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peach
Bon Vivant
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This is an interesting discussion, and I will be scratching my head about it also.
I'm sure I remember talking about something similar in A-Level chemistry, but alas, I'm a biology boy over a chemist and so some of it has drifted
into the darkness.
Given the odd measurements you're getting and the impurities you've found in the materials, it may be worth stopping and harvesting together something
you can be more sure of (as per the long delay in that reply appearing in the AlCl3 thread, which will be followed by some more delays as I do the
same again). If you're putting in so much effort, you'll undoubtedly want a solid result.
I have some questions regarding the practical side of things. I'll bold them to make them more immediately separate.
What resolution is the pH meter, and is it a pool kind or a lab kind? Is it temperature compensating?
I have one that displays to 0.001. That's useless for everyday neutralizations where I can pretty much get what I want by eye (without an outside
indicator), but it does show up some of the more hidden effects of pH. For example, I can neutralize to 7, then sit and watch the meter and I can see
it drifting back up or down to where it started because of all the zeros. Quite often, it can drift a fair way. I can accelerate to a true pH by
watching the drift rate (which won't appear on the .1 meters for a long while).
A prime example of this is cannabis growers (and the guys growing those big watery tomatoes) using hydroponic feed bins. If the solution isn't
buffered, they can set the pH (well below where it should be), walk away, and it'll be back over 7 the next day or two as the pH is still stabilizing
a long time later. It can happen over an entire week, so using tiny solution volumes and molarities would help with that, and make the calorific
measurements harder; ah, nature.
If your meter is reading to .1 and it's the pool type, that effect may be hiding from your on looking gaze, causing you to think it's fully
neutralized prior to it being so.
When I'm going for gold on the pH, it needs to be in an Erlenmyer with a big stir bar (I'm talking 75mm, the biggest that'll fit) running at full
speed and it needs to sit for ten minutes between additions as I approach where I want it; I'll leave it for half an hour sometimes. The meter
functions as a datalogger. When I'm done giving myself cancer from whatever the hell is now lurking in the AlCl3 orange solvent of the flask, I could
plot a graph of it happening.
I refer to it as 'bouncing'.
I'm guessing you've calibrated it against a solid buffer. The pool ones often only have one calibration point, calibrate at or close to your end point
(of coarse). If it only has one point, leaving it at the end point to check the stock solutions will skew the reading it gives for those. NurdRage on
youtube has a video all about accounting for the skew between points on a wonky meter.
Accelerating heat loss
Remember not only does it take energy to warm the liner of the thermos up (and the water it's holding), it's being lost through the thermal bridges.
And the rate will increase as you experiment at higher temperatures, which could squiffy up the tolerance bands. So you may wish to check the bridging
by holding the thermos contents at certain temperatures and watching the energy it's taking. As Nicodem suggests, that could be as simple as a
resistive heater left running for a while with a multimeter on it.
Pure materials
I have bottles of glacial and oxalic acid hanging around the house. The oxalic is a mite treatment for the bees, the acetic I bought ages ago and
haven't used much. The former is only 3 or 6% I think, in sugar solution made up with distilled water (so that'd need cleaning up). The glacial I
think is fairly clean, it does what it's name suggests.
If you're concerned about purities, how much are you talking about using? If it's not a lot, I can put some in the post free, for the good of science.
I'm sure I could pick up some pure materials from the lab supplier for some paypaling to cover the cost. I doubt someone bothered about enthalpies,
entropies, neutralizations, the finer points and in need of technical grade materials is baking crack. I may be about to order up some more things
(for my gelatinous mess experiments), so now may be your chance if you're having problems finding pure materials.
I have a 0.1mg balance that was just recalibrated by Avery Weigh Tronixs (it came out of Air Product's cryogenics lab). If you're having measuring
issues and want the tolerances tighter, and are willing to paypal some £ for the materials I'll have to add to the order, I can premeasure stocks of
them and print you receipts on the weights for dilution.
NurdRage also has a video about manganese dioxide he bought from eGimp, in which he dissolves the manganese, has a look under the microscope, and it's
cut with sand. Tisk tisk...
I'm sorry if I'm teaching you how to suck eggs here, I'm not sorry for trying to help.
[Edited on 23-8-2010 by peach]
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blogfast25
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That's a long post, John, I won't be able to address every single point but I'll try.
Firstly, my enthalpic measurements are of course approximate. I do my best but have my limitations: it's a roomy potting shed, not a state-of-the-art
conditioned laboratory!
