Sciencemadness Discussion Board - Boiling down Sulpuric Acid

Boiling down Sulpuric Acid

Trizocy - 16-2-2015 at 05:26

Hey guys! new guy here so hope i can learn from the forum here

Anyway, i just bought a flask of Drain cleaner which at the back stand 60-100% Sulpuric Acid, and i was wondering if its possible to boil away the water content inside it?

Goal is not to reach 99% (if its simple, then yes) but to reach a % that i know whats in the bottle

because when im working know, i have no idea if its 60% 80% or 100% inside the bottle....

Thanks!!

woelen - 16-2-2015 at 05:49

What color does it have? It should be colorless, or maybe pale brown. Some very low grade drain cleaner acid is black like soot. I consider that useless for chemistry experiments.

Normally, drain cleaner is highly concentrated acid, 92% or better, but not more than 98% (higher than 98% H2SO4 is very rare and certainly not something you will find in a hardware store). Anything at 90% or better can be considered concentrated acid.

Some tests you can do:
- Add a few drops of acid to paper tissue. Does the tissue char and do you get holes in it quickly? If so, then it is well above 90%.
- Add a little acid to table salt. Put a teaspoon full of table salt on a piece of glass and then pour 1 ml of acid on it. Do you see slow bubbling and a lot of dense white fumes? If so, then the acid is well above 90%.

Concentrating sulphuric acid can be done, but you need heat-resistant lab glassware, a glass-walled thermometer with a range up to 250 C, and a suitable hot-plate. You also need to do this outside. Concentrating only makes sense if the above two tests fail. Heat the acid in a glass beaker and put a thermometer in it. Keep on heating until it reaches 200 C or a little bit above that. At that temperature the acid will not simply develop water vapor, but a more distinct dense white smoke. DO NOT INHALE THIS SMOKE, IT IS EXTREMELY CORROSIVE!

Do this experiment on a breezy day, with wind from behind and assure that if the beaker cracks then no disaster occurs to you or your property (your heater will be ruined of course). Protect yourself, the very hot concentrated acid is extremely corrosive and it will burn deep holes nearly instantly in everything! Know what you are doing!

AJKOER - 16-2-2015 at 06:24

I remember an old chemistry text remarking that heated MgSO4 till its red is a substitute for H2SO4 as the anhydrous salt decomposes at 1,124°C. My take on the thermal decomposition:

MgSO4 ---Heat---) MgO + SO3 (g)

You could investigate the use of other sulfates as well, but be mindful that the SO3 vapor is extremely dangerous as it will react with water vapor (or the moisture in your lungs) to form some extremely concentrated Sulfuric acid:

SO3 + H2O ---) H2SO4

Best to rethink what you are attempting even via this possibly marginally safer route to pure concentrated Sulfuric acid (I would also recommend you read some SM's thread of members badly disfigured by acid burns).

[Edited on 16-2-2015 by AJKOER]

blogfast25 - 16-2-2015 at 06:36

AJ, behave! There is nothing safe or easy about decomposing MgSO4. And SO3 is notoriously difficult to dissolve in water, which is why they dissolve it in conc. H2SO4 making oleum and then dilute the oleum.

Concentrating 60 % H2SO4 is not easy though.

Trizocy - 16-2-2015 at 06:40

Hey thanks woelen for fast respond!

Well took a picture of it: http://postimg.org/image/y0nkmq3o9/ (little brown-orange-clear color)

Reaction of tissue : http://postimg.org/image/3qzy5fgnt/
Reaction of salt: http://postimg.org/image/oo6bs56gl/ + lot of fumes

Well got some boro glass, and a mobile hotplate with 1500W, but not a thermometer up to 250 C :/ Will this work out? (also read somewhere that u need boiling stones?)
How % do you think it is??

AJKQER:

Could you send me one of link of the SM members accidents

AJKOER - 16-2-2015 at 06:46

Blogfast, I have since edited my text and while the actual cited reference is valid, I found the comment interesting and remembered it.

Difficult yes, I completely agree (and thank god, I did not want to actually cite more expeditious paths).

gdflp - 16-2-2015 at 06:53

Those all look like good signs, generally sulfuric acid will not react with NaCl unless it is at around 90%+ concentration. In addition, the paper will only char at a similar concentration. The acid looks remarkably clean by drain cleaner standards, it should work fine for most everything except very water sensitive reactions where a very high concentration is necessary(e.g. Fischer esterifications) in which case, I would just buy a small amount of reagent grade sulfuric acid.

Trizocy - 16-2-2015 at 06:57

So we can say its between 90-98% ish?

gdflp - 16-2-2015 at 07:05

Yes. There is some detail here about the second test you did.

[Edited on 2-16-2015 by gdflp]

woelen - 16-2-2015 at 09:26

@Trizocy: Your acid looks very good: both tests show a positve result. It certainly is well above 90% (most likely it is around 95%) and it also looks clean. No need to concentrate the acid further, tampering with it only degrades its quality. Acid of a concentration above 90% already is concentrated, you will have a very hard time concentrating it further. Going beyond 96% is impossible without using oleum or SO3.

Cou - 16-2-2015 at 10:48

SO3 is linked to laryngeal cancer in workers.

Bert - 16-2-2015 at 14:05

Quote: Originally posted by Cou

SO3 is linked to laryngeal cancer in workers.

SO₃ is also linked to a rather quicker death from pulmonary edema...

http://www.rtknet.org/node/236

Trizocy:

If you attempt this, please pay CAREFUL attention to Woelen's advice. Do search on this topic here, it has been extensively posted on.

In particular! Please do not place even your borosilicate glass directly on the heating device. USE A SAND BATH. Years ago, I concentrated several gallons of 30% battery electrolyte using a 4000ml beaker inside a cheap electric deep fryer (purchased for $5.00 at yard sale), first placing an inch or so of clean dry sand in the fryer, setting the beaker with about 3000ml of electrolyte on the sand, then pouring more sand around the beaker's walls up nearly to the top of the fryer before beginning to heat the system.

The sand bath heats the beaker of acid evenly, and helps prevent breakage. The fryer is also DESIGNED to operate for long periods at the temperature levels required, which may destroy some laboratory type hot plates when used with a sand bath.

Trizocy - 16-2-2015 at 15:09

Wowow.. This is great news!

My goal was not to reach 100% or 99% but to know how much sulfuric acid it was in the bottle (%). So i can do some experiment without concern if i put 70% acid to the mixture when its actually 90%

Thanks guys! really helped me out!

Now i know its atleast 90%+ inside it

Quote: Originally posted by woelen

@Trizocy: Your acid looks very good: both tests show a positve result. It certainly is well above 90% (most likely it is around 95%) and it also looks clean. No need to concentrate the acid further, tampering with it only degrades its quality. Acid of a concentration above 90% already is concentrated, you will have a very hard time concentrating it further. Going beyond 96% is impossible without using oleum or SO3.

Thanks woelen!!, really appreciate it!

95%

, can someone else also confirm this?

Volanschemia - 16-2-2015 at 15:22

Quote:

95%

, can someone else also confirm this?

If you have access to a base like sodium hydroxide and a pH indicator that changes colour at around pH 7, you can perform an acid-base titration to determine the exact concentration. If you need to know the exact conc, this is a very good way to go.

Do you know how to do a titration?

Trizocy - 16-2-2015 at 15:28

Hey Austria Scientist!

Well, have ordered a accurate digital PH indicator, so just waiting for it now

Tell me more about : titration?
Thanks

Volanschemia - 16-2-2015 at 16:00

Australian, not Austrian.

Where'd you get your pH indicator from? That would work very well, better than an indicator as long as it's calibrated correctly.

An acid-base titration is where you can determine the molarity of an acid or base, by neutralizing it with it's opposite (acid if it is a base, base if it is an acid). It only works if you know the concentration of one of the reactants. So, you want to get the molarity of H2SO4 which is an strong acid. So you need to neutralize it with a strong base of known concentration. (As a side note, you don't need a strong base but it makes things a bit simpler). I would suggest Sodium Hydroxide (solid drain cleaner) because it is very easy to get.

