I would like to produce/make hydrochloric acid: HCl
- I mixed NaCl and H2SO4 (sodium chloride and sulphuric acid) and got HCL, but
is the formula:
NaCl + H2SO4 - > HCL,
NaCl + H2SO4 - > HCL + HSO4,
NaCl + H2SO4 - > HCL + NaHSO4
or something else and why?
Since I've used first: NaCl and then mixed it with H2SO4, I think it is: NaCl + H2SO4 and because I got: HCl:
NaCl + H2SO4 => HCL, but I am unsure.
I'm trying to learn chemistry alone
PS:
I mixed table salt and hydrogen sulfate with 20 ml of water in the bottle, I turned on the heat, it started to boil, etc. ..
My whole point is how I describe the way I've worked with - in formulas.
Is: NaCl + H2SO4 => HCL , right or wrong and why?1281371269 - 3-3-2011 at 09:56
That can't be right because the two sides don't add up.
NaCl(s) + H2SO4(l) --> HCl(g) + NaHSO4(s)
But that's the easy bit. You need a pretty good setup to actually bubble the gas through water without poisoning yourself.
Edit: Missed out an (s)
[Edited on 3-3-2011 by Mossydie]Mixell - 3-3-2011 at 09:57
Please balance your equations...
The formula is 2NaCl + H2SO4--> Na2SO4 +2HCl.
Although it probably depends of the molar ratios and a few other factors, the reaction listed above mine is also correct.
[Edited on 3-3-2011 by Mixell]hissingnoise - 3-3-2011 at 09:59
The products are NaHSO<sub>4</sub> and HCl
H<sub>2</sub>SO<sub>4</sub> + NaCl--->NaHSO<sub>4</sub> + HCl.
Na<sub>2</sub>SO<sub>4</sub> is only formed by strong heating!
The gas should be bubbled into water using an inverted funnel to avoid suck-back . . .
m4unot - 3-3-2011 at 10:07
Many thanks for all your help friends, but lacks explanation.
and Mr. hissingnoise says: H2SO4 + NaCl--->NaHSO4 + HCl.
I have no idea who is right and who is wrong...hissingnoise - 3-3-2011 at 10:42
Hmmm, All of the above . . .
The reaction can be forced to completion by sustained, strong heating but it's not used in practice for obvious reasons!
peach - 3-3-2011 at 10:54
Mixell is assuming you're going to heat it through to completion.
Hissingnoise and Mossydie are giving you the realistic version of it being run cold. I haven't seen anyone here bothering to heat theirs when
generating it, as it's a lot of effort for not a lot more when table salt is tens of pence per kilo.
You won't get concentrated hydrochloric if any of your hydrogen chloride escapes the water at the end before being absorbed. And hydrogen chloride be
more prone to escaping the higher the concentration gets; as it's solubility decreases.
[Edited on 3-3-2011 by peach]Mixell - 3-3-2011 at 11:12
Well, heating it to completion doesn't decrease the amount of NaCl you use, but the amount of H2SO4 that is consumed in the process.
But yes, heating it isn't a very wise thing to do if you want to make hydrochloric acid (although preparing it in the first place doesn't worth the
efforts, you can buy it in basically any DIY store). m4unot - 3-3-2011 at 11:20
Mossydie says: NaCl + H2SO4--> HCl + NaHSO4
- Since I've used first: NaCl and then mixed it with H2SO4, i agree with: NaCl + H2SO4, and since
the result is or expects to be: HCL
so therefore I think the formula is: NaCl + H2SO4--> HCl, but why HCl + NaHSO4 ?
Can I get an easy explanation.hissingnoise - 3-3-2011 at 11:31
The reactants should be dry; that is ~98% H<sub>2</sub>SO<sub>4</sub> and dry NaCl.
They react immediately and when the reaction slackens some heating is necessary to convert all NaCl to NaHSO<sub>4</sub>.
Ice-cold water is best for absorption . . .
m4unot - 3-3-2011 at 12:15
To everyone who worked on this Conversation and to everyone who participated in it: THANK YOU. It was a pleasure for me.
so therefore I think the formula is: NaCl + H2SO4--> HCl
In that equation, you've just lost a significant number of the atoms on the right to 'nowhere'. 1 sodium, 1 hydrogen, 1 sulphur and 4 oxygens are
missing from the product side.
