Sciencemadness Discussion Board

Titration troubles

Geocachmaster - 26-12-2016 at 19:44

I did my first actual titration today (using proper technique at least, I got a burette for Christmas!:P), and something went horribly wrong.

I bought a 1L bottle of sulfuric acid drain opener at my local hardware store. Not knowing it's concentration, I decided to titrate it. I prepared my titrant by dissolving 3.82g of sodium carbonate in ~15ml of distilled water. One all of the solid had dissolved, I topped it off to 20ml in a graduated cylinder to make a solution that was 1.80 molar. I then prepared the analyte by measuring out 1ml of the sulfuric acid drain cleaner and dissolving it into ~10ml of cold water so the reaction with the carbonate would not be so vigorous. Next, I loaded 10.0ml of the titrant into a burette and added a few drops of phenolphthalein indicator solution to the analyte. I slowly began to add the titrant while swirling the flask and eventually added all 10.0ml. I then refilled the burette and added another 6.4ml before a light pink color emerged and stayed in the flask even with swirling. This brings the total volume of titrant used up to 16.4ml.
Note: my sodium hydroxide is 6 months old so I decided not to use it for the standard solution; it could have a lot of water and sodium carbonate in it.

Now we can do calculations:
1.80 moles a liter x .0164 liters of titrant = .02952 moles of sodium carbonate used
Na2CO3+H2SO4->Na2SO4+CO2+H2O is the balanced reaction taking place
So there were .02952 moles of sulfuric acid in 1ml
.02952x1000=29.52 molar sulfuric acid!
29.52x98.079g/mol=~2.9Kg

One liter of sulfuric acid definitely does not have a mass of 2.9Kg:o. Did I do anything wrong, how can this happen???

Thanks, Tom.

Texium - 26-12-2016 at 20:43

Titrations with sodium carbonate tend to be tricky. I had similar problems when I attempted to use it as a titrant a while back. I think part of the problem is that carbon dioxide can remain in solution which makes the solution stay acidic slightly past the point where you should have stopped titrating.

A bigger factor could be that your sodium carbonate was not fully dried. I assume you dried it in an oven prior to using it, but it can absorb almost twice its molar mass in water as the decahydrate! Whenever you use it in a quantitative situation, make sure to oven dry it, allow it to cool in a desiccator, and then use it fresh.

JJay - 26-12-2016 at 21:04

I've performed a very similar procedure before with no problems, so I'm not sure... a few things, though: you shouldn't count on your burette's absolute measurements to be very precise; rather, you want to look at the differences between measurements. Is your sodium carbonate anhydrous? Is your glassware impeccably clean? Also, you should use only one drop of phenolphthalein with such a tiny amount of solute. Make sure you add the titrant very slowly, especially close to the endpoint - you don't want unreacted sodium bicarbonate and sodium bisulfate to ruin your day. Also, you'll get more precise measurements if you use volumetric glassware at the right temperature, but that's not a source of problems here.

You might also want to try titrating a somewhat larger amount of the acid and/or doing a linear regression with several titrations of different concentrations of solution.

aga - 27-12-2016 at 00:34

Quote: Originally posted by Geocachmaster  
I did my first actual titration today

Nice work !

Getting a good result with titration can be difficult.

Your maths look ok, although 98% sulphuric acid should be around 18[M].

If your sodium carbonate was not dried, it commonly comes as Na2CO3.10H2O (the decahydrate) with a M.W. of 286, the water forming about 63% of that mass.

Adjusting the weight of 3.82g down to 37% = about 1.41g.

Plugging that into the rest of your calculations makes the acid about 11 [M], which is more reasonable.

Titration solutions tend to be around 0.1[M] or less, which helps with the accuracy/reduces the effect of measuring errors.

For example, if your acid sample had been just 0.1ml more (1.1ml) you'd get bang on 18[M] as a result after accounting for the 10H2O.

It's also good to titrate at least 3 times and average the results.

P.S. You probably did well to NOT choose NaOH because if you leave it in a burette overnight, it sticks the valve like superglue ! Lost at least 1 burette that way.

JJay - 27-12-2016 at 01:23

Quote: Originally posted by aga  


For example, if your acid sample had been just 0.1ml more (1.1ml) you'd get bang on 18[M] as a result after accounting for the 10H2O.



Huh? How do you figure?
Quote:

One liter of sulfuric acid definitely does not have a mass of 2.9Kg:o. Did I do anything wrong, how can this happen???


I just noticed that if you happened to use sodium bicarbonate instead of sodium carbonate to do the titration, you'd get very similar results to the ones you observed. Rather good quality sodium carbonate monohydrate is available at most grocery stores at low cost as Arm & Hammer Washing Soda ("the standard of purity").

aga - 27-12-2016 at 01:39

Quote: Originally posted by JJay  
Huh? How do you figure?

Very Badly !