So what you need to do is make a solution of Sodium Hydroxide with a known molarity, I would suggest 5M. I'll leave the calculation of how much NaOH you need up to you. I would suggest making double the volume of H2SO4 you want to titrate. You don't need much, 25mL would be heaps. So make 50mL of 5M NaOH solution. Now this is very important: Make sure you have exactly 50mL. Using a graduated cylinder with 1mL markings would work and a burette would be even better. Place the pH meter in the acid and add the NaOH solution until the pH rises to pH 7. Indicating the acid has been neutralized. You will end up with a solution of Sodium Sulphate according to the reaction below:

H₂SO₄ + 2NaOH = Na₂SO₄ + 2H₂O

Once you are finished, measure how much NaOH solution you have left and subtract it from your starting amount. This is how much you used.

Now you can calculate the molarity of the H2SO4 via the following equation.

MH₂SO₄ = VNaOH / VH₂SO₄ * MNaOH * CH₂SO₄ / CNaOH

Where C = Coefficient, V = Volume, M = Molarity.

Hope you understood all that!

j_sum1 - 16-2-2015 at 16:32

Titration -- a very basic analytic method. It is as accurate as your ability to measure volume and mass. Acid-base titration is the most basic of titrations. Basically the idea is that the more concentrated your acid is, the larger the volume of a base solution will be needed to neutralise it. So, if you prepare some solutions and measure the volume of base needed to neutralise you should be able to calculate the concentration of your acid.

Measure accurately a small amount of your acid. Dilute with water in a flask or beaker. Add some indicator. Phenolphthalein works well for this. You won't see a colour change. (For accuracy it is better to weigh a larger amount of acid, dilute to a known volume in a volumetric flask and pipette a small amount of this for your titration.)
Choose a suitable base that will react with your acid. It must be something of high purity that you can measure accurately. Ie, not sodium hydroxide since that has an unknown amount of absorbed moisture. Anhydrous sodium carbonate would work although it is not ideal since it will release bubbles. Barium hydroxide would work.
The next step depends a bit on what volumetric measuring equipment you have. A burette is designed for the purpose but you can get an approximate result with a measuring cylinder or large graduated syringe.
Based on what you estimate your sulfuric acid strength to be, calculate the amount of your base that is needed to react with that amount.
You need to prepare a standard solution of your base to react with the acid you have measured. You need to do some calculations to decide how concentrated to make this. Ideally, you want the reaction to complete when around half the volume of your burette, measuring cylinder or syringe has been added to the acid, so work out what that is. For accuracy it is better to weigh out a larger quantity and dilute to a large but known and accurately measurable amount. Volumetric flasks are designed for this.
Slowly add your base to your acid until you see a colour change (pink in the case of phenolphthalein.) If you swish your beaker you should see the colour disappear. Then add really slowly until you see the faintest pink colour just stay visible. Record (accurately) the volume of your base that you used.

Calculations are as follows
Calculate the amount of base used in moles.
Calculate the amount of acid in moles that reacted with that base. This will be proportional to the amount of base in some kind of integer ratio. You need to have a balanced chemical equation to do this.
Calculate the concentration of the acid you used.

Useful formulas:
converting mass to moles. n= number of moles. m= measured mass. M= molecular mass.
n=m/M

calculating concentration in moles per litre. c=concentration. n=number of moles. v=volume
c=n/v

calculating concentration when diluting. c and v are concentration and volume as before. 1 and 2 are before and after diluting.
c₁v₁=c₂v₂

That's it from me. You should be able to find a lot of information on titration with a simple search.
[Ninja'd by Aus Scientist on preview]

Trizocy - 16-2-2015 at 16:36

Hahah sorry about the:" Australian, not Austrian." stuff, i think it was the sandman, sparking some dust on my eyes

The PH indicator is bought from alibaba.

Okay, this was interesting, just wanna say that my science skills is not the best, so hope anyone help me out.
Everytime i read it, i understand it more and more

Anyway, (newbie question) about molarity, is that the concentration of the sulf acid in the bottle?
And what is Coeffecient?

You said: "I would suggest 5M of Sodium Hydroxide solution" does this works then:?

100% NaOH powder

25ml total, 23,75ml water and 1.25ml NaOH powder = 5M or 5%, am i on right track

?

Have 10ml ,5ml and even 1ml pippets if that would work

Could you maybe give a calculation example?

Thanks, really appreciate, "Australian"Scientist

aga - 16-2-2015 at 16:44

Right now i have 500ml of H2SO4, and Clear (not brown) from a bottle of Brown Drain Cleaner.

Add 3% OTC H2O2 and stir while heating (it spits a lot) and all the brown goes away.

The Brown is not a problem at all, but it just does not look as good.

Heat some more until the acid starts fuming ; you will definitely know when it is not water leaving the pot anymore, and you have Clear, 98% H2SO4.

aga - 16-2-2015 at 16:56

Quote: Originally posted by Trizocy

Anyway, (newbie question) about molarity, is that the concentration of the sulf acid in the bottle?

This belongs in the Beginnings topic.

'Molarity' refers to the number of Molecules of whatever in a litre.

Google and Wiki : http://en.wikipedia.org/wiki/Molar_concentration

Understanding Molecular Weights and Molar Concentration is a very good thing to do.

Volanschemia - 16-2-2015 at 18:05

Quote: Originally posted by Trizocy

Hahah sorry about the:" Australian, not Austrian." stuff, i think it was the sandman, sparking some dust on my eyes

?

Have 10ml ,5ml and even 1ml pippets if that would work

Could you maybe give a calculation example?

Thanks, really appreciate, "Australian"Scientist

Chemistry is a great hobby to get into, you pretty much never stop learning!
Like Aga said, molarity is the number of Moles of a molecule per Litre of solvent.
1 Mole is Avogadro's number: 6.022*10 to the 23rd power.
So 5M NaOH solution would be a solution that contains 5 * Avogadro's number NaOH molecules per L of solvent (in this case, water).
Molarity is not percentage, which gets a lot of people confused.

To figure out Molarity, you first need to molecular weight of the compound in question. NaOH's molecular weight is 39.9971g/mol. Meaning 1 Mole of NaOH molecules would weigh 39.9971g. So if you were to dissolve 39.9971g of NaOH in 1L of water, you would end up with a 1M solution!

So for 5M we just multiply NaOH's molecular weight by 5, giving us 199.9855g. Let's just say we need 200g. Since we don't need a full Litre, we divide. We're after 50mL. 1000/2 = 500 / 10 = 50. 200/2 = 100 / 10 = 10g of NaOH in 50mL of water to make a 5M solution. Make sure you have an accurate set of scales for weighing the NaOH.

Coefficient is the number beside a reactant or product in a reaction equation. For example in the following equation:

2H₂ + O₂ = 2H₂O

The "2" before the Hydrogen and the Water are the coefficient. So the equation is saying 2 Hydrogens react with one Oxygen to give 2 Waters.

What are the measurement increments on your pipettes? They may work, but I'm not sure about the accuracy of pipettes. A syringe might be better and a grad cylinder would be perfect if you have one.

OK, a calculation example.

I have an some OTC HCl and I don't know the concentration. I have some NaOH solution that I know is 5M.

I measure out 25mL of the HCl in a beaker and add an indicator that will change colour at pH 7 (or in your case, a pH meter). I then add the NaOH in small amounts until the indicator changes colour (or the pH meter hits 7). I then measure how much NaOH solution I used. Let's just say I used 30mL.

