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Acid strength test?

CitricAcid - 15-12-2015 at 19:57

According to a video by NurdRage on making chloroform, chloroform is decomposed by sodium hydroxide to give formic acid. I want to try and produce some formic acid that way,but is there any way to test the concentration of the acid? I know that a good way to test to see whether or not sulfuric acid is concentrated to 98% is to add it to sugar and see whether or not it turns to carbon,but is there any way to test to see whether or not formic acid is concentrated to 85%?

Detonationology - 15-12-2015 at 20:16

Quote:

...the chloroform produced decomposes under strongly alkaline solutions... chloroform reacts with sodium hydroxide to produce formic acid and sodium chloride.

I haven't quite found any data on this quite yet, but I can definitely recommend a method of producing formic acid from anhydrous glycerin and oxalic acid. Here is a link,and TheChemiKid has made a video of this process. Here is an extremely simplified version of the scheme. In step 1, oxalic acid and glycerin react at elevated temperatures to yield glyceryl monoxalate and water. In step 2, the oxalate part of the compound is decarboxylated to yield glyceryl monaldehyde and CO2. The final step is the hydrolysis of the aldehyde to yield formic acid and glycerin, which is regenerated from this process.

[Edited on 12-16-2015 by Detonationology]

Ozone - 15-12-2015 at 21:36

I imagine this would proceed via dichlorocarbene (look up hazards of this). Look out for carbon monoxide and the possibility of phosgene.

I'd be very careful when trying this out. Carbenes are frisky.

O3

CitricAcid - 15-12-2015 at 21:49

Quote: Originally posted by Detonationology

Quote:

...the chloroform produced decomposes under strongly alkaline solutions... chloroform reacts with sodium hydroxide to produce formic acid and sodium chloride.

What is that at the end of the reaction scheme? The formula for formic acid is CH2O2.
Nevermind,I looked it up,the results came back with formaldehyde,I think you made a mistake there

. Thanks for drawing the sequence out though.

[Edited on 16-12-2015 by CitricAcid]

CitricAcid - 15-12-2015 at 21:51

Quote: Originally posted by Ozone

I imagine this would proceed via dichlorocarbene (look up hazards of this). Look out for carbon monoxide and the possibility of phosgene.

I'd be very careful when trying this out. Carbenes are frisky.

O3

How would I get adequate amounts of dichlorocarbene?
"Look our for carbon monoxide and the possibility of phosgene"
I'm assuming it reacts with oxygen in the air to produce phosgene,is there any other alternative?

[Edited on 16-12-2015 by CitricAcid]

UC235 - 15-12-2015 at 22:51

Chloroform does indeed react with aqueous sodium hydroxide. In the presence of a secondary solvent like methanol so that the reaction mixture is homogenous, this reaction can get rather agressive as it is very exothermic.

The reaction proceeds by intermediacy of dichlorocarbene. This is not isolatable. The hydroxide deprotonates chloroform to trichloromethyl anion, which eliminates Cl- to charge-neutral but very reactive dichlorocarbene. This inserts itself into a water O-H bond forming "dichloromethanol" which eliminates HCl to formyl chloride and is hydrolyzed and deprotonated to formate.

As a source of formate, the whole process is massively inefficient. 79ml (119g) of chloroform requires 160g of sodium hydroxide for complete reaction. After boiling down the solution and discarding 174g of NaCl, you would get 68g of sodium formate (in a perfect world, which this is not).

As for turning it back into formic acid, you can't use sulfuric acid unless you want to die of CO inhalation. I'm not even sure what the best approach would be. Phosphoric acid, maybe. The acid forms a high boiling 77% azeotrope at 107C which will be functionally impossible to sepatate from water using distillation on a lab scale. Not counting water, a perfect yield would be 46g. I would be stunned to get half that.

The process using glycerol and oxalic acid posted above will be far simpler and higher yielding.

If you want to know the concentration of formic acid, you could try to find a density table and take very accurate density measurements at the reference temperature. But really, you should learn how to titrate for a definitive answer.

[Edited on 16-12-2015 by UC235]

Detonationology - 16-12-2015 at 04:53

Quote: Originally posted by CitricAcid

What is that at the end of the reaction scheme? The formula for formic acid is CH2O2.
Nevermind,I looked it up,the results came back with formaldehyde,I think you made a mistake there

. Thanks for drawing the sequence out though.
[Edited on 16-12-2015 by CitricAcid]

Thank you for pointing that out. We can all get ahead of ourselves sometimes. The product of this reaction is indeed formic acid, I just forgot one measly O at the end. That's all it takes to screw you in chemistry, I guess.

ave369 - 16-12-2015 at 08:58

For sulfuric acid the test is,whether it carbonizes sugar. For nitric acid - whether it sets a narrow strip of rubber on fire. For the other acids, no strength tests exist.

Boffis - 16-12-2015 at 09:32

Somewhere on the SM Forum I have already posted a patent (try US 3718545) about the azeotropic concentration of formic acid to >98% using heptane or petroleum ether of the correct boiling range. I didn't try it because I found an OTC supplier of 98% formic acid and its cheap.

aga - 16-12-2015 at 14:17

Is there no Titration possible, seeing as formc acid is, erm, an acid ?

Detonationology - 16-12-2015 at 14:27

Quote: Originally posted by aga

Is there no Titration possible, seeing as formc acid is, erm, an acid ?

I think this on the right track...

aga - 16-12-2015 at 14:44

For those that don't know, Titration is really useful and lets you find out how 'strong' an acid or base is by using an indicator, like phenolpthalein.

