indicators for the most part as well.
[Edited on 22-4-2014 by smaerd]
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Magpie
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Quote: Originally posted by aga  |
Titrating a known 6.8M nitric acid solution, i get reasonably close (seeing as i'm a noob) using phenolphthalein, and surprisingly, red cabbage juice,
but an error of at least 50% using the Methyl Orange and Yellow.
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The different indicators have color changes in different pH ranges. You should use the indicator that changes color in the pH range that coincides
with the steepest part of the neutralization curve of pH vs volume of titrant. Therefore the indicator that is appropriate depends on the particular
neutralization that you are doing.
For a strong base/strong acid neutralization like NaOH/HCl phenolphthalein works well, going either way IMO.
If you are titrating a weak acid like acetic acid with a strong base then a different indicator may well be more appropriate.
If you have a biprotic (H2SO4) or triprotic (H3PO4) acid you would use a different indicator depending on the neutralization of interest. For H3PO4
there are 3 neutralizations, all occurring in different pH ranges.
A good book on analytical chemistry should show you some characteristic examples.
The single most important condition for a successful synthesis is good mixing - Nicodem
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Mildronate
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You can make some control sample with buffer solution like your endpoint. You can also use photometer!
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blogfast25
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If you really are going to the expense of 'going electronic' then potentiometric (recorded pH) is far more advisable.
But good old indicators are cheap and easy and only need a bit of practice.
I'm all for good books but the expected end point pH can be calculated easily in most cases, using simple acid-base theory. The case of H3PO4 is a
little more complicated, if you want to use all three neutralisation points.
[Edited on 22-4-2014 by blogfast25]
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Mildronate
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Hm i think photometric titration is more accurate at pH near 0.5pKw=7, ok with pH meter and data logger you can write some titration curves, then make
derivative d(pH)/dC and volla at peak max you have stehiometric point and forget about eye  .
[Edited on 23-4-2014 by Mildronate]
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