I want to use sodium carbonate as a primary standard to determine the molarity of some hardware-store muriatic acid. All of the references I've found
say that the second equivalence point for the titration of sodium carbonate is around a pH of 3.7 and that an indicator appropriate for that pH, such
as methyl orange, or a pH meter should be used to determine the endpoint of the titration.
But when you add, for example, two moles of HCl to a solution of containing one mole of sodium carbonate, you end up with a solution that has some
dissolved CO2 and some Na+ and Cl- ions.
2HCl + Na2CO3 --> 2NaCl + CO2 + H20
By the end of this titration, you are not really titrating the original carbonate ions, you are titrating carbonate ions that were created during the
titration reaction.
CO2 + H2O --> H2CO3
To further complicate the situation, some sources call for boiling the solution when it gets close to the second equivalence point so as to drive off
the CO2 from the titration reaction and then finishing the titration at an equivalence point of around pH 4. Here's an example, starting on page 22:
But once you drive off all the CO2, aren't you just left with a solution of salt water with a pH of 7? (Or maybe a little lower due to CO2 from the
atmosphere?) And so, if you boil away the CO2, wouldn't you have a final equivalence point of around pH 7.
I am confused by all of this. Can anybody point me to a source that deals with my confusion? Thanks!
happyfooddance - 18-2-2018 at 17:02
The difference between a pH of 7 and a pH of 4 is only going to be a drop or maybe two of HCl, and likely won't affect the precision of your titration
as much as, say, the innacuracy of your balance.
I bet if you were to try both methods you outlined above (pH 4 endpoint, and boiling to drive of CO2) you would get the same results if your technique
is okay.
[Edited on 2-19-2018 by happyfooddance]j_sum1 - 18-2-2018 at 17:28
Your bigger problem will be your source of carbonate -- do you have anhydrous or the decahydrate? How much moisture has it absorbed from the
atmosphere? And how would you know?
Your best bet is to begin with sodium bicarbonate and heat it to produce your own fresh carbonate. This is ok as a primary standard.
2NaHCO3 = Na2CO3 + CO2 + H2O
Use this to mix up your standard solution. For titration you want it pretty dilute.
Your equation for the titration is correct:
2HCl + Na2CO3 = 2NaCl + CO2 + H2O
You are also correct to question the role of the CO2 produced: there are a couple of things that can happen to it:
1. It can escape as a gas.
2. It can remain dissolved in the solution as CO2(aq)
3. It can participate in an equilibrium reaction with the water:
CO2(aq) + H2O(l) <=> H2CO3
and
H2CO3 <=> 2H+(aq) + CO32–(aq)
This second equilibrium is not really much of a concern: it is pushed to the left by the presence of carbonate ions already present. And when they
are exhausted you have excess H+ ions appearing also pushing the equilibrium to the left.
This makes sodium carbonate a weakish base. (I would say weak but it is not that weak really. pKa=10.33) You would therefore follow standard
procedures fro titrating strong acid/weak base: employ an indicator with a transition pH slightly acidic: say, pH=6. Methyl Red would be good.
If all you have to work with is phenolphthalein then you won't be too far off. You'll get a slight overestimate of the molarity of your acid. Maybe
by 1-2%.JJay - 18-2-2018 at 19:56
happyfooddance, thanks for your advice. I will try the titration both ways to see if there's much difference.
j_sum1, I have a sample of reagent-grade sodium carbonate, but it's well over 20 years old. I also have some fresh Arm & Hammer Washing Soda,
which claims to be 100% sodium carbonate. I'll try them both to see if there's a difference. My plan is to heat the sodium carbonate at 500F for an
hour to convert it from the decahydrate to the anhydrous and to keep heating if necessary until I get a constant weight. Since baking soda is so easy
to get, I'll also try heating that to get the sodium carbonate.
"ARM & HAMMER™ Super Washing Soda is 100% sodium carbonate and it is used as a laundry booster and general household cleaner. ARM &
HAMMER™ Baking Soda is 100% sodium bicarbonate and has a myriad of household cleaning, personal care, and deodorizing uses, as well as being a
leavening agent. It is important to note that these two are two very different products and cannot be substituted for one another."
JJay, thanks for the Journal of Chemical Education reference. I live near a university that offers public access to most science journals and I will
download the paper this week.
Sulaiman - 19-2-2018 at 10:14
Before you start, decide what level of precision and/or accuracy you require.
If +/- 1% is good enough (it usually is) then just pretend that there are no 'complications' and you should be ok.
I have had difficulty (aka failure) titrating simple mineral acids vs. common bases then back again to 0.1% consistency.
- everything gets complicated.