The reaction equation for HCl/NaOH neutralization is:

HCl + NaOH = NaCl + H₂O

(If there is no coefficient number, the coefficient is one)

Equation:

MHCl = VNaOH / VHCl * MNaOH * CHCl / CNaOH

MHCl = 30 / 25 * 5 * 1 / 1

MHCl = 6

[Edited on 17-2-2015 by TheAustralianScientist]

Trizocy - 17-2-2015 at 01:49

Quote: Originally posted by TheAustralianScientist

Hey

Thanks for the example, really learnt much here

but got some questions:

- Will a 0,00g scales works? or do i need a 0,0000g?
- Well the pipettes looks like this, but i will try to find and buy a burette, also heard that pycnometer is good to measure also, what do u think?
- I see you used HCI as a example this time, does NaOH work on every acid?

Quote: Originally posted by TheAustralianScientist

Molarity is not percentage, which gets a lot of people confused.

I think im still one of them,

. how much concentration is it in 6M HCI? Because u said :"acid-base titration to determine the exact concentration."
But i only know how much Moles in it now.

Just want to say thanks for your effort, really appreciate it

Volanschemia - 17-2-2015 at 02:06

You're welcome. I'm glad I can teach you.

0.00g scales will work fine.
Those pipettes actually look very accurate. The 10mL one would be best I think. You won't need to use it to measure out the 50mL of solution to begin with, a rough amount is fine. But for the titration I think it would work well. Just add the solution in 10mL increments until the pH gets to about 5 and then go slower and slower until you hit 7 exactly. Make sure you record the exact amount of NaOH solution used.
As a side note, make sure your pH meter is calibrated before you measure or it could be very off. Is it just a portable meter or a computerized one?

A think a pycnometer measures relative density, not volume.

NaOH is a strong base which means it is great for titrating against any strong acid (eg. H2SO4, HCl, HNO3, HF) but not perfect for weak acids (eg. Acetic Acid, Citric Acid). For those you want a weak base like ammonia. You can titrate with a strong/weak combination but it requires a different pH indicator.

Molarity is usually more important to a chemist than percentage, but it is easy to convert. First you need the molar mass of the substance. So HCl = 36.46g/mol. We have 6M solution of HCl so 6 * 36.46 = 218.76g of HCl in 1L of solvent. From there it is easy to get percentage. 218.76 / 100 = 0.21876 or 21.876% w/v.

Chemosynthesis - 17-2-2015 at 02:15

Quote: Originally posted by Trizocy

Hey

Thanks for the example, really learnt much here

but got some questions:

- Will a 0,00g scales works? or do i need a 0,0000g?

My opinion is that 0.00g scales are fine, and 0.000g scales are great. Anything more accurate than that gets into analytical equipment, and while nice, is probably excessive and expensive.

Quote:

- I see you used HCI as a example this time, does NaOH work on every acid?

Yes. Arguably any strong base could work on any strength acid. The reason you want to use NaOH is that it essentially completely dissociates. This makes it easy to correlate neutralization of acid with quantity of base for comparisons.

Quote:

I think im still one of them,

. how much concentration is it in 6M HCI? Because u said :"acid-base titration to determine the exact concentration."
But i only know how much Moles in it now.

There is more than one type of concentration. Which one you want to use can vary.

If you want a mass concentration, you need to take moles of your solute and convert to mass of solute. Now divide this by the volume of solvent to find out what mass of substance is in what volume of solvent.

Volume concentrations get a little more complicated because, surprisingly, adding two liquids together, or a liquid and a solid, is often slightly different from the sum of the separate components. For this reason, it's often used to approximate mass concentration due to convenience of pouring rather than taring and weighing.

Once you have those down, you might look into molality... which is different from molarity. But not until you're comfortable.

Edit- oops, Australian Scientist posted an excellent post while I was typing, making this pretty irrelevant. Just listen to him.

[Edited on 17-2-2015 by Chemosynthesis]

Trizocy - 17-2-2015 at 02:19

Hmm, interesting.

"Is it just a portable meter or a computerized one?"
- Its a portable, with temp and ph meter

"Molarity is usually more important to a chemist than percentage"
- May i ask why? would it not be better to say: yup this bottle is 100% acid ? than 281/mols in it

?

Thanks, waiting for the ph meter to arrive, so gonna do a try, will keep you updated

Chemosynthesis - 17-2-2015 at 02:25

Quote: Originally posted by Trizocy

"Molarity is usually more important to a chemist than percentage"
- May i ask why? would it not be better to say: yup this bottle is 100% acid ? than 281/mols in it

Molarity is more important, generally, because that makes reactions easier to express. When you write a reaction, you consider equivalents of reactants, like X molecules of reactant A combining with Y molecules of reactant B. Moles are directly proportional to molecules.

If you wanted to use mass, you'd need to start converting more units again because they are not equivalent in the same way. This takes time and might introduce calculation errors.

Volanschemia - 17-2-2015 at 02:30

Quote: Originally posted by Trizocy

Hmm, interesting.

"Is it just a portable meter or a computerized one?"
- Its a portable, with temp and ph meter

"Molarity is usually more important to a chemist than percentage"
- May i ask why? would it not be better to say: yup this bottle is 100% acid ? than 281/mols in it

?

Thanks, waiting for the ph meter to arrive, so gonna do a try, will keep you updated

Yeah, with a portable indicator, it is a good idea to calibrate it every time you use it. Did you also purchase a buffer solution?

Well, percentage is useful for getting an idea of the concentration at a glance but you need Molarity to figure out how much of a certain reagent you need for a certain reaction.

For example, if you want to make 100g of NaCl from reacting HCl and NaOH (pretty stupid reaction I know, but for an example its nice and easy) you would need to know how many moles of HCl and NaOH are in each solution so you can work out how much of each you need.
So 1 Mole of HCl reacts with 1 Mole of NaOH. So if you had a 5M solution of NaOH and HCl, well that's easy, you just do equal amounts. But if you have 10M HCl and 5M NaOH, you would need a 2:1 ratio. In this context, percentages are useless because a 10% solution of HCl and a 10% solution of NaOH are completely different molarities.

Yeah, keep me updated! I'd like to see how it goes.

[Edit] Heh, Chemosynthesis beat me this time!

[Edited on 17-2-2015 by TheAustralianScientist]

j_sum1 - 17-2-2015 at 03:36

All good stuff here.
Let me reiterate that there are issues with NaOH. It absorbs moisture from the atmosphere and so you never know quite how much you actually have. You could be up to 10% off easily.

However, Give it a shot anyway. It is one thing to talk about it and swap calculations. It is quite another thing to do it.
In any case, it is common to do a rough titration first to get an approximate value; which makes it simpler to do an accurate one later.

Trizocy - 17-2-2015 at 03:47

Quote: Originally posted by j_sum1

All good stuff here.
Let me reiterate that there are issues with NaOH. It absorbs moisture from the atmosphere and so you never know quite how much you actually have. You could be up to 10% off easily.

Woot! up to 10%

Quote: Originally posted by TheAustralianScientist

Did you also purchase a buffer solution?

Nope,

While i was searching for burette i found this one: LINL lets talk about accurate then

but quite expensive

no need for pipettes,burettes etc etc then

[Edited on 17-2-2015 by Trizocy]

Volanschemia - 17-2-2015 at 04:06

If you dry the NaOH in an oven before dissolving it in solution you shouldn't have any troubles.

You need to get a buffer if you want to get an accurate pH measurement. You should be able to get some online pretty easy. Here is an example but not sure of the quality since it comes from china. A liquid would be better and you should be able to get it from a lab supplier if you have one nearby.

You also need to keep your pH meter in a solution of KCl or similar when in storage because it can get damaged if the bulb dries out.

A pipette vs a burette would probably only make a difference of 0.5% if that so don't spend on one unless you really need one.

morganbw - 17-2-2015 at 04:15

Quote: Originally posted by TheAustralianScientist

If you dry the NaOH in an oven before dissolving it in solution you shouldn't have any troubles.

You need to get a buffer if you want to get an accurate pH measurement. You should be able to get some online pretty easy. Here is an example but not sure of the quality since it comes from china. A liquid would be better and you should be able to get it from a lab supplier if you have one nearby.

You also need to keep your pH meter in a solution of KCl or similar when in storage because it can get damaged if the bulb dries out.