They don't go missing, they form into something else. The bisulphate.
The sulphuric is displacing the chlorides on the sodium.
Quote:
but why HCl + NaHSO4
Because what's left of the sulphuric and salt once the chloride has left the salt and one of the hydrogens has left the sulphuric has to go somewhere.
One of the first things you learn in chemistry is that equations need balancing. Some of that can get complicated, but with NaCl + H2SO4--> HCl,
you should easily be able to see by glancing at it that a significant percentage of what you've got on the left is missing from the right; whole
elements are disappearing, e.g. the sodium, sulphur and oxygen.
Quote:
Unbalanced Chemistry Equation:
NaCl + H2SO4 => HCL + Na2SO4
balanced Chemistry Equation:
NaCl + H2SO4 => HCL + NaHSO4
Whilst the top one is unbalanced, it's also for a different reaction.
The second one occurs first, at room temperature. The top one occurs once the mixture is heated, and it occurs via an intermediate.
Cold: NaCl + H2SO4 => HCL + NaHSO4
Then apply heat: NaCl + NaHSO4 → HCl + Na2SO4
[Edited on 3-3-2011 by peach]m4unot - 3-3-2011 at 12:31
Very beautiful research, Thank's peach.hissingnoise - 3-3-2011 at 13:08
In the reaction; H<sub>2</sub>SO<sub>4</sub> + NaNO<sub>3</sub>---> HNO<sub>3</sub> +
NaHSO<sub>4</sub>.
Temperatures routinely reach 120°C and only the bisulphate forms . . .
Mixell - 3-3-2011 at 14:55
Well, not exactly, dinitrogen pent-oxide (nitric acid anhydride) is also formed, some of it attracts water molecules in the gas phase I assume, and
some attracts water molecules at the condensation phase (its also important to mention that some of the dinitrogen pent-oxide decomposes to oxygen and
the quite toxic nitrogen dioxide gas, so the reaction should be performed with a gas tight apparatus or in a ventilated location.
But again, nitric acid can be easily purchased in most of the countries, so performing this reaction only for the sake of receiving nitric acid is
quite useless (and dangerous!), unless you require concentrated nitric acid (95%+ can be reached without too much effort).
Also, nitrogen dioxide anesthetizes the nose even on low concentrations, so if you perform the reaction and you stop smelling the harmful gases, it
doesn't mean that they are gone. Bot0nist - 3-3-2011 at 16:54
@ Mixell
I tried the ole _NO<sub>3</sub> + H<sub>2</sub>SO<sub>4</sub> route a while back. Not fun. I couldn't get any
condensation and a little leak gave me a small taste of orange. Not fun indeed.
I'll stick to my 68%. Even at the prices.
[Edited on 4-3-2011 by Bot0nist]
[Edited on 4-3-2011 by Bot0nist]m4unot - 3-3-2011 at 23:49
People often mention Sir Isaac Newton's laws and Einstein, but forget the father of chemistry, Jaber bin
Hayyan, Ja'far al-Sadiq who was known for many physics and chemistry discoveries and Al-Khwarizmi the
father of Algebra, which is used in mathematics / chemistry equations .hissingnoise - 4-3-2011 at 03:36
Well, not exactly, dinitrogen pent-oxide (nitric acid anhydride) is also formed.
We've drifted a little?
Anyway, N<sub>2</sub>O<sub>5</sub> cannot exist in the presence of water.
In distillation, it is HNO<sub>3</sub> that decomposes to produce NO<sub>2</sub>.
The rest of your post too, sounds a bit mixed-up!
Mixell - 4-3-2011 at 04:13
Almost all of the N2O5 turns into nitric acid (not immediately), but some decomposes too.
My first try at distilling nitric acid failed because I was using rubber stoppers (due to the lack of a proper distillation kit), in the 5 minutes the
reaction ran, the dinitrogen pentoxide oxidized a 0.2 mm layer of the rubber stopper (which isn't attacked by nitric acid).
The N2O5 exists in the gas phase when a production of nitric acid via the nitrate salts and sulfuric acid occurs, when there is a lack of water vapors
due to the dehydration quality of sulfuric acid.