I used the weight of the water instead of the weight of the carbonate ... :(

[Edited on 27-12-2016 by aga]

diddi - 27-12-2016 at 03:00

I am worried by the small quantities used. measuring 1ml accurately is not easy unless you have a proper 1cc pipette. let us have an error of +- .1 ml which for your first titration is probably being generous on the accuracy side. that gives rise to 10% error. if you had made up 10ml of acid into a 500 volumetric flask, and taken a sample aliquot of say 10ml, your error would be somewhere around 1%

then there are the errors in your burette skills. with concentrated titration solutions, small errors and over shoot become enormous. as you reduce the concn, the errors reduce. the endpoint is easier to find and life is easier.

there is also the problems with your suspect hydrated reagent as stated above.

my thoughts: dilute your acid and find a better standard solution. fresh clean dry NaOH is a good choice to use to make a standard solution. and do several runs. for practise you could do 5 times. we look for 'concordant' results. ie results that are the same or very similar. so repeat until you get 3 results within some limit. depending on skills you can make the limit .3ml or .2 or .5 or whatever. any that are outside your range are excluded in calculations.

JJay - 27-12-2016 at 04:15

You could measure out several aliquots of acid using your buret, carefully recording the size of each. It's considered acceptable to estimate the last significant digit, so you should be able to measure to a hundredth of a milliliter if your buret has 0.1 mL graduations. That will easily give you 3 significant digits.

I've run into difficulties doing complex iodometric titrations (which I now believe were caused by my starch suspension)....

I was thinking about buying myself a buret for Christmas....

[Edited on 27-12-2016 by JJay]

unionised - 27-12-2016 at 04:44

I suggest making some fresh, pure sodium carbonate by heating bicarbonate of soda in the oven at about 200C until it has lost all the CO2 and H2O
That way, you know what you are actually weighing out.
Vigorous stirring will help remove CO2 as the titration takes place.

It's also a good idea to make sure that you use most of the volume of the burette in a titration to get the best accuracy.

Neme - 27-12-2016 at 04:51

From what I remember, phenolpthalein is best for titrations with strong base for determinating both weak and strong acids. I'd say for strong acid and weak base would be methyl orange or methyl red better.

Also, best way to remove the carbon dioxide is heating the titration flask.

brubei - 27-12-2016 at 04:58

strong base have to be titrated with strong acid (NaOH) with micromolar phenolphtalein. Commonly by dropping your H2SO4 diluted ten time(#) in 50mL 0.1M NaOH (and 2 or 3 droplet of phenolphtalein). Color change from full pink to incolor near pH = 7 indicate equimolar concentration of NaOH and H2SO4 where:
[H3O+]=[OH-]
C(H2SO4) * V(H2SO4) = 2 (C(NaOH * V(NaOH))
C(H2SO4) = 2((C(NaOH) * V(NaOH) / V(H2SO4))

(#)try to prepare a dilution where your acid concentration will be approximatly 0.1M otherwise your titration will be less accurate if your acid is to much concentrated or too loooong by addind to much volume of overdiluted acid.


[Edited on 27-12-2016 by brubei]

Success!

Geocachmaster - 27-12-2016 at 12:22

Using the advice of several helpful members, I did another two titrations.

This time I used 5ml of acid for each, and I baked the sodium carbonate in an oven for an hour at a little over 200C. There is 90% humidity here right now, so after baking I allowed the carbonate to cool in a large mason jar dessicator over anhydrous calcium chloride. I then prepared the titrant at half the concentration of my previous attempt for more accuracy. Also, after approximately 70% of the theoretical amount of titrant had been added, I added broken glass and began to heat the sulfuric acid solution between addictions (~8ml each) until bubbling ceased. This way, most of the CO2 would be driven off and could not affect the pH. Close to the end, I would heat every 2ml or so.
After the titrations I calculated the drain cleaner to be 17.642 molar for the first and 17.29 molar for the second one. The average of these comes out to the drain cleaner being 93.5% sulfuric acid, by weight. From what I've read, around 93% is typical for many acids, so I am very pleased with the results this time.
My skills can still use some work, so I think I'll practice more, but I am very happy. It's so nice when the numbers work out right:D.

Thanks for all the help, Tom.

aga - 27-12-2016 at 12:26

Quote: Originally posted by Geocachmaster  
Using the advice of several helpful members, I did another two titrations

Congratulations !

Nice work there.

Great to see someone getting the hang of things : you no longer have to rely on what it says on the bottle because you can now measure it for yourself.
Quote: Originally posted by Geocachmaster  
around 93% is typical for many acids

That only works for OTC concentrated sulphuric acid.

Other acids will be different : titrate some and find out !

(and please let us know)

[Edited on 27-12-2016 by aga]

diddi - 27-12-2016 at 18:37

what was the difference in you titration volumes? repeat again to get another value which will be closer to one of the first 2, thus you can improve your value by eliminating the odd one out