A pipette vs a burette would probably only make a difference of 0.5% if that so don't spend on one unless you really need one.

Just drying does not fix it completely. Sodium Carbonate is also formed when exposed to the air.

Trizocy - 17-2-2015 at 05:13

Hmm one question,

I got my sulfuric acid in my garage, will it absorbs moisture from the atmosphere over time?

Molecular Manipulations - 17-2-2015 at 06:08

If the container is airtight, then probably not. Keep a good seal and it will do fine, I put Teflon tape around the threads if it's stored for a long time.

Volanschemia - 17-2-2015 at 18:49

If you put the NaOH in the oven for just five minutes it will be fine.

I don't think you need a result down to the millimole do you? I highly doubt you will get a margin of error above 1% unless you leave the NaOH out without a lid for an hour and I think a percentage within 1% was all you needed am I correct?

Since your Sulphuric Acid already seems quite concentrated its hydroscopy will be fairly minimal. Like MM said, and from my own experience, as long you keep the lid on you shouldn't have any trouble.

Have got looked at pH buffers yet? If you want, you can use a pH indicator like Phenolphthalein with much the same results.

Molecular Manipulations - 17-2-2015 at 18:54

Quote: Originally posted by TheAustralianScientist

If you put the NaOH in the oven for just five minutes it will be fine.

Define "fine". Cause if you think five minutes in an oven at any reasonable temperature will dehydrate sodium hydroxide, you are sadly mistaken.
An hour at 500 deg.C might bring it to 95-97%.

Quote:

Since your Sulphuric Acid already seems quite concentrated its hydroscopy will be fairly minimal.

Umm, say what? The more concentrated the acid, the more hydroscopic it will be. And sulfuric acid is very hydroscopic.

[Edited on 18-2-2015 by Molecular Manipulations]

jsc - 17-2-2015 at 19:04

I hope I caught you before you tried this.

Boiling sulfuric acid is a bad idea. What will happen is that a large amount of it will be vaporized. If you do it indoors, those vapors will then condense on EVERY SURFACE IN THE HOUSE. For example, if you have wallpaper, this will result in large black splotches on the walls, giving it sort of the "cow" look. It will also rust any and all metal in the house. Also, acids have a tendency to eat through clothes, so you wash your clothes and there will be holes in them. Also, if you are in the house while those vapors are there you will be breathing those vapors. Get the picture?

At a bare minimum you need some kind of condenser.

A larger problem is that a drain cleaner will have all kinds of crap in it. If you remove water, you will concentrate the crap as well as the acid. To purify it you have to distill it, a dangerous activity that requires specialized equipment because it has a high boiling point.

You can remove some of the organics by adding a small amount of concentrated hydrogen peroxide (see YouTube).

Volanschemia - 17-2-2015 at 19:12

Fine for a rough titration.
I also don't think that NaOH stored in an airtight container will be 10% water. If that was the case it would be a slush.

My mistake with the Suphuric Acid. You are right there. But if it's kept in an airtight bottle it will be fine. I have had 98% Sulphuric Acid for 6 months now and my last titration showed 97.5% so nothing serious has happened.

Volanschemia - 17-2-2015 at 19:20

Quote: Originally posted by jsc

I hope I caught you before you tried this.

Boiling sulfuric acid is a bad idea. What will happen is that a large amount of it will be vaporized. If you do it indoors, those vapors will then condense on EVERY SURFACE IN THE HOUSE. For example, if you have wallpaper, this will result in large black splotches on the walls, giving it sort of the "cow" look. It will also rust any and all metal in the house. Also, acids have a tendency to eat through clothes, so you wash your clothes and there will be holes in them. Also, if you are in the house while those vapors are there you will be breathing those vapors. Get the picture?

At a bare minimum you need some kind of condenser.

A larger problem is that a drain cleaner will have all kinds of crap in it. If you remove water, you will concentrate the crap as well as the acid. To purify it you have to distill it, a dangerous activity that requires specialized equipment because it has a high boiling point.

You can remove some of the organics by adding a small amount of concentrated hydrogen peroxide (see YouTube).

Yes, don't boil in the house. I don't think he was planning to though, since he's already determined it's above 90%.

Trizocy, what were you planning on using the acid for?

[Edited on 18-2-2015 by TheAustralianScientist]

Cou - 17-2-2015 at 19:21

Quote: Originally posted by jsc

Boiling sulfuric acid also releases SO3, which *could* be a carcinogen, especially with long term exposure such as in factory workers. If you absolutely have to boil sulfuric acid, wear a dual respirator with an SO2 or SO3 filter (not sure if those filters work interchangably). I always use a respirator when working with chlorine gas and it works wonders.

[Edited on 18-2-2015 by Cou]

Molecular Manipulations - 17-2-2015 at 19:25

Jsc, I don't think he was planning on doing this indoors. What you said is mostly correct, if he was to boil down a bathtub full of acid, but a few 100's of mL, not so much.
TAC, not slush by a long shot. Any technical grade NaOH you find is less than 90% pure, even 70% would resemble a dry powder or chips. You do realize that in hydroscopic materials, the water is incorporated in the crystalline structure. Magnesium sulfate hexahydrate has six water molecules per Mg and SO₄ ion, and is about 46% water by weight, but is still dry.

[Edited on 18-2-2015 by Molecular Manipulations]

Cou - 17-2-2015 at 19:30

NEVER MAKE TOXIC GASES INSIDE THE HOUSE. Always use a fume hood or do it outside. And wear a respirator... as fun as chemistry is, I don't think it's worth going to the doctor 5 years later, and finding out you only have 6 months to live because of cancer. Or pneumonia the day after getting too big a whiff of SO2 and SO3.

[Edited on 18-2-2015 by Cou]

[Edited on 18-2-2015 by Cou]

Volanschemia - 17-2-2015 at 19:57

Quote: Originally posted by Molecular Manipulations

TAC, not slush by a long shot. Any technical grade NaOH you find is less than 90% pure, even 70% would resemble a dry powder or chips. You do realize that in hydroscopic materials, the water is incorporated in the crystalline structure. Magnesium sulfate hexahydrate has six water molecules per Mg and SO₄ ion, and is about 46% water by weight, but is still dry.

[Edited on 18-2-2015 by Molecular Manipulations]

As far as I and several others I have conferred with know, NaOH does not hydrate, merely dissolves in the water it pulls in, and it certainly can't be compared to MgSO4.

Molecular Manipulations - 17-2-2015 at 20:10

Sorry but what you said makes no sense.
Who are these "others"? If it absorbs water then it must hydrate. The water is absorbed for a reason and certainly is incorporated in the crystalline structor.
It makes NaOH-2H₂O and NaOH-3.5H₂O.
For it to randomly absorb water without forming a hydrate is beyond ridiculous.
It may not be comparable with magnesium sulfate in the sense that the hydrates are ill-defined and still hydroscopic themselves, but in the sense that the analogy was used it's comparable as anything.

[Edited on 18-2-2015 by Molecular Manipulations]

Amos - 17-2-2015 at 20:21

Quote: Originally posted by Molecular Manipulations

Sorry but what you said makes no sense.
Who are these "others"? If it absorbs water then it must hydrate. The water is absorbed for a reason and certainly is incorporated in the crystalline structor.
It makes NaOH-2H₂O and NaOH-3.5H₂O.
For it to randomly absorb water without forming a hydrate is beyond ridiculous.
It may not be comparable with magnesium sulfate in the sense that the hydrates are ill-defined and still hydroscopic themselves, but in the sense that the analogy was used it's comparable as anything.

[Edited on 18-2-2015 by Molecular Manipulations]

If dry, dusty soil becomes wet, it surely must be due to water being incorporating it into its crystal structure!

I'm with TheAustralianScientist here, I don't think that NaOH hydrates in the same sense as those compounds that have a well-defined crystal structure and incorporate water in a set ratio. I'm hesitant to say more with my admittedly small knowledge on the subject, but if you're going to belittle TheAustralianScientist, why not educate him as well? Pull out the sources.