On pH measurements, I use a regularly calibrated (two points) 'pot grower' model which has served me excellently over the years. 0.1 pH point
resolution. Max. temp. 50 C, but I never use it above 25 C. I treat it like the Queen bee!
For acid base titrations, it's the pH jump that indicates the end point, not some absolute value. Not perfect but more precise than indicators.
Sometimes I use both. For titrations I haven't done before a 'master pH v. volume' curve is recorded to actually see the pH jump clearly and
graphically. Currently I'm having bizarre problems with OTC citric acid. Need to confirm or I'll have egg on my face...
For the neutralisation experiments, pH doesn't even come into it. I standardise the acid solution, use a precise volume, then neutralise with a small
but known excess of alkali. pH is only checked at the end to confirm it is higher than the theoretical end point of the acid (pH of the conjugated
base).
For acid/base standard for the moment I use very carefully twice recrystallised and dehydrated soda. I've perfected that, even if I say so myself. I
don't have purified oxalic acid. If you have some to spare, by all means send me some. Use the U2U to arrange something.
Scales to 0.1 mg? In my dreams! These things from Sartorius cost up to £5,000! I'm investing in a 0.001 g scales. But currently I manage 4
significant digits, that's OK for my purposes. A general lab upgrade is also in the works, considerable investment...
Thermos flask: heat losses are severely limited by the small temperature difference between in and out of the flask: rarely does the temperature in
the flask exceed 25 C. But I get your point. And take the greatest care possible (okay, that's probably an exaggeration ;-) )
Manganese oxide and other oxides: beware of the 'Potters' (pottery grades), many are cut with silica (maybe just sand, of course), from my experience
with various thermites/analysis. For glazing purposes a bit of the old silica is of course a bonus!
My pride and joy from the last years:
http://www.popsci.com/node/30347
Now I'm off to distill some ammonia. The things you've gotta do!
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blogfast25
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Just finished the 'dry' distillation of ammonia from 0.5 mol garden grade ammonium sulphate and 1 mol household NaOH. Just mixed the crystals and
pellets with a minimum of water into an old vinegar bottle, connected with silicone tubing (RC aircraft fuel lines!) to another vinegar bottle with
about 0.5 L of iced DIW.
Gas starts flowing immediately and the reactor was then heated on steam bath. Amazing how the bubbles of NH3 get gobbled up by the cold water, with no
bubbles surviving to the surface of the scrubber water! After 15 min gas evolution more or less stopped.
Resulting solution of pH about 11.8 and titrated as 2.56 w% NH3, which is about 1.5 M. Yield was thus about 75 %, not bad.
Interesting factoid: the silicone tube sits on a bit of brass tubing (yes, RC aircraft fuel line) that sticks through the cork and into the reactor.
That bit of brass, exposed to moist NH3 had gone blue, presumably due to Cu2+ ammonia complex formation...
About half a mol of this ammonia will now be used for a neutralisation experiment.
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peach
Bon Vivant
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You seem to have those things I brought up under control.
The only thing I'd say is that the jump is indeed important, but it also depends on the solution volumes and concentrations being used. I've seen pH's
jump, sit still, then drift a good way from where they're supposed to be as the solutions continue to react over ten minutes or more, for a few
hundred ml's. With litres, and tens of litres, it can still be doing it days or a week plus later, even with gentle changes. Trying to acidify the
garden at present, it's taking a year plus.
I have seen and used that RC silicon tubing before, it's nice stuff considering it's so readily available.
The brass thing doesn't surprise me at all. Whenever things start getting reactive, gas wise, things happen. 
2007, 20kT of ammonia make a break for it at the Seward plant
The rusty stainless scissors from another post, left poking into the HCl(g) generator overnight, after it was empty;
The sink was splotless, prior to the HCl(g) being around, there are green patches and rust appearing all over it. The new chromed tap is going green
at the stem;
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blogfast25
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Regarding pH jumps, the ammonia titrations yesterday were carried out on undiluted, weighed samples (about 2 g) of the 1.5 M NH3, then diluted a
little for the pH probe. Fine end points, no problem. But every time I read the endpoint at 5.0 to 4.3ish, the 4.5 drifted back up to 5 and
beyond! Stirring made it go back down. I suspect NH3 gas in the headspace being absorbed near end point.
Well, my sink isn't quite as bad as yours but then it really is The Kitchen Sink!
Silicone tubing is fine form for impromptu, al fresco type 'apparati' but it's piss poor with acids: falls a part to almost pure silica very quickly.
Whenever things start getting reactive, gas wise, things happen.
If you can get it to react! 