P.S. - I'm a bit drunk at the moment of writing this, so things may not be so clear to the sober individual.
And about the mixed-up part, can be, but this is how I write...trezza - 4-3-2011 at 04:18
in the 5 minutes the reaction ran, the dinitrogen pentoxide oxidized a 0.2 mm layer of the rubber stopper (which isn't attacked by nitric acid).
Wow! You're well on the way to rewriting the entire lit. on chemical physics!
And I was commenting on content rather than style . . .
Mixell - 4-3-2011 at 04:33
Might be =)
But its a name I used in a video game 6 years ago, which is currently my nickname in basically everywhere.
But this is already drifting away from the topic.
And about the distillation of nitric acid, if the author decides to distill nitric acid, please adjust the heat very carefully and heat the vessel
gradually, because you cannot imagine what hell will break loose if you over heat the system (something I learned from a mistake caused by the lack of
patience...).
Edit- I'm not a big fan of sarcasm, and what was the problem with the content in my post, a bit of interesting information regarding safety isn't
completely useless I assume.
[Edited on 4-3-2011 by Mixell]hissingnoise - 4-3-2011 at 04:39
Look Mixell, the fact that rubber is immediately and aggressively attacked by HNO<sub>3</sub> is something everyone who's ever handled
HNO<sub>3</sub> knows by heart!
Mixell - 4-3-2011 at 04:45
I don't know if it was rubber, but in fact I tried to use this stopper with nitric acid (68%) and it worked fine, no visible corrosion was detected.
Just tried it again, it doesn't react with the nitric acid, so it must be some other material and not rubber.
[Edited on 4-3-2011 by Mixell]hissingnoise - 4-3-2011 at 04:48
Quote:
I'm not a big fan of sarcasm . . .
Don't invite it, then!
Mixell - 4-3-2011 at 04:54
Well, then sorry for confusing the name of the material my stopper is made of, but the point is its not rubber and it was aggressively attacked in the
process of distilling nitric acid (presumably by N2O5).
And about the sarcasm part, I prefer first to explain what is wrong and only if it doesn't help- use sarcasm and stuff...
Anyway, going to sleep for a bit, drinking too much beer in the noon is not a wise thing to do.
[Edited on 4-3-2011 by Mixell]hissingnoise - 4-3-2011 at 05:04
Well OK! But what part of; N<sub>2</sub>O<sub>5</sub> cannot exist in the presence of water, are you having difficulty
with?
[edit] Jeeez! I see my post-count has reached 25 . . .
What the fuck am I doing?
[Edited on 4-3-2011 by hissingnoise]blogfast25 - 4-3-2011 at 09:07
Not sure either, HN… Going by Mixell’s earlier posts I’m still fully expecting him to mix XeF4 with AuCl3 in, I dunno… acetone (whatever)?
Just ‘to see what happens!’… He might complain about the stopper being attacked too! Nought queerer than folk...Mixell - 4-3-2011 at 09:12
Oh, never mind, not going to argue about that anymore...
The only thing I'm saying is that N2O5 can exist where there is a phase with a lack of water vapors, just like what happened in my distillation
attempt.
That is all.
[Edited on 4-3-2011 by Mixell]chem101 - 12-9-2011 at 01:24
Hi, can you please answer my question.
"Hydrogen Chloride, HCl, may be prepared in the laboratory by heating Concentrated Sulphuric Acid, H2SO4, with Sodium Chloride, NaCl."
What happens if we use dilute H2SO4 instead???
bbartlog - 12-9-2011 at 05:22
If you use dilute H2SO4, then you end up with HCl in aqueous solution, rather than gaseous HCl being produced. If by 'dilute' you mean 35% or 50%
rather than 95%, this would be a reasonable way of getting the aqueous azeotrope on distillation, i.e. you could get 20% (or was it 21%?) HCl this
way.
If you use more dilute H2SO4 than that, then you end up with relatively dilute HCl and attempts at distillation would have to drive off some water
before you got to the azeotrope.
But the passage you quote is discussing the production of HCl gas, not aqueous solutions, and for that you would want concentrated H2SO4.