And how would you explain hygroscopic polymers, activated alumina, etc? Are they forming whole new crystal structures, or merely adsorbing the water on their surface or within pores?

[Edited on 2-18-2015 by Amos]

Volanschemia - 17-2-2015 at 20:39

Molecular Manipulations, I would be more than happy to admit that I am wrong if you can provide me with evidence that NaOH does indeed hydrate like you are suggesting.

In fact, I would be happy to have learned something new.

Molecular Manipulations - 17-2-2015 at 20:45

Cause adding water to dirt to make mud is the same as it being absorbed from the air...
I'm on my phone now and posting sources is a pain, if you want them, wait till tomorrow - or find them yourself.

Volanschemia - 17-2-2015 at 20:51

Quote: Originally posted by Molecular Manipulations

Cause adding water to dirt to make mud is the same as it being absorbed from the air...
I'm on my phone now and posting sources is a pain, if you want them, wait till tomorrow - or find them yourself.

You didn't mention the hydroscopic polymers...
Anyway, I'm happy to wait until tomorrow for sources.
I will be sticking to my opinions though until they have been proven incorrect by a trusted source of information.

Molecular Manipulations - 17-2-2015 at 20:58

Ah what the hell, one little source couldn't hurt for now.

How's this?
And if that's not enough, this picture is clear proof, I got it online, so it must be true:

image.jpg - 3kB

[Edited on 18-2-2015 by Molecular Manipulations]

Bert - 17-2-2015 at 21:00

Quote: Originally posted by TheAustralianScientist

If you put the NaOH in the oven for just five minutes it will be fine.

Would you care to describe the oven temperature used for that 5 minutes? Also the container material used?

Quote: Originally posted by TheAustralianScientist

Since your Sulphuric Acid already seems quite concentrated its hydroscopy will be fairly minimal.

"Hydroscopy"

HYDROSCOPIC

The word is HYGROSCOPIC. Hydroscopy would be looking into water... using a hydroscope.

Sulfuric acid is QUITE hygroscopic. It's used in desiccators. The more concentrated the acid, the more hygroscopic it will be.

Amos - 17-2-2015 at 21:26

Quote: Originally posted by Molecular Manipulations

Cause adding water to dirt to make mud is the same as it being absorbed from the air...

I never said water was purposefully added to the dirt. Very dry dirt when placed in a more humid environment WILL become more moist, and not because it hydrates, but because water diffuses into anything dry. I was just illustrating how not everything is changing structure and hydrating just because it captures water.

Volanschemia - 17-2-2015 at 21:31

Bert,
For the oven, at the time I was thinking 180C on a glass dish would suffice but I now realize that may not be sufficient. I am still of the opinion that even without drying, NaOH can still be used reliably in a titration.
What do you think about the hydration matter?

With the Acid, my mistake was pointed out in the next post and in the next after that, I acknowledged my mistake.
Maybe you missed that.

I realize now that I spelled 'Hygroscopy' with a 'D' and I apologize for that.
Are you saying however that "The hyGroscopy is minimal" is not a valid sentence? I know for certain that saying "The hydroscopic is minimal" is not correct. Here is a page on Wikipedia that describes "Hygroscopy" as "the ability of a substance to attract and hold water molecules from the surrounding environment."

Molecular Manipulations, thankyou for the reference, I will take the time to thoroughly look through it.

Bert - 17-2-2015 at 22:19

Melt a small amount of Sodium hydroxide (318 C if anhydrous, about 250 C @ 10% water) on a glass surface, one which you do not mind losing. Please report on the condition of the glass afterwards. Draw conclusion on suitability of high temperature/glass containment method for dehydration of this chemical.

Sodium hydroxide is not only hyGroscopic, but deliquescent. It will take up water from the air until it DISSOLVES itself, and then continue to absorb even more water... As mentioned by others, it will also scavenge atmospheric CO₂ and convert itself to carbonate.

In general, if you have not performed an operation, or had hands on experience with a material- but wish offer advice? Please make it clear you discuss it from your reading (and give your references!). Or clearly say what you provide is CONJECTURE, rather than authoritatively state a procedure. Someone who knows even less than you may take your advice as gospel, perhaps with a bad outcome.

[Edited on 18-2-2015 by Bert]

[Edited on 18-2-2015 by Bert]

Volanschemia - 17-2-2015 at 23:27

Quote:

Sodium hydroxide is not only hyGroscopic, but deliquescent. It will take up water from the air until it DISSOLVES itself, and then continue to absorb even more water... As mentioned by others, it will also scavenge atmospheric CO2 and convert itself to carbonate.

What we were discussing was whether NaOH hydrates (ie. integrates water into it's crystal structure).

Quote:

Draw conclusion on suitability of high temperature/glass containment method for dehydration of this chemical.

As I said, if I had to put it in the oven, I would have set the temp 180C, which is below melting point of NaOH.

Sorry if I sound stupid, but I'm not quite sure what that diagram is supposed to mean.

I freely admit there is MOUNTAINS of stuff I still need to learn and would be more than pleased if someone is willing to teach. Meanwhile, I like to pass on what knowledge I do know to others in this forum where I can be corrected if wrong.

Trizocy - 18-2-2015 at 01:54

Wow guys, i am really charmed with all the help i got

Can say this, was on a other high popular science forum and asked about the same question, well the help was not good, felt like reply was "fuck you" in a indirect way

Quote: Originally posted by TheAustralianScientist

I don't think you need a result down to the millimole do you? I highly doubt you will get a margin of error above 1% unless you leave the NaOH out without a lid for an hour and I think a percentage within 1% was all you needed am I correct?

Hey im not working as chemist in pharmaceutical companies, im doing this for fun, and to learn, like a hobby. 1%-2% im MORE than happy. When i was playing with the drain clear, i quess it was 80%, because the description said 60%-100% sulfuric acid, well my mixture dident end good as i want. So the only thing i wanted is to know what % was in my drain cleaner bottle

Quote: Originally posted by jsc

I may be a newbie in chemist, but i was not born yesterday

, have seen 100 of videos on youtube what sulfuric acid can do, so you think i would do it inside? Where i eat and sleep, nono m8. Hmm thought i wrote that i had a portable hotplate, earlier in this thread

But thanks for reminding me of safety!

About the NaOH stuff, (which seems like a war discussion here :p ). if it absorbs water, im okay. As long as the error of the acid concentraion is between 1-2%

But anyway, waiting for my PH meter to arrive, and i will do some test. Will keep you guys updated

And thanks for all the attention and commitment!

Volanschemia - 18-2-2015 at 02:05

Hey!
Sorry about the massive amount of disagreement thats been going on! As long as someone (possibly me) learns something from it, it's worth it.

In my opinion, I don't think the error will be more than 2% easy if you use NaOH, but if anyone else disagrees feel free to correct.

What is the other forum you were on? Yes, there are some great people on here, I have gotten my fair share of help from a lot of people.

Just found a two part video from NurdRage that explains about pH meters really nicely. You may have already seen but if not it is a great video to learn from. (Part 1 and Part 2)

j_sum1 - 18-2-2015 at 02:29

Quote: Originally posted by TheAustralianScientist

Hey!
Sorry about the massive amount of disagreement thats been going on! As long as someone (possibly me) learns something from it, it's worth it.

In my opinion, I don't think the error will be more than 2% easy if you use NaOH, but if anyone else disagrees feel free to correct.

What is the other forum you were on? Yes, there are some great people on here, I have gotten my fair share of help from a lot of people.

Just found a two part video from NurdRage that explains about pH meters really nicely. You may have already seen but if not it is a great video to learn from. (Part 1 and Part 2)

Yes, the error is likely to be more than 2%. That shouldn't be enough to stop anyone attempting a titration to get some practice at the technique. It should be enough to stop someone being confident in the result.

@AusChem, if you have titrated H2SO4 against NaOH in the past and both have absorbed water, then the errors introduced will cancel each other out somewhat. Reporting 98% H2SO4 has dropped to 97.5% is consistent with this. But I wouldn't hold a lot of confidence in that conclusion.