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blogfast25
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OK. Another single data point for the neutralisation of NH4OH (aq) + HCl (aq) --> NH4Cl (aq) + H2O (l)
330.2 g of a 2.56 w% solution of NH3 (0.496 mol NH3) reacted with 330 ml of a 20.4 w% HCl solution (about 10 % HCl excess). TNH3 (in flask) = 18.5 C,
THCl = 21.0 C, Tend (in flask) = 29.9 C.
Estimated value of ΔH = -58.5 kJ/mol of NH3. So a bit higher than the value for NaOH + HCl but considering my own observed standard deviations on
very similar measurements, here for ammonia the value isn't substantially different from the actual neutralisation enthalpy for 1/2 H3O+ (aq) + 1/2
OH- (aq) ---> H2O (l).
I think I might yet try and include one more measurement, for the neutralisation of disodium citrate (say, Na2HCt): HCt 2- (aq) + OH- (aq) ---> Ct
3- (aq) + H2O (l). The third dissociation content of citric acid is very small...
Four hours later the temperature inside the flask had dropped to 28.5 C, more of a drop than I expected...
[Edited on 25-8-2010 by blogfast25]
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blogfast25
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The trouble with citric acid monohydrate...
A week or so ago I created a master 'pH v volume' titration curve for titration of citric acid with standardised NaOH 0.1 M. I chose a citric acid
concentration of about 0.05 M, hoping to titrate the first two equivalence points corresponding to pKa1 = 3.15 and pKa2 = 4.77. I dissolved about 7 g
(accurately weighed) citric acid monohydrate (CAMH) in 750.0 ml of water.
On titration, the first part of the curve between pH = 2.5 and pH = 6.0 is almost linear and uneventful. The sharp first end point then appeared at
pH about 9.
Problem: I used 20.0 ml of the CAMH solution but needed about 30 ml, just to reach the first equivalence point! That indicated a molarity of the CAMH
solution of 0.15 M, NOT 0.05 M. The titration was repeated a few times, with very similar results.
I checked my calcs and weigh-up and couldn't find fault with it and decided to leave well alone for a couple of days.
Well, today I made a 0.1 M solution of CAMH: 15.76 g of CAMH in 750.0 ml of water. The MW of CAMH is 210.14 g/mol, so (15.76 g x 1 mol / 210.14 g) /
0.75 L = 0.1000 M (assuming the product is 100 % pure of course).
Titrated with 0.1 M NaOH (t = 0.981) just one titration of 20.0 ml of CAMH solution required... 59.9 ml of NaOH titrant solution to reach the first
equivalence point!
This would indicate the CAMH solution was 0.294 M, not 0.1000 M, so again 3 times stronger than planned!
I know some will think I'm getting my normalities mixed up with my molarities but that's not the case: I rarely use normality anyway.
And a 0.1 M CAMH solution would require about the same amount of 0.1 M NaOH to neutralise the first CAMH proton anyway:
NaOH + H3Ct ---> NaH2Ct + H2O
I will titrate 20/3 or about 7 ml, to see if I can see all three equivalence points, or at least the first two...
Totally baffled...
Update:
I'm beginning to wonder whether the fact that the three pKa values of 3.15; 4.77 and 6.40 are really quite close together has something to do with it.
The end-point for 0.1 M solutions are approx. 8.5; about 9 and 9.7! Titrating 5 ml of CAHM I ran right past number 2 and 3!
Compared to e.g. H3PO4: 2.148; 7.198 and 12.375 the CAMH values are really close together. Am I titrating all three protons
together???
And looking closer at the pH curve it's fairly clear that's what's happening. The highest pH point is 10.6 but it looks very much part of the pH jump
for equivalence point 1. Yet 10.6 seems already past EQ2 and EQ3! Bugger! The three protons come from fairly three fairly isolated acid groups and
thus are about equally eager, unlike H3PO4...
[Edited on 26-8-2010 by blogfast25]
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blogfast25
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Coulometric calorimetry (aka Electical Compensation Calorimetry)…
Months ago I bought a simple coulometric calorimeter off eBay, an educational tool by Phillip Harris (so old the documentation mentions a telex number
but no email!)
It consists basically of a 200 ml thermos flask (Dewar bottle) with a 12 V - 21 W light bulb fitted so that it can be immersed into the flask. A
second hole in the rubber bung allows for inserting a thermometer. Basic setup:
Connecting the light bulb to a 12 V DC power supply and either a watt meter or an amp meter and a volt meter allows the power output P = I . U (P =
power (W), I in current (A), U = voltage (V)) to be measured. My two identical box standard multimeters:
I use a 12 V stabilised power source that once belonged to a laptop or a computer game (max. P = 48 W)
Measurement of reaction heats in watery medium (and at atmospheric pressure) can then be made as follows. Assume the calorimeter contains m gram of
solution, with a heat capacity of Cp and that the calorimeter has a heat capacity of k, then if Q Joules of heat were released by a reaction, Q = m Cp
ΔT + k ΔT = (m Cp + k) ΔT.