In Summary :-

you stick a measured quantity of your formic acid in a beaker, make an exact concentration of say NaOH solution (or some base that reacts in a known ratio with your acid), add a coloured acid/base indicator to the acid in the beaker, then drip the basic solution into the beaker, drop by drop, until the indicator changes colour.

That is what a Burette is for.

From the volume of the base added you can calculate how many moles of acid there were.

From that and the volume of acid you started with, you can calculate the original concentration (viz Strength) of your acid.

It's actually a lot easier to do than describe.

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Quantitativ...

Idea

CitricAcid - 16-12-2015 at 14:50

Seeing as sulfuric acid reacts with formic acid to produce carbon monoxide,a gas,couldn't an experiment be done that determines the amount of gas produced where sulfuric acid of ,say,98% concentration is added to an equal volume of 98% formic acid? Then,the data from the "control" experiment can be compared to data from experiments where an unknown concentration of formic acid is used with 98% sulfuric acid?
Edit:I forgot, but for the discussion,let's say that 10Ml of 98% formic and sulfuric acid is used,both for the control experiment and for the formic acid with an unknown concentration.

Let's say that I find the amount of gas produced with 98% formic and sulfuric acid,which is represented as X,then I test a sample of formic acid with unknown concentration using 98% sulfuric acid,and the amount of gas produced is Y.

Y/X= Concentration of unknown acid sample (Maybe?).

Of course, the measurement will be off by a few %s since 98% acid was used in the control experiment rather than 100%,but I think the plan is pretty sound.

I got the idea after doing some reading about a setup where hydrogen peroxide and household bleach are used to determine the concentration of the NaClO in the bleach,let me know what you think.

[Edited on 16-12-2015 by CitricAcid]

aga - 16-12-2015 at 14:59

CO is not nice, and gasses are harder to measure than dissolved chemical species in a solution.

When in solution, they don't escape so easily.

CitricAcid - 16-12-2015 at 15:09

Quote: Originally posted by aga

CO is not nice, and gasses are harder to measure than dissolved chemical species in a solution.

When in solution, they don't escape so easily.

Let's say that the control experiment is done in a round bottom flask with vinyl tubing leading into an upside-down graduated cylinder filled with water. By "Amount" I meant the volume of carbon monoxide gas that is made.
Edit: I'm talking about a setup similar to what the maker of the linked video does.
https://www.youtube.com/watch?v=j-PrAczOGb0

[Edited on 16-12-2015 by CitricAcid]

aga - 16-12-2015 at 15:23

Does Titration sound too difficult ?

CHRIS25 used syringes 'cos he had no burette.

Got pretty close to 0.1 accuracy if i recall correctly ...

DraconicAcid - 16-12-2015 at 15:29

And if you can't get phenolphthalein, you can use cabbage juice as an indicator.

aga - 16-12-2015 at 15:41

RED cabbage juice.

Tut Tut Draconic.

Not all Cabbages are created alike.

Edit:

The juice from a red cabbage is a remarkably good pH indicator though.

[Edited on 16-12-2015 by aga]

DraconicAcid - 16-12-2015 at 15:45

Quote: Originally posted by aga

RED cabbage juice.

Tut Tut Draconic.

Not all Cabbages are created alike.

Good point, but I'm not an expert in cabbages. I mean, ew.......they're cabbages.

aga - 16-12-2015 at 15:49

Red cabbages should always be on hand for eating, chemistry, dye, offensive flatulence and even as a weapon if you run out of cannon balls.

Cannot ever have enough red cabbages IMHO.

Ozone - 16-12-2015 at 17:31

The best thing would be HPLC. A BioRad Aminex HPX-87H (300 mm) with a RID, eluted with 0.008N H2SO4 at about 0.6 mL/min will give you formic at about 14 minutes (linear up to about 2000 ppmw, LOQ ~40 ppmw, on-column). Oxalic elutes much earlier.

Barring that, the pKa is 3.75, so the end point would probably be somewhere around 6. So, either methyl red or bromothymol blue should give end points that are pretty close.

With red cabbage juice, you could shoot for purple (lavender, actually, which is the problem), but I've found end points to be hard to really nail with this indicator (there are just too many of them). Also, you have to use it fresh. It tends to turn into brown muck.

If you have a pH meter, though, you're good to go. Just put the probe in there and titrate by volume (say 0.1 N NaOH) and plot the pH vs titrant as you go. The center of the plataeu should be very close to 3.75. The endpoint should also be obvious.

If your resolution is fine enough (as your patience allows), you may also see additional features that will indicate other acidic protons (mixed acids).

If time permits, I'll titrate some tomorrow, for reference.

O3

[Edited on 17-12-2015 by Ozone]

j_sum1 - 16-12-2015 at 17:37

A seasonal alternative to red cabbage juice:
http://www.compoundchem.com/2015advent/2015advent17/

I had no idea.

Ozone - 16-12-2015 at 17:47

Stillage from Elderberry wine, adjusted to various pH, for your enjoyment.

ave369 - 17-12-2015 at 03:59

Guys, we have a whole wiki article about DIY pH indicators!

http://www.sciencemadness.org/smwiki/index.php/Anthocyanin

More to the topic, I have access to heptane and no use for it except stove lighting aid. I mean, of course, distilling Et2O out of Quick-Start car sprays. I may try the experiment with heptane and formic acid.

Concerning titrations, a Janet syringe (a giant syringe which is easy to find in drugstores and medical supplies shops) makes a fine burette if you don't have a real one. Of course, this is a better option, but not the only one: I titrated HCl with normal syringes, filling them with the titrant and emptying them several times.

[Edited on 17-12-2015 by ave369]