Phase diagrams... transport me to a happy place. You can gain a lot of information from a phase diagram if you know how to use it. But more than a two minute explanation. What this particular one shows is that there are a lot of intermediate substances between 100%NaOH and 100%water. In other words, NaOH absorbs water -- so much so that it eventually dissolves in the water it has absorbed. Nice find Bert. Care to share your source?

@Trizocy. I would suggest that you will get much better results with an indicator than you will with a pH meter. Those things can be fickle. If you don't have a buffer solution to store it in you can't expect the readings to be accurate. This is coming from someone who has ordered one recently and is waiting impatiently for it to arrive in the mail.
Phenolphthalein is an excellent indicator for this purpose but if you don't have any then you can perform a very acceptable titration using cabbage juice or squeezing the juice out of some coloured flowers. All you need to do is run some tests to determine what colours you are watching for in acidic and alkaline environments.

Boiling down H2SO4 -- something I have done a bit of recently. Yes, do it outside. Yes expect anything steel in the vicinity to rust like mad. Yes protect yourself against spills, splats and accidental elbow manoeuvres. A couple of months ago I did a beautiful job of cleaning my concrete shed floor with a 1L beaker of boiling H2SO4. (I was a bit peeved at losing my favourite beaker too.) It pays to be careful. Yes, avoid breathing in any fumes. Pull it off the heat when dense white fumes appear. All this is assuming it needs concentrating. Trizocy, it seems like your acid is strong enough for most purposes without further concentrating.

Volanschemia - 18-2-2015 at 03:21

Quote: Originally posted by j_sum1

Hi j_sum. Thanks for your input. What base would you suggest for titrating H2SO4? Makes sense about them both cancelling each other out. Thankyou for clarifying the phase diagram. I will look into them further.

Trizocy, I agree with j_sum, that if you don't have a buffer solution to store, you would be much better off with an indicator. Phenolphthalein is pretty easy to find online so you should be able to find some if you look. Phenolphthalein is great because it has such a distinct colour change so you can't miss the equivalence point.

morganbw - 18-2-2015 at 03:31

Quote: Originally posted by TheAustralianScientist

Fine for a rough titration.
I also don't think that NaOH stored in an airtight container will be 10% water. If that was the case it would be a slush.

My mistake with the Suphuric Acid. You are right there. But if it's kept in an airtight bottle it will be fine. I have had 98% Sulphuric Acid for 6 months now and my last titration showed 97.5% so nothing serious has happened.

Where I worked at we had a large tank of 50% NaOH, it was a heated tank because the 50% NaOH would turn to a solid at cool temps.
You can clearly see this on the phase diagram Bert provided.

[Edited on 18-2-2015 by morganbw]

Trizocy - 18-2-2015 at 03:39

Quote: Originally posted by TheAustralianScientist

What is the other forum you were on? Yes, there are some great people on here, I have gotten my fair share of help from a lot of people.

chemicalforums.com

Quote: Originally posted by j_sum1

@Trizocy. I would suggest that you will get much better results with an indicator than you will with a pH meter. Those things can be fickle. If you don't have a buffer solution to store it in you can't expect the readings to be accurate. This is coming from someone who has ordered one recently and is waiting impatiently for it to arrive in the mail.

http://www.aliexpress.com/item/Free-shipping-New-7-00-9-18-1... + distilled water?
Would this work out?

And do i need ph buffer for every ph or its okay to get 3ph 8ph and 12 ph?

Thanks

j_sum1 - 18-2-2015 at 03:46

Suggestions for titrating H2SO4.
Now I have to confess to never having had to perform a highly accurate base titration. You want a base with the following properties

strong base for a well-defined endpoint
high purity
something that does not absorb moisture (or anything else) from the atmosphere

I believe Ba(OH)2 fits the bill but I am unsure how commonly that is used. [edit] Ba(OH)2.8H2O Barium hydroxide octohydrate [/edit]
The alternative is to use an acid standard to analyse your NaOH. Oxalic acid would work here and I believe that this is a common approach.
This is where the schoolteachers step out of the arena and the real chemists step in. Those with actual experience.

[Edited on 18-2-2015 by j_sum1]

j_sum1 - 18-2-2015 at 03:54

@Tizocy
Distilled water is exactly the opposite of what you want. A minute amount of contaminant (for example the remains of the last thing you measured) will greatly alter the pH of the water. The pH could swing widely.
pH meters have a glass electrode. The should remain wet. And I think that they should remain at a constant pH when stored for any period of time. I am not precisely sure what that pH is or whether it varies from model to model. The one at my work is sitting in a solution of pH 4.5 which came when it was purchased.

A buffer solution is a solution of a weak acid and its conjugate base (or vice versa) that maintains a stable pH even when considerable amounts of acid or base contaminant are added. It works on the basis that the weak acid does not dissociate fully. If some of that acid is neutralised by a base then the equilibrium with the conjugate restores the solution close to what it was originally.

Volanschemia - 18-2-2015 at 04:12

Quote:

I believe Ba(OH)2 fits the bill but I am unsure how commonly that is used.

That would work, but like you said, I don't think it is very commonly available and its also fairly expensive IIRC.

Quote:

And do i need ph buffer for every ph or its okay to get 3ph 8ph and 12 ph?

Just a single pH buffer will suffice (although you may need a certain one for your particular meter). Generally you buy it in liquid form, not sure how those powders work.
(As a side note, generally you don't write "3ph" you write "pH 3".)

Personally I think a pH meter is more trouble than it's worth when you can use something like Phenolphthalein. Do you have access to that Trizocy?

Quote:

The alternative is to use an acid standard to analyse your NaOH. Oxalic acid would work here and I believe that this is a common approach.
This is where the schoolteachers step out of the arena and the real chemists step in. Those with actual experience.

I would be interested in learning more if a more experienced person happens to read this. I will do some reading myself as well.

Trizocy - 18-2-2015 at 04:25

Goal is not to get 99% accurate

NaOH pretty common in stores here, Ba(OH) have no idea where to get that :/

Here is the link of my pH meter

what u think?
http://www.aliexpress.com/item/3-in-1-Bench-Type-digital-PH-...

About the Phenolphthalein, im not sure :S

j_sum1 - 18-2-2015 at 04:32

Quote: Originally posted by Trizocy

Goal is not to get 99% accurate

NaOH pretty common in stores here, Ba(OH) have no idea where to get that :/

Here is the link of my pH meter

what u think?
http://www.aliexpress.com/item/3-in-1-Bench-Type-digital-PH-...

About the Phenolphthalein, im not sure :S

No. But you might want better than 85% accurate.
Ba(OH)2 I don't think is that difficult to get. It is probably not the only (or even best) option. It is just the one that came to mind.

The pH meter looks a beast. Look after it. Find out what buffer solution it needs and get it. The temptation will be to think that the box on the bench is the technical part. It isn't. The probe is what you need to look after.

Seriously, if you don't want to be bothered with buying indicators, have a play with some coloured flowers or vegetables. They are some of the best indicators you can get. After all, litmus is merely extract from a lichen.

Chemosynthesis - 18-2-2015 at 06:17

Quote: Originally posted by TheAustralianScientist

Heh, Chemosynthesis beat me this time!

Still appreciating your post, either way

Quote: Originally posted by TheAustralianScientist

Sorry if I sound stupid, but I'm not quite sure what that diagram is supposed to mean.

I freely admit there is MOUNTAINS of stuff I still need to learn and would be more than pleased if someone is willing to teach. Meanwhile, I like to pass on what knowledge I do know to others in this forum where I can be corrected if wrong.