. purity of reference
. accuracy/precision of scales
. weight compensation for bouyancy
. cleanliness of vessels (no drops adhering to walls)
. precision of burette or similar
. burette compensation for temperature
. with good stirring
. liquid temperature pH compensated for temperature
. absorbtion of atmospheric water or carbon dioxide during measurements
. most importantly of all, before starting, do not forget toCarbon8 - 19-2-2018 at 17:47
Sulaiman, thanks for your tips. And what is the most important thing not to forget?j_sum1 - 19-2-2018 at 18:12
Getting anh sodium carbonate from NaHCO3 is a whole lot easier than from Na2CO3.10H2O + moisture. I would go straight for the bicarbonate.
TBH, most of my titrations are rough as guts sloshing with measuring cylinders to give me an approx value to within 5%. It's been a while since I did
an accurate one. I usually don't need the concentration of my HCl that accurate. I guess, work out what you actually need to know.Sulaiman - 20-2-2018 at 00:57
For sodium carbonate sources it depends upon your location, for example;
eBay .uk has 99.7% anhydrous sodium carbonate 500g, £3.99 incl.p&p
eBay .au seems to have only imported sodium carbonate of unspecified purity at stupid-high prices.
So where j_sum1 would start with bicarbonate, I would start with the carbonate.
(unless I'm in a hurry, in which case I'd use bicarbonate from the kitchen if I have no stock)
Again, consider the requred result,
If I use 99.7% pure sodium carbonate I should not expect results to 0.1% accuracy.
(so maybe the cheaper 99.3% technical grade is sufficient for +/- 1% accuracy)
I could re-crystalise and dehydrate to re-purify
... but I can't be sure how pure, or even if it is purer than the starting material.
Or I can bite the bullet and buy some realatively expensive (£9/100g) >99.9% anhydrous sodium carbonate
- which begins to loose its purity as soon as I open the container.
Some experiments require near stoichiometric mixtures of reactants - probably,
most experiments that I have done usually involve adding a reactant until a certain observable point (e.g. 'no more fizzing')
and/or an excess of one or more reactants,
so unless you are into quantitative chemistry,
precise concentrations are rarely required.
I'm fairly confident that other than for deliberately quantitive analytical experiments,
+/- 1% is MORE THAN accurate enough for almost all hobby purposes.
BUT
sometimes it does matter what the impurities are,
e.g. 1% sugar in salt would not be as noticeable as 1% salt in sugar.
My attempts at maintaining 0.1% absolute accuracy whilst titrating three acids vs. three bases failed,
I am however now confident to better than 0.25%, which is more than accurate enough for my needs.
P.S. Carbon8 - I forget
[Edited on 20-2-2018 by Sulaiman]JJay - 20-2-2018 at 01:03
Triple-recrystallized and calcined sodium carbonate is typically extremely pure according to Perrin's Purification of Laboratory Chemicals. I think
you can find a PDF of an older edition on archive.org.Vomaturge - 20-2-2018 at 16:00
This is a bit off topic, but why is it that acids are titrated with sodium carbonate, but not sodium bicarbonate? Is it because the bicarbonate is not
basic enough?
Also, suppose you just added a measured sodium carbonate solution to the HCl until CO2 gas was no longer observed. Would it stop bubbling at the
equivalence point? Would this technique be able to measure an acid's concentration to within 5%? I know that part of the goal of a titration (or any
measurement, really) is to be accurate, but I can totally see a quick-and-dirty estimation being useful in many cases.
I know these are beginner questions, but they seem related enough I chose to put them here.MrHomeScientist - 21-2-2018 at 06:49
You should also be aware that when companies say their product is 100% "x", it doesn't mean the purity of "x" is 100% (since that's impossible). It
just means the only ingredient is "x", and wherever that came from may have "built-in" impurities. A good example is 100% acetone nail polish remover
- there's only acetone in the bottle, but that acetone is contaminated with several percent water and likely some other trace chemicals depending on
how it was made. None of that matters for it's intended use, but it might make a difference with chemistry. Always look up the SDS for the product
you're getting chemicals from; that should give you a slightly better idea of purity. Even the SDS won't list every impurity, typically.Carbon8 - 21-2-2018 at 11:42
Vomaturge,
Strong acids are often titrated with sodium carbonate because sodium carbonate is considered to be a "primary standard," which means that it can be
conveniently obtained in a highly pure form that is also very stable. Another important characteristic of a primary standard is that it not absorb
water (or, if it does, that the water can be conveniently driven off). The reason for this is that when you use a substance as a primary standard,
it's very important to know as accurately as possible exactly how much of the substance you have in a sample. So, when you measure out a certain
amount of anhydrous sodium carbonate, you want to be certain that there are no water molecules in the sample. Here are two links:
Sodium carbonate is a good primary standard because it can be easily obtained in a pure form, is unreactive as a dry powder, and can be conveniently
dried to the anhydrous form in an hour in an oven heated to 200C.