If we assume that for small values of ΔT, (m Cp + k) = constant or d(m Cp + k)/dT = 0 then if we add more heat (Q’) to the calorimeter by
running current through the light bulb for a known period of time, we know that Q’ = (m Cp + k) ΔT’, with Q’ = I U Δt (Δt time of
running current). It follows that Q = Q’ (ΔT/ ΔT’) or Q = I V Δt (ΔT/ ΔT’). It can be arranged so that
ΔT=ΔT’ but that isn’t strictly necessary.
And if the precise number of moles reacted n is known the Standard Enthalpy of Reaction ΔH at RT can be estimated from ΔH = - Q / n.
Preliminary tests showed the immersed bulb to run at about 12.20 V and 1.70 A, which is indeed about 21 W.
1) Heat of neutralisation of ammonia solution (NH3 (aq)) with citric acid solution:
Using a standardised commercial solution of ammonia, three runs were carried out, each in a total volume of about 150 ml (75 g ammonia solution and 75
g citric acid solution) and neutralising a precisely known amount of NH3 with a small excess of the triprotic citric acid. Values obtained were -
45.3; - 42.9; - 44.3 kJ/mol with average of - 44.2 kJ/mol. So again, despite the combination of a weak acid AND a weak base, not far removed from the
accepted wisdom of -57.3 kJ/mol for strong acid + strong base…
2) Heat of neutralisation of NaOH solution with HCl acid solution:
The first obtained values were much too high and re-standardising showed the HCl solution used was stronger than previously established. The HCl
solution was then meticulously re-standardised and measurements repeated. Except, at the end of the first faulty run the bulb had clapped out! It
turned out to be a box standard car bulb (for the reversing light), so that was replaced and sealed back into place with good quality silicone
sealant.
Then using the newly standardised solutions two runs were made yielding values of - 52.4 kJ/mol and - 54.2 kJ/mol, both somewhat short and not better
than the values obtained with the simpler method described above.
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blogfast25
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Two more runs with HCl + NaOH, this time with stronger solutions (2 M HCl, standardised and slightly stronger NaOH, 75 g of both solution makes for
0.1422 mol of HCl neutralised per run) and slightly longer standing times (before temp. reading). This way the ΔT values has three significant
digits (just over 10 K). Also Q’ was made a littler higher (4’ heating instead of previously 3’).
The two values obtained are - 52.8 kJ/mol and - 54.4 kJ/mol, average - 53.6 kJ/mol, still about 2.1 kJ/mol short of the - 55.7 kJ/mol.
Here’s a resource that provides plenty of ideas for the calibration/standardisation of various methods of calorimetry (pdf):
http://www.iupac.org/publications/pac/2001/pdf/7310x1625.pdf
Most are in the ΔH = +/- 20 kJ/mol range, leading to temperature rises that are too small to be accurately measured with my +/- 0.1 C
thermocouple. Of particular interest could have been the standard molar heat of dissolution at 298 K of KCl(cr), certified by NIST to be + 17.584
kJ/mol (https://www-s.nist.gov/srmors/certificates/1655.pdf?CFID=139...) to a molality of m = 0.111 mol/kg but the estimated ΔT is only about 0.46 K
(17584 x 0.111 / (1000 * 4.19) = 0.46 K)
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blogfast25
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Another run, this time with home brewed HCl. The commercial HCl (23 w%) I’m using is yellow with Fe3+ contamination. To eliminate it as a possible
cause of error, some 90 ml of approx. 2 M HCl were concocted using a basic HCl generator. 15 g technical CaCl2, 50 g H2SO4 (94.5 w%) in the right hand
bottle (on steam bath) and 90 ml of deionised water in the left receiver bottle:
After standardising the obtained HCl solution, yield of HCl based on:
CaCl2 (s) + 2 H2SO4 (l) === > 2 HCl (g) + Ca(HSO4)2 (s)
… was found to be about 78 %.
Using this solution, a calorimetric experiment was then conducted using 0.1872 mol HCl and a small excess of NaOH, in a total mass of about 150 g of
solutions.
One single run gave a value of - 53.7 kJ/mol heat of neutralisation of HCl solution with NaOH solution. This more or less excludes the iron
contamination as a major source of error, as expected.
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