You don't sound stupid at all. I wish more people in the forum were as polite and on point as you two when disagreeing, as it makes everything easier. If one of us thinks it, someone else probably would too. I'll take a crack at Bert's diagram, and see if he was coming it from another direction:

Notice that you look at the dependent, or x-axis, you can start from the far right with nearly 100% pure NaOH. Look at the y-axis, and go to your suggested temperature of 180 degrees. Now find the intercept. Because NaOH is deliquescent, it is still capable of forming an 80% solution essentially unaided at this temperature. So, in an ideal environment, unless your NaOH was below 80% solution, I don't think heating to this temperature would be helpful for drying. In practice, it would eat through your glassware and likely take on CO2 from the atmosphere, which happens to an extent at room temperature. This carbonate formation would obviously introduce error into your titration.

That's how I chose to view it. As for hygroscopy on wikipedia... funny! I always used "hygroscopicity" as the noun form. Now I'm going to be thinking about that all day!

Molecular Manipulations - 18-2-2015 at 06:41

Oh Bert, what would I do without you...? I wonder how long I've been misspelling that word too...

Sorry to continue this sodium hydroxide hydrate discussion, but I might as well finish.
First, just think about it. Sodium hydroxide dissolves exothermically in water, why? I talked a little about this here.

Quote: Originally posted by Molecular Manipulations

Water molecules are attracted to solid sodium hydroxide. Why? Water molecules are polar and sodium hydroxide is made of ions. These two are attracted through the ion-dipole interaction, which is a Van der Waals force.
The positive part of the water molecule (the hydrogens) attract the negative hydroxide ion while the negative oxygen part attracts the positive sodium ion.
But why does sodium hydroxide dissolve exothermically while say ammonium nitrate dissolves endothermically? Because all ionic crystals have lattice structures. When a solid lattice forms, energy is released. When a substance is dissolved, the lattice is pulled apart and destroyed - this requires energy. This is where Van der Waals comes in, there are three Van der Waals forces at play here. The "solvent-solute attraction", the "solvent-solvent attraction" and the "solute-solute attraction".
The relative strengths of each force are critical in determining whether it will dissolve endothermically, exothermically, or not at all.
When the solute-solvent attraction is greater than the solute-solute attraction the substance will dissolve exothermically.
So in a sense you're right about the breaking of bonds (the crystal lattice). But it gets way more complicated than that, I've tried to simplify it as much as possible.

That out of the way, why do you think sodium hydroxide is hygroscopic? For the same reason of course! The solvent-solute attraction is very strong, when a water molecule in the air collides with it, rather than bouncing off like it would with most substances, Van der Waals forces cause it to combine - permanently (unless it's strongly heated). Sodium hydroxide wants to dissociate, but without water, it's solute-solute attraction (the crystal lattice) keeps it in place. So when I posted that it incorporates the water into it's structure, I had actually never seen that before, it just made sense, it has to do that.
Check this out
Quote:

Quote:

Metal cations (being positively charged) attract the lone pairs on water oxygens and form coordinate covalent bonds with water. For example, many divalent cations M2+ can form ions like [M(H2O)6]2+. One or more of the waters can be replaced with anions in the hydrate.

The water is bonded to the ions with covalent bonds, this releases a lot of energy.
That same link also answers Amos's question about activated alumina and maybe even dry dirt!

Quote:

Water molecules can diffuse into the material and become trapped there, especially if the material contains voids that can hydrogen-bond the water in place.
Some porous materials can absorb large amounts of water by a phenomena called capillary condensation. For example, zeolites (complex sodium aluminum silicates) are honeycombed with cavities lined with oxygen atoms. Water in air hydrogen bonds to the outer surfaces of the material. As more water condenses, it tries to spread out and wet as much of the surface as possible to minimize its energy through hydrogen-bonding. That carries the water deeper into the material, making more room for water to adsorb to the outer surface. (They soak up water like a sponge!)

Does this answer your question? I know I couldn't find as good of "proof" that it is incorporated into the structure as I thought I could, but this seems good enough.

[Edited on 18-2-2015 by Molecular Manipulations]

Bert - 18-2-2015 at 10:03

If someone were doing titrations to determine reagent concentration, something like these burettes could be handy- Assuming one had figured out how to make or acquired a standard solution to dispense from such.

https://www.sciencemadness.org/whisper/viewthread.php?tid=61...

papaya - 18-2-2015 at 12:46

What about density measurement? Isn't it precise enough to tell you how concentrated is your acid? Titration has it's problems since you need titration standards (in ampules) to prepare solution of exactly known normality. once I also suffered with this same problem, I even found out an article about usage of borax for titration (when recrystallized properly borax has very well defined "composition"), but never went to test it, maybe anyone else did this?

Cou - 18-2-2015 at 13:56

Quote: Originally posted by Trizocy

Goal is not to get 99% accurate

NaOH pretty common in stores here, Ba(OH) have no idea where to get that :/

Here is the link of my pH meter

what u think?
http://www.aliexpress.com/item/3-in-1-Bench-Type-digital-PH-...

About the Phenolphthalein, im not sure :S

NaOH is easily available in stores like Lowes as drain cleaner. Be careful though... it's better to get reagents from amateur chemistry suppliers such as Elemental Scientific, because buying lots of chemicals from hardware stores is associated with cookery.

gdflp - 18-2-2015 at 15:15

Quote: Originally posted by Cou

If you're really worried, use cash. I've had no issue buying sulfuric acid, sodium hydroxide, toluene, MEK, and acetone all at the same time from Ace.

Trizocy - 18-2-2015 at 16:50

Hey, like i said, im not looking for 99,999999% acid, all i want to know is what concenration acid im playing with :/

Buying pure grade sulfuric acid here in my country is hard, you need to document/school/company. But the drain cleaner contains 60-100% sulfuric acid so easy to buy

Right now, im waiting for my pH meter : http://www.aliexpress.com/item/3-in-1-Bench-Type-digital-PH-...
And im looking for a burette, also buffer soultion. And NaOH which they sell everywhere. This will be okay?

Its okay to have 1-4% errors.

Damn when TheAustralianScientist told me about the NaOH, i thought i would be fun and easy, until the accurate guys came and says everything is wrong,:" you will have to much error"

Well thanks guys

j_sum1 - 18-2-2015 at 19:34

If you think I am discouraging you from having a go then you have misunderstood me. But you do need to realise ahead of time that the result you get is likely to be more than a few percent off.
You might get better success in terms of accuracy if you use sodium carbonate: washing soda. It doesn't absorb moisture the same way. It does have the disadvantage of being a weaker base and so you have to be a bit careful of the endpoint -- whether you use indicator or pH meter. And you get the added irritation of CO2 bubbles in your titration.
If it was me doing it with the equipment you have then I would do both titrations.

Bert - 18-2-2015 at 20:40

Density measurements compared with charts of acid density vs. concentration is an OK way to determine acid %- IF you have an accurate graduated cylinder and an accurate scale, plus are careful to use the cylinder near the calibration temperature. I have made the mistake of using standard temperature calibrated graduated cylinders outdoors in below freezing weather (because I chose to boil down my sulfuric acid outdoors!)

I'd settle for within 3% accuracy in acid concentration measurements in most reaction processes. Besides analytical work, where do you need control tighter than that?

Volanschemia - 18-2-2015 at 23:25

Quote:

NaOH pretty common in stores here, Ba(OH) have no idea where to get that :/

I'm not really sure either. The only place i've seen Barium Hydroxide is on Ebay. Here is a link to a listing from a seller I have bought from but it is a bit expensive.

Quote:

Here is the link of my pH meter

what u think?

Wow! That is a nice meter. Make sure you look after it and it will serve you well.

Chemosynthesis, thankyou for your clarification on the phase diagram. Between yours and j_sum's explanation, I now understand.

Quote:

That's how I chose to view it. As for hygroscopy on wikipedia... funny! I always used "hygroscopicity" as the noun form. Now I'm going to be thinking about that all day!

I like 'Hygroscopicity', it has a nice ring to it. It's very confusing with the whole 'd' or 'g' thing!

Molecular Manipulations, thankyou so much for your effort in procuring that information. It seems we have both learned something from this. So, it's not your standard hydration, but I admit that it is still a hydration and I was wrong.