Sodium bicarbonate is not a good primary standard because it is hydroscopic, and thus different samples that contain exactly the same number of moles
of bicarbonate can contain different amounts of water and therefore weigh different amounts. And it's difficult to drive off the absorbed water so as
to dry sodium bicarbonate to an unchanging weight, because when sodium bicarbonate reaches a temperature of between 50 and 80 degrees C, it begins to
decompose into water, carbon dioxide and sodium carbonate.
This decomposition reaction is the reason that j_sum (above) suggested obtaining a pure sample of sodium carbonate by starting with sodium bicarbonate
and heating it up to convert it to sodium carbonate, which can then be dried to a constant weight and used for the titration.
By the way, when you use sodium carbonate to titrate a strong acid, there are two equivalence points. The first equivalence point is when all of the
carbonate ion has been protonated to the bicarbonate ion and the second equivalence point is when all of the bicarbonate has been protonated to the
carbonic acid form. So, another answer to your question is that acids ARE titrated with sodium bicarbonate during the second part of a sodium
carbonate titration.
As for your second question, about using the end of the bubbling of CO2 as a rough indicator of the second equivalence point, I think that might work.
One way to know for sure is to do the bubble titration in parallel with a normal pH titration and see if they agree.Vomaturge - 21-2-2018 at 12:19
Thanks! That was the answer I was looking for. I was just making sure I'm not missing a key concept on the reaction of bicarbonates with strong acids.
It sounds like nothing is intrinsically wrong with the titration reaction itself when using bicarbonate. However, since baking soda absorbs water and
cannot be reliably dried without decomposition , it would be hard to get a solution of known concentration. That's okay, carbonate is easy enough to
make from bicarbonate.
I'll have to try and compare a pH measurement titration with a gas-release titration. That's the scientific way to find out happyfooddance - 21-2-2018 at 12:36
Sodium carbonate is a good primary standard because it can be easily obtained in a pure form, is unreactive as a dry powder, and can be conveniently
dried to the anhydrous form in an hour in an oven heated to 200C.
I would say, dry in an oven until DRY, an hour might not come close to doing it, especially if you don't know which hydrate you're working with. This
is the reason j_sum suggested starting with the bicarb. The crystal lattice more quickly and reliably dehydrates than some carbonate hydrates. There
are ways around this, of course (think ball mill), but bicarbonate is faster all around. Drying in an oven at 200 C° 'might' work, but I would rely
on a constant weight.Pumukli - 21-2-2018 at 13:26
Potassium hydrogen carbonate is a good primary standard so you can titrate strong acids with this compound too.
Unfortunately it is not sold OTC hence making/purifying sodium carbonate is the amateur friendly route.JJay - 21-2-2018 at 17:43
Potassium hydrogen carbonate is sold by some brewing stores, but it is a little hard to find. Potassium carbonate is more common.
Oh, and sodium carbonate monohydrate is the common commercial form.
This paper details an easy procedure for preparing a sodium carbonate primary standard, explaining how to ensure that it is dry and so forth.
Attachment: 1803x0443.pdf (240kB) This file has been downloaded 334 times
[Edited on 22-2-2018 by JJay]byko3y - 22-2-2018 at 02:19
If you don't drive away carbon dioxide, then you will get a carbonate buffer which slowly changes its ph from 8 to 4 during the titration. Ignoring
the second dissociation of carbonic acid we get:
Ka(NaHCO3) = [H+]*[HCO3-] / [NaHCO3]
[H+] = Ka(NaHCO3) * [NaHCO3] / [HCO3-]
ph = pKa(NaHCO3) + p[HCO3-] - p[NaHCO3]
First pKa of sodium bicarbonate is 6.4.
1 mM of carbonic acid + 10 mM of bicarbonate will give you ph = 6.4 + 3 - 2 = 7.4
5.5 mM of carbonic acid + 5.5 mM of bicarbonate - ph = 6.4
10 mM of carbonic acid + 1 mM of bicarbonate - ph = 6.4 + 2 - 3 = 5.4
ph of 10 mM solution of carbonic acid is approx 4.5.
If i had to perform similar titration I would detect transition at pH = 3.5-4.0, e.g. methyl yellow, methyl orange, ethyl orange. bromophenol blue,
congo red. Those require calibration if you need high precision. If you have a ph meter, then you can record a full titration curve and then deduce
whatever you want.