Quote:

This may not hold true for Norway, but I have NEVER gotten a funny look from buying caustic soda (NaOH) from the supermarket.

Quote:

Damn when TheAustralianScientist told me about the NaOH, i thought i would be fun and easy, until the accurate guys came and says everything is wrong,:" you will have to much error"

Sorry for misleading you Trizocy. It seems that I still have a bit to learn myself...

Quote:

You might get better success in terms of accuracy if you use sodium carbonate: washing soda.

Or Sodium Hydrogen Carbonate (Bi-carb Soda) would work as well. Just be aware though that both Na(CO3)2 and NaHCO3 will take more solution (due to it being a weaker base) and I think (again feel free to correct anyone) that small amounts of Carbonic Acid will remain in solution and may skew results.

Measuring density may be your best option but I have little to no knowledge on the subject, so you would have to get your info from someone else.

Sorry about so many quotes, there were so many things I wanted to address.

[Edited on 19-2-2015 by TheAustralianScientist]

Oscilllator - 19-2-2015 at 00:55

For what it's worth, I use density measurements almost exclusively when trying to determine the concentration of acids and other substances. Sulfuric acid is ideal for a density measurement, because the density of pure sulfuric acid is about 1.8.
A 100ml erlenmeyer flask and a decent scale will tell you the concentration of your acid to within a few %, and you don't have to mess around with titrations and other such shenanigans.

Volanschemia - 19-2-2015 at 01:07

Quote: Originally posted by Oscilllator

For what it's worth, I use density measurements almost exclusively when trying to determine the concentration of acids and other substances. Sulfuric acid is ideal for a density measurement, because the density of pure sulfuric acid is about 1.8.
A 100ml erlenmeyer flask and a decent scale will tell you the concentration of your acid to within a few %, and you don't have to mess around with titrations and other such shenanigans.

Oh wow! I can't believe I hadn't thought of that. That sounds like a fantastic plan of action, and I think you already have some 0.01g scales don't you Trizocy? A grad cylinder or a burette would be better than an erlenmeyer flask though I would think.

Trizocy - 19-2-2015 at 02:06

Quote: Originally posted by TheAustralianScientist

Quote: Originally posted by Oscilllator

Damn now we are back on density again, it was this the previous forum also recommended me to do :/
Yes i got 0,01g and have aldready done some measurement with density, and here is the result:

I got a 1L bottle of sulfuric acid, on the bottle it stand 60-100%, when i measure it with 50 mL i got 93,34g

93,34g * 2 = 186,68g

186,68g / 100 = 1,8668 g/ml density

And here is the density table for it: http://www.sschemical.com/wp-content/uploads/2013/05/Convers...

The highest % here is 1,8391 g/ml, but my result is 1,8668???

Also they recommoned me to buy a pycnometer, which is the only way to be truly accurate it they said.
With only burette,pippette or a cyilnder i could never pour on the same amount like last time.

Had a expriment with density and freshwater

Quote: Originally posted by me

I had a brand new fresh 5L water (think its filter water, and yeah! i know that distilled water and filter water is not same, but this was just a experiment on it, gonna try the destilled water when i got some) so i tryed to measure it 20 times, and here was the result: http://pastebin.com/2xSHGNkq - it comes between 48,14g and 48,95g on 50 ml (20-21°C), (average 48,5905g on 50 mL) My standard deviation was 0.27458.

As you see in the link, almost never got the same number.

Thanks

papaya - 19-2-2015 at 05:47

Trizocy, you got higher density due to some error - rather measuring at different temperature or most likely error in volume measurement. But you know what - if you can't take exact volume of acid, how do you expect titration to work for you?

Bert - 19-2-2015 at 07:09

Trizocy-

If the density of water doesn't work out to be 1g/cc with your equipment, either your equipment or measuring technique needs looking at.

For a start: When using the graduate cylinder, you will observe the "meniscus", a curvature of the liquid's surface due to surface tension. How are you accounting for that?

http://en.m.wikipedia.org/wiki/Meniscus

[Edited on 19-2-2015 by Bert]

Trizocy - 20-2-2015 at 02:31

Hmm i think its my measuring technique which must be wrong

Anyway, if i gonna measure density, then i could just buy pycnometer.

Pycnometer (25ml) on 0,01g scale. Note down the Pycnometer weight
Pour Sulf acid in the Pycnometer and put the top on, let acid overflow on the pycnometer.
Put back on scale, to measure the mass

Pycnometer weight - Pycnometer full of sulf acid = mass of sulf acid

Volume / Mass = Density

?

Then you find what tempture it is, and note that down. And so find specific gravity density sulfuric acid. And lookup where you result is = Profit

?

But one problem, i havent found a legit confirm authority specific gravity, there are too many of them on google, and all have difference measuring :S

blogfast25 - 20-2-2015 at 06:12

Quote: Originally posted by Trizocy

Pycnometer weight - Pycnometer full of sulf acid = mass of sulf acid

Volume / Mass = Density

Density measurements with Pycnometers are usaually done relative to pure water. So also measure Pycnometer weight when filled with water.

W0 = pycnometer weight
W1 = pycnometer filled with acid weight
W2 = pycnometer filled with water weight

Density relative to water = (W1 - W0)/(W2 - W0)

Trizocy - 19-3-2015 at 13:23

Hello!

Let's finish this thread

So... Lets continue.

I got a pycnometer now, so i can try to measure the density now

I got 1 Liter of H2SO4, which on the description it stand 60-100% and what i want is to try to find the specific concentration in the bottle.
So what i have done is to measure several times with a 50ml pycnometer, and a temperature meter.

50ml pyconmeter empty: 34,15g
Average 50ml pycnometer with H2SO4: 126,91g (pour in and out several times, and measure it)

Weight of the H2SO4: 92,76g
Temprature: 7,5 cel to 10,5 cel

So how do i calculate it? Tryed to use Chembuddys calculator, but gave me 185% concentration which is not possible

Fulmen - 19-3-2015 at 13:29

You know both the mass AND the volume of this mass of liquid, don't you see it? Just think about the unit's you have and what you're looking for...

92,76g / 50ml = 1,855g/ml. Then it's just a matter of looking this up in a table.

BTW: Well done, a pycnometer is the proper tool for this. I've always had trouble with getting good measurements from hydrometers.

[Edited on 19-3-15 by Fulmen]

Volanschemia - 19-3-2015 at 15:20

It seems to me that you have some VERY concentrated acid, although I think there may be a little error in the calculation.

Azeotropic (98%) Sulphuric Acid has a density of 1.84g/mL so I would say you have overshot a little, but even so, going off that density reading I would say you have at the very least 97%.

hyfalcon - 20-3-2015 at 04:36

Small amount of dissolved SO₃ would throw this off wouldn't it?

Trizocy - 20-3-2015 at 07:55

Hello guys!

Fulmen:
I aldready know that, but here is the problem, my calculation ending up at 1,855g/ml but the highest concentration of H2SO4 is 1,84g/ml so something is wrong here.

TheAustralianScientist:
How can i do a overshoot when i got a pycnometer?

?

Hyfalcon:
May u deepen more about u talking about?

Thanks all!

Sulaiman - 20-3-2015 at 08:54

Trizocy,
Are the weighing scales calibrated ?

Trizocy - 20-3-2015 at 09:48

Nope

Volanschemia - 20-3-2015 at 18:14

Remember the pycnometer is only instrument 1 of 2 in the density measuring process. The scales are equally important, so if the scales are off, your calculation of density will also be off.

Calibration weights are fairly cheap, you just need to get the one that your scale requires.

papaya - 21-3-2015 at 02:56

Also density data are given for specific temperature and your picnometer volume depends on temperature as well, also impurities in sulfuric acid can affect density, since 1.84 is given for azeo H2SO4, which must contain only water.. Better to titrate it, but here you'll come to primary standart preparation problem, yes analytic work is hard to do at home.

Fulmen - 21-3-2015 at 05:14

Start by testing a few known chemicals like water, ethanol etc, that will give you an idea of the accuracy.